Third World Congress on Oxidation Catalysis

Studies in Surface Science and Catalysis 110 3rd WORLD CONGRESS ON OXIDATION CATALYSIS

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S t u d i e s in S u r f a c e S c i e n c e a n d C a t a l y s i s Advisory Editors: B. Delmon and J.T. Yates

Vol. 110

3rd WORLD CONGRESS ON OXIDATION CATALYSIS Proceedings of the 3rd World Congress on Oxidation Catalysis, San Diego, CA, U.S.A., 21-26 September 1997

Editors R.K. Grasselli

University of Delaware, Newark, DE, U.S.A.

S.T. O y a m a

Virginia Polytechnic Institute, Blacksburg, VA, U.S.A.

A.M. Gaffney

ARCO Chemical Company, Newton Square, PA, U.S.A. J.E. Lyons

Sun Corporation, Marcus Hook, PA, U.S.A.

1997 ELSEVIER Amsterdam - - Lausanne m New York--- Oxford m Shannon m Singapore m Tokyo

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

ISBN 0-444-82772-2 91997 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

3rd World Congress on Oxidation Catalysis - Atom Efficient Catalytic Oxidations for Global Technologies

PREFACE

The 3rd World Congress on Oxidation Catalysis has its roots in the European Workshop on Selective Oxidation held in Louvain, Belgium in 1985. Out of this workshop grew the 1st World Congress held in Rimini, Italy in 1989, the 2nd in Benalmadena, Spain in 1992, and the 3rd being held now in 1997 in San Diego, California, USA. Out of the small core of dedicated and enthusiastic scientists assembled in Louvain in 1985, grew now a broad base of scientists and technologists from academia, industry and government laboratories who are fervently pursuing the subject of oxidation catalysis and are eager and willing to exchange their findings at the current meeting. The overall theme of the 3rd World Congress is "Atom Efficient Catalytic Oxidations for Global Technologies". We chose this theme to stimulate the participants to report their findings with an emphasis on conserving valuable material in their catalytic transformations, as well as conserving energy, and that in an environmentally responsible manner. Progress towards this stated goal is substantial as evidenced by the tremendous response of our community in their participation of quality publications compiled in these Proceedings of the Congress. The subjects presented span a wide range of oxidation reactions and catalysts. These include the currently important area of lower alkane oxidation to the corresponding olefins, unsaturated aldehydes, acids and nitriles. In this manner, the abundant and less expensive alkanes replace the less abundant and more expensive olefins as starting materials for industrially important intermediates and chemicals. In the oxidative activation of methane the emphasis is shifting towards the use of extremely short contact times and newer more rugged catalysts. In the area of olefin oxidations, of particular note are the high efficiency epoxidation of propylene, and new detailed mechanistic insights into the oxidation of ct,~-unsaturated aldehydes to the corresponding unsaturated acids. Substantial progress is reported in the area of the selective oxidation and ammoxidation of substituted aromatics and heteroaromatics. These include higher yields of desired products, new and more durable catalysts, as well as a reduction of undesirable byproducts. The use of oxidation catalysis to produce fine chemicals is experiencing an explosive growth. A plethora of novel approaches are presented which include shape selective epoxidations. Oxidation in confined structures is coming out of its infancy and the use of TS-'I as catalyst is becoming a standard. New approaches are being presented invoking the shape selective character of the nano-environment of peroxytungstates anchored on selected supports. The important areas of combustion, engineering and environmental aspects of oxidation catalysis, as well as their theoretical, computational and modeling approaches round off the program. Not to be overlooked is perhaps the most ambitious, the subject of structure selectivity/activity correlations, an area always worthy of further attention and in depth study. A noble effort thereof has been put forward and is reported here. We are coming ever closer to the ultimate goal of a rational approach to catalyst design and synthesis. There is still ample room for further efforts towards this goal, but the foundations are being formed for a bright and rewarding future of rationally predictive oxidation catalysis. The five featured lectures and seven plenary lectures constitute the general background and overview of the subject matter at hand. The 104 contributed papers and 13 poster manuscripts, summarized in this compendium, probe new avenues to achieve catalytically efficient oxidation reactions for the future needs of mankind in a global

environment. A large number of countries responded to this challenge by their representatives giving oral presentations or posters, and in particular by supplying the written scientific documents contained in this volume. Our sincere thanks go out to all of the contributors. The countries participating in the Congress and contributing to the Proceedings reported here made the 3rd World Congress on Oxidation Catalysis a truly international event, they are: Argentina, Azerbaijan, Belgium, Brazil, Bulgaria, Canada, China, Czech Republic, Finland, France, Germany, Greece, India, Iran, Ireland, Israel, Italy, Japan, Korea, Latvia, Netherlands, New Zealand, Poland, Portugal, Russia, Saudi Arabia, Singapore, Slovenia, South Africa, Spain, Sweden, Taiwan, Thailand, Ukraine, United Kingdom, and United States. We conclude on the basis of the foregoing, that the future of oxidation catalysis is secure and has never been brighter than at this juncture. We are confident, and the Proceedings support this view, that many new and improved selective oxidation processes and catalysts will be discovered and commercialized over the next decade, and that our knowledge towards a rational design of selective oxidation catalysts is within our grasp. With this optimism we look confidently towards the future and to a successful 4th World Congress on Oxidation Catalysis in the next millennium, in the year 2001. Thank you all for partaking in our Congress and for working in an exciting and promising area of catalysis. May success come your way in abundance in the coming years.

Robert K. Grasselli S. Ted Oyama Anne M. Gaffney James E. Lyons

vii

TABLE OF CONTENTS

Preface R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons Featured Lectures F-1

Molecular Mechanism of Heterogeneous Oxidation-Organic and Sofid State Chemists' Views J. Haber

F-2

The Multifunctional Properties of Heterogeneous Catalysts, Active and Selective in the Oxidation of Light Paraffins F. Cavani and F. Trifirb

19

Selective Oxidation of Hydrocarbons Catalyzed by Heteropoly Compounds M. Misono, N. Mizuno, K. Inumaru, G. Koyano and X. H. Lu

35

The Future of Industrial Oxidation Catalysis Spurred by Fundamental Advances B. Delmon

43

F-3

F-4

Plenary Lectures P-1

P-2

P-3

P-4

Molecular Approach to Active Sites on Metallic Oxides for Partial Oxidation Reactions J.C. V6drine

61

In Situ Electrochemically Controlled Promotion of Complete and Partial Oxidation Catalysts C.G. Vayenas and S.I. Bebelis

77

Reductive and Oxidative Activation of Oxygen for Selective Oxygenation of Hydrocarbons K. Otsuka

93

The Selective Oxidation of Methanol: A Comparison of the Mode of Action of Metal and Oxide Catalysts D. Herein, H. Werner, Th. SchedeI-Niedrig, Th. Neisius, A. Nagy, S. Bernd and R. Schl5gl 103

viii

P-5

P-6

P-7

Gold as a Low-Temperature Oxidation Catalyst: Factors Controlling Activity and Selectivity M. Haruta

123

The Selective Epoxidation of Non-Allylic Olefins over Supported Silver Catalysts J.R. Monnier

135

Redox Molecular Sieves as Heterogeneous Catalysts for Liquid Phase Oxidations R.A. Sheldon

151

PartA Structure Selectivity/Activity Correlation A-1

Synergistic Effects in Mulficomponent Catalysts for Selective Oxidation P. Courtine and E. Bordes

177

A-2

Synergetic Effects Promoted by in Operandi Surface Reconstructions of Oxides E.M. Gaigneaux, J. Naud, P. Ruiz and B. Delmon

185

Further Study on the Synergetic Effects between MoO3 And Sn02 E.M. Gaigneaux, S.R.G. Carrazan, L. Ghenne, A. Moulard, U. Roland, P. Ruiz and B. Delmon

197

The Nature of the Active/Selective Phase in VPO Catalysts and the Kinetics of n-Butane Oxidation D. Dowell and J.T. Gleaves

199

Understanding the Microstructural Transformation Mechanism which Takes Place During the Activation of Vanadium Phosphorus Oxide Catalysts G.J. Hutchings, A. Burrows, S. Sajip, C.J. Kiely, K.E. Bere, J.C. Volta, A. Tuel and M. Abon

209

Structural and Catalytic Aspects of Some Nasicon - Based Mixed Metal Phosphates P.A. Agaskar, R.K. Grasselli, D.J. Buttrey and B. White

219

Selective Reactivity of Oxygen Adatoms on Mo(112) for Methanol Oxidation K. Fukui, K. Motoda and Y. Iwasawa

227

A-3

A-4

A-5

A-6

A-7

A-8

A-9

A-10

Mechanistic Studies of Alkane Partial Oxidation Reactions on Nickel Oxide by Modern Surface Science Techniques N.R. Gleason and F. Zaera

235

Structure and Catalysis of LixNi2_x02 Oxide Systems for Oxidative Coupling of Methane T. Miyazaki, T. Doi, T. Miyamae and I. Matsuura

245

Reaction Induced Spreading of Metal Oxides: in situ Raman Spectroscopic Studies During Oxidation Reactions Y. Cai, C-B. Wang and I.E. Wachs

255

A-11

Temperature Programmed Desorpfion of Ethylene and Acetaldehyde on Uranium Oxides. Evidence of Furan Formation from Ethylene 265 H. Madhavaram and H. Idriss

A-12

Active Sites of Vanadium-Molybdenum-Containing Catalyst for Allyl Alcohol Oxidation: ESR Study in situ O.V. Krylov, N.T. Tai and B.V. Rozentuller

275

Lower Alkane Oxidation

A-13

Oxidative Dehydrogenation of Ethane over Vanadium and Niobium Oxides Supported Catalysts P. Ciambelli, L. Lisi, G. Ruoppolo, G. Russo and J.C. Volta

285

A-14

Partial Oxidation of Ethane over Monolayers of Vanadium Oxide. Effect of the Support and Surface Coverage. M.A. BaSares, X. Gao, J.L.G. Fierro and I.E. Wachs 295

A-15

The Ethane Oxidative Chlorination Process and Efficient Catalyst for It M.R. Flid, 1.1. Kurlyandskaya, Yu.A. Treger and T.D. Guzhnovskaya

305

A-16

Oxidative Conversion of LPG to Olefins with Mixed Oxide Catalysts: Surface Chemistry and Reactions Network M.V. Landau, M.L. Kaliya, A. Gutman, L.O. Kogan, M. Herskowitz and P.F. van den Oosterkamp

315

Free Radicals as Intermediates in Oxidative Transformations of Lower Alkanes M.Yu. Sinev, L.Ya. Margolis, V.Yu. Bychkov and V.N. Korchak

327

A-17

A-18

Alternative Methods to Prepare and Modify Vanadium-Phosphorus Catalysts for Selective Oxidation of Hydrocarbons V.A. Zazhigalov, J. Haber, J. Stoch, A.I. Kharlamov, I.V. Bacherikova and L.V. Bogutskaya

337

Active Species and Working Mechanism of Silica Supported MoO3 and V205 Catalysts in the Selective Oxidation of Light Alkanes A. Parmaliana, F. Arena, F. Frusteri, G. Martra, S. Coluccia and V. Sokolovskii

347

Mechanistic Aspects of Propane Oxidation over Ni-Co-Molybdate Catalysts D.L. Stern and R.K. Grasselli

357

Oxidative Dehydrogenation of Propane by Non-Stoichiometric Nickel Molybdates D. Levin and J.Y. Ying

367

Selective Oxidation of Propane into Oxygenated Compounds over Promoted Nickel-Molybdenum Catalysts J. Barrault, C. Batiot, L. Magaud and M. Ganne

375

Oxidative Dehydrogenation of Propane on CeNixOy (0 ~ x due to the diffusion barrier created by the pyrophosphate groups is clearly shown by these arrows (from [ 12]). 5. IRON PHOSPHATES AND HYDROXYPHOSPHATES [16-19] Such catalysts appeared to be potentially important catalysts for the oxidative dehydrogenation of isobutyric acid (IBA) to methacrylic acid (MAA). This reaction is important as a first step to form methyl methacrylate monomer used for plexiglass or altuglass formation by polymerisation. The industrial type catalyst contains iron hydroxyphosphate of uncertain nature and Cs or NH4 as additive and unfortunately necessitates a large amount of water in the feed (namely 10 to 12 mol. H20 per tool. of IBA) to remain stable with time on steam. This makes this process difficult to be develop ,ed industrially. Taking into account the ph ase diagramme Fe203-FeO-P205 it could be possible to select and study several phases wt~ch contain Fe 2 , Fe 3+ or both catio,,s able to insure the redox mechanism necessary for the reaction to take place as it was shown to proceed via a Mars and van Krevelen mechanism [20]. It was also shown that the active phase was a phosphate with iron ion at a two oxidation states +2 and +3, ctFe3(P207)2, composed of trimers of face sharing FeO6 octahedra (as schematized in fig. 3a) separated one from the other by PO4 tetrahedra. The active sites were shown to correspond to a group of two trimers facing each other on their respective layers. The r61e of redox couple and of hydroxylation is shown below:

H20 02 Fe23+Fe2+(P207)2 ~---~ Fe23+Fe 2+ (PO3OH)4 > Fe2+x3+Fel-x2+(PO3OH)4-x(PO4)x 400oC The way the octahedra are connected to each other is an important parameter since it has been shown that the other polymorphic form of Fe3(P207)2 namely J3Fe3(P207)2 was poorly active and selective. In this phase, the central iron atom is in fact occupying a FeO6 trigonal prism (fig. 3b). A study of many different iron hydroxyphosphates exhibiting clusters of face sharing FeO6 octahedra of different sizes has been carried out. It turned out that all catalysts were belonging after catalytic testing to the same solid solution of the type Fe3+4.~Fe2+3x(PO4)3(OH)3-

67

3xO3xwith 0_<x _3

Fig. 4. Electron exchange between iron cations in (FeO6)n octahedra clusters calculated by extended Huckel molecular orbital theory (from [21] and [22]. 6. SUPPORTED OXIDE CATALYSTS Supports are very often used in catalysis for several reasons, namely: i. Dispersion of the active phase in order to increase the surface to volume ratio since heterogeneous catalysis is occuring at the solid surface. ii. Heat transfer: this is a particularly important aspect in oxidation reactions because of the high exothermicity involved. This holds particularly true at industrial scale since the hot spot problem is very crucial and should be monitored with precision within a few degrees to avoid local overactivity and subsequently overheating with presumably irreversible phase transformation of the catalyst. iii. Attrition this aspect is of main importance for fluidized or solid transported beds. One very usually uses silica or carborandum as a binder or as a coating to limitate attrition or to facilitate heat transfer respectively. iv. Formation of new catalytic sites this point will be emphasised below in some examples. One will see that well dispersed species of limited size could be formed and exhibit peculiar catalytic properties. v. Modification of the active phase properties due to its chemical interaction with the support, including epitaxial induced modifications. In some cases the chemical effect of the support will be determining for catalytic properties. For example several oxides (V205, MOO3, RezO7, Cr/O3...) deposited on several oxide supports (SiO2, TiO2, AIzO3) were shown by I. Wachs et al [25] to exhibit very different catalytic properties in methanol oxidation reactions as summarized in table 2.

69

Table 2 Turn over number values in (S)"1 for methanol reaction at 230~ deposited on several supports (from ref. [25]). Oxide support Supported oxide V205 MoO3 CrO3 SiOz 2 39 160 A1203 20 2 Nb205 700 32 58 TiO2 1800 310 300 ZrO2 2300 92 1300 6.1. M o l y b d e n u m

on lwt % metallic oxide

Re207 20 12 1200 170

oxide s u p p o r t e d on silica

Different procedures may be used to prepare such samples as impregnation of silica with a molybdate salt or grafting molybdenum chloride or molybdenum based organo metallic compounds as Mo carbonyls on the hydroxyl groups of silica or at last solid-solid reaction between MoO3 and SiO2 at temperatures near or above 500~ A parameter important for the impregnation method is the pH of the molybdate solution. As a matter of fact the following equilibrium has been well established: MO70246" + 4H20 6-~ 7 M0042 + 8H +. The monomeric tetrahedral MoO42 species is favored at high pH and vice versa for the polymeric heptamolybdate anion. In a study of ammonium heptamolybdate impregnated silica [26], the molybdenum loading was varied up to about 20 wt%. The hydroxyl groups of the silica were observed by infrared and/or UV vis-NIR spectroscopies to decrease with Mo loading and to disappear at 7 wt% Mo loading while two UV bands were appearing at 245 and 340 nm. The former band may be assigned to tetrahedral MoO 24 monomeric species and the latter one to octahedral polymeric (polymolybdate) species. The former band was observed to increase in intensity proportionally to Mo loading at low Mo content and to saturate at roughly 3 wt% Mo. The latter band was observed to appear near 2 wt% Mo, to increase then proportionally to Mo content up to 7 wt% Mo and then to decrease slowly with Mo loading. The formation of the former species was described by: \ / /Si x I

2

-SiI

O

OH

+ MOO]- +2H +

--)

/. O

\ Mo "/ / \ /0 \0 / Si \

+ HzO

Catalytic properties were studied for two reactions namely isopropanol conversion and propene partial oxidation. The first reaction is a test reaction which allows to characterize acidic, basic or redox properties of a catalyst. One gets dehydration to propene or diisopropylether for acid catalyst, acetone for basic catalyst in absence of air and acetone and water for redox type catalyst in presence of air. The experimental results at 100~ clearly show that at low Mo loadings acidic features are favored while redox features are favored at higher loadings. This indicates that monomeric MOO]- species are acidic (presumably as in silicomolybdic acid) while polymeric species exhibit redox properties.

70 The second reaction studied is propene oxidation at 380~ Weak activity was observed for low Mo loading. At higher Mo loading, propanal was the major product while acrolein was also observed. At very high Mo loadings and for MoO3 one gets almost exclusively acrolein. Propanal is known to stem from propene by an electrophilic attack while acrolein corresponds rather to a nucleophilic attack. These results indicate that monomeric MoO~- species does not oxidize propene in our conditions while polymeric (polymolybdate) species exhibit redox properties, with O 2 species being rather electrophilic. Molybdenum oxide exhibits redox properties with 02. species being rather nucleophilic. One can thus realize how the size of the active site is important in oxidation catalysis. This is typical of structure sensitive reaction. 6.2. Vanadium oxide on TiO2 support

Such a catalyst is well known for several reactions, such as o-xylene oxidation to phthalic anhydride and selective catalytic reduction (SCR) of NO by ammonia. The anatase form of TiO2 appears to be better than the rutile form. Such catalysts with 1 and 8 wt% V2Oflanatase was prepared by Rh6ne-Poulenc (S ~ 10 m2 g-l) for an exercise of characterization by 25 different european laboratories. All results are assembled in one issue of Catalysis Today published in May 1994, vol. 20 n~ Surface vanadium species were observed to exist in three different forms: monomeric VO43 species, polymeric vanadate species and V205 crystallites [27], the relative amount of which depended on initial wt % V2OJanatase and on the subsequent selective dissolution treatment. The following conclusions could be drawn considering that a polyvanadate species occupies a ca circular zone of a diameter of 0.38 nm per V atom (i.e. 0.165 nm2) and an isolated monovanadate species 0.66 nm per V atom (i.e. 0.43 nm2): i. Monomeric VO43 species exhibit acidic character with OH groups (Br6nsted acidity) and result in propene formation for isopropanol conversion and in total oxidation for o-xylene oxidation. Maximum coverage equals 0.43 wt% V205. ii. Polyvanadate surface species exhibit redox properties, namely give rise to acetone for isopropanol conversion and phthalic anhydride for o-xylene oxidation. iii V205 crystallites exhibit low activity for o-xylene conversion and high selectivity in total oxidation. Here too the size of the VOx species is of primary importance for the catalytic behaviour. 6.3. Niobium oxide species on silica support [28]

Deposition of NbOx species of different sizes on a silica support has been performed starting from organo metallic complexes as Nb(~I3-C3Hs)4 for a monomer, [Nb(qs-CsHs)H-ta(~5, ~1C5i_h)]2 for a dimer and Nb(OC2Hs)5 for a monolayer. All species were characterized by EXAFS, XANES, FT-I~ Raman, XPS and ESR techniques. Catalytic performances for conversion of ethanol are summarized in table 3. The structure of the species was well characterized particularly the NbO, NbNb, NbSi bonds lengths and coordination numbers. The catalytic data show clearly that dehydrogenation occurs majoritarily on monomeric species and dehydration on monolayer (bidimensional) structure for ethene (intra molecular reaction) and at last on dimeric species for both intra and inter molecular dehydration reaction.

71 Table 3 Catalytic performances of Nb monomeric, dimeric or monolayered species NbOx on SiO2 support for dehydrogenation (---> AA) and dehydration (intra ---> E, inter ~ DE) of ethanol (from ref. 29). Initial rate mmol. min~ g(Nb) ~ Selectivity % Catalyst Total AA E + DE AA E DE 1.25 1.2 0.049 96.1 2.8 1.1 monomer* 0.18 0.004 0.176 dimer* 2.1 24.2 73.7 0.11 0.001 0.106 monolayer** 0.9 99.1 0.0 0.17 0.052 0.118 impregnated* 30.5 20.2 49.3 0.0026 0.0004 0.0022 14.8 46.9 38.3 Nb205 bulk** AA acetaldehyde, E Ethene, DE Diethylether Reaction: Ethanol 3.1 kPa *at 523 K **at 573 K 6.4. Various oxides deposited on different oxide supports The idea is here to determine how a support may modify the catalytic properties of the different oxide species (as inorganic clusters of different size). It has been described above how the size of the deposited oxide species (monomefic, polymeric, bulk-type) results in different catalytic properties. Vanadium oxide has been deposited on several supports as silica, alumina, titania but also zirconia, niobia, zirconium hydroxyphosphates, etc. Its reductibility was studied by reduction with hydrogen at 400~ It was observed that V on TiO2 and yA1203 was reduced rather fast and one reached a consumption such as an O/V atomic ratio near to 1 for TiO2 and 0.65 for A1203 [30]. At variance for silica the reduction was much slower and an O/V ratio of 0.57 was obtained. Such differences were interpreted by the authors as due to different species on the surface: mainly monomefic VO43 species for TiO2, dimefic species V2074"for y.AI203 [30] and V205 crystallites for silica [31 ]. The size of the polyvanadate species in solution corresponds to V3093, V40124"with tetrahedrally coordinated V at pH = 7 and 4.5 respectively and V100286 or VI0028H 5 with octahedrally coordinated V at pH = 2.5. Deposition of such polyvanadates of different sizes was performed carefully on yA1203 support and their initial structures despicted above were shown to remain stable even after calcination in flowing air at 500~ and to be only partly modified after catalytic reaction of oxidative dehydrogenation of propane in the 350 to 450~ temperature range. Moreover in such cases, the selectivity towards propene was observed to be the same at the same conversion level, indicating that in this size domain of polyvanadate species the catalytic selectivity was not changed, only the activity was observed to increase with vanadium loading [32]. In a study of VO 2§ and Cr 3§ cations exchanged or impregnated on the surface of zirconium hydrogenophosphates and compared to vanadyl pyrophosphate and chromium (III) phosphate, the best catalytic properties for ethane to ethylene oxidative dehydrogenation reaction at 500~ were observed when the material present either continuous chains of vanadyl or CrO6 octahedra as in (VO)EP207and et CrPO4 respectively or chains of limited size [33]. By comparison with bulk oxides exhibiting similar environment, it was suggested as for supported VOx and CrOx species on zirconium hydrogenophosphates that the catal~ic performances were related to the stronger basic character ofPO4 3" anions with respect to O anions.

72 7. VMGO CATALYSTS [34-42] Such a system was interesting to consider since it is basically composed of an acidic (V205) and a basic (MgO) oxides. One may then expect either to have well defined phases of given catalytic properties or to have such phases deposited on a basic support (MgO). Three phases are well known namely the orthovanadate Mg3V2Os, the pyrovanadate Mg2V207 and the metavanadate MgV206. The first phase exhibits isolated VO4 tetrahedra (separated by MgO6 octahedra), the second one has comer sharing VO4 tetrahedra and the third one has an octahedral structure. The pyrovanadate phase was found [36] to be the more selective phase at equal conversion with respect to the other two phases for oxidative dehydrogenation of propane at 500~ At variance the Mg3V208 phase (exhibiting isolated VO4 entities separated by octahedral MgO6 entities) was found by H. Kung et al [35] to be more efficient for propene formation in propane oxidative dehydrogenation than the other two phases. A. Pantazidis and C. Mirodatos [37] found recently for the same reaction that propane selectivity was about the same (65 to 75 %) within a large range in V205 content (10 to 80 V205 wt %) and at similar conversion levels. A. Corma et al [36] concluded that tetrahedral isolated VO4 species are more selective for oxidative dehydrogenation of Ca and nC4 alkanes. It then appears than contradictory results have been obtained resulting obviously in contradictory interpretations concerning the role of atom arrangements. The discrepancies between authors may arise from differences between samples, supposedly similar, and from different catalytic reaction conditions particularly the alkane to oxygen ratio values. Taking the samples prepared in our group we have tried to go deeper in the characterization of the catalysts. Reducibly and electrical conductivity measurements [39] have shown that the MgEV207 phase was more easily reduced than the other two while anionic vacancies were created according to: C3H8 q- O2"surf > C3H6 (g) + H20 + [-7 + 2e. Moreover, basicity of surface oxygen ions acts on the ability to create anionic vacancies and thus on the catalytic activity. Moreover if the sample contains more MgO than the above stoichiometries for pure phases one may imagine that its basicity may also play an additional r61e. As a matter of fact propene as a base will be more easily eliminated from the surface of a basic catalyst. The acid-base features of the catalysts were studied by the reaction of isopropanol conversion to propene (acidic feature) and acetone (basic feature) under N2 in the feed and the redox features by the reaction under air in the feed. It was observed at 230~ (table 4 from ref 40) that the pyrovanadate sample was much more basic than the other two pure phases and that excess MgO with respect to crystallized phase stoichiometry induced even more basic character (table 4). As the olefins and to a lesser extent the alkanes are basic one may expect the desorption to be favored by surface basic sites. In other words oxidative dehydrogenation of alkanes is expected to be easier on surface exhibiting basic properties. As a matter of fact the results given in table 5 from ref. 41 show that Mg2V207 which is more basic as shown in table 4 is more selective for olefins in propane conversion and to a lesser extent for n-butane and isobutane oxidation reactions than the other two phases. Such a feature is even more pronounced for the samples with excess MgO at least for propane oxidation, samples which were also shown to present higher basicity (table 4).

73 Table 4 Catalytic data for isopropanol conversion under nitrogen (A) or air (B) in the flow at 230~ alter activation of the sample at 230~ for lh30 under N2 flow, values taken (from ref 40). S t2g"1

Samples

V205 (wt%)

n

MgVzO6 Mg2V207 Mg3V208 MgO 40VMgO" 60VMgO" V205

82.1 70.1 6O.8 0 38 58.5 100

0.1 1.7 0.9 1 J,2 43 19 2.1

Conversion % 7.5 1.5

2.2 0.2 1.2 3.6 64

m= 100 mg; flow rate 18.6 cm 3. min"l 9 Pi= 2.0 x 103 Pa

A Propene %

B

Acetone Conversion Propene % % %

90 43 69 0 14 24 94

10 57 31 100 86 76 6

8

1.7 2.8 0.2 1.7 4.3 4.8

90 40 74 0 17 24 75

Acetone % 10 60 26 100 80 76 25

" mixture ofMgO, Mg3V208 and otMgzV207 as determined by XRD analysis.

Table 5 Catalytic data for oxidative dehydrogenation of propane (A) n-butane (B) and isobutane (C) at 540~ (from ref 41).

Samples MgV206 Mg2V207 Mg3V208 40VMgO 60VMgO

Conversion % 7.4 6.9 8.3 8 8

A Propene" % 15 53 6 65 65

B

COx' % 53 28 94 35 35

Conversion % 10 11

15 15c 42 c

C Dehydr. Conversion i butene % products b 33 14 19 40 9 40 13 10 20 20 23 33 6 13 30

flow rate = 50 cm 3min1, P(C3--)=2 x 103pa, C3 or C4: air = 2:98 "" balance to 100 % corresponds to oxygenates as ethanal, acrolein, propanal and acetic acid. b. dehydrogenates correspond to butenes (major) (but l ene and but 2 ene) and butadiene (minor). c taken at 450~ instead of 540~ because of its high conversion level due to its high surface area. It follows that one may conclude that basic properties as well as atomic arrangements at molecular level play an important r61e in alkane oxidative dehydrogenation reactions. Such a conclusion could also be reached from the study of VMgO catalysts [42] for oxidative dehydrogenation of several alkanes as ethane, propane and butane under similar conditions (see e.g. fig. 5 in ref. 42).

74 8 CONCLUSION Some general conclusions may be drawn from this general presentation: i. Oxidation reactions in gas-phase heterogeneous catalysis usually proceed via Mars and van Krevelen mechanism i.e. involve lattice oxygen ions. Such ions exhibit an electrophilic or a nucleophilic character and therefore present different catalytic properties. As a matter of fact the electrophilic oxygen has been suggested to interact with a double bond or an aromatic ring and the nucleophilic oxygen to interact with a C-H bond in ~ of the double bond or of an aromatic ring. ii. Oxidation reactions are structure sensitive and therefore greatly depend on the local and surface structure of the oxide catalysts. A peculiar fitting between stereochemistry of the solid surface and of the reactant molecule(s) is to be obtained to get the best catalyst. Parameters such as reducibility and reoxidability features of the oxides are very important for catalytic reactions. iii. Active sites for oxidation reactions appear to be molecular "inorganic ensembles" of metallic oxide atoms whose size greatly influences the catalytic properties [2, 3, 43]. In some examples the number of atoms constituting the active sites could be established. For instance double trimers of face sharing FeO6 octahedra are particularly active and selective for oxidative dehydrogenation of isobutyric acid to methacrylic acid on iron hydroxyphosphates; ensembles of four dimers of VO6 octahedra are suggested to be the active sites for butane oxidation to maleic anhydride in vanadium pyrophosphate catalysts. Usually monomeric species as MoO~or VO43" exhibit acidic features and then total oxidation properties. At variance low size polymeric species exhibit better selectivity for many partial oxidation reactions than large size species or bulk-type oxide. It is thus obvious that the size of these r inorganic molecular ensembles )) as active sites is important depending on the reaction considered and the support used. The characterization of such ~ clusters )~ laying on the surface is rather difficult and the above statements for supported oxide catalysts are only qualitative. To better characterize such clusters, techniques as XANES (sensitive to local symmetry and oxidation state of transition metal cations), EXAFS (sensitive to the coordination number and nearby elements distances) and radial electron distribution fRED) of X ray diffraction peaks (sensitive to the nearby elements distances) have to be used. Already data have been obtained for MoO3 and V205 deposited on SiO2, Al203 or TiO2 support and for NbOx species on SiO2 (28, 29). This holds true only if one has one type of cluster species to analyse. The reader interested may read a chapter devoted to XANES and EXAFS techniques by B. Moraweck in the book by B. Imelik et J.C. Vrdrine (1988 and 1994) [44, 45]. iv. Many other examples may have been given. For instance heteropolyoxometallates as H~Mo~VO40 material or zeolite materials as titanosilicalite constitute other materials where the isolated or more or less condensed active species results in very different catalytic behaviour. The case of the titanosilicalite from ENI Co is typical since isolated Ti species are active for phenol oxidation to catechol and anthraquinone while condensed Ti species results in H202 decomposition. v. Oxidation catalysts have to be considered with a dynamical view under reaction conditions. This is related to the Mars and van Krevelen mechanism which involves a redox mechanism and also to the mobility of the oxide lattice. This dynamical phenomenon results in the wetting effect observed under catalytic reaction conditions for multicomponent and supported oxide catalysts [46, 47]. It follows that for many catalysts a certain time on stream is necessary before the catalyst reaches its steady state. It is frequent that in an industrial plant a

75 steady state is reached only after one or two hundreds hours, the catalysts lasting several years before having to be replaced. For simple catalysts as doped vanadyl pyrophosphates used for butane oxidation to maleic anhydride the fight size of the active sites (e.g. tetramers of vanadyl dimers) is monitored by the reactants in catalytic reaction conditions leading to the fight VS+/V4+ ratio, by the preparation procedure to change the material morphology (the (100) face of (VO)2P207 being developed), and by the adequate addition of additive elements which regulate the site size and VS+/V4+ ion ratio. The view of an oxidation catalyst as dynamical under catalytic reaction conditions is essential for our understanding of its functioning. REFERENCES

[1] J.C. Volta, W. Desquesnes, B. Moraweck and G. Coudurier, React. Kinet. Catal. Lett.,12 (1979) 241; J.C. Volta and J.L. Portefaix, Appl. Catal., 18 (1985) 1. [2] J.C. V6drine, J.M.M. Millet and J.C. Volta, Catal. Today, 32 (1996) 115. [3] J.C. V6drine, G. Coudurier and J.M.M. Millet, Catal. Today, 33 (1997) 3. [4] J.M. Tatibouet and J.E. Germain, C.R. Acad. Sci., 290 (1980) 321; J. Catal., 72 (1981) 37. [5] J. Haber, in R.A. Sheldon and R.A. van Santen, Editors, Catalytic oxidation, Principles and Applications, World Scientific, Singapore, 1995, p. 17. [6] J.M. Tatibouet, J.E. Germain and J.C. Volta, J. Catal., 82 (1983) 240. [7] M. Abon, B. Mingot, J. Massardier and J.C. Volta in g: ~" ~

PC2H4=0"65kPa

00_ ~..,~.-c~c_,5 ' r I0~ po 2/Pc2H4

n

15' o,

50

2u~0

Figure 3. Effect of gaseous composition on the open-circuit (unpromoted) catalytic rate of C2H4 oxidation on Pt/YSZ and on the electrochemically promoted catalytic rate when the Pt catalyst film is maintained at VWR = 1V (adapted from ref. [6]).

Figure 3 shows the steady-state effect of constant positive potential application VWR=(=+ 1V) on the rate of C2H4 oxidation on P t ~ S Z as a function of the PO2fPC2H4 ratio. The promotion is much more pronounced under oxidizing gaseous compositions, where a sixty-fold enhancement in catalytic rate and turnover frequency is obtained. The stronger electrochemical enhancement under oxidizing conditions is a general observation in NEMCA studies of oxidation reactions [ 14,15] and is due to the fact that the promoting ionic oxygen species forms, upon 02- supply, only after the coverage of normally adsorbed oxygen is near completion [25,26]. An example of using Na-lY'-AI203, a Na § conductor, as the solid electrolyte to induce NEMCA is shown in Fig. 4. The figure depicts the effect of catalyst potential VWR,

,v,

2O

s

Figure 4. Electrochemical promotion of Pt for CO

t.....

"

"

oxidation using Na-13"-AleO3: Effect of P c o , catalyst potential and corresponding linearized sodium coverage on the rate of CO oxidation at

-0.4

~. "'Q'

---.v.. ~ o.oz -"'i 0

5

T = 3 5 0 ~ and Po2=6 kPa (reprinted with permission from ref. [20]).

83 corresponding coulometrically measured [20,32] Na coverage ONa, and Pco on the rate of CO oxidation on Pt/Na-[Y'-A1203 at 350~ [20]. When the Pt surface is Na-free, the rate goes through a sharp maximum with respect to Pco, indicative of competitive adsorption of CO and O (Langmuir-Hinshelwood type kinetics). For high Pco values the surface coverage of O is very low and thus the catalytic rate is also low. Increasing Na coverage (via negative current or potential application) under these conditions causes a 6-fold enhancement in the rate of CO oxidation due to Na-assisted enhanced O chemisorption. The promotion index of Na, PNa, defined from: PNa=(Ar/ro)/AONa

(4)

takes values up to 200 under these conditions. Higher Na coverages (Fig. 4) poison the rate due to the formation of surface Na-CO complexes [20].

3.2. Selectivity modification: Ethylene epoxidation Electrochemical promotion can be used to modify significantly the product selectivity, of catalytic oxidation reactions. An example is presented in Fig. 5 which shows the effect of catalyst potential and corresponding work function change on the selectivity to ethylene oxide (Fig. 5a) and acetaldehyde (Fig. 5b) of ethylene oxidation on A g ~ S Z at various levels of gas phase chlorinated hydrocarbon moderators [31] (The third, undesirable, product is CO2). As shown in the Figure a 500 mV decrease in catalyst potential causes the Ag surface to change from selective (up to 70%) ethylene oxide production to selective (up~to 55%) acetaldehyde production. The same study [31 ] has shown that the total rate of ethylene oxidation varies by a factor of 200 upon varying the catalyst potential. (Fig. 6)

80

A(e@), eV

-0.4

,

,

-0.z

,

',-

0.0

,

,

0.z

,-

A(e@), eV 8o,-q.4

,

-q.~.

o.o

0,.2

l

8.5% 0 2, 7.8%C , H 4 o O'x T=270~ P=500~cPa 50 ~ ~ k~ ~ 0.0 ppm C2t-I,C12 [~ ~, 9 0.4ppm L ~ \ [] 0.8 ppm 40j" / % 9 1.2 ppm I ~ x~lm 1.6 ppm

5O 9 r~

30

~ ~ , . / ~

I0

/~

o600

~-4oo

T__'52~/0Ooc~8_~

iiii m

a

(b)

z0

A 1.2 ppm [] 1.6 ppm @ 2.0 ppm

-~oo

V~ , mV

6

-~

",s

u

-

-z~o, mY

'

a

Figure 5. Effect of catalyst potential and work function and of the gas-phase addition of various levels of 1,2-C2H4C12 on the selectivity to ethylene oxide (a) and acetaldehyde (b) of ethylene oxidation on Ag supported on YSZ (reprinted with permission from ref. [31]).

84

-0.4

-0.2

0.0

0.2

[

10 1

m

0

M

o G

1-

O

I

7

T=260~C, P = 5 O O k P a

//

a.5~

l~

0.1

~/-/ ~w/" ,'~

-700

o~, 9.8~

9 CO,_

-500

~

I

t

-300 V~ , rnV

I

C,_H,

* CJI40 [] C H a C H O

I

-0.1

I

-i00

T

0.01 !00

Figure 6. Effect of Ag/TSZ catalyst potential and work function on the rates of formation of ethylene oxide (A), acetaldehyde (~) and CO2 (O); T = 260 ~ ; P=500 kPa, Po2 = 17.5 kPa; PCzH4=49 kPa.(reprinted with permission from ref. [31])

Figure 7 refers to ethylene epoxidation on a Ag film deposited on [Y'-A1203 [24] and shows the effect of catalyst potential, VWR, and partial pressure of gas phase chlorinated hydrocarbon moderator on the selectivity to ethylene oxide. For VWR--0 and VWR=-0.4 V the Na coverage on Ag is nil and 0.04, respectively. As shown in the Figure, there is an optimal combination of VWR(0Na) and PCzH4CI 2 leading to a selectivity of ethylene oxide of 88%. This is one of the highest selectivity values reported for the epoxidation of ethylene. Figure 7 provides an example of how in situ controlled promotion can be used for a systematic investigation of the role of promoters in technologically important partial oxidation reactions. Sm~x=88~

90ES

r~ 7S

70

t0 ~s J'(""~ 1"J~~~ PP'~ G ~ ~.a ~ . . r "

~c/2

Figure 7. Effect of catalyst potential and gas-phase 1,2C2H4C12 partial pressure on _ / ~ , ~ 3 the selectivity of ethylene / ' ~ epoxidation on Ag/]Y'-A1203. ~ ~ ~ (reprinted with permission ~:~ ~t ,~,~' from ref. [24])

85 3.3. Aqueous electrolyte NEMCA: An example of NEMCA in an aqueous alkaline solution is shown in Fig. 8. The catalytic reaction is the oxidation of H 2 by gaseous 0 2 on finely dispersed Pt supported on graphite [19]. Similar results have been obtained with Pt black supported on a Teflon frit [34]. The figure shows the effect of positive and then negative current application on the rates of consumption of hydrogen (rH2) and oxygen (ro=2ro2) and on the catalyst-electrode potential. It can be shown easily that at steady-state rH2-ro=I/2F [19,34], as also confirmed by experiment (Fig. 8). The induced increase in the catalytic rate is clearly non-Faradaic with A values up to 7.2 and 9 values up to 3.5. Similarly to the case of solid electrolyte induced NEMCA, a detailed kinetic study [34] has shown that increasing catalyst potential weakens the Pt=O chemisorptive bond and weakens the Pt-H chemisorptive bond. This is due to the electron donor and electron acceptor character of chemisorbed hydrogen and oxygen, respectively. !t

Figure 8. Electrochemical

tO

t ~

-

-

~

~

t

promotion of H2 oxidation on Pt in contact with 0.1 M KOH solution ] [19]. Transient effect of .8 applied positive and negative current on the ~ rates of consumption of 4 _~ H2 (rH2) and ~ ( r o = 2 r 0 2 ) and on Pt

!

9 8 7

,, . . . . 4

"

' 6

/--'4,, / - - '~ ~ ' ~ - - - - - - . 4

,'

t,~ ~

~

":~"23y/I(2F ) ~.[ 0 -t L !

I

" "~"'-

t.2

.... t / . ' ~ O " %

1

~ - .....

[r t t , l i ~ , ,. , I . : , ~ : i't~ . . . .

0 ~ I 2 5 7 9 11 1.3 15 25 2FN/7 t, rain

electrode;Vs.Catalyst_electrodea reversiblepH2 =0.75potentialkpa,Ha

_ rt'0.4

20

P~176 total gas flowrate 280 cm3min -1 at STP(adapted from ref. I19]).

3.4. Electrochemical kinetics and the magnitude of A: Table 1 provides a list of the catalytic oxidation reactions studied so far from the view point of electrochemical promotion and of the measured A, P and Pi values. As shown in this table measured IAI values range from 1 to 3x10 $. It has been shown both theoretically [9] and experimentally [9,14] that the order of magnitude of the absolute value IAI of the Faradaic efficiency A can be estimated for any reaction, catalyst and solid electrolyte from the following approximate expression:

= ro/(I0/2F)

(5)

where r o is the open-circuit catalytic rate and I0 is the exchange current of the catalyst-solid electrolyte interface. The latter can be obtained from standard current-overpotential (Tafel)

86 plots [9,14]. Equation (5) indicates that in order to obtain a non-Faradaic rate enhancement (IAI>I) it is necessary for the intrinsic catalytic rate, r o, be higher than the intrinsic electrocatalytic rate I0/2F.

Table 1 Catalytic oxidation reactions investigated in electrochemical promotion studies (Adapted from Table 3 in ref. [ 14] which provides specific references to each reaction) I. Electrophobic reactions ~/~e~)>0; ~r/bVwR>0; ~r/~I>0; A>0 Reactants

Products

Cata- Electrolyte lyst (Promoting ion)

C2H4, 02 C2H4, 02 C2H4, 02 C2H4, O2 C2H6, 02 C3H6, O2 CH4, O2 CH4, O2 CH4, 02 CO, 02 CO, 02 CO, 02 CH3OH, 02 C2H4, 02 CO,O2 H2, 02 H2, 02 CO, 02 C2H4, 02

CO2 Pt CO2 Rh CO2 IrO2 C2H40, CO2 Ag CO2 Pt C3H60, CO2 Ag CO2 Pt CO2 Pt CO2,C2H4,C2H6 Ag CO2 Pt CO2 Pd CO2 Ag H2CO, CO2 Pt CO2 Pt CO2 Pt H20 Pt H20 Pt CO2 Pt CO2 Pt

YSZ (O2-) YSZ (O2-) YSZ (Oz-) YSZ (O2-) YSZ (02-) YSZ (O2-) YSZ (O2-) YSZ (O2-) YSZ (02-) YSZ (02-) YSZ (02-) YSZ (02-) YSZ (02-) ~"-A1203(Na+) 13"-A1203(Na+) Nafion (H§ KOH-H20 (OH-) CaF2(F-) T102(T1Ox,O-) 9

9

+

2

T(oC)

A

P

Pi

260-450 250-400 350-400 320-470 270-500 320-420 600-750 590 650-850 300-550 400-550 350-450 300-500 180-300 300-450 25 25-50 500-700 450-600

3.105 5.104 200 300 300 300 5 50 5 2.103 103 20 104 5.104 105 20 20 200 5.103

55 90 6 30* 20 2* 70 3 30* 3 2 5 4* 0.25 0.3 6 6 2.5 20

55 90 5 30 20 1 7 3 30 2 1 4 3 -30 -30 5 5 1.5 20

-100 -3-103 -500 -60 -104 -3 -1.2 -105

7 6 6 3 15" 3 8* 8

250

II. Electrophilic reactions ~9r/3(e~).. - 3 Ag (111) single crystal between 600 K and 900 K in order to 0 ~2 prepare the relevant species. NEXAFS [19] can be used as a highly surface-sensitive tool with about 80 % of the information 011 arising from less than 2 nm depth 520 530 540 550 560 570 of the sample. The unique Photon Energy (eV) resolving power of the method for the details of the hybridisation is Figure 2: NEXAFS data of three atomic oxygen evident. The sharp first structure species on silver is the n* resonance arising from transitions from O ls into unoccupied d-sp hybridised states exhibiting a crystal-field splitting fine structure. The broader features above 536 eV arise from the o* ' transitions from O ls into unoccupied sp r states. The significant differences in chemical bonding between oxygen atoms ~ 2 2o and substrate are much more clearly seen than in the more conventional X-ray 0 photoelectron spectroscopy (XPS) in which ((z) and (?) give rise to a chemical shift of s about 1.5 eV without disclosing, however, 6 ~ ~ / ~ ~ the origin of the chemical shift. The absence of multiple n* resonances and of any polarization effects on well ordered Ag surfaces (data not shown, see [19]) prove . ~ 1 together with Raman spectra [25,18] without any doubt the atomic nature of all three species. The peroxo species existing 40o 600 soo at low temperatures and being not Temperature(K) relevant to selective methanol oxidation is difficult to prepare in a spectroscopically Figure 3: Isotopic exchange TDS with 18 02 pure form as its reactivity leads always to of fully oxidized HPA at 1 K/s heating rate. interference with carbonate [17,8] H2CO was obtained in a TPRS run. formation. Observation of reactive oxygen species on oxides is significantly more difficult. Figure 3 shows an isotope exchange experiment with the molecular oxide HnPVMollO40-14 H20, a heteropoly acid (HPA) [26,27] and oxygen 18 in the gas ,

I

,

,

I

I

I

I

I

I

I

I

I

I

'

I

107 phase. The sample stored molecular oxygen together with the crystal water which exchanged readily at low temperatures due to the acidic environment. Attempts to chemisorb 18 02 onto the HPA surface after dehydration were unsuccessful even at atmospheric pressure. In an atmosphere of 18 Oz it was possible, however, to observe a dynamic exchange between gas phase and lattice oxygen as could be seen from the increase of the scrambled oxygen partial pressure with temperature. At 685 K the structure began to collapse giving rise to an increased scrambling rate and at slightly higher temperatures to oxygen evolution from the sample bulk. The correlation with the formaldehyde formation in an atmosphere of methanol is useful for an assignment of oxygen reactivity. A sustained production of formaldehyde was observed only after the constitutional water was lost lasting up to the temperature at which the breakdown of the structure finally gave rise to an increased (but not lasting) activity. The data show that the HPA is only active in this reaction in its dehydrated state and that the useful temperature window at which steady state oxygen activation from the gas phase is observed is limited to the interval between full dehydration and beginning decomposition. The activation of molecular oxygen is likely to be a complex process requiring a significant amount of activation energy to overcome the 493 kJ/mole bond energy in molecular oxygen which can be moderated by the ease at which electrons are transferred from the catalyst to the chemisorbed oxygen molecule. As the whole process needs to be reversible it is of course not possible to use strong reductants such as early transition metals or alkalis as these materials will not be reducible by the organic substrate. Rather inert metals such as silver and gold are a good choice as well as oxides with multiply valent states (defects).

3. Selective Oxidation Pathways The term ,,selective oxidation" can be defined as a set of reactions in which the formal oxidation state of the organic substrate is increased but not to the level of t h a t of CO or CO2. Selective oxidation is then not only the addition of an oxygen atom to the organic molecule , the abstraction of hydrogen via dehydrogenation and oxidative dehydrogenation (or frequently oxidehydation) is also enclosed in ,,selective oxidation". The dehydrogenation is an endothermic process due to the high energy content of C-H bonds without a suitable compensation by the exothermic water formation. For this reason there is a high tendency of the hydrogen product to react irreversibly with the metal-oxygen complex of the active site and to irreversibly damage it. Dehydrogenations can thus only occur with very stable catalysts and under conditions where the reductive chemical potential is modified by the presence of e.g. a large excess of molecular water [28](e.g. in styrene formation over iron oxide). Different redox reactivities with atomic oxygen are the consequence of differing oxygen-substrate interactions. The (a) species is either transferred to the substrate (oxidation) and/or to hydrogen (oxidehydrogenation). The (~,) species reacts as a basic center and is strongly attached to the catalyst which needs not to be a metal such as silver or

108 platinum [11,12] but may well be a partly reduced metal ion in a sub-oxide catalyst [29].

4 . O x i d a t i o n of M e t h a n o l Methanol can be transformed into formaldehyde in principally two different ways known as dehydrogenation and oxidative dehydrogenation. The main side products are hydrogen and water respectively. A cumulation of the mechanistic considerations drawn from model experiments on metal surfaces [3,6,30] is given in Scheme 2. The catalyst surface containing pre-adsorbed oxygen of either ((z) or ('1I) type transforms methanol into methoxy (1). This can react along three main pathways. Reaction (2) is the simple oxidative dehydrogenation leading to formaldehyde and water. Reaction (3) occurs with (~) oxygen and is a thermodynamically unfavorable dehydrogenation path. At long residence times and under excess (a) oxygen a series of complex reactions occurs which begins with the surface oxidation of methoxy (4) to formate. The labile formate + will either decompose into carbon dioxide and water (5) or even into carbon dioxide and hydrogen in the absence of sufficient chemisorbed oxygen. The desorption of formic acid (6) or the formation of dimethyl ether or dimethyl ketone are alternative side reactions. The relative values of the relevant rate constants will control the selectivity to the side reactions although they rarely dominate the overall conversion. The reaction pathways denoted by the intermediates (2) and (4) have been subject to extensive fundamental studies [3,6,31,32] whereas the dehydrogenation pathway was rarely found in such studies. The obvious reason for this problem Scheme 2: Reaction pathways of methanol on metals in the presence of different surface oxygen species. For the is that (?) oxygen cannot be generated from the gas phase numbers se text. under conditions accessible in UHV systems [20]. Only after prolonged use of a single crystal the bulk-dissolved

CH3OH

(1)~O--HH O--CH H

~ HH~1 methoxy HC

(2

H--O HC--O--~

H H

i

H

(3)

(4) 0

(5~0

formate I H--d-- I

/~'~

HC. (6) IH-o

109 species (unintentionally prepared) will segregate to the surface during heat t r e a t m e n t s and create some (7) oxygen. The reactions described in Scheme 2 require the existence of metal sites besides chemisorbed oxygen. The existence of active sites containing both metallic and oxygen centers was imaged directly in atomic resolution on silver saturated with (7) oxygen. An insitu (111) faceted grain of electrolytic silver was used for the image shown in Figure 4. Within the hexagonal a r r a y of silver atoms (normal interatomic distance of 0.285 nm, white) oxygen atoms are intercalated in the top atomic layer of the catalyst. The large contrast arises from the profound change in local electronic work function [19] on going from metallic to anionic atom sites. The acidity of intermediate OH groups formed on metal Figure 4: STM image of Ag (11 l) with (7) catalysts is assumed to be insufficient to oxygen. The interatomic distance of silver interfere with the elementary steps of measured along the white line is 0.285 nm. oxidative dehydrogenation. Oxide surfaces are significantly more For the ,,hole" contrast see text. complex in the range of reactivity of their t e r m i n a t i n g atoms. On HPA strongly acidic OH groups with pka values of mineral acids exist [33]. Surface hydroxyl groups on nominally pure oxides are also known to react acidic with a wide distribution of pka values. ,,Metal" centers a r e , in contrast to frequent colloquial designations, not existent on most catalytic oxides. Cation sites from regular metal-oxygen polyhedra are also difficult to access by adsorbates due to the inherently large size of oxygen ions relative to all cations. Defects in closed-packed surfaces and incompletely coordinated oxide polyhedra (e.g. octahedra) are, however, accessible for chemisorption of oxygen containing adsorbates such as methoxy. These sites are Lewis acids in contrast to metal sites on the metal-oxygen catalysts which are all in contact with the conduction band and hence Lewis bases. A variety of defect types is conceivable and has also been observed experimentally on binary and on complex oxide surfaces [34]. Figure 5 shows an atomically resolved STM image of a FeaO4 (111) single crystalline surface to give an impression about the complexity of a defective oxide surface. The oxygen atoms of 0.26 nm a p p a r e n t size form a regular a r r a y of a hexagonal closed packing as expected from the bulk structure. A hexagonal Moiree pattern indicates slight lattice mismatch between surface and bulk of the oxide film [35]. Point defects of various geometry and the modified local electron density on the adjacent oxygen atoms (lighter contrast) can clearly be seen. A fundamental problem for experimental [36,37] and theoretical [38] studies is our still very limited knowledge about geometry and defect disposition [34] on most oxide surfaces even when they are considered as

110 single crystals by diffraction techniques. The view that oxides contain two types of oxygen sites termed frequently as ,,terminal" (7) and ,,bridging" ((z) is certainly not incorrect [39] but maybe oversimplifying. As a consequence of the oxygen species distribution and the considerable acidity of oxide surfaces alternative reaction pathway for oxi-dehydrogenations have to be taken into account. The cleavage of the C-O bond in methanol becomes now a feasible reaction. The strong acidity allows further dehydration and the formation of dimethyl ether and of the acetal of formaldehyde as a prominent side reactions at low conversions [37, 40]. The kinetics of the oxidative dehydrogenation may exhibit four rate determining processes involving the formation of the methoxy intermediate, the cleavage of the C-O bond, the activation of the methyl hydrogen combined with the desorption of the formaldehyde product or the re-oxidation of the active site under evolution of water. It is the overwhelming conjecture [32] that for alcohol dehydrogenations the methyl activation step should be rate-limiting and neither the re-oxidation [40,41] nor the formation of alkoxy [42,43,44] control the overall kinetics. The methyl protons should be quite acidic considering the bonding situation in methanol and hence should be activated by a basic oxygen site [44]. This is in conflict with the observation [41] that basic oxide surfaces tend to totally oxidize the alcohol (multiple proton abstraction) and that acidic oxides exhibit a correlation between acidity and formaldehyde production. On suitably acidic surfaces the reaction mechanism may be dominated by the electrophilic action of protons on the methanol leading to an ,,embedded" methoxy which is easily activated at the methyl positions. On HPA catalysts this reaction path was claimed to be dominating [27] as long as the catalyst was sufficiently acidic. For the dehydrogenation of the methyl proton a radically different view was developed recently from theoretical Figure 5" STM image of Fe304 (111). The surface is considerations of HPA oxygen-terminated with each oxygen atom being resolved catalysts [38]. A direct 9LEED and ISS were used to verify the termination. ,,metal"-hydrogen bonding interaction between the proton

111 and one Mo 6§ center was claimed from an orbital analysis of an extended Hfickel calculation of transition geometries of this reaction step. It will be necessary to spectroscopically assess the covalency of the ,,metal"-oxygen interaction at the active sites in order to support this chemically surprising interpretation. The complexity of the methanol-oxide surface interaction is responsible for the still inconsistent picture about this reaction which may well proceed via several p a t h w a y s on the same bulk oxide. The reaction scenario in Scheme 3 is a reflection of the site complexity depicted in Figure 5 allowing for a wide spectrum of chemical reactivity of ,,metal"-oxygen groups. This spectrum is directly related Mo to the local geometric structure (,,bond I + CH3OH distance") and it is thus not surprising that a very pronounced structure H3CO sensitivity of the selective oxidation of Ogasphase O--..~ ~--~O methanol was found [36,37,45]. The silver-oxygen system restructures under reaction conditions into a uniformly active surface regardless of its initial O~ljO ~ O~ /O Mo Mo + CH30 H orientation [46,47] and acts as a structure-insensitive catalyst. - H2CO t + Olattice Y " H20 In the following we will investigate H H tO---.. ~ O t the methanol oxidation process over H Mo ] H silver, copper and heteropoly Mo molybdates in order to identify the + CH3OH occurrence of the possible reaction l H2CO pathways from Schemes 2 and 3 on l~3COH polycrystalline surfaces and at atmospheric pressure. The main O~ jO Mo Mo emphasis in these experiments will be on the source of active oxygen. This focus was chosen to better u n d e r s t a n d Scheme 3" Reaction pathways of methanol the involvement of bulk and sub-surface [10,48] chemistry in selective oxidation oxidation over early transition metal oxides. catalysis. Such an u n d e r s t a n d i n g is required when the catalytic performance is compared between series of chemically different systems which are chosen to investigate only one surface property such as acidity. These experiments will naturally only cover a small selection of the problems discussed with the reaction Schemes.

5. A n , , I n t e r m e d i a t e "

Experiment

The surface science studies of methanol oxidation over metals were carried out u n d e r conditions where no (y) oxygen was present on the surfaces. A t e m p e r a t u r e - p r o g r a m m e d reaction experiment was performed [49] in order to

112

prove that the reaction paths (2) and (3) can coexist on a surface. Electrolytic silver was loaded with molecular oxygen at 900 K and 10 mbar for 10 min. The generation of two species of atomic oxygen ((z) and (7) was confirmed by a TDS scan. After reloading the sample was heated in methanol vapor. The results are collected in Figure 6. The TDS trace exemplifies the well separated responses from (a) and bulk-dissolved oxygen (670 K) and from (y) oxygen at 990 K. In the TPRS traces the formaldehyde generation coincides exactly with the desorption peaks for the two oxygen species. This confirms the reactivity of both species and shows that the reaction is of the Langmuir-Hinshelwood type as the formaldehyde knows the bonding of the - ~ IIII" oxygen involved. The trace for carbon dioxide shows a parallel selectivity I ~ I I ~ I with partial oxidation for the (a) state II ~ iI but no total oxidation from the (7) state. 02 I A "'~" /"% This allows to assign the (a) state as the species for oxidative H2 dehydrogenation and the (y) state as the center specific for dehydrogenation. This assignment is supported by the m/e 18 trace exhibiting no intensity at the high temperature formaldehyde signal. The two sharp desorption peaks in the m/e 44 trace above 400 K arise from I I t i 400 600 800 1000 methoxy and carbonate (formate) Temperature[K] species which hardly produce Figure 6: TDS (dashed) and TPRS (full) of formaldehyde. This finding strongly oxygen/methanol on electrolytic silver. Heating underlines the notion that methoxy and rate: 1.3 K/s formate prepared in the surface science experiment as long-lived species [50] are different from methoxy and formate existing as reaction intermediates in the steady state conversion situation. It is speculated that minor differences in the chemical bonding of the chemically identical ad-species decide over their character as stable spectator species or unstable intermediate structures. Vibrational spectroscopic data [44] of methoxy and the analysis of the homologue ethoxy system [51] strongly support this view.

6. E x p e r i m e n t s w i t h Silver and C o p p e r The existence of the two reaction pathways to formaldehyde shown in Scheme 2 is documented in principal with the data from Figure 4. At steady state and at atmospheric pressure the contribution of the two reaction channels needs to be shown independently. In Figure 7 a section through the reaction parameter space of electrolytic silver under isothermal conditions is shown. The molar

113 stoichiometry was varied from oxygen-rich (in the explosion limit!) to methanolrich and the effect on conversion and selectivity was monitored. It is significant t h a t the optimum formaldehyde production was found at slightly methanol-rich conditions precluding the exclusive [2] contribution of oxidative dehydrogenation to the practical conversion. About 50 % of the formaldehyde production is observed at strongly understoichiometric feed compositions indicating t h a t dehydrogenation contributes significantly to the total activity. The selectivity to total oxidation is "-. "~ strongly coupled with 6O the excess of oxygen in the gas phase indicating ',/ \ w , t h a t the total oxidation II / I \ 10 . 20 o 40 of methoxy via formate is a rapid process when f, the surface is oxygen20 covered. The follow-up ,,............................ reaction to formic acid .............. i...................................... m ~I I I (see Scheme3) occurs not 50 1O0 150 200 250 in parallel with the Ratio CH30H : 0 2 (%) production of the Figure 7: Conversion over electrolytic silver at atmosphereic precursor formaldehyde pressure. A reactor holding 5.0 g catalyst in a 6 mm high bed was indicating that the used. The catalyst load was 10 g/h cm2. All data in mole % oxidation of reactive methoxy to formate is a unlikely process and that formic acid m a y be produced from the strongly adsorbed form of methoxy which occurs under conditions of lean reactive surface oxygen. From this only moderately reactive methoxy a finite selectivity leads to carbon dioxide in full accord with Scheme 2 as indicated by the increase of the CO2 production at very methanol-rich compositions. A quantification of the contribution of dehydrogenation to the total activity can be obtained from determination of the hydrogen gas in the tail gas. This value is s o m e w h a t uncertain, as reaction (5) from Scheme 2 will positively contribute and the consecutive hydrogen + oxygen reaction to water will negatively contribute. The data in the inset of Figure 7 indicate the contribution of several effects to the hydrogen yield. The change in slope at 873 K coincides with the occurrence of the (y) oxygen species (see Figure 4). The high t e m p e r a t u r e branch is attributed to the dehydrogenation with (~,) oxygen (reaction 3 in Scheme 2). The low t e m p e r a t u r e branch with low total conversion and significant evolution of CO2 contains significant contributions from formate decomposition to the hydrogen production. Evidence for the presence of several intermediates with significantly different residence times on the surface of silver was gained from a non-steady state conversion experiment. Electrolytic silver after 200 h time on stream was loaded at 870 K with pure oxygen [25,24]. After purging with nitrogen a stream of methanol was injected in the nitrogen stream from time 0 to 100 s and the

114 e v o l u t i o n of products was monitored every 40 s. The normalized response curves in Figure 8 indicate t h a t metallic silver cannot only store oxygen but also activated methanol in CO2 significant quantities X - " 6.4 T= 680 K ~. . . . HOHO~ on its surface. This is unexpected in the light of the failure to analyze the bonding I state of adsorbed methanol on silver at ! ! high temperatures ! I [17] with surface analytical tools where ! molecular methanol was found to desorb I I I 1 above ca. 200 K. The 32 33 34 35 product curves show Time (min) t h a t the Figure 9: Rate oscillations over copper chips. The methanol : oxygen oxygen-covered ratio was 6.4. Set temperature 680 K. SV 250 h~. surface is highly active for total oxidation. With progressing removal of the surface oxygen species [19] the selective partial oxidation wins with most likely a sizable contribution from dehydrogenation which does not require a stoichiometric a m o u n t of oxygen and can thus produce formaldehyde in a 100 situation lean in ~ ~ , ~ k H2CO surface oxygen (after ,, 80 ca. 300 s in Figure 8). The reaction product formic acid / / \ / \ \ sToK 60 from methoxy-toformate oxidation 40 does not form in parallel with the 20 abundance of methoxy (formaldehyde) nor C with the abundance of 0 1O0 200 300 400 500 Time (s) oxidizing surface oxygen. It occurs as final product from an Figure 8: Response of a 100 s pulse of pure methanol over oxygeni n t e r m e d i a t e residing predosed electrolytic silver. SV: 82000 h~. The response function of the non-reacting system (with SiC) was smaller than the CO2 profile, long on the surface as re-adsorption of m e t h a n o l is highly unlikely 300 s after the end of the pulse in a s t r e a m i n g system. The methoxy intermediate is, however, a precursor to this state as its

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115 population vanishes with the removal of the methoxy species responsible for gas phase formaldehyde. The copper catalyst [6,30] offers an additional way of analyzing the contribution of multiple pathways to the practical reaction by studying the regime of oscillatory behavior. For a wide range of compositions temperaturecontrolled rate oscillations [52] can be observed in the methanol partial oxidation reaction. The origin of this unstable activity is the thermally driven interchange from an inactive bare copper metal surface to an oxygen-covered state which is triggered by oxygen being dissolved in the bulk [53]. This state leads easily to a binary oxide which is only active for total 1.5 oxidation. The rate """~ T = 620 K Cu-L3 oscillations for hydrogen, _ ~ O-K o,s,,o SS'* '~t ? 1.4 formaldehyde and CO2 are presented in high temporal resolution in 1.3 j," C u L III '~~~/ Figure 9. The contribution of several . ,,, 1.2 sources to the hydrogen production can clearly be I1 ~ ~ 1.1 seen. In the beginning of each~ oscillation period "I" 1, I I I I hydrogen and CO2 evolve 0 0.1 0.2 0.3 0.4 0.5 from decomposition of CHsOH- Pressure (mbar) formate. The selective Figure 10: Integrated normalized intensities of the first resonances oxidation via oxidative of Cu L II and 0 K for copper foil treated insitu with a mixture of dehydrogenation and via methanol in oxygen gas. dehydrogenation gradually wins over the total oxidation indicating the gradual modification of the copper-oxygen chemical bonding. This interpretation of the oscillations requires again the existence of chemically inequivalent oxygen species present at the copper surface in atomic form. Using X-ray absorption spectroscopy it was possible to substantiate the dependence of the chemical bonding between oxygen and copper on the chemical potential of the gas phase. Figure 10 shows the systematic increase in the spectral weight of the ~* resonance with decreasing oxidation potential of the gas phase (modified by the addition of methanol to the oxygen atmosphere at 50 mbar). The simultaneous change of the Cu L II white line indicates the gradual transition in bonding character from strongly interacting with rehybridised states (maximum in Cu sub-oxide, higher than in Cu II oxide) to non-bonding atomic-like with a weak O 2p-Cu 4s-p interaction. The reaction conditions imply that this state is responsible for the oxidative dehydrogenation. .

.

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116

7. Heteropoly Acids The discussion so far has shown that several surface oxygen species coexist with a bulk-dissolved species in metallic catalysts. The HPA as molecular analogue of complex oxide catalysts possesses three types of oxygen species namely surface terminating oxygen, surface bridging oxygen and bulk bridging oxygen. The latter species are possibly similar to bulk mobile anionic species that replenish defective oxygen sites where an organic substrate has removed an atom from a surface bridge or terminating site [54,55,56]. Aspects of scientific dispute [57,58,59,60,61] are the actual structure of the HPA anion in its active state which may be an opened Keggin cage (,,lacunary form"), the amount of water present under reaction conditions and the location of the excess electronic charges during a catalytic cycle. It has to be mentioned that the structural definition of the reacting HPA is difficult [62,63,64,58] despite its ,,molecular" character. The catalytically most relevant partially cation-exchanged derivatives [65,27,66] of the free acid forms are difficult to structurally assess as the apparently simple cubic X-ray diffraction pattern needs a detailed evaluation to disclose the real structure. For example four different models have been proposed for partly cation exchanged derivatives of H4PVMollO40. Here we concentrate on the free-acid form to avoid any interference with these structural problems. Under reaction conditions for methanol selective oxidation the HPA is nominally in its anhydrous form [63,64,67] which is metastable with respect to irreversible phase separation under prolonged catalytic reaction. Pulse experiments with varying delay times are suitable to follow the reactivity pattern. A single pulse response with a weakly oxidizing feed is presented in Figure 11. The most prominent effect is the strong capacity of the HPA to store the reaction water from the oxidative dehydrogenation reaction. Under reaction conditions the catalyst contains water of insufficient abundance, however, to

117 transform the anhydrous structure into the nominal 2-hydrate form [63,64]. The selective oxidation of methanol can proceed without gas-phase oxygen as seen from the pulse profile for formaldehyde in Figure 11. The dehydration to dimethyl ether occurs only in parallel to the presence of oxygen from the gas phase indicating a change in the surface acidity when the lattice oxygen reservoir is depleted. A series of 100 consecutive pulses was insufficient to consume all lattice oxygen available to the reaction at 573 K revealing the participation of deep volumes of the catalyst which should consist of a molecular crystal. Under reaction conditions the compacted crystal structure of the anhydride seems to allow oxygen transport to the surface. A series of pure oxygen pulses shown in Figure 12 allowed the replenishment of the lattice-oxygen reservoir. In addition, the total oxidation of strongly chemisorbed intermediates was indicated by the formation of a small amount of CO2. A measurable amount of formaldehyde was found only with the very first oxygen pulse indicating t h a t methoxy represents only a small proportion of the organic surface coverage. A significant abundance of water (see Figure 12) was liberated during the oxygen regeneration indicating that the catalyst was able to store hydrogen. This shows that also on the oxide catalyst a significant activity for dehydrogenation occurs which is not detected in steady-state measurements due to the consecutive hydrogen+oxygen reaction. A series of methanol pulses showed t h a t the catalyst converted it with lattice oxygen into formaldehyde and t h a t in a slow process some methanol was stored together with the reaction water. No liberation of the stored methanol with the gradually beginning water desorption was observed. The methanol seemed to be strongly adsorbed and to be converted into reaction products. All these findings show t h a t the bulk of the HPA crystal is actively involved in the catalytic action as it can store water, polar molecules, hydrogen and oxygen. Insitu UV-VIS spectroscopy [57] is a suitable tool to follow geometric structural integrity and electronic structural modification as function of reaction parameters. Figure 13

118 s u m m a r i z e s some results. The spectra exhibit two high energy m a x i m a at 250 and 310 nm characteristic of oligo-molybdate and the intact Keggin-ion (disappears upon opening of the cage structure). In its pristine (yellow) state the catalyst shows no other absorption. Upon h e a t i n g in oxygen or u n d e r inert conditions the loss of hydration w a t e r occurs (see inset) which is accompanied w i t h a slight electronic reduction as seen from peaks at 680 nm and at 1330 nm. These w e a k structures are due to intervalence charge t r a n s f e r bands from oxov a n a d i u m structures and do not occur with vanadium-free HPA. The inset shows the unexpected correlation of reduction and dehydration t r a n s f o r m i n g the HPA into an active state with the partially reduced electronic structure which is also reflected in the broadening of the Mo-O charger transfer bands. Cooling in w a t e r vapor restores fully the initial state indicating the reversibility of the slight reduction. After 723 K in oxygen and w a t e r vapor, the reduction is irreversible as seen by the change of the 310 nm absorption indicating a partial b r e a k d o w n of the Keggin ion. U n d e r m e t h a n o l vapor and conversion to formaldehyde the s a m e spectrum as u n d e r He at elevated t e m p e r a t u r e s can be seen. The intensity of the peaks depends critically on the methanol to oxygen ratio indicating t h a t this intensity can be used to monitor the degree of 1.2 .~ ~ . ~ . ~ bulk reduction of the ~'/~ catalyst or the 1

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600

800 1000 1200 Wavelength (mn)

1400

1600

1800

involvement of lattice oxygen. Figure 14

reports such an experiment in w h i c h the methanol-tooxygen ratio was varied from 10:1 to 1:2 without a significant change in reactivity. The 680 nm

absorption scales

very well with the Figure 13: Insitu UV-VIS spectra of HPA under various conditions, modified oxygen The inset compares the water TDS (thin) with UV absorption at 1330 abundance in the gas nm (thick) revealing reduction during dehydration, phase and is fully reversible in its intensity upon re-oxidation. The time scale for reversible reaction is, however, longer (6000 s) t h a n the waiting times in Figure 14. A slow diffusion controlled process of anion migration through a lattice is a likely explanation for the response characteristics observed. In s u m m a r y , the d a t a confirm the involvement of the bulk of the HPA crystals in the catalytic activity. They support the conjecture t h a t the v a n a d i u m acts as local sink for electrons from the catalytic cycle. The slow and highly activated oxygen m i g r a t i o n plays a practical role in catalytic conversion which is very m u c h more selective with lattice oxygen t h a n with gas phase oxygen. W a t e r is

119

beneficial at least in helping to maintain the structural integrity. At high levels of hydration which are structurally relevant, water suppresses the catalytic activity in this reaction. Thermal load alone leads to a slight reduction of the catalyst which is seen as beneficial for the function as it will enhance the sticking coefficient of molecular oxygen.

7. C o n c l u s i o n s The experiments have shown that mechanistic aspects discussed in the surface-science literature of methanol oxidation are of relevance under highpressure high-temperature conditions. Surface science has provided the techniques and fundamental insight into species present on metal surfaces. The remote reaction conditions do not allow the assignment of the stable intermediates which were spectroscopically characterized as the reaction intermediates. Experiments at conditions closer to practical conversion have revealed a novel oxygen species and cast doubt on the relevance of a stable methoxy species for catalysis. Several species of atomic chemisorbed I A I B I C I B I A oxygen were 0,2identified with surface science tools 0.15 and assigned to specific functions. n" ~" 0.1 These species occur on all catalysts investigated. Their 0.05 -02/MeOH structural disposition is, however, quite 0 20 40 60 80 1 O0 120 140 different in the Time (min) materials studied. The most relevant Figure 14: Evolution of the UV absorption for reduced V centers with different insitu reaction conditions. From times A to C the oxygen to problem for surface science remains to methanol ratio was reduced produce structural details of terminating and defective oxide surfaces which would be needed for meaningful theoretical approaches to the problem of selective oxidation. The key to understand selective oxidation is to characterize the chemical bonding of various atomic-oxygen species interacting with metal ,,ions" in quite distinct ways. Several examples of characterizations were discussed in this text. The degree of d-state interaction of the oxygen 2 p states decides over the covalent or anionic bonding character which translates into a basic or oxidizing function. In the presence of hydrogen the metal-to-oxygen bonding determines further the acidity of hydroxyl groups which can open additional reaction channels by electrophilic activation of the C-O bond. In all three catalyst systems ,

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120 investigated, the bulk plays an active and catalytically relevant role by providing reservoirs of oxygen atoms which control the type and abundance of surface oxygen species.. The reaction intermediates methoxy and formate seem to occur in two forms each. One is very active and carries the reaction under steady-state conditions. The other is more a spectator species with increased binding energy to the substrate which can, however, undergo complex side reactions and can contribute to the selectivity spectrum of the overall process. The external reaction conditions affect sensitively the co-operation of the dehydrogenation and oxidative dehydrogenation pathways leaving room for the speculation to devise a technical process for undiluted formaldehyde production at reasonable conversions.

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3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama,A.M. Gaffney and J.E. Lyons (Editors) 1997 Elsevier Science B.V.

123

Gold as a l o w - t e m p e r a t u r e o x i d a t i o n c a t a l y s t : f a c t o r s c o n t r o l l i n g activity and selectivity M. Haruta Osaka National Research Institute, AIST Midorigaoka 1-8-31, Ikeda, Osaka 563, Japan

The catalytic activity and selectivity of gold for oxidation can be widely tuned by control of three major factors ; selection of type of metal oxide supports, size of the gold particles, and control of the contact structure of the gold particles with the supports. These factors are markedly influenced by the preparation method. Examples are presented which show surprisingly high activities for CO oxidation at temperatures as low as -77~ and excellent selectivities to propylene oxide in the partial oxidation of propylene. It is found that there is a critical particle size of gold around 1-2 nm where the catalytic nature of the supported gold changes dramatically.

1. I N T R O D U C T I O N

Gold has attracted little attention as a catalyst because of its inert character and low melting point (1063~ which causes difficulties in depositing gold on supports with high dispersion. In the past, gold catalysts were considered not to be competitive with other noble metal catalysts in terms of activity [1, 2]. However, we have found remarkably high catalytic activity in the lowtemperature oxidation of CO with some supported gold catalysts prepared by coprecipitation [3, 4]. Since this work catalysis by gold has received growing attention and is currently investigated by many research groups [5]. The catalytic properties of supported gold catalysts appear to change dramatically when different metal oxide supports or different preparation methods are used. From this standpoint, supported gold catalysts behave differently from platinum-group metal catalysts. This paper attempts to summarize major factors which determine catalytic activity and selectivity in supported gold catalysts.

124

2. EXPERIMENTAL 2.1. Preparation of highly dispersed gold catalysts The catalytic properties of Au depend markedly on the preparation method, because it can bring about a large difference in the size of Au particles and the interaction with supports. Until now six methods, including solid-, liquid-, and vapor-phase techniques, have proven effective for preparing active gold catalysts. Amorphous alloys of gold produced by arc-melting can be transformed to highly dispersed gold particles that interact strongly with the oxides formed from the alloy counter metals during catalytic reaction or through oxidizing treatment. A good example is presented by Shibata, et al. for an Au-Zr alloy used to prepare AufZrO2 [6]. Coprecipitation [4] is useful for the preparation of powder catalysts, especially in the case of Au/a-Fe203, Au/CoaO4, Au/NiO, and Au/Be(OH)2, which are extraordinarily active for the low-temperature oxidation of CO. Active AufPiO2 catalysts can also be prepared by coprecipitation, but only in the presence of Mg citrate [7]. Deposition-precipitation [8] is effective for depositing gold with high dispersion on MgO, TiO2, and A1203. This method is applicable to any form of support including beads, honeycombs, and thin films. In principle, support materials are required to have high specific surface areas, most desirably, larger than 50 m2/g. Since gold hydroxide does not deposit at low pH, this method is not applicable to acidic metal oxides having low points of zero charge, for example, SiO2. The surface reaction of [Au(PPh3)](NO3) with the OH groups of freshly prepared metal hydroxides has recently been reported by the group of Iwasawa [9]. This process is especially useful for the preparation of Au/MnOx, which is active for CO oxidation in air. Chemical vapor deposition of organic gold complexes, typically, dimethylgold (m) ~-diketone, is applicable to the widest range of metal oxides [10]. This technique can be easily used with high surface area materials. It can also be used to deposit Au as nanoparticles even on acidic metal oxides, including M CM41 [11]. Co-sputtering [12] is used for preparing clean thin films which are applicable as gas sensors. A target composed of an Au plate and a metal oxide disk is sputtered in 0.4 Pa of oxygen to deposit on a glass substrate heated at 250~ Oblique deposition on a rotating substrate produces columnarstructured porous thin films.

2.2. Catalytic activity measurements Catalytic activity measurements were carried out in a fixed bed quartz reactor of inner diameter 6 ram. The catalyst (usually 200 rag, 125-212 gm fraction) was placed between two glass wool plugs with a bed length of 10-20 ram. For the oxidation of CO and H2, a standard gas containing 1 vol% CO or H2 in air was dried in a silica gel column cooled down to 0~ or -77~ and passed through the catalyst bed at a space velocity of 20,000 h-lml/g-cat.. For the partial

125 oxidation of propylene a mixture of propylene, oxygen, hydrogen, and argon, most frequently, in a volume ratio 1:1:1:7, was passed through the catalyst bed without drying at a space velocity of 2000 h-lml/g-cat.. 2.3. Catalyst characterization Characterization of the catalysts was made using high-resolution transmission electron microscopy (TEM) (Hitachi H-9000), X-ray diffraction (XRD), extended X-ray absorption fine structure (EXAFS), X-ray photoelectron spectroscopy (XPS), temperature programmed desorption (TPD), and Fourier transform-infrared spectroscopy (FT-IR).

3. RESULTS 3.1. CO oxidation Gold supported on metal oxides and hydroxides is active for CO oxidation at temperatures below 0~ irrespective of the type of metal oxide support when gold is deposited as nanoparticles. Table I shows three groups of gold catalysts which are classified in terms of stability, activity, and requirement as the size of gold particles for activity.

Table 1 Classification of supported gold catalysts for low temperature CO oxidation. Group Stability and Support Preparation Size Activity Methods Requirement less stable but A the most active Be(OH)e, Mg(OH)e CP, DP Au clusters at -77 ~C hydroxides Do

-0.00

_

I

10

I

80-

1

I

I

20 30 D i a m e t e r (A) I

I

40

.

50

I

cuboctahedra sahedra ss distribution

~600

-0.10

= r r..,.o r

o,,,=

o

==40-

....

L _

-0.05

Ir

m 20 -

o" c

---.

o-z

;Z>o -0.00

_

I

10

l

I

20 30 D i a m e t e r (A)

I

40

50

Figure 1. Mass fractions of clusters corresponding to (a) the DFA of freshly prepared (1 week old) Au/Mg(OH)2 and (b) aged Au/Mg(OH)2 (3 months). The solid line gives the corresponding non-discrete distribution of mass. The influence of preparation methods on the catalytic properties of supported gold is markedly large, whereas it is marginally small on supported Pt [19]. In Table 3 are listed the mean particle diameters of Pt and Au, rate of CO oxidation at 27~ turnover frequency (TOF) based on surface exposed metal atoms calculated from the mean particle diameter and actual metal loading, and apparent activation energy. TOFs for AudiO2 prepared by depositionprecipitation are larger by about one order of magnitude than those for Pt catalysts and by about 4 order of magnitude than those for AufFiO2 prepared by photochemical deposition. The apparent activation energies also provide an interesting contrast: for all Pt catalysts and A u catalysts prepared by impregnation and photochemical deposition they are in the range of 50-60 kJ/mol, while for Au catalysts prepared by deposition-precipitation they are

128 around 20 kJ/mol.

Table 3 CO oxidation over P t f r i 0 2 and Aufrio2 Metal P r e p a r a t i o n Loading Size Method wt% nm DP 1.0 1.3 IMP 1.0 1.4 Pt FD 0.9 2.4 DP 0.7 3.1 tt 1.8 2.7 Au IMP 1.0 ? FD 1.0 4.6

Rate at 27 ~C mol.s-l.g i 1.4x10 -7 1.9x10 -7 2.4x10 -s 6.9x10 -7 5.5x10 -6 1.7x10 -1~ 1.5x10 -10

TOF at 27 ~C s -1 2.7x10 -3 3.8x10 -3 9.2x10 -3 3.4x10 -2 1.2x10 -1 .... 9.6x10 -6

Ea kJ/mol 49 60 53 19 18 58 56

DP: deposition-precipitation, IMP: impregnation, FD: photochemical deposition

3.2. Partial oxidation of propylene Table 4 shows t h at, in the presence of both oxygen and hydrogen, propylene is converted to propylene oxide(PO) with selectivities above 90% over Aufrio2 [20]. Carbon dioxide is a m a i n by-product. Over P d f r i o 2 and PtfriO2, in contrast, propane is m a i n l y produced. It should also be noted t h a t AufriO2 prepared by impre g n a t i o n is poorly active and produces mainly CO2 at higher t e m p e r a t u r e s . The oxidation of propylene in the absence of hydrogen occurs only at t e m p e r a t u r e s above 250~ to produce a l m o s t exclusively CO2, while the oxidation of hydrogen is r e t a r d e d by the presence of propylene. This implies t h a t the presence of hydrogen not only enhances the oxidation of propylene but

Table 4 Reaction of propylene with oxygen and hydrogen over Au/, Pd/, P t f r i 0 2 catalysts Catalyst Metal T Conversion, Selectivity, % PO PO STY a % yield wt% ~C C~I6 H2 PO acetone C~Is C02 % retool/h/g-cat Aufr i o 2 b 1 50 1.1 3.2 >99 1.09 0.18 Aufrio2 c 1 80 0.2 8.9 70 0.00 0.00 Pdfrio2 c 1 25 57.1 97.7 - 0.4 98 1 0.00 0.00 Ptfri02 c 1 25 12.1 86.6 2 92 6 0.00 0.00 Feed gas: propylene/~e/Ar=lO/lO/lO/70 (vol%), flow rate: 2,000 ml/hr, catalysts: 0.5g. a space time yield, b p r e p a r e d by deposition-precipitation, cprepared by impregnation.

129 also leads to the selective partial oxidation to the epoxide. Dioxygen may be transformed, probably at the perimeter interface of TiO2, through reduction with hydrogen to an active species, possibly hydroperoxo species, that can selectively oxidize propylene adsorbed on the surface of the Au particles. It is surprising that the reaction pathway switches from oxidation to hydrogenation at specific Au metal loadings. Figure 2 shows that while partial oxidation occurs at Au loadings above 0.2 wt%, below this amount the reaction switches to hydrogenation to produce propane[21]. Hydrogenation takes place with higher turnover frequencies at temperatures below 100~ and the selectivity to propane reaches 100 %. This result appears consistent with the result reported by Naito, et al. for Au/SiOe catalyst, where using very low Au loadings, it was found that the hydrogenation of olefin was accelerated by the presence of oxygen [22]. Although the number of gold particles that could be observed by TEM was not large for the catalysts with Au loadings smaller than 0.1 wt%, the main difference between the 0.4 wt% and 0.1 wt% samples appear to be whether the mean particle diameter of Au is above or below 2.0 nm (Figure 3). Taken together these results indicate that gold particles larger than 2.0 nm are primarily responsible for the formation of propylene oxide, whereas those below

F

=

~

I

Au/TiO2

r

2.4 o-e 3

-d m =,,,,,

>" o

"o O

a.

2

C H3CH-CH2 \/

2.2

o

propane C H3CH-CH2

1.8

~O/

1

0.02 0.05

0.1

mNN 0.2

0.4

1

Au Loadings, wt%

Figure 2. Percentage product yields from propylene on Au/TiOz as a function ofAu loadings. Feed gas: propylenelOz/HJAr= 10/10/10/70 in vol%, flow rate: 2,000 ml/h, catalyst :0.5g, reaction temperature: 80~

;5 1.6 1.4

9

C H3C H2C H3

II

' 0:2

' 0:4

' 0'.8

' 0'.8

'

Au Loadings, wt%

Figure 3. Mean particle diameters as a function of Au loadings in Au/ TiOz. 2 nm appears to be critical for switching-over of the products between propylene oxide and propane.

130 2.0 nm catalyze the formation of propane. It is also probable t h a t the presence of Au clusters, smaller t h a n 1.0 nm, which could not be observed by TEM, may contribute to hydrogenation. In order to improve the conversion of propylene while m a i n t a i n i n g selectivity to propylene oxide above 90%, TiO2 or Ti cations were dispersed on high-surface area supports. This i s to obtain a high density of active sites which would be, nevertheless, sufficiently separated from each other. It was expected t h a t higher conversions of propylene and the depression of successive oxidation of propylene oxide to CO2 would result. Both TiO2 deposited on SiO2 (specific surface area, 310 m2/g) and Ti-MCM supports for Au gave slightly improved conversions but with different time-on-stream behavior. TiO2 on SiO2 showed a decreasing conversion with time while Ti-MCM showed gradually increasing conversion reaching a steady s t a t e value after l h [23]. It is likely that the reaction pathways are different for Au supported on TiOJSiO2 t h a n for Au supported on Ti-MCM.

3.3. Spectroscopic

investigations

FT-IR spectroscopy clearly shows that the reactants, CO and propylene, are moderately adsorbed on the surface of Au particles. The presence of water has either no effect or an enhancing effect on the adsorption. The two absorption bands for carbonyl on the Au surface (2106-2118 cm-1 and 2090-2110 cm-1) indicate that there are the neutral and the positively charged surface of Au [24]. The Au atoms at the perimeter interface with the support (TiO2) are probably positively charged and can activate oxygen molecules. The TPD spectra in Figure 4 show that the deposition of Au on TiO2 appreciably increases the intensity of the desorption peak of oxygen. The desorption temperature at around 250~ suggests t h a t the surface oxygen species is weakly adsorbed on the surface. The desorption spectra of CO from AufriO2 show t h a t the amount of CO taken up is increased by Au deposition on TiO2 and t h a t the desorption always takes place as CO2 even at 50~ P r e t r e a t m e n t of the catalyst sample in a stream of He instead of 02 decreases the amount of CO t a k e n up and causes the desorption to take place as CO in the lower temperature region. These results indicate t h a t the uptake of CO involves oxygen adsorbed on the surface of Aufrio2.

4. DISCUSSION The experimental results presented above clearly show t h a t there are three important controlling factors which determine the catalytic properties of gold supported on metal oxides for oxidation. The first is the selection of suitable m e t a l oxide supports. In the complete oxidation (combustion) of CO and nitrogen-containing organic compounds like trimethylamine gold catalysts are

131

Introd. CO pulse

(c)

m I

(d)

CO

(e)

0 Q 0 I=

02 p r e - t r e a t . (b)

x4 I

0

.,

100

I

I

200

300

400

Temperature, ~

Figure 4. TPD for Au/TiO2 (Au loading: 3.3 wt%, DAu: 3.5nm) and TiO2. Pretreatment in O2: (a) Au/TiO2 and (b) TiO2. Pretreatment in 02 or He followed by exposure to CO pulses up to full uptake: (c) Au/TiO2 in 02, (d) Au/TiO2 in He, and (e) TiO2 in 02. Heating rate: 10 ~C/min, rate of flow: 30 ml/min, sample weight: 200mg.

more active than Pt-group metal catalysts. For the combustion of trimethylamine, iron-based metal oxide supports gave the most active gold catalysts. Since iron oxides adsorb trimethylamine strongly, this suggests that metal oxides having a stronger affinity for the reactants provide higher catalytic activity. For CO oxidation, when gold is deposited as nanoparticles, almost all metal oxides can provide activity at temperatures below 0 ~C. In a previous paper [4] we reported that only select metal oxides, the oxides of 3d transition metals of group VIII (Fe, Co, Ni), gave catalytic activity at -77 ~C. However, this was only for the case where gold catalysts were prepared by coprecipitation. Since then, we have succeeded in depositing gold on SiO2 as nanoparticles and have found similar catalytic activities for CO oxidation at temperatures below 0~ Because a simple mechanical mixture of gold nanoparticles with diameter of 5

132 nm in colloidal dispersion and TiO2powder exhibits poor activity [25] and Au deposited on SiO2 exhibits lower activity when calcined at 300 ~C t h a n at 400 ~C [11], a strong interaction between gold particles and the metal oxide support appears to be necessary in the genesis of the low-temperature activity for CO oxidation. A second factor important for activity is the minimization of the size of gold particles. In general, the diameter of gold particles should be smaller than 10 nm for CO oxidation over Au supported on metal oxides. It is very interesting t h a t for CO oxidation over Au/Be(OH)2 and Au/Mg(OH)2 extraordinarily high catalytic activity is observed only when gold clusters of 13 to 55 atoms with icosahedral and fcc cuboctahedral structures are deposited on these metal hydroxides. This requirement as the size and structure of gold needs further investigation. The reaction of propylene with oxygen and hydrogen over Aufrio2 also indicate the existence of a critical size for gold, in this case, at around 2 nm. When gold particles are larger than 2 nm propylene oxide is selectively formed, while, when they are smaller than 2 nm only propane is formed. The results also indicate that hydrogen molecules can be dissociated on the surface of gold in the presence of oxygen. Since propane was not obtained from propylene and hydrogen in the absence of oxygen, oxygen might modify such small gold particles rendering them electron deficient so as to alow them behave like Pd and Pt. The third factor necessary to obtain high activity is to control the interaction of the gold particles with the metal oxide supports. The remarkably large influence of preparation methods on the catalytic properties of supported gold is observed for the oxidation of CO and H2. Over unsupported gold powder and gold deposited on SiO2 by impregnation, the oxidation of hydrogen takes place at lower temperatures than the oxidation of CO, whereas the opposite is the case over gold deposited on metal oxides (including SiO2) by deposition-precipitation and chemical vapor deposition. The turnover frequency of Aufrio2 prepared by deposition-precipitation for CO oxidation is larger by 4 orders of magnitude than that of Aufrio2 prepared by photo-deposition. These differences may arise from different contact structures at the periphery of the gold particles.. While poorly active gold catalysts are composed of spherical gold particles weakly interacting with the metal oxide supports, highly active gold catalysts are composed of hemispherical gold particles attached to the supports on their flat planes. This contact structure presents a longer distance around the perimeter interface of the gold particles. It is likely that the perimeter junction contains the sites for the activation of oxygen for CO oxidation and for the epoxidation of propylene.

133 5. CONCLUSIONS Gold supported on metal oxides and hydroxides has been prepared by several different methods and tested for CO oxidation and partial oxidation of propylene. The conclusions obtained are as follows: 1) The method of preparation has a considerable effect on the catalytic properties of supported gold. Coprecipitation, deposition-precipitation, chemical vapor deposition methods are especially effective for depositing gold as nanoparticles with diameters smaller than 5 nm and with strong interaction with the supports. 2) Over highly dispersed gold catalysts, CO oxidation can take place even at -77 ~C, and propylene oxide can be selectively produced at temperatures around 100 ~C. 3) For CO oxidation over gold supported on metal oxides, the catalytic activity is almost insensitive as to the type of metal oxide support but strongly depends on the strength of interaction with the supports. For the reaction of propylene with oxygen and hydrogen, only titanium-based oxide supports lead to the selective production of propylene oxide. 4) The size effect is dramatic in CO oxidation over Au/Be(OH)2 and Au/Mg(OH)2 and in the reaction of propylene with oxygen and hydrogen over Aufrio2 catalysts. In the latter case, a critical diameter of gold particles is 2 nm; above this value propylene oxide is selectively obtained, while below it propane is favored. 5) The catalytic nature of gold can be widely tuned by the three factors; type of support, size of the gold particles, and the structure of the contact interface between gold and the support.

REFERENCES

1. I.E. Wachs, Gold Bull., 16 (1983) 98. 2. J. Schwank, Gold Bull., 16 (1983) 103, and 18 (1985) 2. 3. M. Haruta, T. Kobayashi, H. Sano, and N. Yamada, Chem. Lett., (1987) 405. 4. M. Haruta, N. Yamada, T. Kobayashi, and S. Iijima, J. Catal., 115 (1989) 301. 5. M. Haruta, Catal. Today, 36 (1996) 153. 6. M. Shibata, N. Kawata, T. Masumoto, and H. Kimura, Chem. Lett., (1985) 1605. 7. M. Haruta, T. Kobayashi, S. Tsubota, and Y. Nakahara, Jap. Pat. 1778730 (1993). 8. S. Tsubota, D. A. H. Cunningham, Y. Bando, and M. Haruta, Preparation of Catalysts VI, G. Poncelet, et al. (Eds.) Elsevier, Amsterdam, pp.227-

134 235, 1995. 9. Y. Yuan, K. Asakura, H. Wan, K. Tsai, and Y. Iwasawa, Chem. Lett., (1996) 755 and Catal. Lett., 42 (1996) 15. 10. M. Okumura, K. Tanaka, A. Ueda, and M. Haruta, Solid State Ionics, 95 (1997) 143. 11. M. Okumura, S. Nakamura, S. Tsubota, T. Nakamura, M. Azuma, and M. Haruta, Catal. Lett., submitted. 12. T. Kobayashi, M. Haruta, S. Tsubota, and H. Sano, Sensors and Actuators, B1 (1990) 222. 13. M. Haruta, T. Kobayashi, S. Tsubota, and Y. Nakahara, Chem. Express, 3 (1988) 159. 14. S. Tsubota, M. Haruta, T. Kobayashi, A. Ueda, and Y. Nakahara, Preparation of Catalysts VI, G. Poncelet, et al. (Eds.), Elsevier, Amsterdam, pp. 695-704, 1991. 15. M. Haruta, S. Tsubota, T. Kobayashi, H. Kageyama, M. J. Genet, and B. Delmon, J. Catal., 144 (1993) 175. 16. D.A.H. Cunningham, W. Vogel, H. Kageyama, S. Tsubota, and M. Haruta, submitted to J. Catal.. 17. W. Vogel, D.A.H. Cunningham, K. Tanaka, and M. Haruta, Catal. Lett., 40 (1996) 175. 18. D.M. Cox, R. Brickman, K. Creegan, and A. Kaldor, Atoms, Molecules and Clusters, 19 (1991) 353. 19. G.R. Bamwenda, S. Tsubota, T. Nakamura, and M. Haruta, Catal. Lett., 44 (1997) 83. 20. T. Hayashi, K. Tanaka, and M. Haruta, Symp. Heterogeneous Hydrocarbon Oxidation, 211th ACS Meeting, New Orleans, March 1996, pp. 71-74. 21. K. Tanaka, T. Hayashi, and M. Haruta, Interf. Sci. Mater. Interconnection, Proc. JIMIS-8, Jpn. Inst. Metals, pp.547-550, 1996. 22. S. Naito and M. Tanimoto, J. Chem. Soc. Chem. Commun., (1988) 832. 23. Y.A. Kalvachev, T. Hayashi, S. Tsubota, and M. Haruta, Proc. 3rd World Congr. Oxid. Catal., September 21-26, 1997, San Diego. 24. F. Boccuzzi, A. Chiorino, S. Tsubota, and M. Haruta, J. Phys. Chem., 100 (1996) 3625. 25. S. Tsubota, T. Nakamura, and M. Haruta, Catal. Lett., submitted.

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 1997 Elsevier Science B.V.

135

The Selective Epoxidation of Non-Allylic Olefins Over Supported Silver Catalysts John R. Monnier Chemicals Research Division, Research Laboratories, Eastman Chemical Company, P.O. Box 1972, Kingsport, TN, 37662, USA.

ABSTRACT The epoxidation of non-allylic, or kinetically-hindered, olefins can be carried out using supported silver catalysts. While epoxidation does occur for unpromoted catalysts, the strength of olefin epoxide adsorption leads to low activity and selectivity, as well as irreversible catalyst fouling. The additon of certain alkali metal salts, such as CsCI, lowers the desorption energy of the olefin epoxide, permitting dramatic increases in activity, selectivity, and catalyst lifetime. In the case of butadiene, the addition of an optimum level of CsCI increases activity and selectivity from approximately 1% butadiene conversion and 50% selectivity for epoxybutene to 15% conversion and 95% selectivity, respectively. Epoxidation of butadiene occurs by addition of dissociatively-adsorbed oxygen to one of the localized C=C bonds to form epoxybutene. The addition of oxygen across the terminal carbon atoms does not occur to any measurable extent. The direct participation of molecular oxygen can be ruled out based both on selectivity arguments as well as the kinetic model for the reaction. The kinetics imply a dual site mechanism. One site, which is unpromoted, serves as the site for butadiene adsorption, while the second site, which is promoted, functions as the site for dissociative oxygen adsorption and epoxybutene formation. Epoxybutene and derivatives represent the beginning of several new families of chemicals that were either not available, or were too expensive, to be considered for large-scale, or even fine chemical, production. More than one hundred chemicals have been prepared so far; several of these are in commercial production at the semiworks scale. 1. INTRODUCTION The gas phase epoxidation of ethylene using molecular oxygen to produce ethylene oxide is one of the most successful examples of heterogeneous catalysis to date. In 1995, over 7.62 billion lb. of ethylene oxide were produced in the US alone(I). In fact, ethylene oxide accounts for 40-50% of the total value of organic chemicals produced by heterogeneously-catalyzed oxidation(2). Thus, while the importance of olefin epoxidation is apparent, efforts to selectively epoxidize higher olefins to their corresponding epoxides using molecular oxygen have been unsuccessful. Many explanations have been proposed for the lack of success in this important area, although the most obvious reason is the reactivity of allylic C-H bonds, which are present in the higher olefins that have been tested. The bond dissociation energy of the allylic C-H bond in propylene is 77 kcal/mole, while the vinylic C-H bond in ethylene is 112 kcal/mole (3). Thus, abstraction of one of the

136

allylic C-H bonds in propylene becomes energetically favorable relative to electrophilic addition of oxygen across the C=C double bond. Once abstraction of hydrogen occurs, formation of an epoxide is precluded. Because of the higher bond strengths of the C-H bonds in ethylene, oxygen addition is the energeticallypreferred route. Previous attempts at higher olefin epoxidation have largely ignored the kinetic reactivity of allylic C-H bonds, but focused on modifying the desirable and unique catalytic properties of silver (the ability to activate oxygen so that it electrophilically adds across a C=C double bond) by promotion or selection of reaction conditions to hopefully suppress allylic hydrogen abstraction while not affecting the addition of oxygen to the C=C double bond. In almost all cases, the only effect has been to poison the silver surface with no substantial improvement in selectivity to the olefin epoxide. In this study we show that it is possible to selectively epoxidize higher olefins to their corresponding epoxides if the olefins do not contain reactive allylic hydrogen atoms. Many olefins, including styrene, substituted butadienes, and norbornene have been selectively epoxidized. Because of the usefulness of 3,4-epoxy-l-butene as a new chemical intermediate, most of the data and discussion will involve the selective epoxidation of butadiene. The catalyst composition and overall kinetics of the reaction will be discussed in some detail. 2. EXPERIMENTAL METHODS

Catalyst preparation has been discussed earlier (4,5) in greater detail. The methods of preparation are typical of those used for supported silver catalysts. In the cases where AgNO3 was used as the silver source, water was used as the solvent. The promoter salts in the appropriate amounts were added to the AgNO3 solution before impregnation and reduction of the AgNO3. The volume of the aqueous salt solution was typically in slight excess of the volume required for incipient wetness. If necessary, excess liquid was decanted before tumble drying at approximately 60-70~ The dried precursor was reduced in a flowing H2/N2 stream at temperatures up to 400~ The various combinations of reduction conditions (gas composition, space velocity, temperature ranges, etc.) have been previously discussed (6). Unless otherwise stated, the silver weight Ioadings for all the catalysts ranged from 12-15%. Promoter Ioadings were varied from a few ppmw to several thousand ppmw. In addition to AgNO3 as the silver source, in some cases silver oxalate, Ag2C204, was used as the silver source. In these cases, a solution of ethylenediamine and water was used to dissolve Ag2C204. Promoter salts were added to the Ag2C204 solution before impregnation and calcination. These catalysts were activated by thermal reduction in flowing air at temperatures up to 300~ The supports used for the various compositions were low surface area, fused alpha-AI203 carriers typically used for supported silver catalysts. Surface areas ranged from 0.2-1.0 m/gm with pore volumes between 0.4-0.6 cc/gm; median pore diameters ranged from 1-10 microns. Most catalysts were prepared on shaped carriers, which were sieved to give a narrow range of granules before evaluation. All catalysts were evaluated at steady-state conditions in a one atmosphere flow reactor using gas flows of He, C4H6, and 02, which were delivered by mass flow controllers. For the kinetic dependency studies, CO2 was also added by a mass flow controller, while different levels of epoxybutene and H20 were introduced from a helium-swept vapor-liquid saturator. In a similar manner, furan, epoxybutene, 2,5-dihydrofuran and crotonaldehyde were added to the feed stream

137

during investigation of the reaction mechanism. Analyses of both the feed and product gas streams were made using in-line gas sampling. Comparison of the feed and product streams permitted accurate mass balances at all reaction conditions. The gas chromatograph employed both thermal conductivity and flame ionization detectors hooked in series to give quantitative analyses of all gaseous products. 3. RESULTS AND DISCUSSION Table 1 Comparison of olefin epoxidation reactions over unpromoted 5% Ag/A120 3. Reaction conditions: T = 250~

catalyst weight = 2~

grams, and flow rate = 50 SCCM of

He/olefin/O 2 = 3/1/1.

Reaction

% Select

CH2=CH2+ 02 EO CO2

53 47

CH2=CHCH3+ 02 CO2 PO CH2=CHCHO

92 0 8

CH2=CHCH2CH3+ 02 CO2 BO CH2=CHCH=CH2

91 0 9

CH3CH=CHCH3+ 02 CO2 2-BO Other

98 0 2

CH2=CHCH=CH2+ 02 CO2 EpBT M Oxirane

25 75

% Conv. 12

3~

2.8

The data in Table 1 summarize catalytic activities for epoxidation of a variety of olefins over an unpromoted 5%Ag/AI203 catalyst. These data illustrate the preferential reactivity at the allylic position relative to addition of oxygen across the C=C bond. While the selectivity to ethylene oxide is typical for an unpromoted catalyst, the selectivities to propylene oxide and butylene oxides are non-existent for propylene, 1-butene, and 2-butene, respectively. In addition to small amounts of the selective allylic oxidation products (acrolein in the case of propylene and butadiene in the case of 1-butene), the only products are those of combustion. However, the results for butadiene reveal it is possible to epoxidize this non-allylic olefin at moderate selectivity and activity. What is not obvious from Table 1 is the short-lived nature of this activity. After 2-3 hours of reaction time, activity and selectivity typically decreased to approximately 14 kcal/mole lower than that for the unpromoted Ag catalyst. Epoxybutene is a very versatile and reactive chemical intermediate that has not been prepared before in large, or even reasonable, quantities. Currently, it is available only from chemical supply houses for prices > $10/gram. The molecule is highly functional, with each carbon atom chemically distinct. It even contains an asymmetric carbon atom at the C3-position. Epoxybutene has two separate and reactive functionalities, a C=C bond and an epoxide group. Each of these groups has rich chemistry, both from a polymer viewpoint and a chemical feedstock perspective. The histogram plot in Figure 7 indicates many different structural isomers are thermodynamically accessible from epoxybutene. In fact, the derivative tree in Figure 8 outlines some of the more than 100 compounds that have been prepared from epoxybutene. Key to this tree of chemical compounds is the thermodynamically favored transition of the family of C4H60 structural isomers from epoxybutene to 2,5-dihydrofuran (by 19 kcal/mole) to 2,3-dihydrofuran (by 7 kcal/mole). Even the formation of cyclopropylcarboxyaldehyde (CPCA) becomes thermodynamically favorable enough at elevated temperature, so that high conversions and high selectivities are observed at temperatures above 350~ In fact, this transition from epoxybutene to CPCA is being commercially practiced to provide a rich variety of chemicals to be used in the pharmaceutical and agricultural sectors (18,19). The various classes of products outlined in Figure 8 include addition of oxygen-containing and nitrogen-containing nucleophiles to epoxybutene to yield an almost endless variety of hydroxy ethers and amino alcohols, respectively, having an extremely wide range of chemical properties as well as boiling points and solvating properties. Addition of H20 to epoxybutene gives both 1,4- and 1,2butenediols, providing a novel means of formation for some interesting glycols.

147

Kcal/Mole 40 30

29.7

20

,04//~o 4.75

-20

t

.,3., ~.,.o-----r~-1s.1 1

I~-cH~

-40

/~,/CHO

1

-21.5

-25.6

Compound

Figure 7. Thermodynamic enthalpies of formation of epoxybutene and related isomers. Enthalpies calculated from CHETAH (Chemical Thermodynamic and Energy Release Evaluation Program) at 25~ and ideal gas conditions.

HO~

O

R

t

OR

OH

t,/

CH ~:~OH

t

X

t

''

RHN~

O

H

1 Figure 8. Abbreviated EpB TM oxirane derivative tree. Compounds in boxes are currently being produced or have been demonstrated in pilot plant operation.

148

Addition reactions to the C=C bond include hydrogenation to give epoxybutane, providing an indirect, yet efficient way to prepare this important epoxide (when coupled with hydrogenation) using molecular 02. Halogenation with CI2 and Br2 form the corresponding dihaloepoxybutanes, which are possible components in flame retardents. Another interesting addition reaction is the selective addition of CO2 to form the highly versatile monomer, vinyl ethylene carbonate, which is similar to ethylene carbonate. From 2,5-dihydrofuran, the most obvious derivative is tetrahydrofuran (THF), formed by the hydrogenation of the C=C bond. The C=C bond is also reactive for hydroformylation and olefin metathesis to add additonal functionality and structure. As stated earlier, the next intermediate, 2,3-dihydrofuran (2,3-DHF) serves as the gateway to CPCA family of chemicals. Hydration of 2,3-DHF produces 2-hydroxytetrahydrofuran, which can be readily converted to gamma-butyrolactone, pyrrolidinone (and N-substituted pyrrolidinones). Finally hydrogenation of hydroxytetrahydrofuran yields 1,4-butanediol in high yields. Yet another demonstration of the versatility of epoxybutene comes from the asymmetric center, which has been converted into a number of four-carbon chiral synthons, such as 3-butene-l,2-diol and various derivatives, in >99% enantiomeric purity (20,21). 4. CONCLUSIONS

The epoxidation of non-allylic olefins, or olefins containly kinetically-hindered allylic olefins, using promoted silver catalysts has been demonstrated. The epoxidation of butadiene to form epoxybutene marks the first example of an olefin other than ethylene to be selectively epoxidized at steady state and commerciallyrelevant conditions using gas phase oxygen and heterogeneous silver catalysts. Epoxidation of higher, non-allylic olefins does occur without the use of an alkali metal salt promoter. However, the olefin epoxide is strongly adsorbed to the silver surface and undergoes a number of side reactions as well as surface fouling. The addition of a promoter, such as CsCI, apparently lowers the desorption energy of the olefin epoxide from the surface, permitting both selectivity and activity to dramatically increase. In the case of butadiene, the addition of an optimum level of CsCI promoter to the silver catalyst increases selectivity and activity from about 50% selectivity and 1% conversion to approximately 95% selectivity and 15% conversion, respectively. Catalyst lifetime also increases from less than 3-4 hours to commercially-relevant periods of time. Epoxidation of butadiene occurs by electrophilic addition of dissociativelyadsorbed oxygen to one of the localized C=C bonds to form the epoxide. The addition of oxygen across the terminal carbon atoms to form 2,5-dihydrofuran does not occur to any measurable extent. When it is observed for unpromoted catalysts, 2,5-dihydrofuran is formed from the isomerization of strongly-bound epoxybutene. The direct participation of molecular oxygen addition to the C=C bond can be ruled out based both on selectivity arguments as well as the kinetic model for the reaction. The kinetics imply a dual site mechanism for butadiene epoxidation. One site, which is unpromoted, functions as the butadiene adsorption site, while the second site, which is promoted, serves as the site for dissociative oxygen adsorption and epoxybutene formation. Under normal reaction conditions, the reaction is zero-order in butadiene pressure and first-order in oxygen pressure (each site is 1/2 order in oxygen pressure). Because of the kinetic inhibition effect of epoxybutene, the reaction at high conversion is negative first order in epoxybutene.

149

Finally, epoxybutene and derivatives represent the beginning of several new families of chemicals that were previously either not available or simply too expensive to be considered for large-scale industrial, or even fine chemical, application. More than one hundred chemicals have been synthesized from epoxybutene, and many more are currently being synthesized and evaluated for a wide variety of applications.

ACKNOWLEDGEMENTS

The author acknowledges the efforts of Peter Muehlbauer and George Oltean for assistance in catalyst preparation and reaction kinetics, David Hitch in reaction engineering, and Jerome Stavinoha and Windell Watkins for many meaningful discussions regarding development of this technology. Steve Godleski and Steve Falling, among many, were instrumental in much of the derivative work involving epoxybutene as a new organic intermediate. REFERENCES ,

2. 3. .

5. 6. 7. ,

9. 10. 11. 12. 13. 14. 15. 16. 17. 18. 19. 20. 21.

Chem. & Eng. News, April 8, 1996, p. 1. R. K. Grasselli and J. D. Burrington, Adv. Catal. 30, 133 (1981). J. A. Dean, "Lange's Handbook of Chemistry," p. 4.25, McGraw-Hill, Inc., New York, 1992. J. R. Monnier and P. J. Muehlbauer, U.S. Patent No. 4,897,498 (1990). J. R. Monnier and P. J. Muehlbauer, U.S. Patent No. 4,990,773 (1990). J. R. Monnier and P. J. Muehlbauer, U.S. Patent No. 5,081,096 (1992). J. T. Roberts, A. J. Capote, and R. J. Madix, J. Am. Chem. Soc. 113,9848 (1991). B. Schiott and K. A. Jorgenson, J. Phys. Chem. 97, 10738 (1993). A. M. Lauritzen, U.S. Patent No. 4,833,261 (1989). M. M. Bhasin, P. C. EIIgen, and C. D. Hendrix, U.S. Patent No. 4,916,243 (1990). Commercial ethylene oxide catalyst graciously supplied for purposes of comparison. J. R. Monnier and P.J. Muehlbauer, U.S. Patent No. 5,145,968 (1992). P. R. Blum, U.S. Patent No. 4,894,467 (1990). S. Hawker, C. Mukoid, J. S. Badyal, and R.M. Lambert, Surf. Sci. 219, L615 (1989). J. T. Roberts and R. J. Madix, J. Am. Chem. Soc. 110, 8540 (1988). N. W. Cant, E. M. Kennedy, and N. J. Ossipoff, Catal. Lett. 9, 133 (1991). W. M. H. Sachtler, C. Backx, and R. A. Van Santen, Catal. Rev.-Sci. Eng. 23, 127 (1981). Chem. & Eng. News, August 21, 1995, p. 7. D. Denton, S. N. Falling, J. R. Monnier, J. L. Stavinoha, Jr., and W. C. Watkins, Chimica Oggi, May 1996, p. 17. N. W. Boaz and R. L. Zimmerman, Tetrahedron: Asymmetry, 5, 153 (1994). N. W. Boaz, U.S. Patent No. 5,445,963 (1995).

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3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 1997 Elsevier Science B.V.

151

R e d o x M o l e c u l a r Sieves as Heterogeneous Catalysts for Liquid Phase Oxidations R.A. Sheldon Laboratory for Organic Chemistry and Catalysis, Delft University of Technology, Julianalaan 136, 2628 B L Delft, The Netherlands 1. INTRODUCTION Catalytic oxidation is widely used for the conversion of petroleum-derived hydrocarbons to commodity chemicals [1 ]. Moreover, in fine chemicals manufacture there is increasing environmental pressure to replace traditional stoichiometric oxidations with inorganic reagents such as dichromate and permanganate, with cleaner, catalytic alternatives which do not generate excessive amounts of inorganic salts as byproducts [2]. Catalytic oxidations in the liquid phase generally employ soluble metal salts or complexes as the catalyst. However, solid catalysts have several advantages compared to their homogeneous counterparts, such as ease of recovery and recycling and amenability to continuous processing. Moreover, siteisolation of discrete redox metal centres in inorganic matrices can lead to oxidation catalysts with unique activities and selectivities by virtue of the fact that oligomeriTation of active oxometal (M = O) species to inactive ~t-oxo complexes is circumvented. One approach to designing stable solid catalysts with unique activities is to incorporate redox metal ions or complexes into the framework or cavities of zeolites and related molecular sieves. So-called redox molecular sieves [3, 4], unlike conventional supported catalysts, possess a regular microenvironment with homogeneous internal structures consisting of uniform, well-defined cavities and channels. Confinement of the redox active site in these channels and/or cavities can endow the catalyst with unique activity as a result of strong electrostatic interactions between acidic and basic sites on the internal surface and the substrate or reaction intermediate analogous to interactions with acidic carboxyl and basic amino groups of amino acid residues in the active site of (redox) enzymes. Indeed, the analogy with enzymes can be taken even further: f'me-tuning of the size and hydrophobicity of the redox cavity (see later) can provide unique, tailor-made catalysts by influencing which molecules have ready access to the active site, on the basis of their size and/or hydrophobic/hydrophilic character. Hence, application of the terms 'mineral enzymes' and zeozymes to such catalysts is appropriate [5]. Up until the late seventies attempts to develop redox molecular sieves were mainly limited to the ion-exchange approach (see later). This situation changed dramatically with the discovery, by Enichem scientists in 1983 [6, 7], of the unique activity of titanium silicalite-1 (TS-1) as a catalyst for oxidations with 30% aqueous hydrogen peroxide. Following the success ofTS-1, interest in the development, and application in organic synthesis, of redox molecular sieves has increased exponentially and has been the subject of several recent reviews [8-11 ]. It has even provoked a revival of interest in another approach to producing redox molecular sieves: the so-called ship-in-a-bottle method [ 12-15]. Why use a redox molecular sieve? Although the major motivation stems from the expectation of producing a catalyst with unique activity we note that, in many cases, a stable,

152 recyclable solid catalyst exhibiting the same activity/selectivity as its homogeneous counterpart would be a useful objective. Thus, even trace amounts of 'heavy metals' in aqueous effluent are undesirable which means that catalysts should be recyclable. A good case in point is chromium: the 'chromofobia' which is currently in vogue imposes an essentially zero emission constraint on this metal. 2. OXIDATION MECHANISMS A conditio sine qua non for understanding oxidations of organic substrates with 02, H202 and RO2H catalyzed by redox molecular sieves is a thorough appreciation of the elementary mechanisms of oxidations in the liquid phase [ 1]. One can safely assume that the elementary steps involved in liquid phase oxidations (see Figure 1) do not change fundamentally when the metal catalyst is located in a molecular sieve; only that the relative rates of these steps may change substantially.

Free r a d i c a l c h a i n a u t o x i d a t i o n initiation

RH

diff.controiled

R. + 0 2 RO 2 9+ RH ROaH + M n§ RO2H + M(n-1)+ ROe + RH RO 9+ RO2H 2 RO 2 9 \ 2 CHO 2 9 /

rate iim. slow

(1)

RO 2

(2)

ROaH + Re

(3)

~- RO 2, + M(n+)+ + H +

fast fast fast

termination

R,

=

(4)

RO, + Mn'IOH

(5)

ROH + R,

(6)

ROH + RO 2 9

(7)

2 ROe + 0 2

(8)

~HOH + ~ C = O + ' 0 2 / /

(9)

Catalytic oxygen transfer RO2H + S

M

~--

ROH + SO

(10)

M+SO

(11)

2M--O

(12)

Mars-van Krevelen mechanism M~O+S 2M+O

2

Figure 1. Oxidation mechanisms

153 One aspect which sets oxidation apart from other reactions, e.g. hydrogenation and carbonylation is the fact that there is almost always a reaction (free radical chain autoxidation) in the absence of the catalyst (Reactions 1-3). Moreover, (transition) metal ions which readily undergo a reversible one-electron valence change, e.g. manganese, cobalt, iron, chromium, and copper, catalyze this process by generating alkoxy and alkylperoxy radicals from RO2H (Reactions 4-6). From the viewpoint of selectivity this ubiquitous competing free radical chain autoxidation is often, but not always, something to be avoided: In principle, it can be circumvented by employing H202 or RO2H as the terminal oxidant in an oxygen tranfer process (Reaction 10). Particularly in the context of fine chemicals such reagents can be economically viable. However, as noted above, variable valence metals can also be expected to catalyze the homolytic decomposition of RO2H and H202 via the so-called Haber-Weiss mechanism (Reactions 4 and 5; R = alkyl or H). Moreover, subsequent free radical chain decomposition of RO2H via reactions 7 and 8 leads to the formation of dioxygen which returns the reaction to an autoxidation manifold (Reactions 1-3). In the case of secondary alkylperoxy radicals, termination via the Russell mechanism [16] can even lead to the formation of singlet dioxygen (Reaction 9). Obviously competition between reactions 6 and 7 will depend on the relative concentrations of substrate (RH) and oxidant (R'O2H). In many cases both homolytic and heterolytic pathways afford the same products, e.g. alcohols and ketones from hydrocarbons, which means that results have to be interpreted with care. Certain elementary tests for homolytic pathways need to be performed, e.g. inhibition by a radical scavenger such as Ionol indicates a free radical chain mechanism and loss of yield on flushing with an inert gas suggests the intermediacy of dioxygen in reactions with H20~ or RO2H. More sophisticated 'reality tests' can also be performed to demonstrate the intermediacy of alkoxy radicals in oxidations with RO,H [17]. The 'holy grail' of oxidation chemistry is the design of catalytic systems capable of mediating oxygen transfer from dioxygen, without the need for a sacrificial reductant, i.e. a Mars-van Krevelen mechanism [18] in the liquid phase. Indeed, the confinement of substrate molecules in the micropores of molecular sieves might be expected to create quasi gas phase conditions conductive to such a mechanism at the expense of free radical chain autoxidation (we note, however, that a mechanism involving two metal centres as shown in eq. 12 would be difficult to achieve in a molecular sieve). The active oxidant in catalytic oxygen transfer processes may be an oxometal (M = O) or a peroxymetal (MOOR) species (Figure 2) [19]. It will be readily appreciated that catalytic oxygen transfer may be considered as a special case of the Mars-van Krevelen mechanism. Peroxometal mechanisms are favourable when the metal in its highest oxidation is both a Lewis acid and a weak oxidant;e.g, early transition metal ions with do configurations (Mo v~, W vl and Tirv). The oxidation state of the metal ion does not change during the catalytic cycle; catalysis being due to the Lewis acid character of the metal ion. Strong (one-electron) oxidants, exemplified by later and/or first row transition elements such as Crv~, Mn m, Co tu and Fe I", favour oxometal pathways and/or Haber-Weiss type homolytic decomposition of RO2H. Vanadium(V), being a strong Lewis acid and a reasonably strong (one-electron) oxidant exhibits all three types of activity.

154

( RO,,H

ROM n*

....=

-ROH

v

M.*

=

MoVI

b)

M(n*2)+

-

Zr w

(a)1

\O/

I

[

Ti w

S

M" * OR

RO

V v

WVX

ROMn* + SO

0 =

V v ' C r vl , F e v , R u vx

M .§

Co m

C r vI

Mn m

Vv

Fe m'

E~

1.82

1.48

1.51

1.0

0.771

I I

I

1

TiIV

MoVX

WVl

0.08

0.2

0.03

Figure 2. Peroxometal (a) vs oxometal Co) pathways in oxygen transfer Hence, the oxometal pathway is more complicated than one might assume from Figure 2. Conversion of peroxometal to oxometal species can involve either homolytic or heterolytic cleavage of the O-O bond (Figure 3). If'leak~e' of RO. (or HO- in the case of H202) occurs, the reaction is transferred to an autoxidation manifold (Figure l). Competition between homolytic and heterolytic pathways for oxometal formation will be influenced by many factors, e.g. ligands surrounding the metal, solvent, etc. In short, catalytic oxidations with 02, H202 and RO2H are, from a mechanistic viewpoint, intricately interwoven, including various homolytic and heterolytic pathways. Examples of oxidative transformations involving heterolytic, pemxometal pathways are olefin epoxidation, sulfoxidation and oxidations of various nitrogen compounds. In contrast, allylic and benzylic oxidations and (cyclo)alkane oxidations are typical of oxometal and/or free radical autoxidation pathways, which are difficult to distinguish. Alcohol oxidations may involve peroxometal or oxometal pathways. There are few cases, e.g. stereoselective olefin epoxidation with TiW/RO2H (R ffi H or alkyl) must involve a heterolytic, peroxometal pathway, which are unambiguous.

155

M.~...01~0 R hemolysis= I!(,.~)+.O 1 diffusion....=RO, =

recombination Mn+(~ b,,'f~'ORheterolysis /?~") ~,,~-o/ = 0 OR Figure 3. Hemolysis vs heterolysis ofperoxomelal complexes

3. REDOX MOLECULAR SIEVES: GENERAL CONSIDEI~TIONS 3.1. Structures and composition of molecular sieves Zeolitcs and zeotypes are crystalline oxides comprising comer slmfing TO4 tetrahedra (T = Si, AI, P, etc.) and consisting of a regular pore system with diameters of molecular dimensions, hence the term molecular sieve. Zeolites refers to altmfinosilicates (T = Si and AI) and zeotypes to molecular sieves having analogous sm~tures but with a different elemental composition, e.g. aluminophospha~s (AIP0's; T = AI and P) and siticoaluminophosphates (SAPO's; T = Si, AI and P). Molecular sieves are categorized on the basis of their pore diameters into small pore (< 4 A~, medium pore (4-6 A), large pore (6-8 A), extralarge pore (8-14 A) and mesoporous (15-100 A). The pore system may be one, two or three dimensional. This can be important in the context of catalysis as a few obstructions in a one dimensional pore system would seriously impede access to a large proportion ofthe catalytic sites while in two or three dimensional pore systems alternative diffusion paths can be found. A molecular sieve having a particular topology is described by a mnemonic three letter code [201. The AFI structure (7.3 ~), for example, is one dimensional while the FAU structure (7.4 A) consists of three orthogonal channel sytems (7.4 A) intersecting in larger cavities (13 A), so-called supercages, and molecules can travel in all three directions. Other selected examples are given in Table 1.

156 Table 1 Pore dimensions and dimensionalities of molecular sieves Pore size Small Medium

Structure type LTA

Trival name

Pore Ringdiameter [/~] size

Zeolite A

4.1

8

3 3 3

MFI MEL AEL

ZSM- 5, TS-I ZSM-11, TS-2 AIPO4-11

5.3 x 5.5 5.3 x 5.4 3.9 x 6.3

10 10 10

MOR BEA FAU

6.5 x 7.0 7.6 x 6.4 7.4

12 12 12

7.4 x 6.5

12

2

LTL AFI

Mordenite Zeolite beta Faujasite Zeolite X or Y Hexagonal faujasite Linde type L AIPO4-5

7.1 7.3

12 12

1 1

VFI CLO

VPI-5 Cloverite

12.1 13.2

MCM-41

ca. 40

Large EMT

i

Extra large Mesoporous

Dimensionality

'

18 20

1

I I

1

3 1

The framework of molecular sieves is not completely rigid and incoming molecules are able to induce slight structural changes. Hence, ca. 10% should be added to the pore diameters given in Table 1 to obtain the limiting sizes of molecules having access to the pores. Their well-defined pore systems combined with their capacity for at least small substrate-induced structural changes enable molecular sieves to recognize, discriminate and organize molecules with a precision of < 1 A [21 ]. This capability to organize and discriminate molecules with high precision endows them with shape selective properties [22], analogous to enzymes. Hence, also by analogy with enzymes one would expect the highest activity to be observed with the best fit, i.e. when the dimensions of the substrate are comparable to those of the micropores. The recently discovered mesoporous (alumino)silicates, e.g. MCM-41, consist of a regular array of uniform one dimensional pores with diameters in the range 15-100 A and have properties intermediate between those of amorphous SiO2 and A1203 and microporous sieves [23]. This has considerably extended the size of molecules that can be adsorbed: the immobilization of enzymes in MCM-41 has even been achieved [24]. 3.2 A c i d i t y a n d h y d r o p h o b i e i t y

In zeolites the different valencies of Si (tetravalent) and AI (tdvalent) produce an overall negative charge for each A1 atom which is balanced by an alkali (alkaline earth) cation. Exchange of these cations by protons affords strong Brvnsted acids, comparable in strength to sulfuric acid. The acid strength increases with decreasing Al content, while the number of acid sites decreases. Substitution of tfivalent aluminium in the zeolite framework by tetravalent

157 species, e.g. Si or Ti produces (metallo)silicalites with an electrically neutral, hydrophobic framework. Likewise, substitution of silicon by phosphorus produces the electrically neutral, hydrophilic aluminophosphates (AIPO's) or acidic SAPO' s.

3.3 Types of redox molecular sieve We can distinguish three types ofredox molecular sieve on the basis of the method of synthesis: (a) ion exchange, Co) framework substitution and (c) encapsulation. As noted above, the negative charges of zeolite and SAPO frameworks are compensated by exchangeable cations, so that redox cations can be introduced by direct ion exchange. However, the high mobility of the exchangeable cations results in a marked propensity for leaching. Moreover, ion exchange is not applicable to neutral molecular sieves such as AIPO's and silicalites. Framework substitution of A1, Si or P by a redox metal ion leads, in general to more stable redox molecular sieves (Figure 4). So-caUed isomorphous substitution, in which the metal ion is coordinated tetrahedraUy by four oxygen atoms should be possible when the r~io~/roxys~ ratio is between 0.225 and 0.414 [25]. It should be noted, however, that the oxidation state of the metal and, hence, structure and charge of the framework, may change substantially when an as-synthesized material is calcined. For example, Cr-substituted sieves generally contain Crm in the as-synthesirzd material but on calcination this is transformed to Cr vl. Since the latter contains two extra-framework Cr = 0 bonds it can only be anchored to a surface defect site rather than isomorphously substitutexl. By the same token, as-synthesized CrmAPO contains a neutral frwnework (Crm replaces AIm) while in the calcined CrVIAPOthe fraraework contains one Brznsted acid site per Crv~. Hydrophobicity

I

205 !

I AI.

I I

ALPOs

I

M-APO

T

I !

SAPOs

Redox MetaI,M

! s ~)2

I [

Y

Fe

Ti

Zr

III

IV

IV

III

IV

IV

If

ZEOLITES

l

Ti-beta

V

IV I IV IV I

V

Cr

9

ii T

TI-ZSM-5

Sn I I

']

m

i

J IM-SILICALITES I

M-ZEOL

-TAPSO -CrAPSO

-VAPO -CrAPO -CoAPO

As-synthesized Calcined

~

Mn

Co

III

II

II

VI

III

III

TS-1 VS-1 CrS-1

I

Figure 4. Types of redox molecular sieves and oxidation states of the metal ions

158 Another approach to creating novel redox active molecular sieves involves the encapsulation of transition metal complexes in intrazeolitc space: the so-called ship-in-abottle method [ 12-15]. Encapsulated metal complexes should, in principle, ideally combine the advantages of homogeneous and heterogeneous catalysis. Molecular sieves containing supercages, e.g. FAU and EMT (hexagonal faujasite) are ideal for encapsulation as substrate molecules have ready access via the micropores (7.4 A) to the metal complex which is trapped in the supercages (13 A) An advantage of encapsulation is that it allows for the synthesis of redox molecular sieves containing elements that are too large to be incorporated by isomorphous substitution (see earlier). 3.4 Choice of solvent Redox molecular sieves have one important advantage over other heterogeneous catalysts: it is possible to influence which substrate molecules approach the active site by a suitable choice of molecular sieve and solvent. The molecular sieve can be viewed as a second solvent which extracts the substrate molecules from the bulk solvent. Which molecules are selectively extracted depends on the size and hydrophobicity of the micropores and of the substrate. Highly siliceous molecular sieves, such as silicalite-1, are hydrophobic and will selectively adsorb apolar hydrocarbon substrates. This is especially important in hydrocarbon oxidations where the primary products - alcohols, aldehydes and ketones - are polar molecules and are more susceptible to oxidation than the hydrocarbon substrates. Hence, by using a redox molecular sieve it is, in principle, possible to obtain much higher selectivities to primary oxidation products than with analogous homogeneous catalysts. The selective oxidation of alkanes with H202 in methanol solvent, in the presence of TS-1 is, presumably, a manifestation of this effect. Catalytic oxidation with H202 in homogeneous solution are generally strongly inhibited by the water present in the H202 or formed in the reaction. Hence, another important advantage of hydrophobic redox molecular sieves, such as TS-1, is that they are not deactivated by water owing to the preferential adsorption of the hydrocarbon substrate and H202. In general, a hydrophilic solvent, e.g. acetone or methanol, is used to create a single liquid phase with aq. H202 although it has been claimed that this is not necessary [26]. Hydrophilicity of redox molecular sieves increases with increasing aluminium content. Hence, high-alumina zeolites, AIPO's and SAPO's are strongly hydrophilic and selectively adsorb hydrophilic substrates. In this case a hydrophobic solvent should be used to facilitate the adsorption of substrates. Furthermore, it should be noted that the incorporation of AI in silicalites or Si in AIPO's generates Brensted acid sites which may catalyze undesirable sidereactions (see later).

4. SYNTHESIS AND CHARACTERIZATION 4.1 Framework-substituted molecular sieves The so-called hydrothermal synthesis of molecular sieves involves allowing an aqueous gel, containing a source of the framework building elements (AI, Si, P) and a structuredirecting agent (template; usually an amine or a tetraalkylammonium salt) to crystallize in an autoclave, under autogenous pressure, at temperatures ranging from 80 to 200 ~ [27]. Crystallization times can vary from several hours to weeks. Redox molecular sieves are similarly prepared by adding a source of the redox metal ion to the synthesis gel. The as-

159 synthesized material is calcined at ca. 500 ~ to remove the template. As noted above, this can lead to oxidation of the redox metal ion to a higher valence state. In the synthesis of silica-based materials a mineralizer (OH, F) is required to regulate the dissolution and condensation process, i.e. synthesis is generally carried out at high pH. In contrast, (redox) aluminophosphates are crystallized from gels prepared by mixing an alumina slurry with a solution of the redox metal ion in aq. H3PO4 and the template, i.e. synthesis occurs at low pH. Titanium-substituted silica-based molecular sieves, in particular TS-1 (MFI), have been the most intensively studied [6, 7, 9]. This generally involves controlled hydrolysis of a mixture of Si(OEt)4 and Ti(OEt)4 in the presence of the template, the tetrapropylammonium cation in the case of TS-1. Many workers have experienced problems in TS-1 synthesis and the various pitfalls haven been reviewed [9]. Small amounts ofNa § or K § orginating from commercial samples of the template suffice to prevent the substitution of Ti into the framework. Similarly, the presence of F leads to the formation of octahedral extra framework titanium. Following the success of TS-1 a variety of Ti-substituted molecular sieves were prepared by hydrothermal synthesis (Table 2) [28-32]. Furthermore, various redox metals, e.g. V, Cr, Mn, Fe, Co, Cu, Zr, and Sn, have been reportedly incorporated into silicalites, zeolites, A1PO' s and SAPO's [8-11 ] and the list is still increasing. Alternatively, framework substitution can be achieved by post-synthesis modification of molecular sieves, e.g. via direct substitution of A1 in zeolites by treatment with TiCI4 in the vapour phase [34] or by dealumination followed by reoccupation of the vacant silanol nests. Boron-containing molecular sieves are more amenable to post-synthesis modification than the isomorphous zeolites since boron is readily extracted from the framework under mild conditions [35]. Synthesis of framework-substituted molecular sieves via post-synthesis modification has the advantage that it is applicable to commercially available molecular sieves which have already been optimized for use as catalysts.

Table 2 Titanium-substituted molecular sieves Material

Template

Pore size (A)

Ref.

Ti-silicalite- 1 (TS- 1)

Pr4NOH

5.3 x 5.5

6,7

Ti-silicalite-2 (TS-2)

Bu4NOH

5.3 x 5.4

28

Ti-ZSM-48

H2N(CH2)sNH2

5.4 x5.1

29

Ti-beta

Et4NOH

7.6 x 6.4

30

Ti-MOR

none

7.0x6.7

31

Ti-APSO-5

C6HIINH2

7.3 x 7.3

32

Ti-MCM-41

CI6H33(CH3)3NOH

ca. 40

33

Ti-HMS

Cl2H25NH2

ca. 40

33

160 A veritable arsenal of techniques has been mobilized to provide information regarding the structure of redox molecular sieves [9-14]. X-Ray powder diffraction (XRD) provides an immediate check for crystallinity and structural type. X-Ray absorption fine structure spectroscopy (EXAFS) and X-ray absorption near edge spectroscopy (XANES) give further insights into coordination geometry and bond lengths. Infrared and Raman spectroscopy have been used to identify characteristic features, e.g. the 960 cm -~ bond attributed to the Si-OTi stretching vibration in TS-1 [9]. Diffuse reflectance UV-Vis spectroscopy (DREAS) and EPR provide useful information regarding the oxidation state of the metal. Other techniques that are regularly applied are MAS-NMR, X-ray photoelectron spectroscopy (XPS) and scanning and transmission electron microscopy (SEM and TEM). Finally, BET surface area measurements and adsorption experiments are indispensible for checking the structural integrity of the molecular sieve. 4.2. Molecular sieve-encapsulated metal complexes Three different approaches are used to achieve encapsulation: a) Intrazeolite complexation (flexible ligand method) b) Intrazeolite ligand synthesis c) Metal complexes as templates for zeolite synthesis In the first method the metal complex is assembled in the zeolite cavities by allowing the metal-exchanged zeolite to react with ligands that are small enough to access the micropores. The metal complex, once formed, is too large to diffuse out. For example, bis- or trisbipyridyl complexes of FeII, Ru n, Mn H, Co t~and Cun have been encapsulated in zeolite Y (FAU) [ 12-15, 36], Metal-Salen and related Schiff's base complexes have been similarly encapsulated in faujasites [12-15, 37, 38]. However, in this case there is virtually no difference in kinetic diameter between the complex and the free ligand and metal-Salen complexes are readily leached by protic solvents, such as ethanol [ 12]. In the second method the ligand itself, constructed by intrazeolite synthesis, is too large to exit the supercages via the micropores. Most examples of this type pertain to FAUencapsulated metallophthalocyanines, first reported by Romanovsky and coworkers in 1977 [39]. They are prepared by first introducing the metal into the zeolite and then adding 1,2dicyanobenzene, which reacts at elevated temperatures to form the metallophthalocyanine in the supercages. Different methods have been used to introduce the metal ion: via ion exchange or as a metal carbonyl or metallocene [12]. In the former two cases the phthalocyanine ligands are largely metallated but many uncomplexed metal ions are also present. In the rnetallocene method, in contrast, there are no uncomplexed metal ions present but a large proportion of the encapsulated phthalocyanine ligands are metal-free. By using substituted 1,2-dicyanobenzenes encapsulated analogs of substituted metallophthalocyanines can be prepared [ 12]. In the template method the zeolite is allowed to crystallize around the metal complex which is assumed to act as a structure directing agent, i.e. the bottle is built around the ship. This allows for the encapsulation of well-defined complexes without contamination by the free ligand or uncomplexed metal ions (see above). The method is restricted to metal complexes that are stable under the relatively harsh conditions of temperature and pH involved in hydrothermal synthesis. Balkus and coworkers [14, 40, 41] used this approach for the encapsulation of metallophthalocyanines in faujasite. However, in order to fit into the faujasite supercages the phthalocyanine ligands are strongly deformed and Jacobs has

161 expressed some doubt [ 12] regarding the structure-directing capability of a template that requires initial deformation. The characterization of zeolite-encapsulated complexes is by no means simple and the same battery of techniques (see earlier) has been brought to bear [14] as with frameworksubstituted sieves. 5. CATALYTIC OXIDATIONS- FRAMEWORK-SUBSTITUTED MOLECULAR SIEVES 5.1. Ti, Zr, Sn and V Titanium(IV) silicalite (TS- 1), the first example of a framework-substituted redox molecular sieve, catalyzes a variety of synthetically useful oxidations with 30% aqueous hydrogen peroxide under mild conditions (see Figure 5). Examples include phenol hydroxylation [42], olefin epoxidation [43], cyclohexanone ammoximation with N113/I-I202 [44], secondary alcohols to ketones [42], primary amines to oximes [45], secondary amines to hydroxylamines [46], sulfides to sulfoxides [47] and alkane oxygenation [48]. The remarkable reactivity of TS-1 is believed to be largely due to site-isolation of tetrahedral titanium(IV) in a hydrophobic environment. The latter ensures that hydrophobie substrates will be adsorbed from a reaction medium containing large amounts of water.

OH

O

NOH

11

1 : 1 o:p

~'

R~/ "- C H 2

L~

o NH a

R2CHOH + ~.

[ :

/R,S

I~-1

+

~ J

R2NH ~ , " =_ R2NOH

~RR'CHNH, RR'CHOH

RzSO

RR'C-'-NOH RR'C~O

Figure 5. TS-1 catalyzed oxidations with H20~

162

The TS- I catalyzedhydroxylationof phenol to a I:I mixture of catcchol and hydroquinonc has been commercialized by Enichem. Similarly,the arnmoximation of cyclohcxanone is being developed commercially as a low-saltalternativeto the conventional process for the production of cyclohexanone oxime, the raw materialfor nylon-6. The reaction involves initialTS-I catalyzedoxidationof NHa by H20~ to give NH2OH. The factthat bulky kctones such as cyclododecanone undergo ammoximation is consistentwith subsequent reaction of N H ~ O H with the ketone substratetaking place outsidethe molecular sieve.The method has been used [49] for the conversion of p-hydroxyacctophenonc to the corresponding oxime which is the precursorof the analgesicparacetamol (Reaction 13). NHCOCH

NOH

1

NH3/H202

(13)

TS-I OH

OH

OH

TS-1 is a particularly active catalyst for olefin epoxidation [43], even unreactive olefins such as allyl chloride being smoothly epoxidized at temperatures close to ambient. Relative reactivities of olefin substrates are completely different to those observed in analogous homogeneous systems. Owing to the steric restrictions of the micropores of TS-1 only straight-chain olefins are smoothly epoxidized. Cyclohexene is completely unreactive (see Table 3). Similarly, in contrast to homogeneous titanium catalysts, TS-1 shows no enhanced reactivity towards allylic alcohols indicating that there is no coordination through the OH group. Table 3 TS-1 catalyzed epoxidations with H202a Olefin

T (~

conv. (%)

Epoxide sel. (%)'

H202

Propene

40

72

90

94

l-Pentene

25

60

94

91

1-Hexene

25

70

88

90

Cyclohexene

25

90

10

n.d.

Allyl chloride

45

30

98

92

Allyl alcohol

45

35

81

72

.

_

t (min)

.

.

.

.

.

_

"MeOH solvem; olefin~202 molar ratio = 5; data taken from M.G. Clerici and P. Ingallina, J. Catal., 140 (1993) 71.

163 The solvent of choice is methanol which gives higher rates than aprotic solvents [43]. This is attributed to the formation of a titanium(IV)-hydroperoxo comples (I) in which coordination of the alcohol promotes oxygen transfer to the olefin (Figure 6). Coordination of the alcohol becomes increasingly difficult with i n ~ g steric bulk, consistent with the observed decrease in rate methanol > ethanol > tert-butanol.

sio

\

SiO --Ti--OSi / sio

SiO

R

SiO __XTiJ~H SiO

/C~~C~'~H

H202

SiO HOSi

\ ~- S i O - - T i

OH

ROH

sio/ \o /

c--c

sio \

/o\ c--c

=- SiO - - / T i - SiO

OR

Figure 6. Mechanism of TS-I catalyzed epoxidation Similar heterolytic mechanisms can be envisaged for other nucleophilic substrates, e.g. ammonia, amines, sulfides, phenols, alcohols. With alkanes or aromatic hydrocarbons, on the other hand, homolytic mechanisms, with possible involvement of HO. radicals, would seem more likely. A titanium(IV)-silicalite-2 (TS-2) catalyst having the MEL topology gives similar reactions to TS-I [50]. However, the scope of TS-1 and TS-2 catalyzed oxidations is limited to the relatively small molecules which can access the micropores (5.6 x 5.3 A and 5.3 x 5.4 A, respectively). This stimulated several groups to investigate the incorporation of titanium into larger pore sieves. Thus, Corma and coworkers [30] s u ~ e d in incorporating titanium in zeolite beta. The resulting titanium~silicoaluminate, Ti-BEA, catalyzed the oxidation of larger substrates such as cyclohexene and cyclododecane [51]. However, owing to the Br~rnsted acidity of Ti-BEA, the major products of olefm oxidation were glycols and glycol monomethyl ethers resulting fTom ring opening of the epoxide by H20 or MeOH, respectively. We subsequently showed [52] that ring opening could be suppressed by simply neutralizing the Brvnsted acid sites by ion exchange with alkali metal ions (see Table 4). Recently, aluminium-free titanium-substituted beta was synthesized, using a different template, and shown to catalyze the epoxidation of olefms with H20,, albeit with some ringopening [53].Ti-BEA also catalyzes epoxidations with TBHP [54], in contrast to TS-1 which cannot accommodate the transition state for epoxidation with the bulky TBHP. Ti-APSO-5, which also contains Brznsted acid sites afforded the diol as the main product in the oxidation of cyclohexene with H202 while with TBHP the epoxide was formed in 79% selectivity [55].

164

Table 4 Titanium-catalyzed epoxidations Olefin

Catalyst

Oxidant

, Conv. ]

Res

Selectivity (%)

(%) /

m

1=Hexenea

m

H202

TS-I Ti-beta

Cyclohexene

i

TS- 1 Ti-beta

,,....

98 80

H202 i

.,

,

epoxide

glycol(ether)

96 .12

4 88

nn

,,

:

l

m

,

,,=,l,

l

0. Nonetheless, the NASICON structure provides for some desired V site isolation, however, apparently not complete and hence not sufficient to achieve our desired catalytic goal. Another observed fact is, that the Nb-Ti-V-P-oxide under investigation shows an amorphous overlayer via TEM which is enriched in vanadium. The (V/P)surface > (g/P)particle" One can reason that at the temperature of 900 ~ required to obtain the NASICON structure, the more

225 40

/ / / / / / /

O

E

30

LIJ r'n

.

ICl >...

/

"l"-9 20 Z
HzCO (g)+ 1/2H2(g)). Excess oxygen adatoms on Mo(112)-p(1• which were not incorporated into the p(1• structure, enhanced the selectivity to formaldehyde from 50 % to 90 % and lowered the activation energy of the methanol oxidation. Such oxygen adatoms were more reactive than the oxygen atoms of the p( 1 • structure and reacted with the methoxy species to form H20 by the oxidative dehydrogenation mechanism (CH30(a) + 1/20(a)~ H2CO(g) + 1/2 H20(g)). In a constant flow of methanol, the reaction proceeded several cycles but was deactivated by C(a) accumulated on the surface. The selective oxidation of methanol in flow conditions of CH3OH and 02 successfully proceeded on the Mo(112)-p(1• surface without deactivation.

1. I n t r o d u c t i o n

Control of reaction paths on catalyst surfaces by optimizing the structure and electronic properties is a key issue to be solved in surface science. Iron/molybdenum oxides are used as industrial catalysts for methanol oxidation to form formaldehyde selectively. The iron /molybdenum oxide catalyst consists of Fez(MoO4)3 and MOO3, and shows kinetics and selectivity similar to those of MoO3 for methanol oxidation [1]. It suggests that Mo-O sites play an important role in the reaction. MoO3 has a layered structure along a (010) plane, but the (010) surface is not reactive because it has no unsaturated Mo site [1]. On Mo metal surfaces such as (100) [2,3] and (112) [4], major products in methanol reactions were H2 and CO. Therefore, we considered that partial oxidation of Mo sites is needed for the selective oxidation of methanol. We have reported that methanol reaction pathways on Mo(112) could

228

van der Waals snhere \

oxygen adatom \ \ [1 TO]

oxygen atom

ii, ,!ii

!,, s,,

I

......

, il I

....

iiii,.

....

ti!i!,-

It

p(1 x2)-O

I

0 adatom + p(1 x2)-O

Figure 1. Models for oxygen-modified Mo(112) surfaces. be controlled by modification of the surface by oxygen atoms [4-6]. Formaldehyde (H2CO) was a major product during temperature-programmed reaction (TPR) of methanol on a Mo(l12)-p(l• surface (0o=1.0), while CH4, HE, C(a), and O(a) were the products on surfaces with lower coverages of preadsorbed oxygen. Besides, the reaction on Mo(112)p(1• proceeded without formation of H20 (CH30(a)--~ HECO(g)+ 1/2HE(g)). We have suggested that formaldehyde was formed due to suppression of C-O bonds of methoxy intermediates by selective blocking of the second-layer Mo atoms with the oxygen atoms. Oxygen modification of metal surfaces have been examined on methanol reactions. On some metal surfaces such as Cu(110) [7], Cu(111) [8], Cu(100) [9,10], Ag(110) [11], Ru(001) [12], Rh(111) [13], and Fe(100) [14], oxygen atoms enhanced the formation of methoxy intermediate by extracting the hydroxyl hydrogen of methanol to form OH(a). Another effect of oxygen modification was to stabilize the methoxy species as observed on Ni(110) [15], Mo(100) [3], and W(112) [16]. On Fe(100) surface, the stabilization of methoxy by oxygen atoms resulted in a change of selectivity [ 14,17]. The Mo(112) surface has a ridge-and-trough structure, where the top layer Mo atoms form close-packed atomic rows along the [ 111] direction separated by 0.445 nm from each other and oxygen atoms are expected to occupy quasi-3-fold sites composed of one second-layer and two first-layer Mo atoms [18]. Left-half of Figure 1 shows a model of the p(1• surface (0o= 1.0) which was proposed on the basis of LEED patterns and CO titration experiments [ 19]. Every second Mo row is coordinated by oxygen atoms on both sides(Mo2c), while the other Mo rows are not directly coordinated (MONc). This structure preserves adsorption sites on MoNc for admolecules such as CO, methanol, and ammonia. Selective blocking of the second-layer Mo atoms by oxygen atoms suppressed bondbreaking of C-O or N-H, resulting in m

CH30 [4-6] or NHx (x~2) [20] species persisting on the surface up to 500 K. In this study, we show that excess oxygen adatoms on Mo(112)-p(1• which are not incorporated into the p(1 • structure, enhance the selectivity to formaldehyde and lower the activation energy of the methanol oxidation. The selective oxidation of methanol in a flow of CH3OH and 0 2 successfully proceeds on Mo(112)-p(1• without deactivation.

229

2. Experimental The experiments were performed in an ultrahigh vacuum chamber which was equipped with a low energy electron diffraction (LEED) optics, which was also used for Auger electron spectroscopy (AES), and a quadrupole mass spectrometer (QMS) for TPR. A Mo(112) sample was cleaned by cycles of At+ sputtering and annealing to 1300 K. The sample could be cooled to 150 K by liq.N2 and resistively heated at linear sweep rates of 0.5-15 K/s. The cleanliness and surface order were checked by AES and LEED. The clean surface exhibited a sharp and well-contrasted p(1 • 1) LEED pattern, indicating that the surface preserves the bulk truncated structure. The p(1 • structure (0o-=1.0) was prepared by exposing the clean surface to 2 L (1 L=1.33• Pa.s) of O2 at 300 K and subsequent annealing to 600 K [19]. The oxygen coverage was monitored by AES and LEED patterns.

3. Results and Discussion 3.1. Temperature-programmed reactions (TPR) TPR spectra from the Mo(112)-p(1 • surface exposed to methanol at 200 K (Figure 2A) showed simultaneous desorption of HzCO, CI--I4, CO, and H2 at 560 K. We showed that the composition of the species remaining on the surface above 400 K was C:O:H = 1:1:3 and that the hydroxyl hydrogen recombinatively desorbed as H2 below 400 K [4]. We considered, therefore, that the intermediate was methoxy species as supposed on other metal surfaces [2,3, 7,8,11-13,15,17,21-23] or on MoO3 [1,24]. Methoxy species were also observed on oxygen-modified Mo(110) by X-ray photoelectron spectroscopy (XPS) [25] and on oxygenmodified Mo(100) by high-resolution electron energy loss spectroscopy (HREELS) [3]. The A, r~

I

'

,

'

I

' .,

,

'

~--.,.,~_

_~.._,..~

I

16 amu (CH4)

-

~

I

'

I

'

I

'

i

I

'

I

16 amu (CH4) 28 amu

(CO)

2 ainu (H2) r

._~ O9u~ ~

32 amu (CH3OH)

~;

30 amu (H2CO)

200

400

600

T/K

800

1000

18 amu (H20) x2 30 ainu (H2CO) 200

,

I

400

,

I

,

600

T/K

I

800

,

I

1000

Figure 2. TPR spectra after exposing (A) a Mo(ll2)-p(lx2)-O surface and (B) a Mo(112)p(l•

surface with 0.15 ML of oxygen adatoms to 4 L of CH3OH at 200 K. The heating

rate was 5 K/s.

230 Table 1 The product distribution in TPR for the methanol reaction around 560 K on the Mo(112) surfaces modified with oxygen. Product s peci es

Y ield / ML p( 1x 2)-O

Oxygen adatoms (00' = 0.15)

surface (00 = 1.0)

+ p(1 x2)-O surface (00 = 1.0)

H2(g )

0.10

0.01

H2CO(g )

0.09

0.06

H20(g ) CO(g)

0 0.02

0.04 0 . 6 having t h e h e x a g o n a l s t r u c t u r e b e c o m e a c a t a l y s t which has high selectivity to t h e O C M reaction. However~ even in the Table 1 C a t a l y t i c a c t i v i t y of t h e oxidative coupling of m e t h a n e over LixNi2_xO2. Catalyst NiO Li0.3 Ni 1.rO2 Li0.6 Ni 1.4 02 Li0.rNil.3 02 Li0.9Nil.lO~ LiNiO2

React. Temp. (K)

CH4- Conv. (%)

C2-Select. (%)

Structure

993 1053 993 1053 993 1053 993 1053 993 1053

21.4 33.8 31.7 33.6 33.3 34.1 37.2 36.1 32.2 42.5

32.0 27.3 40.0 25.4 54.1 20.7 51.4 25.1 53.7 63.6

Cubic C u b ic Cubic Cubic Cubic Hexagonal Cubic Hexagonal Cubic Hexagonal Hexagonal

case of these selective catalysts~ its selectivity d r o p s a b r u p t l y at a high t e m p e r a t u r e at 1053K as shown in Table 1. T h e e x p e r i m e n t a l result over Li0 7Nil.302 is shown Figure 5. T h e s t r u c t u r e of LixNi2_xO2 ( 0 . 7 < x < 0 . 9 ) t r a n s f o r m e d from

250

---Hexagonal type --~

Cubic type

v

70

I!!:: :::i(

60 k,

iii l i ! )...e...~--o--'~ 'i ,ii,i[l

50

! !

9-. 40 ..~ -~ 3O rfl ',.,, ~ 20

,

.

,

,I

.......

:~-~*~_~:

l

a. ~**

~ ~i~,!

D

,

i~i:: ::!

,

l:L::::i

10 00

0.0

0.2

0.4 0.6 X (LixNi2_xO2)

Figure 4. The relationship between LixNi2_xO2 ( 0 < x < l ) catalyst 9 100

A

0.8

C2-selectivity

1.0

and the structure of

O2-conv. O..._o

80

t=,

A

60 .~

C2-select.

.,~

~9

4020

yo / O / o / CH4-conv.

O

.._...O 900

950

1000

1050

Temperature / K

Figure 5. The temperature d e p e n d e n c e catalyst.

o f methane oxidation o v e r L i 0 . ? N i 1 . 3 0 2

251 a hexagonal t y p e to a cubic t y p e at higher t e m p e r a t u r e t h a n 973K. It should be noted t h a t the C2-selectivity may be correlated to the crystal s t r u c t u r e of these catalysts. The catalytic activity of the O C M reaction should be predictable by the n a t u r e of the oxygen species on the catalyst surface. In addition, the electronic states of the surface oxygen could be largely related to the bulk s t r u c t u r e in the crystal.

3.3 A c t i v e sites for t h e O C M r e a c t i o n A typical p h o t o e l e c t r o n spectroscopy e x p e r i m e n t was conducted as follows for u n d e r s t a n d i n g of the catalyst surface. The characterization of the surface oxygen species on these oxides were investigated using XPS and UPS. Figure 6 displays representative typical X P S spectra in the O ls region for Lil.0Nil.002 (Fig. 6-a), Li0.3Ni1.702 (Fig. 6-b) and Ni304 (Fig. 6-c). The O ls spectra in Fig. 6 reveal the presence of at least two different types of the near-surface oxygen species on these catalysts. Two structures, denoted by the characters A and B, were observed at 529.4 and 531.3 eV in the XPS O ls s p e c t r a of Ni304, and at 529.2 and 531.4 eV in those of Li0.3Nil.702. The peak A located at 529.3+0.1 eV could be a t t r i b u t e d to 0 2 - species which the catalytic n a t u r e must be nonactive for the O C M reaction [8]. It seems reasonable to u n d e r s t a n d t h a t the main existence of the 0 2 - species on Ni304 oxides are nonselective for the O C M reaction. A n o t h e r O ls peaks (peak B) at higher binding energy in these oxides may be associated with the catalytic

A

Ols

~

(c)

(b)

m

(a) I

538

9

I

536

,

I

,

I

,

I

.

I

,

i - - ,

530 528 526 Binding Energy / eV

534

532

I

524

,

I

522

Figure 6. X P S O ls spectra of LiNiO2 system. (a) LiNiO2; (b) Lio3Nil.rO2; (c) Ni3Oa.

252 activity for the O C M reaction. T h e peak intensity ratio A / B of the Li0.3Nil.rO2 was given to be 36/64 in the last column on Table 2, and the oxygen species of the peak B is reasonably existence on the Li0.3Ni1.rO2 catalyst surface. B u t this catalyst was low activity for the O C M reaction as described previous section. In t h e case of LiNiO2 with the active and high C2-selectivity catalyst for the O C M reaction, a peak B' in the XPS s p e c t r u m of LiNiO2 was observed at Table 2 The O ls level of the binding energy of LixNi2_xO2 catalysts. Crystal

Structure

Ni304 Li0.3Nil.rO2 Li0.vNil.302 LiNiO2

Cubic Hexagonal Hexagonal

Ols (eV) Peak A 529.4 529.2 529.1 529.2

Ols (eV) Peak B 531.3 531.4 531.6 531.9

Intensity ratio A/B 63/37 36/64 21//79 11//89

531.9 eV to higher binding energy level of the peak B and largely related to the selective O C M reaction. The peak B' was also observed at 531.8 eV in the X P S s p e c t r u m of LiCoO2 which exhibit high .activity for the O C M reaction [9]. The slight difference in the binding energy between the peak B and the peak B' m a y be often a b e t t e r indication of the n a t u r e of oxygen species r a t h e r t h a n the absolute value of the binding energy. Thus, the a p p e a r a n c e of the oxygen species is a t t r i b u t a b l e to the high C2-selectivity for the OC]~I reaction. C

r~

m

(b)

(a) "

20

I

15

"

"

1'0

I

5

"

'

EF=()

Binding energy / eV Figure 7. U P S s p e c t r a of the solid solution LixNi2_xO2. (a) Lio.3Nil.rO2; (b) LiNiO2.

253 T h e results of X P S m e a s u r e m e n t for t h e solid solution series LixNi2_xO2 are s u m m a r i z e d in Table 2. T h e peaks B of Li07Ni1302 w i t h a i n c o m p l e t e hexagonal s t r u c t u r e was also observed at 531.6 eV to slight higher b i n d i n g energy. T h e i m p o r t a n t point to note t h a t the crystal s t r u c t u r e of these c a t a l y s t also largely affected to t h e electronic states of the surface oxygen. Figure 7 shows U P S s p e c t r a of Li0.3Ni1702 and LiNiO2. These s p e c t r a were m e a s u r e d at t h e incident hp of 40 eV w i t h reference to EF as t h e zero on the e n e r g y scale. Several structures~ d e n o t e d by the c h a r a c t e r s A-C~ were observed at 2.0, 6.0, and 10.5 eV in those of LiNiO2 (Fig. 6-a)~ a n d at Eb ---- 2.2, 6.0 and 11.0 eV in t h e U P S s p e c t r a of Li0.3Nil.702 (Fig. 6-b), T h e s t r u c t u r e located at Eb -- 2.0~2.2 eV can be assigned the 3d b a n d of nickel oxide~ a n d the b r o a d peak at 6.0 eV in the valence b a n d can be assigned to the O 2- species associated w i t h NiO [10]. It was suggested from these results t h a t the bulk s t r u c t u r e largely affected t h e surface oxygen species which play an i m p o r t a n t role for the selective of t h e O C M reaction. Figure 8 shows the location of oxygen ions, Ni and Li cations viewed t w o - d i m e n s i o n a l l y along the vertical direction of the cubic closest packing of oxygen ions of LixNi2_xO2 series. As l i t h i u m c o n t e n t was increasingly d o p e d , t h e lattice s t r u c t u r e t r a n s f o r m e d from a cubic t y p e (Fig. 8-a) to a hexagonal t y p e (Fig. 8-c) t h r o u g h a interm e d i a t e t y p e (Fig. 8-b) b e t w e e n these s t r u c t u r e . T h e oxygen species in the cubic lattice is placed in a h o m o g e n e o u s charge d e n s i t y w i t h a r a n d o m mixed Li ~~ Ni 2 ~ and Ni 3~ . On the o t h e r hand, the oxygen species in a hexagonal lattice is placed in a anisotropic charge density b e t w e e n m o n o c a t i o n Li ~ layer and t r i c a t i o n Ni 3~- layer. It is inferred from the relationship b e t w e e n t h e i r s t r u c t u r e of LixNi2_xO2 a n d t h e selectivity of the O C M t h a t Li + and Ni 3~ cations t w o - d i m e n s i o n a l l y configurated as shown in Fig. 8-c is the i m p o r t a n t

0

0

o

Ni

9

Li

9

Ni

or

Li

Figure 8. Active oxygen species of the solid solution series LixNi2_xO2. (a) Cubic type; ( b ) I n c o m p l e t e hexagonal type; (c) H e x a g o n a l type. role for the oxidative reaction. T h e f o r m a t i o n of active O- species occurs in t h e lattice oxygen and on t h e catalyst surface. It is p r o p o s e d t h a t an active O - species of the c a t a l y s t surface e x t r a c t s a h y d r o g e n from m e t h a n e , and accordingly t h e coupling of m e t h y l radicals take place.

254

CONCLUSIONS In this work, we present the results for LixNi2_xO2 ( 0 < x < l ) catalysts of the O C M reaction. In order to u n d e r s t a n d the surface catalytic active species, the characterization of the oxygen species of LixNi2_xO2 oxides is examined by the combined approach of XPS and UPS. From the X P S O ls and U P S spectral analysis, the main oxygen species of LixNi2_xO2 (x ethanol > 2-butanol. 4. DISCUSSION The thermal spreading of metal oxides over oxide supports has been intensively investigated over the past decade and much is currently known about this process [ 1,5]. The driving force for the thermal spreading of metal oxides is related to the lower surface free energy of crystalline oxides such as V205 and MoO3 compared to crystalline oxide supports such as TiO2, S n O 2, A1203, etc. This process is analogous to the wetting of one solid by another induced by the forces of surface tension in order to lower the surface free energy of the system [2]. The low Tamman temperatures o f V 2 0 5 and M o O 3 (345 and 397.5~ respectively) are responsible for the efficient spreading of these metal oxides at temperatures of 400-500~ Furthermore, the spreading kinetics of the metal oxides are (1) dependent on the structure and morphology of the oxide support, (2) enhanced over well-developed crystal planes and (3) dependent on the specific gaseous environment (oxidizing vs. reducing or wet vs. dry) [5]. Under oxidizing conditions and elevated temperatures, the spreading of crystalline V205 and M o O 3 is initiated spontaneously at the metal oxide-support interface and subsequent migration occurs by surface diffusion of the metal oxides via vacancies or unoccupied sites in the two-dimensional metal oxide overlayer. Amorphous metal oxide phases are suggested as a transient form between the crystalline metal oxides and the two-dimensional metal oxide overlayers. Moisture enhances the surface diffusion of the metal oxides [1 ]. Under mildly reducing conditions, the spreading of crystalline metal oxides is significantly retarded due to the much higher Tamman temperatures of the

262 corresponding reduced crystalline metal oxides [1,5]. The present thermal treatment experiments in air revealed that extensive dispersion of MoO3 occurred at 400~ and essentially complete dispersion took place at 500~ for the loose powder physical mixture of 4% MoO3/TiO2. In contrast, very little dispersion was observed for comparable thermal treatments for the loose powder physical mixture of 4% VzOs/TiO 2. The observation that the kinetics of MoO 3 disperion are faster than the kinetics of V205 dispersion were also previously observed [12]. The lack of V205 dispersion by the thermal treatments is somewhat surprising and may be related to the structure and morphology of the titania support employed in the present investigation. The form of the physically mixed metal oxide was also found to significantly affect the dispersion kinetics due to the presence of significant mass transfer limitations in the catalyst pellet relative to the loose powder. The presence of mass transfer limitations in catalyst pellets or wafers typically employed for Raman and IR studies is welldocumented in the literature [ 13]. The present studies demonstrated that significant dispersion of M o O 3 o r V2 05 o n a titania support could not be achieved at temperatures of 500 ~ with the physically mixed oxides in the form of a catalyst pellet. Thus, dispersion of M o O 3 and V205 on oxide supports at much lower temperatures for physically mixed catalysts in the form of pellets can not be due to thermal spreading and must occur by another mechanism. The in situ Raman studies clearly demonstrate that spreading of MoO3 and V205 over different oxide supports in the form of catalyst pellets readily occurred during methanol oxidation at temperatures as low as 230~ Such low temperatures, which are below the Tamman temperatures of these oxides and the temperatures required for thermal spreading in the catalyst pellet (above 500~ implies that thermal spreading is not involved in the spreading mechanism taking place during methanol oxidation. This suggests that a strong interaction between the gas phase components and the crystalline metal oxide phases may be occurring. Formaldehyde is the major selective oxidation reaction product and is known to interact very weakly with metal oxides such as molybdates and vanadates, and adsorbed formaldehyde is readily displaced by the presence of moisture and methanol [14,15]. Moisture interacts strongly with molybdates [ 14,15] and vanadates [8], but the thermal spreading experiments did not result in significant dispersion of the crystalline metal oxides in the physically mixed catalyst pellet. The interaction of carbon dioxide with molybdates and vanadates is extremely weak and adsorption is usually not even observed at room temperature [16,17]. In contrast to these gaseous components, the interaction of methanol with molybdates and vanadates is very strong and is much stronger than moisture since adsorption of methanol can displace adsorbed moisture [14,15]. Furthermore, methanol oxidation over crystalline M o O 3 and V205 results in the deposition of molybdena and vanadia at the exit of the reactor, which possesses lower temperatures [ 18]. This observation suggests that methanol is able to strongly complex with Mo and V present in crystalline MoO3 and V205 to form volatile Mo(OCH3) n and V(OCH3)n complexes. The alkoxy complexes of vanadia and molybdena are well known and are liquids at room temperature possessing high vapor pressures. Thus, the low temperature dispersion of metal oxides over oxide supports during methanol oxidation is due to the formation of volatile metal-methoxy complexes that result in vapor phase transport of the oxides. The dispersion mechanism may also occur by surface diffusion of the metal-methoxy complex, but no such information is currently available. The absence of Mo and V deposits at the reactor exit during methanol oxidation suggests that either surface diffusion or readsorption of the volatile M-alkoxides is also taking place. In summary, a new phenomenon of reaction induced spreading of crystalline metal oxides on oxide supports is observed in the present investigation at temperatures much lower than that required for thermal spreading via

263 solid-state reactions, 200-250~ vs. 400-500~ Thermal spreading depends on the Tamman temperature of the crystalline metal oxide phases and reduced metal oxide phases possess very high Tamman temperatures which significantly retard migration [1,5]. However, essentially complete dispersion ofV205 on TiO2 was observed during methanol oxidation even though the in situ Raman spectra revealed that the vanadia was reduced under the reaction conditions (see Fig. 3). Essentially complete dispersion of M o O 3 o n TiO2 was also observed after treatment of the catalyst in an oxygen-free methanol environment. Thus, the oxidation state of the metal oxide does not appear to influence the kinetics of reaction induced spreading of crystalline metal oxides. Reaction induced spreading of MoO3 on oxide supports during oxidation of higher alcohols is significantly reduced relative to methanol oxidation (methanol > ethanol > 2-butanol). The reduced migration kinetics is most probably related to the stability and reactivity of the various alcohols. The rate determining step during the oxidation of alcohols to their corresponding aldehydes or ketones involves breaking the alpha C-H bond of the alkoxides (the carbon bonded to the alkoxy oxygen), and the stability of this bond is related to the number of additional carbon atoms coordinated to the alpha carbon: stability decreases with increasing number of carbon atoms coordinated to the alpha carbon [ 15,19]. Thus, the methoxy complex is more stable than the ethoxy complex, and the 2-propoxy complex is the least stable among these alkoxy complexes. The greater stability of the Mo-methoxy complex most probably is responsible for the greater volatility and spreading observed during methanol oxidation. The current findings that reaction induced spreading of metal oxides on oxide supports can occur during oxidation reactions at very low temperatures have important implications for commercial applications as well as fundamental studies. The oxidation of methanol to formaldehyde is industrially conducted with F e 2 ( M o O 4 ) 3 . M o O 3 catalysts that contain excess MOO3. The strong interaction between methanol and the MoO 3 component results in the stripping of the molybdena from the catalyst and its deposition as MoO3 crystalline needles at the bottom of the reactor where the temperatures are somewhat cooler. This volatilization phenomenon is responsible for catalyst deactivation (loss of activity and selectivity) and pressure build-up in such commercial reactors [20]. The opposite behavior is observed during methanol oxidation over MoO3/SiO 2 catalysts at 230~ The strong interaction of methanol with Mo and the weak interaction between surface molybdena species and the silica support results in agglomeration and crystallization of the surface molybdena species to beta-MoO3 particles during methanol oxidation [21 ]. A very important consequence of reaction induced spreading of metal oxides during alcohol oxidation is that the catalyst preparation method of many supported metal oxide systems is not critical since the same surface metal oxide species will form during reaction (especially methanol oxidation) [ 12,21 ]. Furthermore, the possibility that reaction induced spreading occurs during oxidation reactions over catalysts composed of physical mixtures needs to be very carefully investigated in such systems before other mechanisms are proposed to account for observed reactivity patterns [22]. 5. CONCLUSIONS A new phenomenon of reaction induced spreading of crystalline M o O 3 and V205 on oxide supports is observed during methanol oxidation at temperatures much lower than that required for thermal spreading via solid-state reactions, 230~ vs. 400-500~ The migration of the metal oxides appears to proceed by the formation of volatile M-(OCH3) . complexes and is not

264 influenced by the oxidation state of the metal oxide (both oxidized and reduced metal oxides are readily dispersed). The kinetics of reaction induced spreading of metal oxides during alcohol oxidation is much slower for higher alcohols because of the low stability of the corresponding M-alkoxides compared with the more stable M-methoxides. REFERENCES

1. 2. 3. 4. 5. 6.

H. Knoezinger and E. Taglauer, Catalysis, 10 (1993) 1. J. Haber, T. Machej and T. Czeppe, Surf. Sci., 151 (1985) 301. D. Honicke and J. Xu, J. Phys. Chem., 92 (1988) 4699. Y. Xie and T. Tang, adv. Catal., 37 (1990) 1. J. Haber, T. Machej, E. M. Serwicka and I. E. Wachs, Catal. Lett., 32 (1995) 101. F. Cavani, G. Centi, E. Foresti, F. Trifiro and G. Busca, J. Chem. Soc., Faraday Trans., 1, 84 (1988) 237. 7. J.-M. Jehng, H. Hu, X. Gao and I. E. Wachs, Catal. Today, 28 (1996) 335. 8. G. Deo and I. E. Wachs, J. Catal., 146 (1994) 323. 9. H. Hu and I. E. Wachs, J. Phys. Chem., 99 (1995) 10911. 10. M. A. Vuurman, I. E. Wachs and A. M. Hirt, J. Phys. Chem., 95 (1991) 9928. 11. G. Went, S. T. Oyama and A. T. Bell, J. Phys. Chem., 94 (1990) 4240. 12. T. Machej, J. Haber, A. M. Turek and I. E. Wachs, Appl. Catal., 70 (1991) 115. 13. Y. Cai and I. E. Wachs, to be published. 14. W.-H. Cheng, J. Catal., 158 (1996) 477. 15. W. Holstein and C. J. Machiels, J. Catal., 162 (1996) 118. 16. K. Segawa and W. K. Hall, J. Catal., 77 (1982) 221. 17. A. M. Turek, I. E. Wachs and E. DeCanio, J. Phys. Chem., 96 (1992) 5000. 18. G. Deo, H. Hu and I. E. Wachs, to be published. 19. W. E. Farneth, R. H. Staley and A. W. Sleight, J. Am. Chem. Soc., 108 (1986) 2327. 20. R. Pearce and W. R. Patterson, Catalysis and Chemical Processes (Wiley, New York, 1981) p. 263. 21. M. Banares, H. Hu and I. E. Wachs, J. Catal., 150 (1994) 407. 22. P. Ruiz and B. Delmon, Catal. Today, 3 (1988) 199.

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

Temperature

Programmed

Desorption of Ethylene

265

and A c e t a l d e h y d e

on

U r a n i u m Oxides. E v i d e n c e o f F u r a n F o r m a t i o n from Ethylene. H. Madhavaram and H. Idriss Materials Chemistry, Department of Chemistry, The University of Auckland, Private Bag 92019, Auckland, New Zealand The reactions of acetaldehyde and ethylene have been investigated on the surfaces of UO 2 and UO3 by temperature programmed desorption (TPD). On UO 2 two molecules of acetaldehyde undergo reductive coupling to C4 olefins. This is due to the fluorite structure of UO 2, which can accommodate large numbers of excess oxygen, up to UO2.25. The vacant surface oxygen of UO 2 were titrated by N20 adsorption and were equal to 1.86 x 1019 molecules/g, representing an apparent surface area of vacant oxygen sites of 1.9 m2/g. On the other hand, ethylene-TPD on [3 UO3 indicated the desorption of acetaldehyde (490 K). In addition, an unexpected product was also observed. This product was identified as furan (C4H40, m/e 68, 39) which desorbed at ca. 550 K with a carbon selectivity of ca. 40 %. Furan formation from ethylene on UO 3 requiring the formation of one carbon-carbon bond and of one carbon-oxygen bond, is most likely accompanied by oxygen depletion from the UO 3 surfaces and subsequent reduction of U cations into lower oxidation states. The observation of furan from ethylene shows that one may obtain oxygenated products with a high carbon number from ethylene (a relatively abundant feed stock) via one single step. 1. INTRODUCTION Oxidation-Reduction reactions very often track the cation oxidation states of oxide materials [1 ]. Changing the oxidation state of a given cation is accompanied by structural change (such as from a rutile or anatase TiO 2 to corundum Ti203 or from orthorhombic V205 to rutile VO 2, i.e., changing of the coordination numbers of metal cations [2]). Another way of changing the oxidation states of cations is by creation of surface defects, where the surface looses its ordered structure [1, 2]. In the case of oxide materials several factors affect oxygen depletion (or in other words reduction of cations) the most important are the mass difference between the cation and the anion, the bond energy, and the formation of ordered or semi ordered cluster defects. While the mass difference is essentially important in the case of reduction via particle bombardment (see Sigmund theory [3]), bond energy and surface structure are most likely the dominant factors during

266 chemical reduction. The investigation of the effects of changing the structures and oxidation states of oxide materials is crucial to the understanding of their catalytic properties. The uranium oxides system is a good candidate for this investigation due to its presence in different stable and metastable structures - the main product of oxidations of uranium metal are UO2, U407, U308, U409, and UO 3 - as well as the presence of a wide range of oxidation states (from +2 to +6) [4]. The main reason for this wide range of oxidation states in U oxide (and the early actinides in general) is relativistic effects [5], which is simply a mass correction for the core electrons that lead to greater shielding of the higher orbitals, or in other words a decrease in the ionization potential and work function. Another important feature of some phases of U oxides is their possibility of accommodating large numbers of interstitial oxygen atoms in clusters, [2:2:2] clusters [6], without changing the crystal structure, such as UO 2 to UO2+x (x up to ca. 0.25); U cations in proximity of these clusters may have higher oxidation states than +4. There is also another implication to the ease of removing electrons form the outer shells, one can change the surface from one electronic state to another by reduction or by oxidation. Recently we observed that H2-reduction as + +4 well as Ar -sputtering of U308 resulted in surface cations exclusively in a U oxidation state [7]. This net change in oxidation state is unlike what one observes on early transition metal oxides such as Ti [8] and V [9] oxides. In addition, and most important, U oxides are known as good catalysts (or as components of catalysts) for serval industrial reactions such as olefins and paraffins ammoxidation [10-13], hydrocarbons dehydrogenation [14], and very recently for total oxidation of chlorinated compounds [15]. It has also been observed that U308 is active in C-C bond formation reactions such as the formation of isobutene from acetone [7]. Despite these technological importance fundamental understanding of the reactivity of U oxides surfaces towards organic reactions is lacking due to the very few amounts of work interested in this system, the most complicate oxide known [16]. This work is devoted to the understanding of the oxidation of CH2=CH 2 on UO 3 and the reduction of CH3CHO on UO 2 by temperature programmed desorption (TPD). Surface and bulk characteristics were investigated by X-ray Photoelectron Spectroscopy (XPS) and Xray Diffraction (XRD) as well as by N20 adsorption.

2. EXPERIMENTAL

TPD at atmospheric pressure was performed using a fixed-bed reactor interfaced to a high vacuum chamber equipped with a Spectra Vision quadrupole mass spectrometer (base pressure ca. 1 x 10.7 ton). The mass spectrometer is multiplexed with an IBM PC which is equipped with a programme (RGA for windows) that allows the monitoring of 12 masses simultaneously at a cycling rate of ca. 5 s. Catalysts were loaded into the reactor and heated under dry air (or hydrogen) for 2 hours at 800 K (or 10 hours, in the case of hydrogen reduction) prior to reaction. After cooling to room temperature (under H2 or air), the carrier gas was displaced by He (ultra pure) before adsorption of acetaldehyde or of ethylene. Acetaldehyde was placed in a saturator at room temperature. In order to obtain surface

267 saturation, dosing of acetaldehyde was performed while monitoring its m/e 29. A decrease in the signal (due to adsorption) followed by signal restoration is indicative of surface saturation (in ca. 2 minutes). In the case of ethylene dosing was obtained upon exposure for 5 minutes. The catalyst was then purged with He for ca. one hour at room temperature in order to remove traces of the reactant in the TPD line as well as weakly adsorbed molecules on the surface of the catalyst. The gas flow was introduced into the chamber through an -1 interface which consists of a leak valve differentially pumped to 10 torr during operating -6 conditions. A constant pressure of ca. 5 x 10 torr was maintained during all TPD runs. During the purging the m/e 29 (the highest m/e of acetaldehyde) or m/e 27 (for ethylene) were monitored and the TPD started when no change in this m/e signal was observed (after ca. one hour of purging time). The ramping rate during TPD was kept fixed at 15 K/min. The fragmentation pattems of each product were checked in order to identify unambiguously the reaction products by the method described previously [17]. This involved: (a) the separation of the desorption peaks into different domains of temperatures, (b) the analysis of the fragmentation pattern of each product separately, (c) starting from the most intense fragment for each product (m/e 29 for acetaldehyde, for example) and subtracting the corresponding amount of its fragmentation until the majority of the signals were accounted for. XPS was performed using a KRATOS XSAM-800 model with a base -9 pressure of ca. 10 torr. U (4f), O(ls), C(ls) and Ar(ls) (in the case of the Ar-ion sputtered samples) regions were scanned each run. Unreduced samples were loaded into the system without further treatment. Ex situ reduced samples (using H 2 at the same conditions as for TPD) were exposed to air (although, under oxygen free N 2 flow) for about 30 to 60 s, at room temperature, before introduction into the XPS chamber. Ar-ion sputtering was m /

performed using a direct beam KRATOS ion gun at a pressure of ca. 5 x 10 torr. Mg Ko~ radiation was used at 170 W. Collection of spectra were performed at a pass energy of 38 eV. Sample charging up to 5 eV occurred under X-ray irradiation. Binding energies were calibrated with respect to the signal of adventitious carbon (binding energy at 284.7 eV). No charging was observed with UO 3 samples. XRD data were collected using a Phillips 1130 generator, and a Phillips 1050 goniometer. XR radiation was achieved using a Cu tube (broad focus) (Kc~; X = 1.514 A) at 44 kV and 20 mA. N20 titration was performed in a pulse reactor. Pulses of N20 were introduced into the reactor at 480 K. A thermal Detector at the end of a Porapack Q column allowed the monitoring of N2 and N20 peaks. The absence of formation of N 2 accompanied by total restoration of N20 signal was indicative of total titration. This took about 4 pulses of l ml each (1 atm.) per 1 g of UO 3. UO 3 was prepared from a uranium nitrate solution by precipitation with NH 3 at pH 9. After filtration and drying at 373 K over night the powder was calcined at 673 K for 5 hours. XRD indicated a pure 13UO 3. Polysrystalline U30 8 (from BDH No. 26216) was used as received.

268 3. RESULTS 3.1. Surface and bulk characterisation of UO 2 and UO 3.

XRD spectra of UO3, U308 and H2-reduced UO 3 are presented in Figure 1. H2-reduction of UO 3 at 800 K for 10 hours resulted in transformation of the monoclinic phase of [3 UO3 into the orthorhombic fluorite structure of UO 2, although some orthorhombic ~ U308 is also present. Similar results were observed from H2-reduction of U308 [7], with complete transformation of o~U308 to UO 2, however. Table 1 Titration of oxygen vacancies by N20 adsorption on unreduced UO 3 and H2-reduced UO 3 (UO2). Reactor temperature 480 K, BET surface area of UO 3 = 33 m2/g, 1 g of catalyst, reduction temperature 773 K, 16 hours. Pulse number

N 2 (molecules)

N20 (molecules)

H2-reduced UO 3 (UO2) (1 g) 1 1.37 1019

1.32 1019

2

0.33 1019

2.4 1019

3

0.14 1019

2.6 1019

4

0.02 1019

2.75 1019

5

negligible

2.78 1019

total

1.86 10

19

Unreduced UO 3 (1 g) 1

no reaction

2.78 1019

XP spectra of UO2, U308, and UO 3 were analysed elsewhere [7]. A brief description is +

given here. Figure 2 (adapted from ref. 7) shows the XPS U4f region of UO 3, Ar -sputtered UO 3 and H 2 reduced U308. Three important points need indication. First, XPS U4f7/2 and U4f5/2 of UO 3 (spectra a) are higher in binding energy than those in spectra b and c. Second, spectrum a does not contain satellites while both spectra b and c contain satellites + at 386.5 and 397.5. Third, the XPS U4f peak positions of Ar -sputtered UO 3 as well as of H2-reduced U30 8 are those of UO 2 and UO2+ x respectively (see Table 1 in ref. 7), clearly indicating that one can shift the cation oxidation state from one position to the other (from +6 in UO 3 to +4 in UO2). This is unlike early transition metal oxides where, although they + are sensitive to H 2- or Ar - reduction, the resultant surfaces still contain considerable

I

XRD H2 - reduced U 0 3

'3O8

Ar+-sputtered UO, El

a,

u

El

L

U

d\x,j H,-reduced U,08

I

I

I

I

20

40

60

80

28 Figure 1. XRD of U03, U30s, and H2 - reduced U03 (mainly U02).

WOJ

400

395

390 385 380 375 Binding Energy (eV)

Figure 2. X P S of U03, Art -sputtered UO3 (U02) and H2 - reduced U30x (UOz,x)

270 amounts of stoichiometric phases. This unique characteristics of U oxides affect its chemical reactivity (see below), particularly with regard to oxidation-reduction reactions. The pulse method of N20 was investigated on UO 3 and H2-reduced UO 3 (UO2) (Table 1). This method is successful for titration of the surface area of Cu ~ and Ag o catalysts [ 18] and we wanted to try it to titrate oxygen vacancies instead of using oxygen in order to avoid formation of multilayers of dioxygen on the surface. N20 dissociated on UO 2 but not on UO 3 (Table 1).The dissociation reaction is activated, below ca. 425 K no dissociation occurred. A temperature of 480 K was observed as optimum were N20 dissociated non catalytically (catalytic decomposition occurred at ca. 525 K and above). From Table 1 one may estimate the total surface area of potential vacant sites which may abstract oxygen 2 19 from oxygenated compounds. Assuming that one m contains 1 x 10 atoms, N20 titration data indicated a surface of ca. 1.9 m2/g, or about 6 % of the total BET surface is composed of oxygen vacancies. 3.2. Acetaldehyde-TPD on UO 2.

Figure 3 and Table 2 present the desorption products during TPD after acetaldehyde adsorption on UO 2 (H2-reduced UO3). Table 2 Carbon yield and carbon selectivity of products formed during TPD after acetaldehyde adsorption at room temperature on UO 2 Product

Desorption Temperature (K) Acetaldehyde (m/e 29) 390 Propane (m/e 39) 610 Butadiene (m/e 54) 540 butene (m/e 56) 673 Ethanol (m/e 45) 415 CO 2 (m/e 44)

730

Carbon Yield (100%) 65.9 12.2 6.3 0.9 0.7 14.0

Carbon Selectivity (100 %) 35.8 18.5 2.6 2.0 41

Serval reactions occurred evidenced by a complex desorption products. First, acetaldehyde (m/e 29, 15, 44) desorbed at 390 K followed by traces of ethanol at 415 K (2 % of carbon selectivity, table 2). Three other products were observed. Butadiene and butene desorbed at 540 and 673 K respectively with a combined carbon selectivity of 21.1%. This reaction pathway follows a reductive coupling mechanism which has been observed previously on the surfaces of TiO 2 single crystal and powder [19-21]. The formation of C4 olefins is a clear example of the capacity of UO 2 surfaces to abstract large amounts of oxygen from surface carbonyls, via pinacolates [ 19], as follow

271

2 CH3CHO + 2 U

+4

- Vint.vac.

)

CH3CH=CHCH 3 + 2 U

+4+x - Oint.

Vint.vac.: interstitial oxygen vacancy, Oint.: interstitial oxygen. 3.3. E t h y l e n e - T P D on U O 3

Figure 4 and Table 3 show the desorption products during ethylene-TPD on ~ UO 3. Table 3 Carbon yield and carbon selectivity of products formed during TPD after ethylene adsorption at room temperature on UO 3 Product

Desorption Temperature (K) Ethylene (m/e 28, 27) 400-700 Acetaldehyde (m/e 29) 480 Furan (m/e 68, 39) 550 CO 2 (m/e 44)

above 800

H20 (m/e 18)

ca. 500

Carbon Yield (100%) 85.7 8.3 6.0

Carbon Selectivity (100%) 58 42

not calculated

In addition to ethylene desorption in a large temperature domain, acetaldehyde was clearly observed evidenced by its m/e 15, 29 and 44 (Table 3). The formation of acetaldehyde from ethylene indicates the facile removal of surface oxygen on UO 3 and shows its high reactivity towards oxidation of olefins. It is important to note that during TPD there is no regeneration of surface sites (in contrast to a steady state oxidation reaction with oxygen). This reaction requires a subsequent reduction of surface cations as follow

CH2=CH 2 + U

+6

-O

~

CH3CHO + U

+4

+ VO

(480K)

Vo: surface oxygen vacancy In addition, another important product was observed, furan (C4H40, m/e 68 and 39) at 550 K with comparable yield to acetaldehyde (42 % carbon selectivity). Thus, furan formation indicated that U surfaces are also active for C-C bond formation in their oxidised form, in addition of being an active C-O bond formation catalyst. The key route to this reaction is the formation of C4 olefin (most likely butadiene) which in its turn reacts with the surface oxygen to give furan as follow

CH2=CH 2 + CH2=CH 2 + O 1

CH2=CH-CH=CH 2

+ H20

cthylene/UO

acetaldehyde/U02

G

0 0-

-I--,

E4 k

0

v1 Q)

ct1i;inol x 40 0

pl-opane x 5

Q)

k k

0

0

5 300 400 500 600 700 800

300 400 500 600 700 800

273 CH2=CH-CH-CH 2 + 201

C4H40 (furan) + H20

(550 K)

Ol: lattice oxygen

Tow further points are worth mentioning. Firstly, XRD of the used ]3 UO3 (after TPD) indicated a mixed phase materials composed mainly of [3 UO3 and o~ U308. TPD of ethylene on this used UO3 (which have been regenerated under a dry air flow at 473 K for 90 minutes) showed a furan yield very similar to that on pure 13UO3 [22]. This result (which is under further investigation) may indicate that o~ U308 is also active towards this oxidative coupling reaction. It is important to mention that U308 contains substantial amounts of U +6 cations (together with U +4 or U +5 cations [7]). Secondly, in order to understand the reaction mechanism, TPD after butadiene adsorption at room temperature on ]3 UO3 was also investigated. Furan was clearly observed together with maleic anhydride [22]. This last point reinforces the above reaction mechanism.

4. CONCLUSIONS The oxidation of ethylene has been investigated on polycrystalline 13 UO 3 surfaces. Two oxygen containing products were observed: acetaldehyde, and furan. Furan desorption which requires a C-C bond formation, most likely is formed via dimerization of two adsorbed ethylene molecules followed by cyclization with available surface oxygen. Both the formation of acetaldehyde and furan from ethylene on UO 3 are clear examples of the ease of removing oxygen atoms from UO 3 surfaces. The reduction of acetaldehyde was also investigated on UO 2. Two molecules of acetaldehyde couple together to make a symmetric olefin: butene (which undergoes further dehydrogenation to butadiene). This is similar to what has been observed on TiO 2 and CeO 2 surfaces before [ 19-21, 23 ]. These complex chemical pathways indicate the richness of the U oxides system and open routes to further investigations.

References 1. M.A. Barteau, Chem. Rev., 96 (1996) 1413 and references therein. 2. V.E. Henrich and P.A. Cox, The Surface Science of Metal Oxides, 1994, Cambridge University Press, and references therein. 3. P. Sigmund, Sputtering by Ion Bombardment: Theoretical Concepts. Topics in Applied Physiscs, 47 (1981) 9. 4. C.A. Colmenars, Prog. Solid State Chem., 9 (1975) 139. 5. M. Pepper and B.E. Bursten, Chem. Rev., 91 (1991) 271. 6. R.J.D. Tilley, Defect Crystal Chemistry, Blakie, Glasgow and London, 1986.

274 7. H. Madhavaram, P. Buckanan and H. Idriss, J. Vac. Sci. Technol. A, 1997, in press. 8. H. Idriss and M.A. Barteau, Catal. Lett., 26 (1994) 123. 9. H. Poelman, L. Fiermans, J. Vennik and G. Dalmai, Surf. Sci., 275 (1992) 351. 10. K.M. Taylor, US Patent No. 3,670,006 (1972). 11. R.K. Grasselli and R.C. Miller, US patent No. 4010188 (1977) 12. R.K. Grasselli and D.D. Suresh, J. Catal. 25 (1972) 273. 13. D.D. Suresh, M.J. Seely, J.F. Brazdil and R.K. Grasselli, US Patent No. 4855275 (1989) 14. J.M. Hermann, J. Disdier, F.G. Freira and M.F. Portela, J. Chem. Soc. Farad. Trans., 91 (1995) 2343. 15. G.J. Hutchings, C.S. Heneghan, I.D. Hudson and S.H. Taylor, Nature, 384 (1996) 341. 16. F.A. Cotton and G. Wilkinson, Advanced Inorganic Chemistry, 1972, Wiley, New York, third edition. 17. H. Idriss, K.S. Kim and M.A. Barteau, J. Catal., 139 (1993) 119. 18. M. Boudart and G. Djega-Mariadassou, Kinetics of Heterogeneous Catalytic Reactions, 1984, Princeton University Press. 19. H. Idriss, K.G. Pierce and M.A. Barteau, J. Am. Chem. Soc. 116 (1994) 3063. 20. H. Idriss, K.G. Pierce and M.A. Barteau, J. Am. Chem. Soc. 113 (1991) 715. 21. J.E. Rekoske and M.A. Barteau, Ind. Eng. Chem. Res., 34 (1995) 2931. 22. H. Madhavaram and H. Idriss, work in progress. 23. H. Idriss, C. Diagne, J.P. Hindermann, A. Kiennemann and M.A. Barteau, J. Catal., 155 (1995)219.

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

275

Active Sites of Vanadium-Molybdenum-Containing Catalyst for Allyl Alcohol Oxidation: ESR S t u d y i n s i t u . O.V~Krviov. N~.uen Tien Tai. B.V.Rozentuiler N.N.Seme~ov Ln.stitute o! ( . h . m l c a ! Physics, Russian A c a d e m y of Sciences, ul.Kosygina 4, Moscow, 117884, Russia Allyl alc.r)hol oxidation into acrolein on the rhombic phase of molybu,:J~uJJ~ oxide in,-,u_ified with v a n a d i u m oxide has been studied by the Kne~nou ,Kl~leLic ,1 . , aim, , t,y: ESR of V . I + ions in sit'u. It was showl,, that active sites for ti~is reaction are V 'r ions situated in the bulk of the catalyst. ~w n e a r its s~rface, but n~:.,t at the surface. Fast diffusion r)f ,?iectro_ns and a more slower diffusion of oxygen ions b e t w e e n the s,~rface and t;he bulk occur d u r i n g the reaction. AI .

.

.

.

.

.

.

.

.

.

.

!1

u.

' 'roduction. A widespread opinion exists about the m e c h a n i s m s of oxidative catalytic reactions, that active sites for these reactions are surface coordinatively u n s a t u r a t e d transition metal ions. which can be r e d u c e d and reoxidized dllring the reaction. O u r studies of sevel-a] oxidation reactions [1,2] by .tt',S.R .in situ have shown t h a t the active sites can be situated not on!v at the surface, but in the bulk of the eata!y~t. We havc~ studied CO oxidation over the paraelectric phase of BaTiO3. , which contained about 0.01% Mn z+ ions as a n a t u r a l impurity. It t u r n e d .... J. ~u,o,L, 1 . _ ~. ~~ A~2+ ,_,u~., at 400-~t~0 K the active sites for CO oxidation are ..,,., ions. .Al(,r~g with adsorption and catalytic m e a s u r e m e n t s , we have proved using in situ ESR s ! ~ c t r a studies of Mn 2+ ions t h a t t h e y are situated in the bulk of BaTif) 3. It was found out that CO adsorption proceeds in exact accordance with the stoicbiometric equation, w h e r e C)~2- is a c,]rface c,x,,~en it,-' '~, Mr, 4+ + CO + O~'~- ---~ Mn ~+ 4- CO2,

(1)

i.e. adsorption of one CQ molecule results in a p p e a r a n c e of one Mn 2+ ion. On the contrary, 02 adsorption decreased the i n t e n s i t y of the ESR spo.otrum in accordance with the s t n i e h i n m e t r y

276 a~,~ ":,-,rf-'-! ~

~: u+_"__> 2Mn~ +-+ zuo.+"-' -, _ ,,

:.~,

tzl

i o. t:tvo Mn zl- inn~ d+sapp~.ar on a d s o r p t i o n of o n e nxyg~.n mnloe~iP. (In e x p o s u r e to t h e 2CO4-O z s t m e b i o m e t r i e m i x t u r e ~.he c a t a l y t i c r e a c t i o n is pr.,..-.eeeding t h r e u g h t h e r e d o x m e c h a n i s m f o l l o w s E q u a t i o n s (1) a n d {2). The 99 ..... - ~+.- o f m , , , . ~ , , . ~ o u r c o x i d a t i o n ( E q u a t i o n 2) is a l m o s t 2 ordcr~" of iiia,~iii|,tid~, N* u , , '~ J " t h a n t,ia~ ' ' oI~ t h e r e d u c t i o n ( E q u a t i o n l ) . In tht ~, 9 "........ " ' - "

.

.

established. cal,

.

. . . . . . .

.

The

.

.

lOW

reaction

proceeds

lu

predominantly

"

t2.oliCuIJtI

over

the

" "~t 'L.IUII

i8

oxidized

a J v s++. t.

~|'ho.~o invo.~tigntions wPro. e o n t i n ~ e d in m+r stud.les of a iiyi aieohoi ]r,+,:~ .'.~er,,l,~in m , e r t b e h e x a g o n a l pb_a~e of MoO a m o d i f i e d b y ! w ! % of Vz() 5 [2 ~]. T h e reacti~;n "" "~ '~

~,,,II~O11

:' .--~ \

+' '-'

.

Lot

ii2,k)

p r o c e e d s at 400-470 K w i t h i O0% s e l e c t i v i t y . P r a c t i c a l l y all t h e V r h,t~s rneam+t'o+d w i t h ~ S R i n si~.u w ~ r e in t h e b u J k of the, e a t a j v s t . VorT.r!atjoD of o~.e V 4.b i o n d u r i n g r e d u e t i o n of t h e catalyst. ]_s a c c o m p a n i e d b y t h e disapl.x~aranee of o n e allyl alcohol m o l e c u l e ,

. ,-, . + .'--:.i.,*~ . -,** ....§

+ [ ]~ + C : + 1 1 4 0

(+~,~

+ 119.O

it w a s s h o w n [3] t h a t V 4+ ions o b s e r v e d in E S R s o e e t r a a r e u n i f o r m l y c i i s t r i b u t e d o v e r t h e wiaole b u l k of h e x a g o n a i MoO~. (Ipym,~ito. offoet~ havo. bo~n o b s e r v e d on o x i d a t i o n of t h e c a t a l y s t b y ~,y.ygen. T h i s p r o , _ ' e s : ~ e.9....n+ be d e s c r i b e d b y f h e s t o i e h i o m e t r i e e q t m t i n n : q'~r4-t._t_

I 1

.l.. , q

., ~'tI1)',~-t--

/3-~

tV,~

AL 435 K Lhe t'aLe or Lhe s t a t i o n a r y c a t a l y t i c r e a c t i o n is proport"I O I l ~ l ' i:o t h e c,.Jneentration of v a n a d i u m ions a n d to t h e o x y g e n p r e s s u r e in th~ mixture

. . . .

k.;z.i '

J

,~ .... /

*~, i.e., at ~,~e- t e m p e r a ~++...4 at t h e t e m p e r a t u r e s ,,,+s,,~' 1.{~+1...... +l , ~ n.- a~-nn ~ - 6 0 0 ~r l;ui~'~ ot" at+ai,ioij~ti- s ~ t M y t i c oxidMiui,, t h ~ h e x a g o n M p h a s e of ~+ ***vO,,~ i~ [l'Hill

o I slt~n:~e.u

inl, o

the

more

s~ame, +

~

1

rhotnbie

phase 1

[6]. Thi~-

puper

is

277

devoted to the investigation of the mechanism of allyl alcohol oxidation into acrolein on the rhombic phase of MoO~ modified with V20~.

Experimental Samples of V-Mo-oxide catalysts have been prepared by mixing (NI-I4)6MoTOz4 and NH4VO.3, drying for 4 hours at 470 K, and calcination at 550-80~2 K Catalysts with 0.2, 0.5, 12, ,~.0, and 3.0 wt% of VzO5 were prepared. Their phase coml~)sition has been d e t e r m i n e d on the in~ t r u m e n t DRON-2 with Fe-K..~-emission. ESR s t ~ c t r a have been registered with the spectrometer EPR-V constructed in the Institute of Chemical Physics and equipped with an a t t a c h m e n t for the t e m p e r a t u r e regulation of the ampol!le with the s'ample directly :in the resonator of the spectrometer. Accuracy of the . is 0.2 ~ Calibration of the spectrometer has been t e m p e r a t u r e ~e"ulation, .~. Clone with the help of solutions of stable nitroxyl radicMs in benzene. g-values and IIFS constants were d e t e r m i n e d by comparison with a Mn24-/MgO standmxl. The absolute error in the d e t e r m i n a t i o n of the spin concentration by means of double integration is '_*50,%, the relative one is +2%. The error in g-value is 0.001. and that in ItFS constants is 0.1 Gs. The V4+ ions concentration was d e t e r m i n e d via double integration using the second parallel component of the ESR signal. A flow microcatalytic set-up has been combined with the ESR-spectrometer. Catalyst samples were placed into a flow- reactor which was at the same time an ampoule for the ESR studies. Gas mixtures He+O~, He+CsHsOII. and tIe§ have been prepared with the help of special 4- and 6-way valves. Gas analysis have been performed on the Carbowax column of the gas chromatograph, the length of the column was 1 m, and its t e m p e r a t u r e was 300 K. The catalytic reaction was studied at 380-540 K directly in the heated resonator of the ESR-spectrometer, at higher t e m p e r a t u r e s it was studied outside of the resonator. '

}

.

.

9

.

.

Results and discussion. Paramagnetic centres in V-Mo-oxide catalyst. Each of the observed ESR signals consists of 24 ItFS component~ due to interaction of an unpaired electron with the v a n a d i u m nucleus (s!V I=7/2. p=99.7o~). Their dependence on t h e , t e m p e r a t u r e of calcination of 2%V2Os/MoO 3 is shown in Fig.l The si.~nal tt (gx=gy=l.95r gz=l.908) corresponds to V4+ in the hexagonal phase of MoO~, the signals A, B, and C

278

~} 3"V.,(reI.uni tz )

Figure 1. De~mndence of the intencity of ESR signa]s observed in V-MoO~ catalyst on the heating temperature.

,K 593 623 673 723 873 q23 correspond to V4+ in the rhombic phase. It is seen that the rhombic phase only exists at temperatures higher than 600 K. This result was also confirmed by the XRD study. The signals A (gx=1.976, gy=1.974, gz=l.921) and B (gx=1.974, gy=l.970, g~.=1.928) are characterized by additional tIFS, which is typical for the interaction between unpaired electron of vanadyl and the nitrogen nucleus l'14N. I=l, p=99.63%). A calculation of the orbitals with an unt~aired electron [7] shows that the A signal corresl~mds to the V imide complex, where the NH group occupies the tx.,sition of one of the oxygen ligands around vanadyl. The B-signal a p p e a ~ at a higher t e m w r a t u r e (700 K). A calculation shows, that this signal corresponds to the interaction of the unpaired electron of vanadyl with a NO molecule. Both, the A and B complexes, are formed during preparation of the samples from arr lmonium molybdates and vanadates. The signal C (gx=l.971, gy=1.969, gz=1.872) s essentially different. It is fi~rmed at high temtmratures (800-900 K). Every component of its additional ftFS consists of one intensive line in the centre and 6 lines of equal intensity, separated by equal distances. The intensities of these lin,::s are in the ratio of 100:5.1. The appearance of this signal can be explained by the formation of non-stoichiometric phases in MoOa, the so called Magnelli phases [81, where the MoOs octahedrons are connected by planes and edges, but not by apexes. Such a non-stoichiometry contracts distances between the metal cations. Interaction of an tlnpaived electron of V4~ of the first octahedron with the molybdenum nucleus ('~,97Mo, I=5/2, p=25.18 %) of the adjacent octahedron connected by an edge with the fi.,~'t one gives additional HFS. The signals A, B, and C were observed at all concentrations of VzOa. Fig.2 shows the dependence of the n u m b e r of V 4+ ions and of the ESR line width on the VaO~ content. The maximal concentration of V 4+ ions

279

al:ter the catalyst reduction by allyl alcohol has been obsex"qed for 2% o[ V205. The line width increases monotonically with increasing percent of V~.O5 The V 4t ions are distributed uniformly and separately in the catalyst bulk with the increase of VzO 5 up to 2%. At higher V205 concentrations the signal intensity increases and the lines are broadened, because of strong mutual interaction of the V 4+ ions. At higher VzO 5 concentrations clusters of V 4+ ions are formed, and the n u m b e r of V 4+ ions observed in ESR diminishes. ,......

'~ ,3

500~ for the high vanadium content catalysts. It was reported [16] that at low loading vanadia interacts preferentially with the most basic hych'oxyl groups present on the titania surface. At high vanadium loading, most of Ti-OH groups are replaced by new Bronsted acid sites which give rise to a NH~ band whose intensity increases with the vanadium content. This result is in agreement with the substitution of strong Lewis acid sites of TiO2 with weaker Bronsted acid sites due to V-OH groups observed in the vanadium containing samples. The TPD profiles of ternary catalysts are also different from those of binary catalysts, showing that the presence of niobium results in a different distribution of acid sites, especially at low vanadium loading. However, both xV6Nb/Ti (I) and (II) give rise to a lower temperature signal. It is noteworthing that the presence of niobium oxide enhances the acidity of the bynary catalysts, more strongly at high vanadium loading. The amount of desorbed NH3 from the ternary catalysts is slightly affected by the preparation method, especially at high vanadium loading (Table 1).

3.3. Laser-Raman spectroscopy. In Figure 2 Laser-Raman spectra of 6Nb, 1V/Ti, 6V/Ti, 1V6Nb/Ti and 6V6Nb/Ti catalysts are reported. The spectrum of 6Nb/Ti shows a narrow peak at ca. 990 cm 1 which can be attributed to the double bond Nb=O of both tetrahedral and octahedral NbOx species [21]. A broad band, with lower intensity and centred at ca. 920 cm -1, probably due to Nb-O-Nb bridges in octahedrally coordinated species, is also present [21]. 6V/Ti catalyst shows a narrow peak at 1030 cm 1 due to the double bond V=O and a broad band at 915-920 cm 1 attributed to V-O-V bridges in polycondensed species [17]. The band at 915-920 cm -1 is absent in the spectrum of 1V/Ti sample, where tetrahedral isolated species prevail. The spectrum of 6V6Nb/Ti catalyst shows no remarkable difference from that of 6V/Ti; moreover there is no evidence of a contribution at 990 cm 1 related to Nb=O bonds. However, in the spectrum of 1V6Nb/Ti catalyst a broad signal at ca. 990 cm -1 indicates the presence of Nb=O bond. Its intensity is negligible if compared to the corresponding signal in the spectrum of 6Nb/Ti. The disappearence of the Nb=O signal in the ternary catalysts could suggest an interaction between the two oxide phases. The formation of V-O-Nb-O-V bridges can be suggested, or the grafting of vanadium onto niobium oxide phase can be hypothesized.

3.4 Catalytic activity tests The possible occurrence of homogeneous reactions was tested by performing experiments in the absence of catalyst under the same reaction conditions of the catalytic tests. No ethane conversion was observed up to 700~ In the activity tests the oxygen conversion was kept always A1203 > ND205 > SiO2. The low activity of 4VCe and 5VNb catalysts, despite their reducibility (4VCe) and acidity (5VNb) may be due to structural transformations by reaction of vanadia with the underlying oxide at the high temperatures required for ethane oxidation. Concerning selectivity, the more reducible oxide support systems show a high selectivity to deep oxidation (CO). 4VCe shows high selectivity to CO at low ethane conversion. The acidic supports, alumina and niobia, also yield CO as the main oxidation products. Only silica-supported vanadium oxide shows higher selectivites for ethylene. Acetaldehyde and formaldehyde are also produced son 12VSi and 12VSi-H20. The relevance of V-O-V bonds can be evaluated for the performance of V205/A1203 and V205friO2 at different surface coverages. Alumina supported vanadium oxide shows increasing TOF numbers for oxygen, CO and CO2 as

303 vanadium oxide loading increases up to monolayer coverage. At monolayer coverage, where the (V-O-V) / (V-O-Support) ratio is expected to be highest, the TOF's of ethane and ethylene decrease, but TOF of oxygen, CO and CO2 increase. This could be indicative of the higher reducibility of surface polymeric vanadium oxide species with respect to isolated surface vanadium oxide species (4,8), which appears to lead to a less active and selective catalyst. A similar trend is observed for VTi series: at monolayer coverage, ethane and ethylene TOF numbers decrease. For the titania-supported vanadium oxide catalysts, the TOF's for oxygen, CO and CO2 do not increase at vanadia monolayer coverage as in the case of the VA1 series. On the contrary, they decrease slightly, but isolated surface vanadium oxide species on titania are more reducible than isolated surface vanadium oxide species on alumina. This may account for the higher TOF(oxygen)/TOF(ethane) ratio observed on the VTi series. This ratio becomes closer for VTi and VA1 series at monolayer coverage, where both series are expected to show s higher reducibility of the surface vanadium oxide species. The ternary V205/TiO2-SiO2 catalyst shows interesting structural and catalytic properties. Surface vanadium oxide species preferentially coordinate to titania sites in the TiO2/SiO2 supports (8). However, the use of a titania-silica support prepared so that titanium oxide is highly dispersed and strongly interacting with silica support results in titania with different characteristics to pure titania. The titania-silica support used here has 20% of the titanium atoms in tetrahedral coordination as determined by XPS and no crystalline aggregates of titania are formed, as determined by Raman spectroscopy (10). The V=O mode observed for the dehydrated 10V5TiSi catalyst is at 1036 cm -1, much closer to that of silica-supported vanadium oxide than to that of titanium-supported vanadium oxide (Table 1). The surface vanadium oxide species are isolated (100 % dispersion) and must also have a different coordination environment (probably anchored on both, titania and silica sites) that yields an activity similar to that on 12VSi but more selective, since no CO2 is formed and the selectivity of ethylene increases. The lower selectivity of oxygen -containing products suggest that vanadia species on the highly dispersed titania-on-silica supports may be less reducible than on the pure constituting oxide supports. 5. C O N C L U S I O N S The surface vanadium oxide species on silica, water-treated silica, alumina, ceria, titania, zirconia, niobia and titania-silica have been characterized and studied for the selective oxidation of ethane. The terminal V=O bond does not appear to be directly involved in the reaction (no correlation with TOF). However, the bridging V-O-V or V-O-Support bonds appear to critical for the oxidation of ethane. The nature of the V-OSupport bond is determined by the specific support. Bonding to a reducible support metal ion yields active catalysts (e.g. 6VTi and 4VZr). Acidic supports show some activity, but much lower than the reducible ones. The silica support is not reducible and does not possess acidic sites and shows the lowest TOF numbers. However, silica-supported vanadium oxide catalysts possess the highest selectivity. The very low activity of 4VCe and 5VNb could originate from a reaction of vanadia with the underlying support. The surface coverage increases

304 the (V-O-V) / (V-O-Support) ratio. Polymeric surface vanadium oxide species are more reducible than isolated surface vanadium oxide species in the presence of butane (15). If we assume a similar trend of reducibility with ethane than with butane, if turns out that more reducible surface vanadium oxide species are less active and selective. This effect is more evident for the VA1 series than for the VTi series, since the isolated surface vanadium oxide species on alumina are much less reducible than on titania. All the catalysts that show higher reducibility, either due to its interaction with the support or due to its surface polymerisation show lower selectivity. The surface vanadium oxide species have a different environment for 10V5TiSi catalyst, which yields an activity similar to that of 12VSi but is more selective. Further research is going on to fully understand the environments of vanadia sites in this catalyst.

ACKNOWLEDGEMENTS This research has been partially funded by the Fundaci6n Caja de Madrid (Spain).

REFERENCES "

.

3. "

5. 6. .

"

.

10. 11. 12. 13. 14. 15.

E. A. Mamedov, and C. Cortes Corberfin, Appl Catal A : General, 127, 1 (1995) G. Bond, and S. Flamerz Tahir, Appl. Catal., 1 (1991) G. Deo, I. E. Wachs, and J. Haber, Critical Reviews in Surface Chemistry 4 (3/4), 141 (1994) I. E. Wachs, and B. M. Wechkhuysen, Appl. Catal. in press (1997) S. T. Oyama, and G. A. Somorjai J. Phys. Chem., 94, 5022 (1990) J. Le Bars, J. C. Vedrine, and A. Auroux, S. Trautmann, and M. Baerns, Appl. Catal tk" General 88, 179 (1992) M. Merzouki, B. Taouk, L. Tessier, E. Bordes, and P. Courtine, in "New Frontiers in Catalysis" (Guczi et al., Eds.), p. 753. Elsevier, Amsterdam, 1993 I. E. Wachs, J.-M. Jehng, G. Deo, B. M. Weckhuysen, V. V. Guliants and J. B. Benziger, Catal. Today, 32, 47 (1996) J. -Mirn Jehng, and I. E. Wachs, Catal. Letter, 13, 9 (1992) X. Gao, M. A. Bafiares, J. L. G. Fierro. and I. E. Wachs, unpublished results G. Busca, Mater. Chem. Phys., 19, 157 (1988) H. Eckerdt, and I. E. Wachs, J. Phys. Chem., 93, 6796 (1989) J. Hanuza, B. Jezowska-Trzebiatowska and W. Oganowski, J. Mol. Catal., 29, 109 (1985) G. Deo, and I. E. Wachs, J. Catal., 146,323 (1994) J. Haber, A. Kozlowska, and R. Kozlowski, J. Catal, 102, 52 (1986)

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

305

The ethane o x i d a t i v e c h l o r i n a t i o n process and efficient catalyst for it M.R. Flid, I.I. Kurlyandskaya, Yu.A. Treger and T.D. Guzhnovskaya Scientific Research Institute "Syntez", 2 ,Ugreshskaya str., P.O. Box 56, Moscow, 109432 Russia Formation of the mixed cement-containing systems within the range of low copper concentrations with addition of alkali metal dopants as well as catalytical properties of these systems in the ethane oxidative chlorination process have been investigated. Based on the obtained data the efficient and stable copper-cement catalyst has been worked out. This catalyst will assist in the development of a new technology of the vinyl chloride production from ethane. The basic parameters of the ethane oxychlorination process have been determined : at 623-673K, time-on-stream 3-5s and reactant ratio of C2H6: HCI: :02 = 1:2:1 the conversion of ethane is more than 90% and the total selectivity to ethylene and vinyl chloride is 85-90%.

1.1ntroduction The gas-phase catalytic process for oxidative chlorination of ethane to vinyl chloride according to overall equation C2H6 + HC1 + 02 = C2H3C1 + 2H20,

(I)

proceeds in two consecutive kinetically independent reactions: (1) the oxidation of hydrogen chloride to chlorine and (2) the chlorination of ethane. This process is promising for developing a rational technology of vinyl chloride production, because ethane utilized in it is a cheap hydrocarbon raw material [ 1,2 ]. The process is conducted at high temperatures, and ethane converts to vinyl chloride due to a combination of consecutive and parallel radical-chain and heterogeneously catalyzed reactions: oxidation, chlorination, and dehydrochlorination. The contributions of homogeneous and heterogeneous reactions to the overall rate of chlorinated hydrocarbon conversion depends on the temperature ranges at which the reaction proceeds. The process as the whole may be represented by the following schematic diagram [3]: CO + C02

C2H6 - -

t

t

C2H5CI~2H4CI2 C2H4 ~

C2H3C1

t

~2H3C13

t

~2H2C14

--~ C2H2C12 --~ C2HC13

CO + CO2

t

~2HCl5 --~ C2C14

--

C2C16 (II)

306 from which it follows that the major products of ethane oxidative chlorination are ethyl chloride, 1,2- and 1,1-dichloroethanes, 1,1,2-trichloroethane, chloroethylenes, and carbon oxides as the products of deep oxidation. At relatively low temperatures (623m723K), the reaction mixture consists mainly of chloroorganic saturated compounds [3-6]. The situation changes dramatically with raising the temperature. Figure 1 demonstrates the effect of the temperature on the oxidative chlorination of ethane over the well-known conventional salt CuC12--KC1/silica gel copper-containing catalyst. 80

1

70

.

'

~

-

"

60,

-~

3

~,9 50

L 0

,'.o

40

2

30 "0

o

5

20 10

4

"'

0 723

773

823

Temperature,K

Figure 1. The effect of temperature on the ethane oxidative chlorination process (silica gel as the support, copper content of 6.0 wt %, potassium content of 4.0 wt %, reactant ratio C2H6 : HC1 : O2= 1 : 1 : 1, x = 3 s). 1 is the conversion of ethane; 2 is the yield of oxidation chlorination products; 3.4, and 5 are the yields of ethylene, deep oxidation products, and vinyl chloride, respectively ( x is time- on-stream ). Thus, in the presence of traditional catalytic systems, the yield of vinyl chloride to converted ethane does not exceed 35%. The total yield of vinyl chloride and ethylene ranges up to 80%. It was shown [3,5,6] for saturated compounds ethane, ethyl chloride, 1,2dichloroethane, and 1,1,1-trichloroethane that the observed conversion rates are satisfactory described by the equation r, = k,. Pi" Pci2 0"5

(1)

The observed rate constant in equation (1) in this case decreases in the order C2H6 > > C2H5C1 > C2H4C12. The activation energies for the transformations of saturated (130 kJ/mol) and unsaturated compounds (40--90 kJ/mol) differ dramatically; as a consequence, the yield of chloroalkenes increases with temperature. Oxidative chlorination of ethane gives rise to considerable amounts of carbon oxides. The overall rate of these side reactions is described by the empirical equation rco + c02 = ki. pi" Po_~"Pcl2~

(2)

Unsaturated compounds make the dominant contribution to formation of carbon oxides. Whereas the introduction of one chlorine atom into ethylene molecule results in a 7m 8-fold increase in the observed rate constant of deep oxidation, the further increase of chlorine content in molecule diminishes the oxidation of chloroalkenes.

307 It is essential that the reactions of saturated compounds exhibit zero orders with respect to both oxygen and hydrogen chloride and proceed kinetically independently of one another. For the unsaturated compounds, the conversion rates represent complex functions of the reaction mixture composition. Under the conditions when the reaction exhibits zero order with respect to hydrogen chloride, the kinetics of unsaturated compounds oxidative chlorination is described by the equation: 2ko2 po2" ki'pi ri

(3)

--

2ko2 "po2+ L-k,.p, where index i relates to unsaturated compounds. The process of ethane oxidative chlorination imposes heavy demands on the catalysts. The conventional salt supported catalysts are composed of Cu, K, Ca, Mn, Co, Fe, Mg, and other metal chlorides containing various additives; these salts are precipitated on alumina, zeolites, silica gel, and other supports. Catalytic systems that represent solid solutions of iron cations in the lattice of the o~-A1203 and a-Cr203 phases doped with cations, such as K, Ba, Ce, and Ag are also known [7]. The activity of the known catalytic systems and, especially, their selectivity to vinyl chloride are insufficient. In addition, the known catalytic systems tend to rapid deactivation because of gumming and carbonization of their surfaces. The main problem that determines the possibility for industrial utilization of the process is the creation of highly efficient, stable, and selective catalytic systems performing at relatively low temperatures. This problem was alleviated due to the development of a new generation of heterogeneous catalysts based on high-alumina cements and intended for the synthesis of chloroorganic compounds, l These catalysts fortunately combine the properties required in industry and genetically intrinsic to cements thermal stability, high mechanical strength, and basicity of the surface, which prevents its carbonization with the possibility of imparting the system special properties desired in a particular process [8]. The mechanism of the ethane oxidative chlorination process is distinguished by the fact that the catalyst accelerates primarily the reactions of hydrogen chloride oxidation and dichloroethane dehydrochlorination. This necessitates the modeling of cement catalytic system with the surface carrying active sites capable of catalyzing both reactions mentioned. The analysis of the known and our own experimental data indicated that the properties required may be offered by a copper-containing cement-based catalytic system modified with alkali metals. In this catalyst, copper-containing active sites catalyze the oxidation of hydrogen chloride, whereas the activity of the catalyst in the dehydrochlorination reaction is determined by the acid--base surface properties, which are inherent to cements with different phase compositions. The development of this catalytic system made it necessary to investigate the formation process of mixed cement systems within the range of low copper concentrations and with addition of alkali dopants and determination of the correlation between properties of the obtained catalytical systems and their activity in the ethane oxychlorination process.

I 'Fhe catalysts based on high-alumina cements were developed in collaboration with Prof. V.I. Yakerson (Zelinsky Institute of Organic Chemistry, Russian Academy of Sciences, Moscow, Russia) and Prof. E.Z. Golozman (Institute of Nitrogen Industry, Novomoskovsk, Russia).

308

2.Experimental. The catalysts were prepared by chemical mixing of high-alumina cements ( technical calcium aluminate- talyum) or cement-based supports (calcium aluminates with the developed surface area and various CaO/A1203 ratio - galyumin or galyumin C) [9] with the sources of copper and alkali metals in water--ammonia or ammonia ---carbonate media; the mixing was followed by the drying and thermal treatment of the samples obtained. A comprehensive study on the formation of cement catalytic systems was performed by X-ray diffraction, thermal analysis, electronic diffuse-reflectance spectroscopy and IRspectroscopy. Table 1 presents characteristics of some of the investigated catalytic systems. Table 1 Characteristics of copper--cement catalysts No

Sample

Support

Preparation

Phase composition

Ssoec,

conditions

without thermal with thermal m2/g treatment treatment at 673K 1 CuO-K20- Galyumin 348K water-gibbsite, C3AH6, KC1, CaCO3, CuO, 130 CaO-A1203 ammonium CaCO3, CHA, CuO, ,/-A1203,C12A7 solution KC1

2 CuO-Cs20- Galyumin 348K CaO-Al203

C

water-- CsC1, CaCO3, CuO, CsC1, CuO, ammonium C3AH6 CaCO3(calcite), solution C12A7, CaCO3 (aragonite)

15

Kinetic measurements were made at 623 - 773K using a circulatory flow installation. Reactions were studied in the fixed bed catalyst. Time - on- stream was varied within the range 1,5 - 10s at a reactant ratio of C2H6: HCI: O 2 = 1:1 +3,3, : 1 +1,4. Air was used as a source of oxygen. The grains of the catalyst were 0,25 - 0,5 mm in size. The gas was fed at a volumetric flow rate of 600 h - ~ . The catalytic systems were preactivated with a hydrogen chloride nitrogen mixture at 573-623 K. The analyses were based on the chemical methods (determination of hydrogen chloride and chlorine) and the gas chromatography.

3.Discussion and results. The stage of chemical mixing of the catalysts preparation involves the hydration of cements with forming C3AH6 (C is CaO, A is A1203, and H is H20), gibbsite, and calcium carboaluminate as well as the exchange processes with forming CaCO3 and copper hydroxoaluminate (CHA). The depth and the rate of hydration as well as the distinctions in the exchange processes are determined by the type of cement-containing agent. The stage of thermal treatment involves the formation of C12A7, ,/-alumina, and solid solution of aluminum and copper oxides, which is followed by the precipitation of the excess of highly dispersed copper oxide and by the formation of copper aluminate spinels with various degrees of disorder.

309 Thus, copper-containing phases can occur both as free oxide and as the forms bound with the matrix of the support; the concentrations of bound forms increase with temperature and with the duration of chemical mixing. The estimation of the depth of interaction revealed that not only the implantation of the Cu 2. ions into the matrix lattice with forming isolated ions is possible, but the formation of small surface clusters (CuO)• with highly covalent Cu--O bonds. The distribution of catalytically active component between the free oxide, clusters, and ions implanted into the matrix lattice depends both on the conditions of formation and on the composition of the catalytic system as well as on the type of cement-containing agent. As it was shown in [ 10], cement-containing matrix exerts a strong modifying effect on the active copper-containing sites. At equal concentrations of the active component, the activity of copper-containing sites incorporated into the copper--cement catalyst is higher than that of the supported salt catalysts. When the concentration of copper and surface concentration of copper-containing sites are decreased, specific catalytic activity of coppercontaining centers sharply increases. So, at an extended specific surface of the copper-cement catalyst, high catalytic activity to the oxidation of hydrogen chloride can be accomplished even at a low concentration of active copper-containing component provided that the latter is bound with the matrix of the support. The surface area of cement catalysts, which carries aluminum- and calcium-containing oxide fragments, exhibits pronounced acid--base properties. These properties can manifest itself as a catalytic activity to the reactions of dehydrochlorination, which proceed via the formation of donor--acceptor complexes between the substrate and acid or base sites at the catalyst surface. The existence of different calcium aluminate phases in the aluminum-calcium catalysts was proved by diffuse-reflectance IR spectroscopy. The presence of these phases is responsible for the complex structure of the catalyst surface. At the surfaces of these catalysts, calcium ions with lower coordination numbers can occur together with the ions octahedrally surrounded by oxygen anions. These ions can act as balance cations in the structure of C12A7, being responsible for the existence of specific terminal hydroxyls and Lewis acid sites bound to calcium. At the surfaces of galyumins, bridging hydroxyls exhibiting somewhat stronger acid properties are present along with terminal hydroxyl groups. The hydration of galyumin surface can supposedly be attended with the weakening of the A1--O--M bond (M = A1 and Ca) resulting in the appearance of additional strong adsorption sites[8]. The enrichment of surface layer in galyumin C with Ca )-+ ions at the increase of CaO/ AL203 ratio is essential for reducing the yield of deep oxidation products and preventing the carbonization of the surface. The data on the state of copper-containing phases and acid--base properties of active sites occurring at the surface of mixed cement systems, which were presented above, enable us to conclude that these catalysts can be employed in the oxidative chlorination of ethane. It ~is known that the chlorination of ethane with chlorine formed in the oxidation of hydrogen chloride proceeds by a heterogeneous--homogeneous mechanism [3]. This is why the efficiency of cement catalysts was studied separately by the examples of Deacon reaction and dichloroethane dehydrochlorination reaction. It was found that for galyumin-based cement system, the variation of copper content within 8--25% (in terms of CuO) virtually does not affect the rate of chlorine formation. For the oxidation of HC1, the rate constant is 1.2.10 -3 mol HC1/g cat.h. This value is comparable with the rate constant of HC1 oxidation in the presence of copper-containing salt catalysts. The

310 introduction of potassium chloride into a copper--cement system results in a 1.5-fold rise of the rate constant for the HC1 oxidation. Thus, the activity of copper--cement catalysts in Deacon reaction is comparable with that of commonly used salt catalysts. Systematic investigations on the performance of cement-containing catalytic systems with various chemical and phase compositions in the reaction of 1,2-dichloroethane dehydrochlorination with forming vinyl chloride C2H4CI 2 -4

C2H3C1 +

HC1

(III)

revealed that the catalytic activity of these catalysts in the process under consideration is high. At the constant composition of the reaction mixture, the maximum reaction rate was accomplished with using a cement system whose specific surface is 130 m2/g. Thus, at 623K and time-on-stream of 3.8 s, the reaction rate was 0.42--0.46 mol of vinyl chloride per litre.hour. This value is more than two times higher than the reaction rate accomplished with using a well-known supported salt catalyst CsC1--SiO2. It was also shown that the presence of copper in cement catalytic systems does not affect the activity of the catalyst in the dehydrochlorination reaction (see Fig. 2).

70 60

"6 ,..

0

50 40

P,

~

~ L

30

C 0

0

20

0

[

523

.

.

.

.

.

r

. . . . . . . . . . . . . . . . . . . . . . .

573

' ....

623

" ........

673

Temperature,K

Fig. 2. The conversion of 1,2-dichloroethane in the dehydrochlorination reaction at various catalysts as a function of temperature, z = 8 s; 1- galyumin (Sspec = 130 m2/g), 2- galyumin with a dopant of copper (8 wt % in terms of CuO); 3- CsC1/silica gel. Thus, cement-containing systems provide the conversion of dichloroethane to be increased to more than 70% even at 673K. An important positive factor is that vinyl chloride molecule is stable at this temperature. At 673K, the side reaction of vinyl chloride dehydrochlorination with forming acetylene proceeds slowly, acetylene does not form, and the reaction is not complicated by the formation of a number of by-products, for example, of perchloroethylene. Thus, the above-made supposition about bifunctional character of copper--cement catalytic systems was confirmed in the investigations of their activity in the above-mentioned reactions.

31l The oxidative chlorination of ethane as a whole was studied by using of the cementcontaining catalysts with a specific surface of 130 m2/g (sample 1) and 15 m2/g (sample 2). Copper concentration was kept constant and equal to 8 wt % in terms of CuO (see Table 1). It was found during the investigations that when the temperature was raised from 623 to 773 K, the conversion of ethane somewhat increased, and sample 1 exhibited better activity in comparison with that of sample 2. At the moderate temperatures (623--673K), an extended specific surface of sample 1 was favorable for increasing the yield of target unsaturated compounds: ethylene and vinyl chloride. The further temperature increase led to a decrease in the process selectivity because of a noticeable increase in the yield of deep oxidation products, CO• The effect is more pronounced for sample 1 (see Fig. 3). 100 F

...........................................................................................................................................................................................................................

...,,-

90

-,--I'"

.

.

.

.

I-'-

""

""

-6

.~

__,a

-----4

80 7O

.__

..... -1 b

......._-

I

6o 5O

t-

40

o

(..)

30 20 ~C

1~ L 0 ,"F:

623

-.

-

--T

,

673

'"'-

. w

.

723

.

.

.

.

_

.

.~C

._~

773

Temperature,K

Fig. 3. The conversion of ethane and the yields of reaction products for catalysts 1 and 2 as functions of temperature. Time-on-stream of 3 s; 1; the reactant ratio of C2H6 : HC1 : 02 - 1 : 2 : 1;-- - catalyst i ; catalyst 2; a is the conversion of ethane, b is the total yield of ethylene and vinyl chloride to converted ethane, c is the yield of deep oxidation products COx. The dependences shown in Fig. 3 reveal that employing a catalyst with a larger specific surface area with rising temperature would, probably, lead to the deep oxidation of vinyl chloride and, to a lesser extent, of ethylene, resulting in a decrease in the total yield of ethylene and vinyl chloride. A certain increase in the overall yield of CO• products, which was observed for catalyst 2, is accompanied with an increase in the total yield of ethylene and vinyl chloride. This suggests that saturated chlorinated h y d r o c a r b o n s - ethyl chloride and 1,2-dichloroethane m are oxidized predominantly and that the rate of oxidation is lower rate compared to that of the dehydrochlorination of these compounds. Thus, the decrease in specific surface of the catalyst involves a noticeable drop of the yield of deep oxidation products, whereas the yields of vinyl chloride and ethylene remain high. We see little reason in the further cut of the specific surface, because the rate of catalytic dehydrochlorination therewith decreases.

312 The results obtained circumstantially testify that the dehydrochlorination and oxidation reactions proceed at different active sites. It is likely that the oxidation of chlorinated hydrocarbons proceeds at the copper-containing sites. This agrees with the data we obtained in the oxidative chlorination of ethylene [ 11 ]. Taking into account the fact that the value of specific surface is a crucial factor in the choice of catalyst, the further investigations we conducted with using catalyst 2. Both the time-on-stream and the reactant ratio are important chemical engineering parameters affecting the characteristics of the process. It was found that the increase in the time-on-stream at T = 673K can improve both the conversion of ethane and the yield of ethylene. The total yield of chloroorganic products therewith decreases, but the concentration of vinyl chloride passes through a maximum. We also observed an increase in the yield of deep oxidation products COx (see Table 2). Table 2 The effect of time-on-stream on the oxidative chlorination of ethane Catalyst- 8 wt % CuO/cement; T = 673K; reactant ratio C2H6 : HC1 : O2 = 1 : 2 : 1. No.

~, s

Reactant conversion, % C2H6

HC1

O2

Yields scaled to converted ethane, % C2H4C12

C2H3C1

C2H4

COx

1

1.5

80.6

36.0

91.2

25.0

34.1

34.6

2.6

2

3.2

87.5

31.7

89.5

19.4

36.8

40.1

3.2

3

5.6

89.1

30.4

88.6

11.3

38.2

43.5

4.0

4

7.9

90.9

30.0

87.9

10.1

35.6

46.8

6.5

5

10.0

92.7

29.1

86.0

8.2

33.0

47.6

9.7

We can suppose on the strength of the data listed in Table 2 that at the short times-onstream, the major contribution to the formation of deep oxidation products is made by saturated chlorinated hydrocarbons: 1,2 dichloroethane and ethyl chloride. On increasing timeon-stream to more than 6 s, we observed a sharp increase in the yield of deep oxidation products together with the decrease in the yield of vinyl chloride. It is likely that at the longer times-on-stream, the rate of deep oxidation of vinyl chloride would increase and become higher than the rate of dichloroethane dehydrochlorination. Taking into account this fact, we believe that the optimum time-on-stream assuring the best total yield of ethylene and vinyl chloride would be 3--5 s. It was shown in the investigations that the ratio of initial reactants also essentially affects the process. It was found that the excess of hydrogen chloride is favorable for improving the selectivity of the process with reducing the yield of deep oxidation products. At 673K and the reactant ratio of C2H6 : HC1 = 1 : 1, the yield of COx ranges from 6 to 7%; at the reactant ratio of C2H6 : HC1 = 1 : 2, the corresponding yield is 3-----4% (see Table 2). A positive factor is that the carbonization of the catalyst therewith decreases. On the other hand, the increase in the excess of HC1 to ethane up to 3 : 1 involves the decrease in the yield of unsaturated hydrocarbons due to the inhibition of the dehydrochlorination of 1,2dichloroethane and ethyl chloride with hydrogen chloride. The excess of oxygen increases the conversion of ethane mainly due to its oxidation: the yield of carbon oxides increases by 1.8-2 times. Thus, the optimum reactant ratio to provide the best yields of the target products is C2H6 : HC1 : O2 = 1 : 2 : 1.

313 Perfect stability of copper-containing cement catalysts in the oxidative chlorination of ethane was confirmed by their performance for 1500 hours without any decrease in the catalytic activity.

4.Conclusions The results obtained substantiate that the utilization of copper---cement catalysts offers promise for the synthesis of vinyl chloride from ethane at law temperatures in a single step. The proposed efficient and stable copper-cement catalyst will assist in the development of a new technology for the production of vinyl chloride from ethane. This technology is lowwaste and balanced in raw materials with meeting modem requirements of ecological safety. It would be appropriate to conduct the process of vinyl chloride production from ethane, hydrogen chloride, and oxygen in a fixed bed of copper---cement catalyst modified with alkali metals, for example, at 623--673K, time-on-stream of 3--5 s, and reactant ratio of C2H6 : HC1 : 02 - 1 : 2 : 1. Under these conditions, the conversion of ethane is more than 90%, and the total selectivity to ethylene and vinyl chloride is 85-90% at the yield of deep oxidation products COx no more than 3--4%.

REFERENCES

1. Yu.A. Treger, V.N. Rozanov, M.R. Flid, L.M.Kartaschov, Usp. Khim., 57,No 4(1988) 577 2. H.Rigel, H.D.Schindler, M.C.Sze. Chem.Engng.Progr.,.69, Nol0, (1973) 89 3. E.I. Gel'perin, Yu.M. Bakshi, A.K.Avetisov, A.I. Gel'bschtein, Kinet. Katal., 19, No 6 (1978) 527. 4. A.J.Magistro, P.P.Nicholas, R.T.Carrol, J.Organ. Chem., 34 (1969) 271 5. E.I. Gel'perin, Yu.M. Bakshi, A.K.Avetisov, A.I. Gel'bschtein, Kinet. Katal., 20, No 1 (1979) 129. 6. E.I. Gel'perin, Yu.M. Bakshi, A.K.Avetisov, A.I. Gel'bschtein, Kinet. Katal., 24, No 3 (1983) 633. 7. M.M. Mallikarjunan and S. Zahed Hussain, J. Sci., Ind. Res., 42 (1983) 209. 8. V.I. Yakerson, E.Z. Golosman. React. Kinet. Catal. Let., 55, No2 (1995) 455 9. V.I. Yakerson, E.Z. Golosman. Scientific Bases for the Preparation of Heterogeneous Catalysts. VI Intern. Symp.Preprint. 3 Poster Session II Louvain-la-Neuve (Belgium), (1994) 105 10. I.I. Kurlyandskaya, I.G. Solomonik, E.D.Glazunova, E.A.Boevskaya, Yu.M.Bakshi, E.Z. Golosman,V.I. Yakerson, Khim. Prom-st, Moscow, No. 6 (1996) 368. 11. M.R. Flid, I.I. Kurlyandskaya, I.G. Solomonik, M.V.Babotina, Khim. Prom-st, Moscow, No. 6 (1996) 364.

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3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

315

Oxidative Conversion of LPG to olefins with Mixed Oxide catalysts: Surface Chemistry and Reactions Network M.V.Landau a, M.L.Kaliya a, A.Gutman a, L.O.Kogan a, M.Herskowitz a and P.F. van den Oosterkamp b aBlechner Center for Industrial Catalysis and Process Development, Ben-Gurion University of the Negev POB 653, Beer-Sheva 84105, Israel Tel. (972-7)-6472141, Fax.(972-7)-6472902 bKinetics Technology International (KTI) B.V., POB 86, 2700 AB Zoetermeer,The Netherlands,Tel.31 (79)-3531453, Fax.31 (79)-3513561 The catalyic performance of three mixed oxide catalytic systems V-Mo-, V-Mg and RE-LiHalogen (RLH) in LPG oxidative conversion was measured at different O2/LPG ratios, temperatures and WHSV. At high LPG conversions V-Mo-based catalysts yielded low olefins selectivity and high LPG combustion (CB), V-Mg - medium olefins selectivity by oxidative dehydrogenation (ODH) route and medium LPG CB selectivity, while RLH catalysts displayed high olefins selectivity by ODH and cracking (CR) routes at low CB. TP-reaction experiments and the effects of oxygen partial pressure on catalytic performance indicated a dynamic interaction of surface oxygen in the ODH, CB and CR routes. ESCA and TPD measurements detected three types of surface oxygen with different nucleophility and bonding strength. Their distribution correlated with LPG conversion selectivities. A correlation between catalysts acidity, the surface exposed metal cations concentration and the productivity by the CR route was derived. The surface basicity was also significant in olefins productivity by the ODH and CR routes. The selectivity of LPG oxidative reactions were attributed to different intermediates formed on the surface as a result of interaction of C3-C4 paraffins with oxygen atoms of different nucleophility. Both the redox balance of surface metal cations and the acidity-basicity balance are proposed to be significant. 1. I N T R O D U C T I O N Catalytic oxidative conversion of low paraffins into olefins, a potential alternative to steam cracking, is one of the attractive optiopns that could decrease the process temperature, minimize the coke deposition at the reactors walls and increase the olefins productivity. Various catalytic processes for oxidative production of ethylene, propylene and butylenes have been published. A review of the published results measured with individual C2-C4 paraffins [1] allowed to select three most efficient oxide catalyst systems for the study: V-Mo- [2], V-Mg- [3] and Mg-RE-LiHalogen (Mg-RLH) [4]. Comparison of their performance in LPG oxidation showed that V-Mocatalyzed mainly the full paraffins CB, V-Mg- displayed average olefins selectivity producing a large amount of butadiene while the RLH - containing oxide systems showed the highest olefins selectivity at high LPG conversions producing substantial amounts of C2-C3 olefins by CR and ODH routes [ 1] The purpose of this work was to study the states of surface oxygen and relate them to the catalytic performance of selected catalysts: V-Mo, V-Mg and RLH.

2. EXPERIMENTAL

Preparation of Catalysts. V-Mo-catalysts were prepared according to procedure described in [2]. Ammonium metavanadate and paramolybdate were dissolved separately at 70~ third solution containing all the other metal components in form of nitrate salts was mixed with the first two evaporated by mixing. The catalyst material was crushed, sieved, dried at 120~ and calcined at 350oc for 5 h. V-Mg samples were prepared by mixing the MgO obtained by decomposition of Mg(NO3)2 or Mg(OH)2 (with addition of SiO2 or TiO2 powders in some cases) with water solution of ammonium metavanadate (containing metal nitrates in some cases), evaporation the suspension to dryness, dried at 120~ and calcined at 550~ for 6 h. The RLH- catalysts were prepared via an aqueous slurry containing LiNO3, NI-I4-halogen salt, Dy-oxide and the second

316

metal oxide (MgO,Ce-oxide or transition metal oxide). The water was evaporated, the paste dried at 130~ resulting solid was crushed,sieved and calcined at 500oc for 2h and at 750oc for 16h. Catalysts testing. A tubular titanium reactor 17 mm ID and 250 mm length supplied with the central thermowell was designed to test the catalysts over wide range of temperature and various feed compositions. Hydrocarbons - 25wt.% n-C4H10- 25wt.% i-C4H10- 50wt.%C3H8 (LPG artificial mixture) or its components, oxygen and nitrogen were fed separately by mass flow controllers (Brooks Instrument) and mixed in preheater at 450oc. The reactor was inserted into Carbolate tubulat oven, uniformly heated over a length about 50 mm. 1-5 g catalyst diluted with quartz pellets at 1:3 ratio was loaded between layers of quartz pellets. Axial temperature gradient in the catalyst layer during the tests was less than 5~ Homogeneous LPG oxidation in titanium reactor filled with quartz pellets at temperatures lower than 600oc was less than 5 wt.% conversion. The analysis of the reaction products excluding water was performed on line with GC HP-5890 that contained four columns - 45/60 Molecular Sieve 13X, 10 ft x 1/8"; 50 m x 0.53 mm Plot A1203; 80/100 Hysep Q 4 ft x 1/8" and 1 ft x 1/8", with internal switching valves and two detectors TCD and FID controlled by ChemStation analytical software. Selectivity was defined as wt of olefins in product divided by the wt of converted LPG feed. Catalysts characterizations. The catalysts composition was measured by energy -dispersive Xray (EDAX) - JEM-35, JEOL Co., link system AN-1000, Si-Li detector. The surface area was determined using BET method (ASTM 3663-84). Phase composition was measured by XRD in conventional, automated Philips PW 1050/70 diffractometer equipped with a long, fine focus Cu anode tube, 40 kW, 28 mA, a scintillation detector and a diffracted beam monochromator. The phase identification was carried out according to JCPDS-ICDD powder diffraction cards. PHI 549 SAM/AES/XPS apparatus with double CMA and Mg Ka X-ray source has been used for X-ray Photoelectron Spectroscopy (XPS) measurements of the catalysts. After recording general survey spectra, high resolution scans were taken at pass energy (25 eV) for the O ls peaks. The spectral components of O signals were found by fitting a sum of single component lines to the experimental data by means of non-linear least-square c.urve fitting to Gauss-Lorentz shape function using software provided by instruments manufacturer for peaks deconvolution. Care was taken to protect the calcined fresh samples from the contact with atmosphere by pressing them into 10 mm disks and transfering to the ESCA analytical chamber. The quantitative distribution of oxygen atoms with different O ls characteristics as well as total atomic surface concentrations of oxygen were calculated by conversion the peak areas into atomic compositions taking in account the sensitivity factors of all detected elements. Binding energies were referenced to the carbon ls line at 284.5 eV. The TPD and TP-reaction measurements were carried out in AMI-100 Catalyst Characterization System (Zeton-Altamira) equipped with quadrupol mass-spectrometer (Ametek1000). 3. RESULTS AND D I S C U S S I O N

3.1. Phenomenological description of observed catalytic effects Table 1 presents the olefins selectivitiy and productivity measured catalysts at about 30% LPG conversion. The measurements were temperature and O2/LPG ratio, keeping the LPG conversion constant by olefins selectivity is determined by a few basic components increasing in

with all the tested oxide carried out at constant varying the WHSV. The following sequence:

V-Mo- (5.1-8.4%) < V-Mg- (39.2-55.0%) < RLH (67.0-79.0%). The nature of promoters or components in RLH catalysts affected mainly the olefins productivity. The significance of different reaction routes is apparent in Table 2 that compares the CR and CB selectivities measured with selected representatives of the three catalyst groups: V-Mo-Nb-SbCa (Cat.A), 0.07V2Os-Mg(Cat.B) and Mg-Dy-Li-C1 (Cat.C). It also includes the results obtained with a catalyst that yielded a higher olefins productivity where the RLH composition was supported on a transition metal oxide (TM-RE-Li-C1, Cat.D). LPG was almost fully combusted on the V-Mo-catalyst. V-Mg-catalyst converted LPG mainly by ODH and CB routes with about equal efficiency. RLH catalysts enhance the ODH and CR routes with relatively low CB. Table 3 comoares the catalvtic oerformance of M~-Dv-Li-C1 catalyst in oxidation of individual LPG

317

components. All the hydrocarbons were converted mostly by ODH and CR routes, with CR selectivity increasing in the sequence: propane < n-butane < i-butane, so that the contribution of cracking products to total olefins yield was 55-65%. Figure 1 presents the olefins selectivity as a function of LPG conversion. Such plots are commonly used for comparison of low paraffins oxidation catalysts [5,6]. The V-based catalysts showed strong decrease in olefins selectivity with increasing conversion ( more expressed with V-Mo-) normally found with ODH catalysts [5.6], while the selectivity of RLH catalysts was almost independent on LPG conversion. Table 1 Compositions and performance of the tested catalyst belonging to the three selected groups Catalyst group

V-Mo

V-Mg

Catalyst composition

S.A., m2/g

0.09V205 0.74MOO3 0.02Nb205 0.02Sb203 0.13CaO 0.05V205 0.83MOO3 0.12CaO 0.17V205 0.83MOO3

Phase composition Olefins Olefins sel.*),% product. g/gCat., h

14

Sb204, Nb205, 8.4 SbNbO4 [Mo4011]O 5.1 MoO3,[Mo4011]O, 7.5 VMoO14

6 10

0.07V205 0.93MgO 0.07V205 0.93MgO 0.05V20 5 0.79MgO 0.16SIO2

60 100 90

0.05V205 0.94MGO 0.006TIO2 0.004Cr203 0.07V205 0.88MGO 0.05Li20 0.06V205 0.79MGO 0.05Li20 0.1C1

55

MgO, Mg3V208 MgO, Mg3V208 MgO, Mg3V208, Mg2SiO4 MgO, Mg3V208

57 52

Mg-RLH 0.8MgO 0.09Li20 0.002Dy203 0.1C1 0.7MgO 0.09Li20 0.002Ce203 0.21C1 0.39MgO 0.43Ce203 0.003Dy203 0.08Li20 0.1 C1 0.88MgO0.01Li20 0.001Dy203 0.1I 0.82MgO 0.1Li20 0.004Dy203 0.08Br 0.77MgO 0.09Li20 0.005Dy203 0.14F *) T = 585~

V-Mo catalysts T = 500~

0.03 0.028 0.027

44.9 44.3 43.5

0.15 0.15 0.14

39.2

0.13

MgO, Mg3V208 --

55.0 54.5

0.18 0.18

20 18 19

MgO,LiDyO2,Li20 -MgO, CeO2

77.3 79.0 78.5

0.08 0.1 0.1

-15 20

MgO,LiDyO2,Dy203 MgO, DyOBr,Dy203 MgO,LiDyO2,Li20

82.0 70.0 77.3

0.25 0.16 0.02

O2/LPG = 1; LPG conversion -- 30%

Table 2 Performance of selected representatives of the three catalyst groups in oxidation of LPG *) Catalyst

A

B d

a

b

C c

d

a

b

D

a

b

c

c

d

a

b

c

d

8.4

91.6

-- 0.03 44.9 55.1 3.1 0.15 77.3 22.0 39.00.08 74.7 28.0 36.2 1.03

*) Testing conditions as in Table 1; LPG conversion -30%; a-olefins selectivity,%, b -combustion selectivity,%, c -cracking selectivity (C1+C2),%, d - olefins productivity, g/g Cat.h A scheme of LPG reactions is proposed in Fig.2 to show the possible low paraffins transformations according to main three routes. It is based on measured products distributions and

318 90

.....

e

9 O'ql~

r

:

9 4~

9

4P,

41,~

9

;>

O r

,,, 30

0

.

.

.

.

I~

I-

0

A

I. . . . . . .

20

40 LPG c o n v e r t ; I o n ,

A

,.A

~.. 60

%

Figure 1. Olefins selectivity vs. LPG conversion plots for all the testcd catalysts

"~

C2H6

CzH4~

C3H 8 ~ - - . E . _ ~ - C " H ? ~

9 i-C4Hzo

~

8. ~ ,.,,

~

n-C4HI0 . ~ ' " 7.

~ *oa

~.

~

, O ~ , ,...~6 ~._,0.

~.~" .n__,_

.

tg.1

~

L ,-,-

7-- co~_

n-C4H~ ~~ cis,trans-C4H8 t6. H24-O2 - - ~ H20 ~7.CH4 + O 2 ~ CO + H20 ~ i-C4H8 l,.C H 4+O 2 ~ CO 2 + H 2

Cn_xH2(n.x),CH4,Cn_xH2(n.x)+2,H20

CRI 1.3,6,8,15,19 ODH CnH2n+2 CB

CO

,.~ ,o2

r.- CnH2n,H20

2,4,7,14 5,9,10,11,12,13,16,17,18

CO,CO/,H20,Hz Figure 2. Reactions network in oxidative conversion of LPG

8O

319

Table 3 Performance of Mg-RLH catalyst C in oxidation of LPG components *) Paraffin

n-Butane

i-butane

a

b

c

d

a

b

72.2

26.2

50.3

0.13

76.1

20.1

c

propane d

a

58.4 0.19

78.2

b

c

d

21.5

60.1

0.18

*) Testing conditions as in Table 1" Hydrocarbons conversion -- 30%. a - olefins selectivity,%, b combustion selectivity,%, c - cracking selectivity,%, d - olefins productivity, g/g Cat.h kinetic studies. The molar amount of hydrogen detected in products was higher than the amount of olefins could produce without a change in the number of carbon atoms while the amount of consumed oxygen was lower than needed for combustion of hydrogen stochiometrically. Therefore reactions like 5, 10, 13 and 19 in Fig.2 were included in the reactions network assuming production of hydrogen as a result of partial combustion. 3.2. Surface oxygen role

in

oxidative conversion of light

alkanes

Lattice oxygen in metal oxides reacted in catalytic cycles is replenished by reoxidation [7-9]. The effect of O2/LPG ratio on the catalytic performance of three selected catalysts shown in Fig.3 indicates that oxygen from gas phase is a reactant in all the three routes of catalytic conversion. Molecular oxygen could react with adsorbed hydrocarbons or oxygen bonding and activation at the catalysts surface could be nesessary.

90

/ ?

60

Catalyst A

~

80

3

90

atalyst B

1

2 ~

3

60

40

3O

30

0 0

0.5

1

1.5

2

OxygenlLPG molar rallo

2.5

--

0

I

0.5

--I .....

1

t-

1.5

OxygenlLPG molar ratio

t---

2

----t--

0

I

0.4

0.8

1.2

OxygenlLPG molar rallo

Figure 3. Effect of O2/LPG ratio on performance of selected catalysts in LPG oxidation at 585~ 1-LPG conversion, 2 - olefins selectivity, 3 - oxygen conversion, 4 - cracking selectivity Three consecutive runs of n-C4H 10-TP-reaction experiments were carried out with selected catalysts A,B,C and D. 25 cm3/min mixture 9%.vol. n-C4H10-He was fed to the reactor of the AMI-100 Catalysts Characterization System containing 3 g catalyst after heating to 200oc in He flow. Then the temperature was gradually increased at 5OC/min up to 600oc (Cat.A), 750oc (Cat.B,C) and 800oc (Cat.D). After reaching the required temperature, the gas flow was switched to He, catalysts were purged for 1 h,cooled to 200oc. Then the procedure of the first run was repeated. Before the third run, performed at the same conditions, the catalysts were reoxidized in 5%vol Oa-He flow at 550oc for 2 hours with subsequent cooling to 200oc. During the n-C4H10-TP-reaction runs the concentration of n-C4H10 in effluent gas as well as

320

concentrations of C4H8 (ODH product), C2H4,CH4 (CR products), CO2, H20 and H2 (CB productg) were monitored by MS. TP-reaction spectra for butane consumption (similar in shape for all catalysts) is shown in Fig.4a. It could be divided into three parts reflecting different catalysts performance as the temperature increases: I - no butane consumption at low temperature, II increasing butane consumption by fresh and reoxidized catalyst and no consumption with reduced catalyst, III- increasing butane consumption in all the runs that could be a result of other reaction routes (e.g. homogeneous reactions with oxygen evoluted by oxides decomposition). In the second series of TP-reaction experiments, the temperature during butane flow was changed in a ramp mode: it was increased in the same way as in previous series up to value a little higher than it corresponded to the end of the part II and kept constant for 1 hour. In this case (Fig.4b) the butane concentration spectra with fresh and reoxidized catalysts showed a minimum as a result of gradual conversion of surface oxygen while the reduced catalyst did not display defined peaks. The concentrations of all the other compounds monitored by MS displayed a -

(a)

F~gure 4. n-C4H10-TP-reaction spectra recorded with Mg.RLHcatalyst B

321

maximum over the same time pcricxt. The normalized MS peaks intensities lot butane, butylene, ethylene, methane, water, carlx~n dioxide and hydrogen measured at maximum butane consumption lot catalysts A,B,C and D are presented in Fig.5. The peaks normalization was done separately for every experiment, so their relative intensities shown in Fig.5 for different catalysts could not be compared. In all cases the products distribution with fresh and reduced catalysts were close to those measurcd in steady-state experiments, excluding high CO2 evolution with the fresh RLH catalysts. Reducing the V-Mg and RLH catalysts in butane flow almost fully depressed their ODH and CB activity shifting the products distribution in the direction of CR and dehydrogenation while the V-Mo- catalyst in reduced form produced the same CB products as fresh and reoxidized form with lower efficiency. These results are evident for the need for adsorbed oxygen species in the reaction cycles producing products according to the three main conversion routes detected in steady state experiments. Then the differences in performance of the three selected catalysts groups in LPG o~dation is probably caused by different states and concentrations of the surface lattice oxygen atoms. It is widely accepted 18-101 that the ability, of surface oxygens Os to react with hydrocarbons and the type of reaction depend on the distribution of Os among the different species: O2(gas) ~ O2(ads) w," O2-(ads) ~-*'20-(~s) w-~'202-(lattice). The performance of the most strongly bonded lattice oxygen that could be removed at high temperatures by the reaction with hydrocarbons in catalytic cycles is governed by their nucleophilicity being directly related to the effective negative charge and bonding strength [8-11]. Those characteristics together with surface concentrations of different oxygen forms for selected catalysts A-D were measured by TPD and ESCA.

Fi~zure 5. n-C4Hlo-TP-reaction products distribution with catalysts A-D at butane consumption

322

3.3. Surface chemistry characterizations The TPD experiments were carried out with 3 g catalysts A,B,C and D in He flow 25 cm3/min monitoring by MS the evolution of 02, CO2 and H20 over a temperature range 200-800oc ( for VMo- catalyst 200-600oc), heating at 5~ The results presented in Table 4 showed that only V-Mo- catalyst contains comparatively weakly bonded oxygen that could be partially desorbed at the temperatures used in steady-state catalytic tests. The oxygen bonding strength corresponding to Table 4 He-TPD of fresh catalyst Catalysts

Desorbed species:

02

A B C D

H20

CO 2

a

b

c

a

b

c

a

b

c

>600 680 705 ND

>20 40 50 ND

480 560 650 ND

ND 700 720 480

ND 500 80 8

ND 510 640 420

ND 710 720 580

ND 30 300 90

ND 580 600 550

a - Temperature of peaks maximum,~ : b - normalized MS peaks intensity c - Temperature of initial product desorption, oC; ND - not detected the temperatures of initial oxygen desorption and its maxima, increased in catalysts sequence: A D ( Table 6) expose more metal cations and chlorine that behave as electron acceptors. Thus increasing the amount of highly nucleophilic oxygen atoms in the same row as electron donors should be accompanied by substantial changes in catalysts acidity-basicity. Those characteristics were measured by NH3- and CO2-TPD after saturation the catalysts samples with corresponding gases at 40oc. The results are shown in Table 8. V-Mo- catalyst displayed the lowest acidity corresponding to the lowest metal cations concentration but about 50% of the acid sites were strong desorbing ammonia at >250~ The other catalysts contain few strong acid sites but the total acidity strongly increased in the sequence B250oC/total -

-

-

A

B

C

D

30 0.5

81 0.002

140 0.04

340 0.03

0.5 0

7 0.3

4 0.4

4 0.6

to the absence of highly nucleophilic oxygen species. The basicity of other catalysts was about one order of magnitude higher: V-Mg and Mg-RLH catalysts displayed about equal distribution between strong and weak basic sites while at the surface of catalyst D the relative amount of strong basic sites was more than twice higher. It corresponds to apperance of highly nucleophilic oxygen species and increasing their nucleophility from C to D (Table 6).

325

3.4. R o l e of d i f f e r e n t

s u r f a c e s p e c i e s in catalytic cycles

Comparison the surface characteristics of selected representatives of the three catalysts groups with their catalytic performance in LPG oxidation show: i - at temperatures less than 600oc all three LPG oxidative conversion cycles - ODH, CR and CB, are controlled by interaction of hydrocarbons with surface lattice oxygen atoms OI and OII, that form surface OH-groups being removed by dehydroxylation before reoxidation, as indicated from the results of TP-reaction and TPD experiments discussed in the part 3.2. ii - combination of OII oxygen species with low nucleophilicity (basicity) bonded to easy reducible metal cations (V,Mo) with acid sites leads to CB increasing with increased acid sites strength; it was indicated by direct correlation between olefins selectivity measured with catalysts A-D (Table 2) and parameter [ 1-OII/Ototal] (Table 6)reflecting decrease of CB selectivity with decrease of the fraction of OII in all the surface oxygen atoms and furthermore by substantial increase of the strong acid sites concentration from V-Mg to V- Mo (Table 8). iii - combination of OI oxygen species with high nucleophility (basicity) bonded to hardly reducible cations (Mg,RE) with weak acid sites leads to ODH and CR increasing with increased basicity of OI atoms, as indicated by comparing changes in the fraction of OI (Table 6) and their basic strength (Table 8) from catalyst A to catalyst D with CR and olefins selectivities of those catalysts presented in Table 2. iiii - the efficiency of CR conversion route increases with increased weak acidity of the catalyst as indicated from the direct correlation between CR productivity of A-D catalysts that could be easily estimated from the data of Table 2 and surface concentrations of metal cations and chlorine given in Table 6. Based on this information two different modes of paraffins activation are assumed, leading to CB or ODH-CR products depending on catalysts surface chemistry that are consistent with generally accepted models [8-11]. V-Mo- catalyst containing strong acid (electron-acceptor) sites, easy reducible cations andweak nucleophilic (proton-acceptor) oxygen atoms could adsorb the hydrocarbon molecule as a result of hydride-ion abstraction by acid sites. Reduction of metal cation with splitting of one of metal-oxygen bonds and stabilizing the proton and carbanion in form of OH and alkoxy species: CnH2n+2 ?-

O H!1

/z,,

CnH2n+ 1 +

OH OCnH2n 1 (n-l)+ M e - O_ MIe (m--l; . _

0

I] m+

n+Me 0 lvle !

! !

| i

!

RLH catalysts do not contain strong acid sites and easy reducible metal cations but have strongly nucleophilic (proton-acceptor) oxygen atoms and weak acid sites. The hydrocarbon molecule could be adsorbed as a result of proton abstraction by strongly nucleophilic lattice oxygen without splitting the metal-oxygen bond and stabilization of proton and carbcation in form of OH and alkyl species: CnH2n+2 O~ CnH2n+ I O - M e (n'l)+ - 0 - M e n+ - 0 ~ ! i

! i

CnH2n+l O - M e (n-l)+ i i

H O - M e n+ - 0 | i

The subsequent transformations of alkoxy radicals containing strong C-O bonds at the surface of V-Mo- catalyst with weakly bonded oxygen atoms yields preferentially formation of full CB products with some hydrogen evolution, while alkyl radicals stabilized on acid sites at the surface of RLH catalysts as a result of C-H bonds polarization in the strong field of metal-

326

oxygen ion pairs should be preferentially transformed to olefins as a result of further hydrogen abstraction (ODH) or CR. The fraction of CR products in olefins depends on catalysts acidity increasing the lifetime of alkyl radicals on the catalyst surface.The V-Mg- catalyst contained the both types of surface oxygen OI and OII in about equal amounts (Table 6) displaying average acidity and basicity (Table 8) and including easy reducible (V) as well as hardly reducible (Mg) metal cations. As a sequence it showed an average catalytic performance. In both cases the catalytic reaction cycle became closed as a result of dehydroxyation of catalysts surface and further oxygen adsorption-insertion in the oxide lattice that in case of V- or V-Mo-containing catalysts is accompanied by increasing of metals oxidation extent. In addition to further reacting of alkyl and alkoxy intermediates at the catalysts surface with dynamic lattice oxygen they could be desorbed into gas phase and react there homogeneously with gas oxygen as it was demonstrated in [ 13] for V-Mg-catalyst. Testing the RLH catalysts in fixed-bed reactor with void fraction of catalysts layer varied from 28 to 43% showed that this route became significant at temperatures higher than 590oc but no substantial changes in products distribution were observed. 4.

SUMMARY

The RLH-based catalysts display high olefins selectivity at high LPG conversions producing olefins by oxydative dehydrogenation and oxidative cracking. The last charactristics allow them to produce ethylene from LPG that is the main product of steam cracking. Supporting the RLH system at different carriers affects mostly the catalysts productivity. The RLH-based catalysts display about 50% olefins yield with productivity per reaction volume close to steam cracking. The high selectivity of RLH-catalysts to olefins is a result of a definite combination of surface oxygen state, oxygen / metal cations ratio, redox properties of metal cations and acidity-basicity balance. Further studies are needed in order to understand the role of the support and the proper functioning of RE-Alkali-Halogen systems in oxidation of low paraffins. REFERENSES

1. M.V.Landau, M.L.Kaliya, M.Herskowitz, P.F.van den Oosterkamp and P.S.G.Bocqu6, CHEMTECH, 26, No.2 (1996) 24. 2. J.H.McCain, US Patent No. 4 524 236 (1985). 3. H.H.Kung and M.A.Chaar, US Patent No. 4 777 319 (1988). 4. C.J.Conway, D.J.Wang and J.H.Lunsford, Appl.Catal., 79 (1991) L 1. 5. F.Cavani and F.Trifiro, Catal.Today, 24 (1995) 307. 6. S.Albonetti, F.Cavani and F.Trifiro, Catal.Rev.-Sci.Eng., (1996), 413. 7. P.Mars and D.W.van Krevelen, Chem.Eng.Sci.(Special Suppl.), 3 (1954) 41. 8. A.Bielanski and J.Haber, "Oxygen in Catalysis", Marcel Dekker, Ink., New York, 1991. 9. V.D.Sokolovskii, Catal.Rev.-Sci.Eng., 32, No. l&2 (1990) 1. 10. G.Centi, F.Trifiro, J.R.Ebner and V.M.Franchetti,Chem.Rev., 88 (1988) 55. 11. H.H.Kung, Ind.Eng.Chem.Prod.Res.Dev., 25 (1986) 171. 12. J.Zi61kowski and J.Janas, J.Catal., 81 (1983) 298. 13. X.D.Peng, D.A.Richards and P.C.Stair, J.Catal., 121 (1990) 99. 14. J.C.Fuggle, L.M.Watson and D.J.Fabian, Surf.Sci.,49 (1975) 61. 15. H.H.Kung, Adv. in Catal., 40 (1994) 1. 16. T.Ito, J.X.Wang, C.H.Lin and J.H.Lunsford, J.Amer.Chem.Soc., 107 (1985) 5062. 17. D.Wang, M.P.Rosynek and J.H.Lunsford, J.Catal., 151 (1995) 155.

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 1997 Elsevier Science B.V.

327

Free Radicals as I n t e r m e d i a t e s in Oxidative T r a n s f o r m a t i o n s of L o w e r Alkanes. M. Yu. Sinev, L. Ya. Margolis, V. Yu. Bychkov, and V. N. Korchak Semenov Institute of Chemical Physics, Russian Academy of Sciences 4 Kosygin street, Moscow 117334, Russia Catalytic oxidative transformations of lower alkanes attract the attention as possible ways to transfer these substances into more suitable chemicals - olefins and oxygenates (alcohols, aldehydes, acids, etc.) - and to involve them into the industrial use as raw materials for chemical and petrochemical synthesis. However, the yields of desirable products reached up to date are not sufficiently high. The progress in the studies of intrinsic mechanism of catalytic partial oxidation of lower alkanes is not sustainable either. We believe that these two facts are correlated and that the analysis we performed in the present work can brighten up some important details of the mechanism of catalytic oxidation of lower alkanes*. Experimental facts and theoretical concepts existing in the literature indicate that the formation of free radicals plays an important role in a number of catalytic oxidation reactions [1-5]. In the present paper we analyze the contribution of free radicals to several oxidative transformations of lower alkanes over oxide catalysts. Based on the thermochemical data and on the results of kinetic simulations it is shown that the observed reaction kinetics and product compositions in the mentioned above processes are determined by a set of interdependent heterogeneous and homogeneous reactions of free radicals, i.e. they should not be considered as "spectators" taking part in side reactions, but as principal intermediates causing the main features of lower alkanes oxidation and design of catalysts. 1. ACTIVATION OF ALKANE MOLECULES Alkane molecules do not have any specific "reactive centers", like functional groups or multiple bonds. This means that their activation can be carried out only by the bond dissociation or charge transfer processes. It is evident that the more energetically favorable is the first step of activation, the higher is the probability of its contribution to the overall reaction. In other words, using the value of the energy expenditures in different elementary steps ( Eex = AH - Est, where AH is the overall enthalpy change and Est is the energy of stabilization of the activated molecule or its fragment on the catalyst surface) one can estimate which one of them is more feasible. Such an analysis may be performed on the basis of the available thermochemical data (see, for example, [6]). The possible processes and corresponding values of Eex which we have considered are given below: * This study was carried out under the financial support of the Russian Foundation for Basic Research (research grant No. 96-03-32440)

328 (i) homolytic dissociation of C-H bonds accompanied by the formation of surface OH-group and free radicals [O] + RI-I => [OH] + R

(1)

Eex = DR_H where [O] - strong oxidizing surface center having a high affinity to the hydrogen atom; D i - energy of corresponding bond dissociation; (ii) heterolytic C-H bond dissociation with a proton abstraction on a strong basic center [ 0 2-] + RH => [ 0 2- ... H +] + R-

(2)

Eex = DR_H + IH - IRwhere I i - ionization potentials of corresponding particles; (iii) heterolytic C-H bond dissociation with a hydride-ion abstraction on a Lewis acidic site [M n+] + RH => [M n+ ... H-] + R +

(3)

E e x - DR_H + I R - I H(iv) ionization of alkane molecule [h+] + RI-I => [h+ ... e-] + RH +

(4)

Eex = IRH where [h+]

- hole center;

[h+ ... e-] - trapped electron. It is easy to demonstrate that the sign and the magnitude of energy changes in all these processes depend on the compensation of Eex by binding the fragments of the activated molecule to the surface centers. The values of Eex given in Table 1 show that the energy which has to be compensated is minimal for the process (1). On the other hand, one may assume that the energy of the O-H bond formed in this process is comparable to Eex, i.e. the second fragment (free radical R) should not be bound to the surface in order to compensate the energy expenditure. This assumption is in a good agreement with calorimetric measurements, according to which in the case of oxide catalysts active in oxidative coupling of methane (OCM) and oxidative dehydrogenation of ethane [7,8], as well as in total oxidation of alkanes [9] the O-H bond strength ranges from 250 to 470 kJ mo1-1. On the contrary, in the case of heterolytic C-H bond dissociation the energy expenditure is so large that its compensation requires the binding of both fragments, which must occur in a

329

Table 1 Energy expenditure on the activation of lower alkane molecules Energy expenditure ( kJ mol "1 )

Molecule

reaction (1)

reaction (2)

reaction (3)

reaction (4) .

CH4

431

1630

1308

1250

C2H6

410

1615

1183

1120

C3H8

398

1609

1162

1078

n-C4HI0

393

1605

1154

1037

iso-C4H10

389

1601

1120

1016

single reaction step. This requires the presence of paired centers with specific configurations and energy relations. Such a mechanism including the simultaneous abstraction of H + and H ions from n-butane as a first step of maleic anhydride formation was suggested by Trifiro et al. [ 10] to explain the unique properties of V-P-O catalysts. One may assume, however, that due to steric restrictions the smaller an alkane molecule, the lower is the probability of such a synchronic mechanism. The process (4) seems to be improbable because it requires the existence of the hole centers with an electron affinity comparable with ionization potentials of alkane molecules. The above analysis shows that the formation of free radicals in the interaction of alkane molecules with the surface of oxides may prove to be energetically preferable as compared to any other mechanisms of their activation. Furthermore, this process requires only one type of single active centers and it proceeds in a single step. The combination of these factors may render this process the most favorable. This conclusion is experimentally confirmed by the correlation between the concentration of strong oxidative sites and the catalytic properties of the oxide catalysts (see, for instance, [ 11,12]); the Polanyi-type relation between the activation energy for the oxidative dehydrogenation of alkanes in their interaction with oxides and the energy of O-H bonds formed simultaneously [13-15]. The difference in the reactivities of C1 - C4 alkanes is mainly caused by the difference in the Eex values. Estimations based on the data of Table 1 show that the difference in rate constants at 700 - 1000 K between methane and butanes over the same catalyst can exceed 103. The H-atom affinity in the case of efficient catalysts for methane activation should be the highest. As a result, if the O-H bond strength is high enough to compensate the energy expenditure in the reaction (1), the process of active sites regeneration (reoxidation) becomes more impeded and the difference in the optimal reaction temperatures for different alkanes can reach 100 K or more. ff the catalytic oxidation of alkane molecule starts with the formation of a free radical on the surface of an active catalyst particle and its escape to the gas phase, the complete reaction network includes both homogeneous and heterogeneous steps of the transformation of primary (CnH2n+l) and secondary radicals. Since all these processes are sufficient for the formation of the final products, the analysis of the influence of different factors on the -

-

330 selectivity of a complex heterogeneous-homogeneous process can be carried out by considering the elementary reactions of free radicals. 2. ELEMENTARY REACTIONS OF FREE RADICALS 2.1 H o m o g e n e o u s

reactions

The main types of primary gas-phase transformations of free radicals formed in the reaction (1) are - recombination CnH2n+l + CnH2n+l (+ M)=> C2nH4n+2 (+ M*)

(5)

- H-transfer and elimination (ifn >__2) CnH2n+l + CnH2n+l => CnH2n+2 + CnH2n

(6)

or CnH2n+l + 0 2 => CnH2n + HO2

(7)

or

CnI-I2n+l => CnH2n + H

(8)

- oxygen molecule capture CnH2n+l + 0 2 CnH2n+102

(9)

- oxidation Cnn2n+l + 0 2 => CnH2nO + O H

(lO)

or CnH2n+l + 0 2 => CnH2n+lO + O

(ll)

Reactions (5)-(8) and (10) lead to the formation of stable molecules (hydrocarbons and aldehydes). Subsequent reactions of peroxy- (CnH2n+lO2) and oxy-radicals (CnH2n+lO) formed in reactions (9) and (11) lead to the formation of oxygenates (alcohols, aldehydes, etc.), carbon oxides, and/or olefins. The fractions of radicals transformed into different fmal products depend on the reaction conditions (temperature, oxygen pressure) and on the number of carbon atoms in the alkane molecule. For example, the stability of peroxy radicals decreases with increase of the number of carbon atoms in the alkyl fragment, that is why the probability of total oxidation via their subsequent transformations decreases from methane to

331 butanes. The higher temperature also decreases this probability, due to the shift of the equilibrium (9) towards alkyl radicals, increasing the fraction of radicals transformed into the products of coupling and dehydrogenation. However, if the temperature increases beyond some certain value, the fraction of oxygen containing products increases again because of more sufficient contribution of reactions (10) and (11). In particular, the low efficiency of reactions of metyl radicals with 02 molecules likely causes the existence of a temperature "window" for the OCM process at 900-1100 K. The development of chains in the gas phase leads to the acceleration of the secondary radicals formation, as well as to the additional conversion of the initial reactants and to the shifts of product selectivities. 2.2. H e t e r o g e n e o u s r e a c t i o n s

As we have mentioned above, if the catalyst pores are sufficiently narrow, i.e. the species diffuse through them in Knudsen or transitional re~mes, the contribution of heterogeneous reactions of free radicals to the overall reaction rate and selectivity may be predominant. The main types of elementary reactions between radicals and surface sites proposed elsewhere [ 15] are the following: - H-atom transfer, for example [O] + CnH2n+l => [OH] + CnH2n

(12)

- O-atom transfer, for example [O] + CnH2n+l => [ ] + CnH2n+~O

(13)

- radical capture [0] + CnH2n+~ => [OCnH2n+q

(14)

Let us consider the possible role of these reactions in the formation of final products. If n > 2, the successive reactions (1) and (12) lead to the formation of the desirable product in the case of oxidative dehydrogenation processes. The possible contribution of alkoxy radicals to the formation of reaction products is already mentioned above. In this section we should emphasize that the relative probability of reactions (13) and (14) depends on the properties of the catalyst (oxygen binding energy E[ol) and on the reaction temperature: the higher the temperature and the lower the E[o], the more probable is the reaction (13). In this case one may expect an increase of selectivity of partial oxidation to oxygenates. The fate of radicals captured by the surface sites with the formation of the alkoxy groups depends on the number of carbon atoms in the alkane molecule, as well as on the properties of the catalyst surface. According to the data obtained by Aika and Lunsford with the use of IR spectroscopy and TPD [ 16], in the case of MgO (an oxide with very high E[ol) the methoxy groups decompose forming CO and H2, but in the case of higher alkoxides the formation of corresponding olefins takes place.

332 Taking into account the whole set of homogeneous and heterogeneous reactions, one may conclude that depending on the target product the requirements to the catalyst and to the reaction conditions should be different: if we wish to increase the yield of oxidative dehydrogenation products, we have to increase the temperature and to use the catalysts with higher E/ol. The rigidity of these requirements increases as the number of carbon atoms in alkane molecule decreases due to the increasing strength of C-H bonds and stability of peroxy-radicals. On the contrary, the lower the temperature and oxygen binding energy Eiol, the higher is the probability of the oxygenate production. We have to notice, however, that these requirements are contradictory if the olefm is an intermediate for the further formation of oxygenates. In this case it is substantial that the catalyst contains the active sites of different types: one of them (strongly-bound oxygen with high H-atom affinity) is responsible for the formation of free radicals, and the second one (with lower Etol) supplies O-atoms for the insertion into the organic molecule. The efficiency of oxides containing vanadium, molybdenum, tungsten, and similar cations as catalysts for partial oxidation of hydrocarbons to oxygenates is likely due to the division of functions between oxygen species of different types (terminal M n+ = O and bridge Mn+-O-M n+) which these cations are able to form. An analogous co-operation of oxygen species is likely to take place in the case of multicomponent catalysts: if one oxide phase actively produces free radicals which react subsequently with weakly-bound oxygen species on the surface of another component, the total rate of the final product formation will be higher as compared to that measured over each individual oxide. This explanation is alternative (or complementary) to the so-called "oxygen remote-control mechanism" discussed by Weng and Delmon [17], according to which a synergy in catalytic action may be caused by the transfer of active oxygen species between two oxide phases with different donor-acceptor properties. 3. SIMULATIONS OF SURFACE-ASSISTED FREE-RADICAL PROCESS The preliminary analysis of elementary reactions of free radicals in the presence of an active catalyst demonstrates that the heterogeneous generation of primary radicals initiates the homogeneous processes. In their turn, both primary and secondary free radicals affect reciprocally on the surface active sites. A kinetic model which considers the heterogeneous and homogeneous transformations as interdependent and presumes that all the particles (stable molecules and free radicals) present in the gas phase oxidation undergo both homogeneous transformations and interactions with active sites of the catalyst surface was discussed elsewhere [18]. This model was previously used to simulate the OCM reaction in a quasihomogeneous system. Taking into account that the subsequent fate of the radicals formed in the reaction (1) depends, in the general case, on the relation between the number of collisions with other particles in the gas phase and with the surface and also on the nature, concentration and reactivity of the surface centers, we utilized the approach proposed in [18] to simulate the heterogeneous-homogeneous oxidation of methane in combination with mass-transfer in the gas-solid system. The reaction space was considered as a gas volume of a varied thickness (L) exposed to a flat surface with a varied concentration of active sites (C). The results of simulations of the reaction accompanied by the one-dimensional masstransfer directed normally to the surface are given in Fig. 1-2. If C = 0, a self-acceleration

333

d[CH4]/dt, -

100

C

nmol/s.

5"10

-

-

16

m-2

5 " 1 0 15 hi-2

10-

5 " 1 0 14

-2

m

i

0.1

I

I I

I I

I

I

i

i

I

1 I

-3

t

I

-2

I

I I

I

I 1 I

I

I

i

-1

I

I

0

I I

I

I

I

I

I

i

I

log t (s.)

Figure 1. Methane conversion rate as the fimction of time at different concentrations of active sites on the surface (1000 K, 1 atm., CH4 902 = 10 91, L - 1xl0 -4 cm)

W(s)/W(tot)

log(lO 4 to.t) gas r e a c t i o n -4

0.8

-3

0.6 -2 0.4>

0

o _

i

i

i 4

0

-6

-5

-4

-3

-2

-1

"-

0 log L ( c m )

io

Figure 2. Effect of gas volume thickness on the fraction of the rate of heterogeneous reaction in total conversion and on the time of 1 0 % oxygen conversion ( 1 0 0 0 K , 1 atm., CH4 9 0 2 1 0 " 1, C = 5 x 1 0 1 6 m 2)

334 typical for chain reactions with branching was reproduced. At C = 5x1016 m -2 a kinetic behavior of that kind disappears, and the process becomes "linear", i.e. its rate reaches the maximum at t = 0 and then declines due to the consumption of the reactants. At intermediate C values the gradual transformation of kinetic behavior from a self-acceleration type to a "linear" type takes place. The effect of the gas volume thickness on the contribution of the surface reaction to the overall kinetics is presented in Fig.2. At L = 10 nm the time of 10% conversion characterizing the rate of reaction is -104 times less than in the case of a homogeneous gas reaction and the fraction of the rate of heterogeneous reaction Ws in the total conversion rate Wtot is nearly 1. At increasing thickness of the gas volume the fraction of the heterogeneous reaction and the rate of overall process both decline. However, even at L = 1 cm, the reaction occurring in the gas volume still experiences the influence of the surface taking part in the radical reactions. Taking into account that the specific surface areas of the oxide catalysts usually used for partial oxidation range between-~1 and 20 m2g1, the characteristic size of solid crystallites is -~10-5-104 cm and the role of heterogeneous reactions of radical species in the catalyst pores is likely predominant. According to the results of kinetic simulations, if the specific surface area of the catalyst is more than-~1 m2g-1, the most of radicals formed in the reaction (1) undergo the reverse transformation into the initial alkane molecules: [OH] + CnH2n+l--> [O] + CnH2n+2

(-1)

This means that, although the surface of pores is much larger than the outer surface of the grains, the contribution of the latter to the formation of the final products can be sufficient due to the lower probability of the reaction (-1) for the radicals formed outside the pores. The grain size which usually ranges f r o m - 0.1 mm to few centimeters, makes it possible for the homogeneous chain reaction to develop in the free volume of the catalyst bed (see Fig.2). Recently we observed the effect which supports the conclusion about the substantial role of the radical reaction outside of the catalyst grains. When a very efficient OCM oxide catalyst (10% Nd/MgO) was placed into the reactor together with an inactive metal filament (Ni-based alloy) the sharp increase of conversion accompanied by the selectivity shift from oxidative coupling to the formation of CO and 1-12 was observed [19]. Since the metal component has a low activity also with respect to ethane oxidation, this behavior is not due to successive oxidation or decomposition of C2 hydrocarbons on the metal surface, but should be attributed to the reactions of methane oxidation intermediates. Almost total disappearance of ethane (which is a product of CI-I3 radicals recombination) and acceleration of the apparent reaction rate by the addition of an "inert" material indicate that the efficiency of methane oxidative transformations can be substantially increased if the radicals have a chance to react outside the zone where they formed and the role of reaction (-1) decreases. Although the second (metal) surface is not active enough to conduct the reaction of saturated hydrocarbon molecules (methane and ethane), the radicals generated by the oxide can react further on the metal surface. As a result, the fraction of the products formed from methane activated in the reaction (1) increases, and the formation of the final reaction mixture of different composition takes place.

335 4. CONCLUSIONS 1. The most energetically favorable process of lower alkanes activation over oxide catalysts is a homolytic C-H bond dissociation with the formation of free radicals. The difference in energy expenditures for the formation of free alkyl radicals cause the difference in reactivities between C1-C4 alkanes. 2. The main factors determining the efficiency of different oxides as catalysts for lower alkanes oxidation are the H-atom affinity of strong oxidizing surface sites and the oxygen binding energy. These thermochemical factors cause the rates and directions of free-radical reactions and, as a result, the catalytic activity and selectivity to certain products. 3. The total rate of reaction and the selectivity to different products (olefins, oxygenates, carbon oxides) depend on relative efficiencies of different transformations of free radicals in the gas phase and in the heterogeneous steps, as well as on the transport phenomena. REFERENCES

1. 2. 3. 4. 5.

V.M. Polyakov, Usp. Khim., 17 (1948) 351. D. J. Driscoll, K. D. Campbell, and J. H. Lunsford, Adv. Catal., 35 (1987) 139. J.H. Lunsford, Langmuir, 5 (1989) 12. T.A. Garibyan and L. Ya. Margolis, Catal. Rev., Sci. Eng., 31 (1989-1990) 35. M. Yu. Sinev, L. Ya. Margolis and V. N. Korchak, Usp. Khim. (Russ. Chem. Rev.), 64 (1995) 373. 6. V.N. Kondratiev (ed.), Chemical Bond Dissociation Energies, Ionization Potentials and Electron Affinities, Handbook, Moscow, Nauka, 1962 (in Russian). 7. V. Yu. Bychkov, M. Yu. Sinev, V. N. Korchak, E. L. Aptekar' and O. V. Krylov, Russ. Kinet. Catal., 30 (1989) 1137. 8. M. Yu. Sinev, V. Yu. Bychkov, V. N. Korchak, and O. V. Krylov, Catal. Today, 6 (1990) 543. 9. V. Yu. Bychkov, M. Yu. Sinev, Z. T. Fattakhova, and V. N. Korchak, Russ. Kinet.Catal., 37 (1996) 366. 10. G. Centi, F. Trifiro, J. R. Ebner, and V. M. Franchetti, Chem. Rev., 88 (1988) 55. 11. D. J. DriscoU, W. Martir, J.-X. Wang, and J. H. Lunsford, J. Am. Chem. Soc., 107 (1985) 5062. 12. M. Yu. Sinev, V. Yu. Bychkov, Yu. P. Tulenin, B. V. Rozentuller, and A. M. Rajput, 9th Soviet- Japanese Seminar on Catalysis, Novosibirsk, Nauka, 1990, p. 75. 13. A. A. Bobyshev, V. A. Radtsig, Russ. Chem. Physics, 7 (1988) 950. 14. A. A. Bobyshev, V. A. Radtsig, Russ. Kinet. Catal., 29 (1989) 638. 15. M. Yu. Sinev, Catal. Today, 13 (1992) 561. 16. K.-I. Aika and J. H. Lunsford, J. Phys. Chem., 81 (1977) 1393. 17. L. T. Weng and B. Delmon, Appl. Catal., 81 (1992) 141. 18. M. Yu. Sinev, Catal. Today, 24 (1995) 389. 19. Yu. P. Tulenin, M. Yu. Sinev, and V. N. Korchak, 1 lth Int. Congress on Catalysis, June 30 - July 5, 1996, Baltimore, ML, USA, Programme and Book of Abstracts, P-275.

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3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

337

Alternative m e t h o d s to prepare and m o d i f y v a n a d i u m - p h o s p h o r u s catalysts for selective oxidation o f h y d r o c a r b o n s . V.A.Zazhigalovla, J.Haberlb, J.Stochlb, A.i.Kharlamov 2, i.V.Bacherikovala and L.V.Bogutskaya Ia Ukrainian-Polish Laboratory of Catalysis: a) Institute of Physical Chemistry, National Academy of Sciences of Ukraine, Pr.Nauki 31, Kyjiv-22, 252022 Ukraine b) Institute of Catalysis and Surface Chemistry, Polish Academy of Sciences, Niezapominajek, Krakow, 30-239 Poland 2 Institute for Materials Science Problems, National Academy of Sciences of Ukraine, Kryjanovski 3, Kyjiv, 252680, Ukraine Among numerous compounds formed in vanadium-phosphorus-oxide system, vanadyl pyrophosphate is known to be an efficient catalyst for C4-C5 paraffins partial oxidation [1 ]. Typical process of its synthesis can be represented by a following scheme: ROH, Reductant V 2 0 5 -k- H3PO4

> V O H P O 4 . 0 . 5 H 2 0 . . . . . 5> ( V O ) 2 P 2 0 7

It has been established that the properties of vanadyl pyrophospate are strongly dependent on its biography, i.e. the preparation method, presence of overstoichiometric phosphorus and additives [1-4]. Therefore, considerable effort of the researchers was directed to optimization of the synthesis technique and in recent studies also, non-traditional methods for the catalysts preparation were considered [5-10]. It has been shown by us [9,10] that mechanochemical treatment is a perspective method to modify the properties of the precursor VOHPOa.0.5H20 and thus, to influence the catalyst prehistory. The present paper deals with the possibilities of mechanochemical and barothermal treatments applied at different stages of the catalyst synthesis: the initial reactants, the precursor and the final catalyst. 1. E X P E R I M E N T A L

V205 (purefor analysis) and H3PO4 (pure) were used as initial reagents. The synthesis of the precursor of VPO catalysts was carried out according to the procedure described in [11 ], starting from V205 and H3PO4 in butanol medium in the presence of organic reducing agent. The solid product, after filtration, was heated stepwise up to 300 ~ in vacuum (total time was 60 h). The activation* of the precursor i.e. its transformation into the (VO)2P207 phase was performed in a reactor, just before the catalytic test, with a gas reaction mixture consisting of 1.7% Call10 in air. The activation was carried for 72 h at the temperature gradually rising up to 440~ Witht the exception of samples activated by means of mechanochemistry

338 Mechanochemical treatment was applied at different stages of the synthesis described by scheme 1: to the starting reagent V205, precursor VOHPO4.0.5H20, final product (VO)2P207 and mixtures of the powder VPBiO precursor + La203. It was carried out in a planetary mill at 3,000 rpm. La203 was prepared prior to the milling, by decomposition of La2(CO3)3.xH20 (Aldrich) in an inert gas flow. The solids for the treatment were either suspended in ethanol or water or used without any dispersant (dry milling). Barothermal synthesis and treatment were carried out in a stainless steel autoclave lined with internal Teflon glass (V - 20 cm3). The "barothermal" procedure included its both well known variants named hydrothermal synthesis (in the presence of water) and organothermal one (in the presence of organic compounds) as well as synthesis without any solvent. For the synthesis the powders of starting compounds and phosphoric acid with/without solvent were loaded into the Teflon glass. For modification of VPBiO, precursor grains (D = 5 mm, L = 5-6 mm) were placed in the glass mold in the reactor, and the modificator was located in the space between the glass and autoclave walls. Different temperatures and times of treatment were applied in these experiments. The method of the barothermal treatment was described in details in [ 12]. Phase composition of the samples was analyzed using DRON-3M X-ray diffractometer with Cu Ka radiation. The specific surface area (SsA) of the samples was measured by BET method on Gasochrom-1. Thermal analysis was carried out with the thermoanalytical instrument Derivatograph Q-1500 D (system F.Paulik-J.Paulik-L.Erdey) in helium atmosphere at a heating rate of 10 K/min. The surface composition was examined using VG ESCA-3 X-ray photoelectron spectrometer (A1 Kal.2). The spectra were calibrated against C ls (284.8 eV) line as the standard in the binding energy determination. Jeol-100 CX transmission electron microscope and Nanoscope scanning tunneling microscope were used for the investigation of morphology. Details of the measurements and data processing are given in [ 11 ]. Catalytic properties of the synthesized samples after activation were examined in the hydrocarbon-air reaction mixture in reactions of the oxidation of: i) n-butane (1.7 vol. % in air) to maleic anhydride, ii) butene-2 (1.6 vol. % in air) to maleic anhydride, iii) n-pentane (1.2 vol. % in air) to maleic and phthalic anhydrides, and iv) propane (1.8 vol. % in air) to acrylic acid. Catalytic tests were performed in the flow system with GC control of the reaction products. 2. RESULTS AND DISCUSSION

2.1. Mechanochemistry 2.1.1. Mechanochemical modification of the initial reagents for synthesis of VPO catalyst To prepare VPO precursor (P/V = 1.15) samples of the V205 reagent untreated (V205-R) and after wet (in ethanol, V205-E) and dry (V2Os-D) milling were used. Some properties of these solids are given in Table 1. It has been established that mechanochemical treatment increases the specific surface area of V205 and produces V 4+ ions. The latter phenomenon is indicated by an appearance of the low-energy contribution in the XPS spectrum of V 2p electrons. The STM study [13,14] showed that after the mechanochemical treatment in ethanol the change of V205 texture took place due to an anisotropic plastic shift deformation along the planes parallel to (001). This leads to an increase of the relative exposure

339

Table 1. Properties of V2_Qs_before and after its mechanochemical treatment Sample

V205-R V2Os-D V2Os-E

SSA

Treatment Medium Time, min.

m2/g

Air Ethanol

3.8 13.8 8.8

XRD R*

12 30

1.33 0.85 4.33

W*

3.5 7.5 3.5

XPS Binding energy O 1s V2p(1) V2p(2)

v(1)/

531.0 531.1 530.6

0 0.09 0.37

516.3 516.1

517.8 517.6 517.6

V(2) .

* R - the ratio of I(001)/I(~10) indexes intensity, W - WHPM - width at the half of the (001) reflection of the latter over the surface of V205 (see Table 1) while an average size of the particles remains almost unchanged [13]. On the contrary, dry milling of V205 by chaotic destruction of the crystals produced smaller particles, which is reflected by the increase of the XRD peaks width (Table 1). The degree of surface reduction of vanadium pentoxide (given by the content ratio V(1)/V(2) in Table 1) was much higher after the treatment in ethanol as compared to dry milling. No special features of the synthesis of VPO-D from VzOs-D were observed as compared to traditional VPO-R compound synthesis, both lasting 12-16 h. But synthesis of VPO-E precursor using VzOs-E proceeded at a much higher rate and in the presence of organic reducing agent was completed in 1h (sample VPO-E 1). Another preparation using V2Os-E was carried out without reducing agent and with smaller amount of the solvent (sample VPO-E2). In this case the time needed for the full formation of the precursor phase was about 2.5-3.0 h. Some properties of the prepared VPO precursors are listed in Table 2. Table 2 Properties of VPO precursors prepared from V2Q5 treated mechonochemically Sample

SsA

DTA

XRD*

Tendo Texo I0.570/I0.329 VPO-R VPO-D VPO-E1 VPO-E2

mZ/g

~

~

20.2 19.5 4.6 12.1

448 435 468 448

500 485 505 495

/I0.293 75/46/100 100/40/90 73/45/100 100/35/76

XPS _Binding energy, eV O 1S V 2p P 2P

(P/V)s

532.3 532.2 532.4 532.3

2.08 2.00 3.30 2.05

517.5 517.4 517.5 517.5

133.6 133.6 133.7 133.7

*Intensity ratio for reflections at d = 0.570, 0.329 and 0.293 nm It follows from the data in Tab.2, that reduction of the V205 particles size results in VPOD precursor in increased intensity of the 0.570 nm peak attributed to the exposure of (001) plane containing the vanadyl groups. Also some decrease of the temperatures, at which the amorphous (Tendo)and the crystal (Texo) phases are formed in the course of vanadyl pyrophosphate preparation, was observed (see [10, 15] for details on phase transformations). The change of V205 texture during its mechanochemical treatment in ethanol leads to the synthesis of VPO-E 1 precursor with low specific surface area and unchanged texture. It can be

340 assumed that freshly formed microcrystalls of the precursor, during the fast synthesis, rapidly grow into agglomerates. This sample shows also an increased temperature of amorphization and of crystallization during the formation of vanadyl pyrophosphate and the increased surface P/V ratio. Modification of the conditions of synthesis, consisting in some deceleration of the process, allows the preparation of the sample (VPO-E2) with larger specific surface area. Moreover, it has favourable morphology with high exposure of (001) crystallographic plane. Table 3 shows the catalytic properties of these samples. One can see from the data, that all catalysts synthesized on the basis of the mechanochemically treated V205 show an increased selectivity towards maleic anhydride and higher specific rate of n-butane and npentane oxidation as compared to those obtained in traditional synthesis. The best effect in the improvement of selectivity can be reached by increase of the relative exposure of (001) plane at the VOHPO4.0.5.H20 surface which is known to be transformed into (200) plane of (VO)2PzO7. The low paraffins conversion over VPO-E samples at the given reaction conditions can be directly connected with their low specific surface area. The comparison made between samples VPO-E1 and VPO-E2 shows that the precursor synthesis using VzOs-E needs to be optimized in order to improve the catalytic performance. Nevertheless, the results clearly demonstrate that mechanochemistry is obviously a promising method for the pretreatment of initial reagents in order to synthesize efficient VPO catalysts of paraffins oxidation. Table 3 Properties of VPO catalysts prepared using mechanochemically treated V205 in reactions of paraffins oxidation. Sample

n-Butane oxidation X, % SMA,% W. 104

n-Pentane oxidation X,% SMA,% SPA,% W. 104

VPO-R VPO-D VPO-E1 VPO-E2

73 79 25 63

52 59 17 43

58.5 68.5 60.0 69.0

1.13 1.45 1.60 1.59

21 35 33 30

12 10 4 6

0.92 1.17 1.32 1.29

Note: X - paraffin conversion, SMA, SpA- selectivity to maleic and phthalic anhydride, respectively, W - specific rate of the paraffin oxidation, mol/h m 2. For n-butane T = 425 ~ GHSV = 3000 h -1 and for n-pentane T = 320 ~ GHSV = 1800 h -~ 2.1.2. Mechanochemical modification of the VPO precursor Previously we showed [ 13,16] that the nature of dispersant had a quite remarkable effect on the properties of VPBiO precursor which was subjected to a short-time mechanochemical treatment (up to 10 rain.) and the best result was obtained when ethanol was used. Here we will describe the influence of the time of treatment on the properties of VPO precursor (P/V = 1.07). As can be seen from Table 4, the longer is the mechanochemical treatment (up to 20 min) the higher are both specific surface area of the precursor and relative intensity of its (001) crystallographic plane reflection. The latter observation can be explained by anisotropic plastic deformation of the crystals arising from their layered structure. The observed increase of the surface P/V ratio is in a good agreement with the mechanism of the VOHPO4.0.5H20 phase transformation [ 15].

341 Table 4. Properties of VPO precursor after its mechanochemical treatment in ethanol Sample Time SsA treatm.,

min

mZ/g

XRD

I0.570/ I0.293

XPS Binding energy, eV O 1s V 2p P2p

n-butane oxidation* (P/V)s

SSA1

X

S

VPO 6.0 74/100 531.7 517.5 133.9 1.43 9.4 62 61 VPO10 10 8.8 95/100 531.6 517.4 133.8 1.62 12.3 68 65 VPO20 20 14.2 100/78 531.8 517.4 133.7 1.80 17.2 73 70 VPO30 30 8.0 ** 532.1 517.3 133.7 1.92 8.5 77 74 VPO60 60 6.4 *** 532.2 517.3 133.7 1.84 6.5 70 68 *T=440~ GHSV=3200h 1, S S A 1 - specific surface area after catalysis (m2/g), X - n-butane conversion (%), S- selectivity to maleic anhydride (%); the samples VPO30 and VPO60 did not need activation prior to catalysis, **Amorphization of the sample, very weak peaks at d = 0.328, 0.305 and 0.285 nm ***All reflections ofvanadyl pyrophosphate are present, the most intense one is at d = 0.313 nm Continuation of the mechanochemical treatment leads to amorphization of the sample followed by the formation of vanadyl pyrophosphate phase. It should be however noted that even after 60 min. of the treatment this compound is not well crystallized and the specific surface area of the final catalyst is much smaller than that of a sample obtained by in situ activation of a precursor after 20 min of its mechanochemical treatment. It follows from the results presented in Table 4, that the samples obtained by mechanochemical treatment of the precursor become more active in the reaction of n-butane partial oxidation so that the hydrocarbon conversion and selectivity to maleic aL!hydride increase. The sample converted into vanadyl pyrophosphate by means of the mechanochemical treatment turned out to be more efficient than that activated with the reaction mixture. The most interesting is the sample after 30 min. treatment which is "half-activated" and consists of the amorphous phase. The active component forming directly in the catalytic mixture without long activation procedure gives rise to the most active and selective catalyst for n-butane oxidation.

2.1.3. Mechanochemical promotion of VPBiO precursor Recently [ 13, 16] we have shown the possibility of efficient promotion of VPO precursor with bismuth compounds. The present paper reports new results on promotion of VPBiO precursor with lanthanum compounds (previously a similar catalyst was shown to be active in tetrahydrofuran formation [17]). Table 5 compares the properties of traditionally prepared VPBiLaO sample (by simultaneous introduction of bismuth and lanthanum additives in the course of the synthesis of VPO precursor) with that (VPBiO-La-M) produced by the mechanochemical treatment of VPBiO precursor and lanthanum oxide powders. The treatment in the latter case was carried out for 10 min. in ethanol medium. It can be seen that introduction of lanthanum by both methods leads to an increase of the catalytic activity. At the same time, its introduction in the course of the synthesis of VPBiO precursor causes a decrease of the selectivity to partial oxidation products in both investigated reactions: oxidation of butane and propane. This negative effect on selectivity was also observed when the catalyst was prepared by means of mechanochemistry but to a much lesser degree. As a result, the latter catalysts were more efficient and produced higher yield of the desired product (see Table 5 data). It can be noted that for the traditional sample a higher value of V 2p-electrons binding energy is observed suggesting that the oxidation degree of vanadium can be in this case higher than in the traditional samples.

342 Table 5 Properties of VPO precursor with additives bismuth and lanthanum. Sample*

VPBiO VPBiLaO3 VPBiLaO5 VPBiO-La3M VPBiO-La5M

XPS Binding energy, eV** P/V Bi/V La/V O 1 s V 2p P 2p

Oxidation n-Butane*** Propane**** X SMA Y X SAg Y

531.5 531.8 531.7 531.6 531.5

48 56 61 55 58

517.4 517.9 517.9 517.5 517.5

133.7 1.58 0.12 134.0 1.61 0.17 0.027 133.9 1.88 0.14 0.063 133.8 1.73 0.09 0.023 133.6 1.95 0.08 0.058

69 61 54 68 66

33 34 33 37 38

40 49 60 47 59

32 21 17 30 28

13 10 10 14 16

*Number in the sample name represents the atomic ratio (La/V).100 ** B.E. Bi 4f = 159.9 160.2 eV, La 3d = 836.5-836.7 eV *** T = 420 ~ GHSV = 3000 h "l, X - n-butane conversion (%), SMA - selectivity to maleic anhydride (%), Y - yield of maleic anhydride (mol. %); **** T = 435 ~ GHSV = 1500 h -1, X - propane conversion, %, SAA - selectivity to acrylic acid, %, Y - yield of acrylic acid, mol. %

2.1.4. Mechanochemical modification of vanadyl pyrophosphate Vanadyl pyrophosphate was prepared by heating the VPO precursor in an inert gas flow at 550 ~ for 24 h. Its characteristics are listed in Table 6. One can see from TSM pictures (Figure 1a) that the catalyst prepared in this way is composed of quite large aggregates including crystals of geometrically-regular shape. Mechanochemical treatment disintegrates them to produce much smaller particles of different shape (Figures 1 b,c,d). It is noteworthy, that in the case of the treatment of (VO)2P207, there is no dependence of the morphology on the dispersant nature at variance with the case of the VPO precursor treated similarly [13]. An increase of the intensity of the reflection at d = 0.387 nm corresponding to (200) plane of vanadyl pyrophosphate was observed by XRD to be the only structural change occurring at the treatment in ethanol. As a result, the catalyst after treatment shows an increase of both activity and selectivity in n-butane partial oxidation. The specific rate of the hydrocarbon conversion decreases after the treatment which is believed to be due to non-proportional growth of the number of active centers and the specific surface area. Table 6 The properties of vanadyl pyrophosphate after mechanochemical treatment Treatment

SSA

Solvent Time, min. m2/g Water Ethanol

10 10 10

* T = 440 ~

4.6 9.2 5.4 7.1

XRD

I0.387/ I0.313 82/100 95/100 92/100 100/73

XPS Binding energy, eV O lS V 2p P 2p 531.5 531.6 531.8 .

517.5 517.4 517.5 . .

133.8 133.6 133.8 .

n-Butane oxidation* SMA, W-10 4 %

P/V

X, %

1.22 1.34 1.42

68 76 69 77

GHSV = 1500 h l , W - rate of n-butane oxidation, mol/h.m 2

60 58 62 69

1.44 0.86 1.19 1.08

343

a.

b.

Figure 1. Transmission Scanning Microscope pictures of vanadyl phyrophosphate, a) initial, and after mechanochemical treatment: b) in water, c) in ethanol and d) dry milling. 2.2. Barothermal synthesis 2.2.1. Barothermal synthesis of VPO catalysts Soon after publications of J.Johnson and A.Jacobson [18,19] hydrothermal synthesis has begun to be applied to different vanadium phosphates synthesis. In the present work an attempt has been undertaken to use organothermal (with n-butanol addition) synthesis and that without any solvent. It has been established that in hydrothermal synthesis starting from V205 and H3PO4 it is possible "depending on the synthesis temperature and duration" to obtain VOPO.2H20 and VOPO4.H20. When VO2 is used the mixture of vanadium (IV) and (V) compounds such as 13VOPO4 and VOHPO4.0.5H20 can be formed. Their catalytic performance in n-butane and butene-2 oxidation (better for samples HS-1) is worse than that of the catalysts prepared by other methods.

344 Organothermal synthesis with the use V205 leads to the formation of [3-VOPO4 (OS-1). The latter compound shows low activity in paraffin oxidation, but is a quite efficient catalyst for olefin oxidation (Table 7). When VO2 was used for organothermal synthesis, an unknown compound was obtained at low temperature and/or short time of the synthesis. Continuation of the synthesis led to the formation of VOHPO4.0.5H20 (OS-2). In the synthesis using V205 in solvent-free conditions, 13-VOPO4 was found to be the only product (AS-1) but quite a high temperature and long time were needed for the reaction to be completed. In the case of the use of VO2, a new compound was formed. Table 7 Properties of the barothermally synthesized samples Butene-2 oxidation*** SMA,% W 104

n-Butane oxidation** W 104

Sample*

X, % SMA,% TS-1 TS-2 HS-1 OS-1 OS-2 AS-1

68 63 35 75 28

60 60 20 64 21

X, o~

1.44 1.11 1.02 1.76 0.48

76 78 74 71 70

71 55 69 51 72

3.22 2.86 3.31 3.14 3.06

* T S - 1 , -2 - catalysts prepared following traditional methods for n-butane and butene-2 oxidation, respectively **T = 440 ~ GHSV = 1500 h -l', ***T = 380 ~ GHSV = 3600 h -1', X hydrocarbon conversion, SMA - selectivity to maleic anhydride, W - rate of hydrocarbon oxidation, mol/h.m 2 2.2.2. Barothermal

modification

of VPBiO

precursor

The VPBiO precursor treatment was performed with n-butanol and phosphoric acid vapours as the perspective media for the treatment of VPO [ 12]. The obtained results, sortie of which are listed in Table 8, show that the higher is the temperature and the longer is the treat ment the lower becomes the specific surface area. At the same time, it should be noted that Table 8 An influence of VPBiO precursor barothermal treatment on its properties Sample*

VPBiO VPBiOnbl VPBiOnb2 VPBiOnb3 VPBiOpal VPBiOpa2 VPBiOpa3

Treatment Time T, h ~ -

6 10 6 14 10 6

-

250 250 300 200 250 300

SSA

XRD I0570/ m Z / g I0293 12.5 11.2 10.0 9.3 12.0 10.2 8.7

100/98 100/78 100/76 100/39 100/76 100/49 100/45

XPS Binding energy, eV O ls V 2p P-2p Bi 4f 531.5 531.7 531.6 531.4 531.1 531.4 531.3

517.4 517.5 517.6 517.1 517.1 517.3 517.3

133.6 133.9 134.0 133.5 133.4 133.5 133.4

159.9 160.2 160.1 159.7 159.7 159.8 159.7

*nb, pa- treatment with n-butanol and phosphoric acid, respectively

(P/V)s (Bi/V)s 1.47 1.87 2.08 2.07 2.23 2.21 1.98

0.08 0.12 0.10 0.10 0.10 0.11 0.11

345 such treatment leads to some change of the structure of VPBiO precursor as it can be seen from XRD results. An increase of the relative intensity of the reflection at d = 0.570 nm attributed to crystallographic plane having vanadyl groups is namely observed. Enrichment with phosphorus is found at the surface of the particles after treatment. Similar effect could be expected at the treatment with phosphoric acid but looks unusual in the case of the treatment with alcohol. The catalytic properties of the samples are presented in Table 9. The data indicate that the barothermal treatment favours an increase of the selectivity in paraffins oxidation. Moreover, the treatment in n-butanol also leads to the growth of catalytic activity of the samples. In the case of the treatment with phosphoric acid vapours, the catalytic activity remains almost unchanged, due to effect of water steam as described in [ 12]. Table 9 Catalytic properties VPBiO precursor after barothermal treatment Sample

SSA 1

m2/g VPBiO VPBinbl VPBinb2 VPBinb3 VPBipal VPBipa2 VPBipa3

14.6 12.0 9.8 9.0 11.6 9.7 8.5

n-Butane oxidation* X, % SMA,~ 49 50 57 57 48 50 45

68 72 75 77 73 81 76

n-Pentane oxidation** X,% SMA,% SPhA,% 67 -

32 -

17 -

61 65 65

39 42 36

14 18 21

SSA1 - specific surface area after catalysis * T = 400 ~ GHSV = 2400 h -1 *** T = 420 ~ GHSV = 1500 h -1

Propane oxidation*** X, o~ SAA,o~ 26 28 29 31

48 55 59 58

-

GHSV = 3000 h -1 **T = 350 ~

Concluding, it should be emphasized that mechanochemical and barothermal methods have been shown to be advantageous as alternative technologies for preparation and modification of VPO catalysts for partial oxidation of saturated hydrocarbons. ACKNOWLEDGMENT This study was supported in a part by ISF (Grants UBI000 and UBI200) and in a part by Scientific Research Committee (Poland) Grant No 3T09A 08010. The authors thank Prof. V.G.Iljin, Dr. G.A.Komashko and V.E.Yaremenko for assistance in some experimental work.

REFERENCES 1. 2. 3.

G. Centi (editor), Vanadyl Pyrophosphate Catalysts, Catal. Today, 16 (1993) 1-147. G.J. Hutchings, Appl. Catal., 72 (1991) 1. F. Cavani and F. Trifiro, Preparation of Catalysts VI, Stud. Surf. Sci. Catal., Elsevier, Amsterdam, 91 (1995) 1.

346 4.

V.A. Zazhigalov, V.M. Belousov, G.A. Komashko, A.I. Pyatnitskaya, Yu.N. Merkureva, A.L. Poznyakevich, J. Stoch and J. Haber, Proc. 9th Int. Congr. Catal., Chem. Inst. of Canada, Ottawa, 4 (1988) 1546. 5. P.F. Miquel and J.L. Katz, Preparation of Catalysts VI, Stud. Surf. Sci. Catal., Elsevier, Amsterdam, 91 (1995) 207. 6. P.M. Michalakos, H.E. Bellis, P. Brusky, H.H. Kung, H.Q. Li, W.R. Moser, W. Partenheimer and L.C. Satek, Ind. Eng. Chem. Res., 34 (1995) 1994. 7. V.V. Guliants, J.B. Benziger and S. Sundaresan, J.Catal., 156 (1995) 298. 8. P.F. Miquel, E. Bordes and J.L. Katz, J.Solid State Chem, 124 (1996) 95. 9. L. Bogutskaya, V. Zazhigalov, M. Misono and T. Okuhara, Japan-FSU Catal. Seminar'94, Catal. Sci and Techn. 21 Century Life, Tsukuba, Japan, (1994) 202. 10. V.A. Zazhigalov, J. Haber, J. Stoch, L.V. Bogutskaya and I.V. Bacherikova, Appl. Catal. A, 135 (1996) 155. 11. V.A. Zazhigalov, J. Haber, J. Stoch, A.I. Pyatnitskaya, G.A. Komashko and V.M. Belousov, Appl. Catal. A, 96 (1993) 135. 12. V.A. Zazhigalov, I.V. Bacherikova, V.E. Yaremenko, I.M. Astrelin and J. Stoch, Teoret. Eksperim. Chem., 31 (1995) 206. 13. V.A. Zazhigalov, J. Haber, J. Stoch, L.V. Bogutskaya and I.V. Bacherikova, 1 l th Int. Congr. on Catal. - 40th Anniversary, Stud. Surf. Sci. Catal., Elsevier, Amsterdam, 101B (1996) 1039. 14. V.A. Zazhigalov, J. Haber, J. Stoch, A.I. Kharlamov, L.V. Bogutskaya, I.V. Bacherikova and A. Kowal, Solid State Ionics (in press). 15. C.C. Torardi, Z.G. Li and H.S. Horowitz, J.Solid State Chem., 119 (1995) 349. 16. J. Haber, V.A. Zazhigalov, J. Stoch, L.V. Bogutskaya and I.V. Bacherikova, Catal. Today (in press). 17. V.A. Zazhigalov, J. Haber, J. Stoch, G.A. Komashko, Ai. Pyatnitskaya and I.V. Bacherikova, New Develop. Select. Oxid. II, Stud. Surf. Sci. Catal., Elsevier, Amsterdam, 82 (1994) 265. 18. J.W. Johnson, A. Jacobson, J.F. Brody and S.M. Rich, Inorg. Chem., 21 (1982) 3820. J.W. Johnson and A. Jacobson, Angew. Chem., 95 (1983) 442. D

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 1997 Elsevier Science B.V.

Active species and working mechanism of silica supported MoO3 and catalysts in the selective oxidation of light alkanes

347

V205

A. Parmaliana% F. Arena% F. Frusteri b, G. Martra c, S. Coluccia r and V. Sokolovskii d aDipartimento di Chimica Industriale, Universitb, di Messina, Salita Sperone 29, 98166 S. Agata (Messina), Italy blstituto CNR-TAE, Salita S. Lucia 39, 98126 S. Lucia (Messina), Italy CDipartimento di Chimica Inorganica, Chimica Fisica e Chimica dei Materiali, Universit/l di Torino, Via P. Giuria 7, 10125 Torino, Italy aDepartment of Chemistry, University ofWitwatersrand, Johannesburgh, P.O. Box 106, South Africa The catalytic performance of a series of silica supported (2-7 wt%) MoO3 and (2-20 wt%) V205 catalysts has been comparatively evaluated in both the partial oxidation of methane to formaldehyde (MPO) and the oxidative dehydrogenation of propane to propylene (POD) in the range 500-800~ and 500-525 ~ respectively. V205 acts as a promoter of the reactivity of the SiO2 for both MPO and POD, while MoO3 plays a generally negative effect on the catalytic functionality of the SiO2 surface acting however as a promoter in the MPO at T>650~ A direct relationship between the density of reduced sites of low-medium loaded silica based oxide catalysts and reaction rate in both MPO and POD strongly suggests the occurrence of a concerted reaction mechanism involving the activation of gas-phase 02 on the reduced sites of the catalyst surface. The nature of the active sites has been defined on the basis of a complete characterization of the surface and redox features of MoO3/SiOz and V205/SIO2catalysts. 1. INTRODUCTION The disclosure of the mechanism of a catalytic reaction leads to the identification of the active sites being then the basis for the design of more effective catalytic systems. Generally, a mechanistic model is adequate for describing the pathway of a class of reactions on a class of catalysts, however this rule is not completely valid for the partial oxidation of light alkanes on bulk and/or supported oxide catalysts [ 1]. Since the excellent performance of supported MoO3 and V205 based catalysts in the selective oxidation of light (C~-C3) alkanes, during the last decade a considerable research effort has been directed to ascertain the working mechanism of such oxide systems as well as the nature of the active sites and the origin of the oxygen involved in the formation of reaction products. However, no definitive conclusions have been still provided about these issues and therefore a much deeper investigation of the reaction dynamic mechanism and the correlation with the nature of the surface is necessary [2]. Several factors, such as the nature of the support, the oxide loading and the reaction conditions control the formation, the coordination and the stabilization as well as the catalytic action of the

348 various surface oxide species. On this account, we have evaluated the catalytic behavior of silica supported MoO3 and V205 systems in both the selective oxidation of methane to formaldehyde (MPO) [3,4] and the oxidative dehydrogenation of propane to propylene (POD) [5] disclosing that V205 acts as a promoter of the reactivity of the SiO2 surface while the action of MoO3 strictly depends upon the kind of the silica support [3-5]. The aim of this paper is to provide a correlation between the catalytic pattern of differently loaded silica supported MoO3 and V205 catalysts in MPO and POD reactions with their surface and redox features in order to highlight the nature of the active surface species in the selective oxidation of light alkanes. 2. EXPERIMENTAL 2.1. Catalysts

Differently loaded (2-7 wt%) MoO3/SiO2 and (2-50 wt%) V205/SIO2 catalysts were prepared by incipient wetness impregnation of a "precipitated" silica (PS) support (Si 4-5P Grade, Akzo Product; S.A.BET, 400 m2.g~) according to the procedure described elsewhere [3]. The list of the catalysts along with their composition and BET surface area values are reported in Table 1. Table 1 List of catalysts Code Chemical composition

(wt %) PS VPS 2 VPS 5 VPS 10 VPS 20 VPS 50 V205 MPS 2 NIPS 4 NIPS 7 MoO3

SiO2 2.0% V205/SIO2 5.3% V205/5iO2 10.1% V205/5i02 20.8% V2OJSi02 50.8% V2OJSi02 V205 2.0% MoO3/SiO2 4% MoOa/SiO2 7% MoO3/SiO2 MoO3

S.A.BET

(m~.~-') 400 260 230 200 190 160 5 300 190 75 2

2.2. Catalytic measurements

Catalytic measurements in MPO were performed by Temperature Programmed Reaction (TPR) tests [4] using a conventional flow apparatus and a linear quartz microreactor connected on line with a Thermolab (Fisons Instruments) Quadrupole Mass Spectrometer (QMS) for continuous scanning of the reaction stream. TPR tests were run in the T range 400-800~ by using 0.05 g of catalyst, a heating rate (13) of 10~ "1 and a reaction mixture He/CH4/O2 in the molar ratio 7:2:1 flowing at 50 STP cma.min-1. Catalytic measurements in POD have been performed at atmospheric pressure in the range 500-525~ using 0.25 g of catalyst sample diluted with same sized SiC (1/5, l/vol) and a reaction mixture in the molar composition CaH8:O2:N2:He=2:l:l:6 flowing at the rate of 100 STP cmS-min'l[5]. All the tests have been carried out at GHSV of 1,700 h"1 (STP m3C3H89m'aeat-h'l).

349

2.3. Catalyst characterization Temperature Programmed Reduction (H2-TPR) measurements were performed using a linear quartz gradientless microreactor and a 6% H2/Ar mixture flowing at 60 STP cm3 rain~ according to the procedure described elsewhere [6]. High Temperature Oxygen Chemisorption (HTOC) measurements were performed in a pulse mode according to the procedure described in detail elsewhere [6]. Reaction Temperature Oxygen Chemisorption (RTOC) measurements in the range 500650~ were performed in a pulse mode after treatment of the samples in the C3HffO2/He or CH4/Oz/He reaction mixture. 02 pulses (Vp,l,o, 4-10.8 mol 02) were injected into the carrier gas until saturation of the sample was attained, the density of reduced sites (p, 1016 sites.gc,t~) being calculated by assuming a chemisorption stoichiometry 02:reduced site = 1:2 [3,5]. Diffuse Reflectance UV-Vis DR UV-Vis spectra of differently loaded V2OflSiO2 samples, calcined in situ in 02 at 600~ were obtained by a Perkin Elmer Lambda 19 spectrophotometer, equipped with an integrating sphere. 3. RESULTS and DISCUSSION

3.1. Catalytic activity Methane partial oxidation (MPO). The catalytic activity of differently loaded MPS and VPS samples in the range 500-800~ expressed in terms of normalized specific surface activity, NSSA (NSSA=SS&/SSAps, where SSAi and SSAps are the specific surface activity of the catalyst i and bare PS support, respectively), is compared in Figure 1A. 1.0

100

A

9 t

,-...

09 z o

o

0.0

I O I

co o d3

-0.5

0 0

t

80

0.5

I

0

60

q

*

d 9

40

1 9

20

t

%

o

-1.0

s60

660'

' '760'

T e m p e r a t u r e (*C)

" "860

P ~ 6-~

" ~6-~

" ~~-,"

' ~'6 o" i6,"

'

CH 4 conversion (%)

Figure 1.Normalized Specific Surface Activity (NSSA, sSmcat/SSmps; SSA=molcH4 co,v..m-2.s-1) vs. T (A) and HCHO selectivity (SHcHO) VS. CI-I4 conversion (B) in the range 500-800~ for differently loaded MPS and VPS catalysts. Legend: PS (*); MPS 2 (A); MPS 4 (r'l); MPS 7 (0); VPS 2 (&); VPS 5 (1); VPS 10 (e) and VPS 20 (~). At any loading, V 2 0 5 promotes the reactivity of the PS at any T (log NSSA>0) according to the following reactivity scale: VPS 5 > VPS 10 > VPS 2 = VPS 20 > PS. By contrast, MoO3 exerts a negative effect on the SSA of the PS carrier (log NSSA700~ The increasing trend of the NSSA with T results more enhanced for the MPS 7 sample (Fig. 1A).

350 The HCHO selectivity as a function of CH4 conversion in the range 500-800~ for MPS and VPS catalysts is shown in Figure lB. For MPS 2 and MPS 4 catalysts, the trend of HCHO selectivity with CH4 conversion (Fig. 1B) is similar to that of the unloaded PS, whereas for the MPS 7 sample a significant improvement in HCHO selectivity at conversion levels lower than 3% is observed. By contrast, at the same level of CH4 conversion, a progressive decline in HCHO selectivity with V205 loading occurs on VPS catalysts. Medium loaded MPS 7 and highly loaded VPS 20 samples are the most and the least selective systems respectively. Oxidative dehydrogenation of propane (POD). The predominant products in the POD over silica based oxide catalysts in the range 500-525~ were propylene and carbon oxides. Ethylene and acetaldehyde along with a considerable amount of acrolein and traces of propionaldehdye are formed on PS [5]. The addition of MoO3 and V205 to the SiO2 support implies a higher selectivity to propylene and correspondingly a lower production of COx, a slight cracking activity and a significant decrease in the amount of oxygenates. Table 2 shows a detailed comparison of the activity of bare and differently loaded MPS and VPS catalysts in terms of propane conversion, propylene selectivity, reaction rate and C3I-I6 productivity values. The catalytic functionality of the SiO2 surface is strongly promoted by the addition of V205, while it is depressed by MOO3, the extent of this effect rising with the MoO3 loading [9]. Table 2 Activity of bare and promoted silica catalysts in the oxidative dehydro[genation of propane Catalyst T C3Hs conv. C3I-I6 s e l . Reaction rate STYc3H6 (%) (].tmol.gc.tl.s "1) (g.kgc.t'l.h "1) (~ (%) PS 5OO 0.9 37 0.49 27.5 525 1.9 37 1.04 57.7 MPS 2 500 0.8 62 0.44 41.4 525 1.3 58 0.71 62.5 MPS 4 5OO 0.7 69 0.39 40.8 525 1.4 59 0.75 67.1 MPS 7 500 0.2 80 0.09 6.5 0.4 525 85 0.18 11.6 VPS 5 500 7.8 60 4.25 386.4 525 13.3 55 7.25 602.8 VPS 10 5O0 9.8 51 5.34 413.0 525 177 41 9.65 600.0 VPS 20 500 7.8 27 4.25 174.0

3.2. Redox properties of MoO3/SiO2 and V2Os/SiO2 catalysts under reaction conditions In previous papers we disclosed a direct relationship between catalytic activity in MPO of low-medium loaded silica based oxide catalysts and oxygen uptake under steady state reaction conditions pointing out that such property governs the catalytic behavior of MPO catalysts [3,7,8]. Then, in order to find out whether such a relationship is also valid for POD reaction, the density of reduced sites (p) of various silica based MoO3 and V205 catalysts in both MPO and POD has been evaluated and related with their catalytic activity. The direct relationships between reaction rate in MPO and POD and p, shown in Figure 2A, well account for the opposite effects exerted by MoO3 and V205 on the activity of the bare PS carrier in both MPO [3,4,6-8] and POD [5]. Indeed, V205 effectively promotes the activity of bare PS, and allows

351 for the stabilization of a higher density of reduced sites owing to its easier "reducibility" [6] whereas MOO3, being essentially unreducible under reaction conditions [6], depresses the activity of the PS carrier because of a negative physical effect linked to the partial coverage of the silica surface own active sites [3,6,8]. However, though the above relationships (Fig. 2A) explain the origin of the catalytic action of low-medium loaded (___5wt%) MPS and VPS systems, they do not account for the activity of highly loaded (>__10wt%) VPS catalysts in MPO. 15

III

~ 9 12 O .m 0~

15

A

t

"~ 12

I

9

t

o

.m

MPO

~

,

E

:

"~

0

,

|

2

|

4

.

|

6

,

|

;A

MPO

',

\

l n,

,, j,d, 'f'f

.

\

!

~ lS~...... "'POD |

,,.,..,,.

!

o

i"

~

l

9

E

J

l t

n,'

$

.

,

.

.

8 10 12

9 ( 1016 siteS'gcat "1)

3

!

0 " qd.o

' l'0b.o

P ( 1016 siteS'goat "1)

Figure 2. A: Relationships between reaction rate and density of reduced sites (9) in MPO (600 and 650~ and POD (500 and 525~ for low (___5wt%) loaded MPS and VPS catalysts. B: Relationship between reaction rate and p for differently (0-20 wt%) loaded VPS catalysts in MPO at 650~ Legend: PS (*); MPS 2 (A); MPS 4(!-I); VPS 2 (A); VPS 5 (11); VPS 10 (o) and v e s 2o (~).

Indeed, the activity of VPS catalysts reaches the maximum for VPS 5 sample, while the density of reduced sites increases steadily up to a loading of 20% (VPS 20) resulting in the peculiar volcano-shaped relationship between P and reaction rate shown in Figure 2B. This singular behavior is to be related with the capability of highly loaded V2Os/SiOz systems to "loose" constitutional oxygen [6] under reaction conditions which, however, is ineffective towards the selective oxidation of light alkanes [5,8,9]. 3.3. Surface structures of MoO3/SiO2 and V2Os/SiO2 catalysts

The H2-TPR profiles of differently loaded MPS and VPS catalysts in the range 200-1200~ compared with those of bulk MoO3 and V205 respectively, are shown in Figure 3A and B, while the values of oxygen uptake (HTOC) and oxide dispersion (O/Me) are listed in Table 3. A wide and convoluted band of H2 consumption starting (To, ,~d) at T ranging between 435 (MPS 7) and 486~ (MPS 2) and spanning a T range of 500-600~ accounts for the stoichiometric reduction of MoO3 to Mo Oin all MPS catalysts. The spectrum of the low loaded MPS 2 catalyst features a low rate of H2 consumption up to ca. 800~ thereafter the reduction rate increases sharply giving rise to a main reduction peak with maximum at 934~ (Fig. 3A, a). The increase in the MoO3 loading from 2 to 4 wt% (MPS 4) causes a marked shift of To,red to lower T (436~ and a concomitant enhancement of the reduction kinetics at lower T ("I-

/

"-o

~b ,.. 10

Q2 ~--i

0

100

*

200

Ai

300

500

600

700

800

~(rrirt) Figure 4 : Relative hydrogenation activity at 423 K under isoprene + helium flow versus time of CeNio.2 reduced at 573 K under H2.

Figure 5 9Hydrogen H* content as a function of treatment temperature under H2 of (o) CeNi0.2 and (11) CeNi0.5.

Besides, for each treatment temperature the relative hydrogenation activity can be reported as a function of the relative hydrogen H* content of the solid, and no proportionality is obtained (Figure 6). In fact, the H* consumption kinetic by the hydrocarbon is a complex phenomenon, in particular, the diffusion of the H* species from the >to the ~<surface >) of the solid has to be taken into account [24, 25]. For treatment temperatures higher than about 447 K, anionic vacancies are created in the CeNix catalyst by the loss of H20 (OH groups), as it is shown by thermogravimetry (Figure 3 1 , 7). After a treatment under H2 at 473 K the solid CeNi0.~ contains 17.5 10- mol.g of H

388 species, and this hydrogen storage has been correlated to the creation of anionic vacancies in the solid [22].

|

0.8 A

0.6

E E

CI

>- 0.4 -r

-8

0.2 .'

i

i

0.5

1

I.r rel.

Figure 6 : Relative hydrogenation activity at 423 K under isoprene + helium flow versus the hydrogen H* species concentration of CeNi0.5 reduced at 423 K under H2.

-12

I

273

I

473

"

TEMPERATURE (1~

673

Figure 7 9 Thermal treatment under H2 of CeNi0.5 (a) and CeNio.2 (b) followed by thermogravimetry.

A great analogy exists between the results presented in this study and those obtained in the laboratory on copper based oxides [24, 25] and molybdenum based sulfides [26, 27] which have been found to be hydrogen reservoirs. As a matter of fact different studies published on these catalysts deal with the relations existing between the active site unsaturation degree and the dienes hydrogenation activity and selectivity, as well as the existence of hydrogen H* reservoirs. It has been shown that the first hydrogen species introduced in the diene during the hydrogenation reaction is a hydride species coming from the solid [28]. Indeed, it has been proposed that under helium+alkadiene, the titrated H* species are for one half H" species located in anionic vacancies and the second halfH + species (coming from OH groups) [24-27]. These species are inserted in the solid during the activation under H2 : O 2" M n+ D + H2 -9 OH-M n§ H" (with D 9anionic vacancy). 3.3. Oxidative dehydrogenation of propane on hydrogen reservoirs. The oxidative dehydrogenation of C3H8 has been performed on previously, at different temperatures under H2, in-situ reduced catalysts. On Figures 8 and 9, propene yield and selectivity are reported as a function of propane conversion, respectively on pretreated under

389

H2 at 473 K CeNi0.2 and on pretreated under H2 at 433 K and 473 K CeNi0.5. The two treatment temperatures (433 and 473 K) lead respectively to the creation of a slight and a large hydrogen reservoir as shown previously in Figure 5. The ODH of propane obtained on the H2 pretreated CeNi0.5 at 433 K is equivalent to that obtained on the untreated solid. Besides, at 648 K, an optimum yield of about 6.9 % can be obtained on the reduced CeNi0.5 at 473 K, while at the same temperature, a maximum yield of 5.4 % is obtained on the untreated solid (Figure 8). As shown on Figure 9, a similar effect is observed on CeNi0.2. Clearly, the presence of the large hydrogen reservoir (i.e. treatment at 473 K under H2) lead to a beneficial effect on the propene yield.

100 1

6

100

8O

5

80

A

......~-.~ ......2 ~ i i 2 ~ : : 0

9

0

I

20

9

A

4

=~

3

~4o

9

5

4~ A

.

9 "'o.~

2

2O

"..~................~....

20

6

9

I

40

g=~::::::=..e ,

CONVERSION (%)

I

60

1

1 0

Figure 8 : propene yield (I, o) and selectivity (D, o) as a function of propane conversion on CeNi0.2 not treated (I, [3) and previously treated under H2 at 473 K (., o).

0

I

0

I

"

I

20 40 CONVERSION

9

I% )

I

60

Figure 9 : propene yield (m, o, 0 ) and selectivity ([2, o, 0) as a function of propane conversion on CeNi0.5 not treated (I, [2) and previously treated under H2 at 433 K ( 0 , 0) and at 473 K (., o).

The existence of the large hydrogen reservoir can be correlated to a partially reduced catalyst, as it has been published previously [22]. So, the results obtained here confirm that for the ODH of propane, the catalyst works in a partially reduced state and that a redox mechanism is involved.

4. DISCUSSION. The catalysts CeNix possess this character of being hydrogenation and oxidation catalysts in agreement with their redox properties. Knowing that in previous studies it has been shown by work function measurements that propane reacts with O 2 species located at the surface of various oxide catalysts during the activation of the alkane [23], and taking into

390 account the results obtained in the present study, a hydrocarbon activation model can be proposed. By analogy to the heterolytic dissociation of hydrogen, a heterolytic dissociation of the alkane involving the abstraction of a H" species from the hydrocarbon can be envisaged on a low coordination site involving an anionic vacancy. This active site, 1, indicating that the active phase is =SbVO4 with a surface enriched with Sb. Characterisation of syntheses in the A1-Sb-V-O system allowed the identification of a trirutile-like phase with the composition All_xSbVxO4 (0 < x < 0.5). The synthesis of this phase, which is active and selective to acrylonitrile, requires excess of aluminium. Over a fresh preparation with A1, Sb, V and W the activity and the selectivity to acrylonitrile increased considerably with time-on-stream. This behaviour shows that the active rutile phase is formed in situ, and EDX analyses gave the average composition A10.1Sb0.8V0.7W0.404. The highest yield to acrylonitrile that was observed for the three systems was 37 % and was obtained over the A1Sb-V-W-oxide. 1. I N T R O D U C T I O N In patents [ 1] it has been indicated that for propane ammoxidation =SbVO4 is an active phase of rutile type in catalysts with Sb, V and W on an alumina-rich support. It has been found that when the amount of antimony in the synthesis is in excess of the stoichiometric ratio, i.e. with Sb:V > 1, a catalyst is produced which is more selective than the pure =SbVO4 phase for acrylonitrile formation. An Sb:V ratio in the range 2-10 was reported as being optimal [1-3]. Over Sb-V-O catalysts with excess Sb and without A1 and W the yield achievable to acrylonitrile is 10-12 %, to be compared with 3-4 % for a corresponding sample with an Sb:V ratio equal to one. It was shown by adding alumina to SbsVOx that the yield to acrylonitrile was further increased up to about 25 %, i.e. by a factor more than two. Another key element in the catalysts of rutile type is tungsten. A composition Sb5VWOx supported on A1203 with 20 wt.% of SiO2 was reported to give a yield to acrylonitrile of 40 % at almost 70 % propane conversion [ 1]. Considering the yield to acrylonitrile, the data show a remarkable step-wise improvement when progressively adding excess Sb, A1 and W to --SbVO4. In the present paper the formation and structure of the active phases in the three systems Sb-V-O, A1-Sb-V-O and A1-Sb-V-W-O will be described and compared considering various characterisation results. It will be shown that there are significant differences between the systems concerning the composition of the active rutile phase, the stage at which the active structure is formed, and catalytic performance.

414 Moreover, it will be demonstrated that data can be rationalised in terms of the site isolation principle, which has been found applicable in other catalyst systems with antimony [4].

2. EXPERIMENTAL 2.1. Preparation of samples =SbVO4 (2.0 m2/g) was prepared by heating an equimolar mixture of V205 (Riedel-deHahn, p.a.) and Sb203 (Merck, p.a.) in air at 800 ~ for 18 hours. The same method was used to prepare a sample with excess antimony oxide (Sb:V = 2), having a specific surface area of 1.4 mZ/g. Another Sb-V-O sample with Sb:V = 2:1 was prepared from a slurry of Sb203 in water solution with NH4VO3 (Merck, p.a.), which was heated under reflux before drying and final calcination at 610 ~ [3]. The specific surface area was 3.6 m2/g. For investigation of phase formation in the A1-Sb-V-O system, weighed amounts of AI(OH)3 (Riedel-de-Ha~n, p.a.), Sb203 and V205 were mixed and ground together and finally heated at 680 ~ for 4 days with one intermittent grinding. A sample, which was used as catalyst, and with the atomic ratio AI:Sb:V = 21:5:1 was prepared from the same reactants following a slurry method which has been described in detail elsewhere [5]. The final calcination was performed at 610 oC and the resulting specific surface area of the sample was 157 m2/g. A catalyst with the nominal composition AI:Sb:V:W = 21:5:1:1 was prepared according to the slurry method [1,5] starting from AI(OH)3, Sb203, V205 and (NH4)6W12(OH)2038 (Strem Chemicals, p.a.). The specific surface area of the sample was 121 m2/g after the final calcination at 610 ~

2.2. Activity measurements The activity measurements were performed at 480 ~ using a plug-flow reactor made from glass. Dilution of the catalyst with quartz grains was necessary to have isothermal conditions. The composition of the inlet flow to the reactor was propane 14.3 vol.%, ammonia 14.3 vol.%, oxygen 28.6 vol.%, water vapour 7.1 vol.%, and nitrogen 35.7 vol.%, corresponding to the stoichiometric ratio between propane, ammonia and oxygen for acrylonitrile formation. Propane and the products CO, CO2, propylene, acrylonitrile and acetonitrile were analysed on a GC equipped with a Porapak Q column, a sample valve, a methanisation column for conversion of the carbon oxides to methane, and an FID detector. Analyses of the conversion of ammonia and the formation of HCN were performed using titrimetric methods [6]. Analysis of conversions and selectivities were performed with time-on-stream after heating the catalyst in air to the reaction temperature and subsequently switching to the reactant composition. The influence of propane conversion on the product distribution was studied varying the amount of catalyst at constant flow rate.

2.3. Characterisations Specific surface areas were measured with a Micromeritics Flowsorb 2300 instrument, using adsorption of N2 at liquid N2 temperature. The samples were degassed at 350 ~ For X-ray powder diffraction (XRD), the samples were crushed and mounted on adhesive tape. Films were recorded using a Guinier-H~igg focusing camera with CuKo~I radiation and with Si as internal standard. Energy dispersive X-ray microanalysis (EDX) was carried out in a transmission electron microscope JEM-2000FX fitted with a Link AN10000 analysis system. The phases were first identified by electron diffraction and thin edges were then analysed using a beam approximately 500/k in diameter and an acceleration voltage of 200 kV. High resolution transmission electron microscopy was performed inoa JEM-4000EX instrument operated at 400 kV and possessing a structural resolution of 1.6 A. Samples were lightly ground in methanol and the dispersion was then transferred to copper grids covered with a holey carbon film. In the microscope, thin crystals positioned over the holes in the carbon film were examined by diffraction and imaging techniques.

415 3. R E S U L T S

AND DISCUSSION

3.1. T h e S b - V - O system Stoichiometric SbVO4 does not form when heating an equimolar mixture of V205 and Sb203 under various conditions, but two [7] or three [8] compositional series with SbVO4 as an end member were used to characterise the rutile-type phases that were produced. M6ssbauer data indicated Sb 5+ and consequently mixed V3+/V 4+ in this rutile [7]. By heating equimolar mixtures of V205 and Sb203 at 800 ~ in flowing gas with various O2/N2 ratios, we have recently demonstrated that the data can be interpreted in terms of a single cation deficient series, namely Sb0.9V0.9+xD0.2-xO4 (0 < x < 0.2), where D denotes cation vacancies [9]. In Fig. 1 are the electron diffraction pattems of selected samples from the series. The pattern of the oxidised end member Sb(V)o.9V(IV)o.sV(III)o.IDo.204 (Fig. lc) shows beside the basic rutile lattice, satellite diffraction spots which are due to ordering of the cation vacancies. An approximate supercell 2~/2a,2,4T2b,4c was observed for a partial oxygen pressure equal to or higher than that for air. For the reduced end member Sb(V)o.9V(IV)o.2V(III)o.9Do.oO4 the diffraction pattern in Fig. 1d shows a different rutile superstructure (~/2a, a/2b,2c), resulting from cation ordering [10]. At intermediate pressures of oxygen (0.1 < P(O2) < 0.21 atm) only diffraction spots from basic rutile are observed (Fig. l a). When the synthesis was carried out with excess of vanadia, we observed [9] formation of a solid solution series Sb0.9VI.IO4 - VO2, i.e. Sb0.9-yV1 l+yO4 with 0 < y < 0.7. The V-rich end composition, which is approaching V(IV)O2, becomes Sb(V)0.2V(IV)I.6V(III)0.204. The diffraction pattern of the V-rich structure showed a basic rutile lattice with strong diffuse intensity between Bragg spots. Investigation of the Sb-V-oxide by XRD showed only reflections from the basic rutile cell.

Figure 1. Electron diffraction patterns of rutile structures recorded with the electron beam parallel to the [100] direction. Superlattice reflections are indicated by arrows. (a): intermediate --SbVO4; (b): (A1,V)SbO4; (c): oxidised =SbVO4; and (d): reduced =SbVO4.

416 Figure 2 shows catalytic data for =SbVO4 and a sample with the atomic ratio Sb:V = 2:1, which both were prepared by solid state synthesis. XRD showed the latter sample to consist of =SbVO4 and Sb204. The plots for =SbVO4 of the propane conversion and the selectivities to propylene, acrylonitrile and C 1-2 degradation products against time-on-stream show no variation with time. At =5 % of propane conversion the selectivities to propylene and acrylonitrile are 54 % and 9 %, respectively. The sample prepared with Sb:V = 2:1 shows at the same conversion level a modest decrease in propane conversion with time-on-stream. Simultaneously the selectivity to acrylonitrile increases from 22 % up to 26 %, while the selectivities to propylene and degradation products correspondingly decrease. The increase in the selectivity to acrylonitrile can possibly be associated with our previously reported finding that migration of antimony from Sb204 to the surface of =SbVO4 occurs during the catalytic reaction [3]. We have, moreover, reported that samples with excess of antimony are slightly less reduced than the pure =SbVO4 after use in the ammoxidation [3,11]. However, these findings cannot fully explain the differences between the data for =SbVO4 and the Sb:V = 2:1 sample in Fig. 2, since in fact the data show that the sample with excess antimony is already initially much more selective than the pure =SbVO4 to acrylonitrile formation. In contrast to previous results [ 12], the XRD data in Table 1 for the samples show no difference in the dimension of the unit cell of the rutile, which for both agrees with the composition Sb0.9Vo.904. 70

70

"~

60

v

60

"= o

50

~= o

50

m

40

o~

40

"o

tO

~ t-

oO

9

9

~-

e.-

30

30 " k'~L--~__ ~ A A A .

o m

'-

20

ID tO

10

..

..

--

==

o

=.

20

10 o----o

0 0

i 1

i 2

i 3

i 4

i 5

T i m e - o n - s t r e a m (hours)

i 6

0

7

0

I

i

i

I

I

I

I

1

2

3

4

5

6

7

Time-on-stream

(hours)

8

Figure 2. Propane conversion (O) and the selectivity to propylene (@), acrylonitrile (11) and degradation products (A) vs. time-on-stream over (a): Sbo.9Vo.904; and (b): solid state preparation with the atomic ratio Sb:V = 2:1. Table 1 Unit cell parameters for the =SbVO4 in solid state preparations with two nominal Sb:V ratios c/a Sb:V atomic ratio a (/k) c (/k) 1:1" 1:1, fresh 1:1, used 2:1, fresh 2:1, used

4.625(4) 4.6228(5) 4.6247(4) 4.6228(8) 4.6251(12)

3.040(2) 3.0385(4) 3.0397(4) 3.0387(6) 3.0369(9)

0.657 0.657 0.657 0.657 0.657

*The average unit cell as determined for one sample using four different methods; from ref. [ 13].

417 Thus, the apparent effect of excess antimony in Sb-V-O preparations to enhance the selectivity to acrylonitrile cannot be explained by a change of the bulk composition of the =SbVO4 phase. To gain further insight into this matter a mechanical mixture (Sb:V - 4:1) with pure --SbVO4 and o~-Sb204 was heated for 8 hours at 800 ~ The fractions of particle size were 0.150 - 0.425 mm and 0.100 - 0.150 mm for =SbVO4 and ~-Sb204, respectively, allowing separation of the two phases by sieving after the co-calcination. Activity data are presented in Figure 3 for the mechanical mixture after the heat treatment and for untreated =SbVO4, not calcined with o~-Sb204. We have previously reported that a simple mechanical mixture of the two phases, which was not submitted to the co-calcination step, does not show improved catalytic performance as compared with the =SbVO4 phase [ 11]. The present data in Fig. 3 for the co-calcined mixture, on the contrary, show considerable improvement. The selectivity to acrylonitrile at --24 % propane conversion is only =2 % over =SbVO4, while it is substantially higher (42 %) for the calcined mixture. Consequently, calcination of -SbVO4 with (~-Sb204 has produced a catalyst which gives less degradation and is more selective to C3 products. Data obtained for each phase from the mixture, after separation by sieving, are as well included in Fig. 3. These data show that the co-calcined =SbVO4 phase alone is responsible for the high selectivity to acrylonitrile, though it is less active than the untreated =SbVO4. Although it is known that both o~-Sb204 and ~-Sb204 have low activity and are not selective to acrylonitrile formation [5], the co-calcined Sb204 phase has become active and selective. XRD data indicate that the reason can be the o~-Sb204 during the co-calcination has transformed partly into Vdoped Sb204 [14], which has been reported to be selective to acrylonitrile formation [5]. Comparison of the data for the co-calcined mixture with those measured for each of its constituents shows that the propane conversions over the latter samples are almost addable, and the data also indicate that the propylene formed over the V-doped Sb204 is converted to acrylonitrile over the =SbVO4. When the co-calcined phases after separation are used in consecutive beds the corresponding data in Fig. 3 show almost no difference as compared with the phases are being used in form of a mixture. This fact indeed shows that the activated =SbVO4 is a major phase, determining the catalytic performance in Sb-V-O catalysts with excess antimony oxide. ~. 100 v

~>9

80

-~

60

"0 c

~

['9 [~ 9 I-]

sel. propylene sel. acrylonitrile propane conversion N H 3 conversion

40

c

o

~ c 0

o

20

0

I

a

i

b

[

c

d

i

e

f

Figure 3. Conversion of propane and ammonia, and the selectivity to propylene and acrylonitrile in propane ammoxidation over (a): =SbVO4; (b): mixture obtained by co-calcination of particles of =SbVO4 with c~-Sb204; (c): =SbVO4 obtained by co-calcination with ff-Sb204 and subsequent separation by sieving; (d): V-doped Sb204 obtained by co-calcination of =SbVO4 with otSb204 and subsequent separation by sieving; (e): co-calcined samples in two serial beds with =SbVO4 in the first and V-doped Sb204 in the second bed; and (f): co-calcined samples in two serial beds with V-doped Sb204 in the first and -SbVO4 in the second bed. The co-calcination was performed at 800 ~ for 8 hours. Amounts: =SbVO4 1.0 g, and Sb204 1.9 g.

418 The data described conclusively demonstrate that the active phase in the antimony-rich side of the Sb-V-O system is formed during the calcination step in the catalyst synthesis. The active phase is -SbVO4, possibly with a surface which is enriched with antimony. In the present work the enrichment can be due to extraction of vanadium by the Sb204 phase, which is in contact with the =SbVO4 surface. However, at lower calcination temperatures where no V-doped Sb204 can form, the enrichment with antimony possibly is due to migration of antimony species from o;-Sb204 to the surface of =SbVO4. For such inference we have previously presented support [3,15]. The present results can clearly be rationalised in terms of isolation of the vanadium centres at the surface is needed to a suitable degree to have a surface which is selective to acrylonitrile, agreeing with the idea on site isolation that originally was presented by Callahan and Grasselli [16]. For a Sb-V-O catalyst prepared by the slurry method [3] and with the atomic ratio Sb:V = 2:1, the selectivity to propylene as a function of the propane conversion is plotted in Fig. 4a, and the corresponding plots of the selectivity and the yield to acrylonitrile are in Fig. 4b. The data are representative for Sb-V-O preparations. The highest yield to acrylonitrile obtained per single pass is 11% and is achieved at about 40 % of propane conversion. The corresponding selectivity passes through a maximum of 37 % at lower conversion (25 %). --

sel. p r o p y l e n e S b - V - O

-"

- -m-- sel. p r o p y l e n e AI-Sb-V-O - o- - sel. p r o p y l e n e A I - S b - V - W - O 60 , , , , ,

I II

"

I i,

[]

o

". \ \

o

0

20 40 60 80 C o n v e r s i o n of p r o p a n e (%)

. . n - - y i e l d acn. A I - S b - V - O - o - -yield acn. A I - S b - V - W - O i

6o ~ 40 "8 g

~ 30

0

b t

~ 40 ".~,

20

i

._. 5O

-1z

0

..

--a---yield acn. S b - V - O

80

a

.;

4o

sel. acn. Sb-V-O

- - m - - s e l . acn. AI-Sb-V-O - o - - s e l . acn. A I - S b - V - W - O

~9

,_~

_

"~

20

0

o)

10 0

.o

~

""

"O

..

/ /

i

-/-I~

,' 0

~ 20

40

60

80

100

Conversion of propane (%)

Figure 4. (a): The selectivity to propylene vs. the propane conversion in propane ammoxidation over Sb-V-O with Sb:V = 2:1, A1-Sb-V-O with AI:Sb:V = 21:5:1, and A1-Sb-V-W-O with AI:Sb:V:W = 21:5:1:1. The conversion of propane at =70 % ammonia conversion is indicated for Sb-V-O (A), A1-Sb-V-O (!-3), and A1-Sb-V-W-O (O). (b): The selectivity and the yield to acrylonitrile over the same samples as in (a). 3.2. The AI-Sb-u system Syntheses in the A1-Sb-V-O system were performed starting with mixtures of AI(OH)3, Sb203 and V205 in various ratios, which were calcined in air at 680 ~ [5]. Characterisation with XRD, electron diffraction and EDX revealed the formation of crystalline 8-A1203, o~Sb204, V205 and A1VO4, as well as two rutile related phases Sbo.9Vo.904 and (A1,Sb,V)204. No crystalline A1SbO4 was observed to form at 680 ~ since its formation requires higher synthesis temperature [ 17]. The 8-A1203phase, which was undetectable with XRD, consisted of 5-10 nm sized crystals and was identified using electron diffraction and EDX. The XRD analyses showed that oc-Sb204, A1VO4 and V205 are formed close to their ideal stoichiometry, though due to some decomposition of A1VO4, V205 appears also in the Al-rich side of the system. In Fig. 5 are the fields of formation for Sbo.9V0.904 and (A1,Sb,V)204 depicted for

419 different starting compositions. Formation of Sb0.9Vo.904 occurs over a wide range of starting compositions, while (A1,Sb,V)204 appears as product exclusively in the Al-rich part of the system. o

o

§

/.: " v v ok*

v v-vv o\O

§

/ 20,.s

O%Sb

.\0

.\o

0\0

ok*

Figure 5. Fields of formation in air at 680 ~ of Sb0.9V0.904 (left triangle) and (A1,Sb,V)204 (fight triangle) as determined by X-ray diffraction. The circles are marked at the starting composition (at.%) of the synthesis. Three samples synthesised under various conditions and containing (A1,Sb,V)204 were selected for elemental analysis. Crystals of (A1,Sb,V)204 in each sample were first identified by electron diffraction and were then analysed using EDX. The data are plotted in Fig. 6. Sb204

SbV

V204

bO4

AI203

Figure 6. EDX analysis of ruffle-related (A1,Sb,V)204 crystals in three representative samples, which were prepared by calcination of mixtures with AI(OH)3, Sb203 and V205 in various ratios. O: starting composition AI:Sb:V = 9:9:2, calcined at 900 ~ for 6 days; O: starting composition AI:Sb:V = 6:3:1, calcined at 680 ~ for 4 days; and +: starting composition AI:Sb:V = 1:2:1, calcined at 1000 ~ for 4 days. In spite of a slight overestimation of the antimony content, the plot shows that a solid solution series between A1SbO4 and SbVO4 exists (A13+ ~ V 3+) and can be formulated as All_xSbVxO4 with 0 < x < 0.5. The series evidently does not correspond to substitution of both A13+ and Sb 5+ with V 4+, i.e. a solid solution between A1SbO4 and V204. The gap in the

420 composition between =SbVO4, or more precisely Sbo.9Vo.904, and A105SbV0504 is reasonable, since the former phase is essentially a V 4+ compound, while the "latter is a V 3+ compound. Comparison of the field of formation for (A1,Sb,V)204 in Fig. 5 with its range of composition in Fig. 6 shows that at 680 ~ (A1,Sb,V)204 is not formed in syntheses which start from compositions corresponding to the phase composition. It appears that the formation of Sb0.9V0.904 is kinetically favoured and that a large excess of alumina (> 60 at.%) is required to obtain (A1,Sb,V)204 as the sole rutile structure. (A1,Sb,V)204 likewise Sbo.9Vo.904 has a rutile-related structure and for both phases only the basic rutile-type reflections are observed with powder X-ray diffraction. Selected area electron diffraction, on the other hand, admits the easy distinction between (A1,Sb,V)204 and Sb0.9V0.904 (see Fig. 1). The diffraction pattern in Fig. l b of (A1,Sb,V)204, or (A1,V)SbO4, shows a 3-fold supercell of trirutile-type [ 18]. We have reported that (A1,Sb,V)204 is the active phase for propane ammoxidation in Al-rich A1-Sb-V-O catalysts [5]. Figure 7 shows for a slurry preparation with AI:Sb:V = 21:5:1 the dependence of the catalytic performance with time-on-stream. The propane conversion and the selectivities show almost constant behaviour, indicating that the active structure is formed in the synthesis of the catalyst. Characterisation with FTIR, FT-Raman, XPS and X-ray diffraction before and after use in ammoxidation showed no difference [5]. The selectivities to propylene and acrylonitrile together with the yield to acrylonitrile are shown in Fig. 4 as a function of the conversion of propane. At about 30 % propane conversion, the selectivity to acrylonitrile passes through a maximum of 45 %. The yield to acrylonitrile reaches 20 % at around 55 % propane conversion and 75 % ammonia conversion. Further increase is limited by the complete consumption of the oxygen in the stream. 60

i

o-----o

!

o

r f~

OU

1

~

I

i

I

I

o--

>,

._> 45

o

_ei-~"~ 60

(9 (9

o

o

o

(9 (/)

"0

,-

-o 40

30

-

t-

o

$

(9 1 5 -

> cO

O

O 0

0

I

I

I

2

i

3

Time-on-stream (hours)

Figure 7. Propane conversion (O) and the selectivity to propylene (O) and acrylonitrile ( I ) vs. time-on-stream over a slurry preparation with the ratio AI:Sb:V = 21:5" 1. 3.3.

2O

t-

(..)

The

AI-Sb-V-W-O

0 ~

0

10

20

Time-on-stream

30

40

(hours)

Figure 8. Propane conversion (O) and the selectivity to propylene (O) and acrylonitrile ( I ) vs. time-on-stream over a slurry preparation with the ratio AI:Sb:V:W - 21"5:1" 1.

System

No detailed characterisation of this catalyst system regarding phase composition has previously been reported in the literature, though kinetic data have been given [2,6]. A freshly charged catalyst with the nominal composition AI:Sb:V:W = 21:5:1:1 showed a considerable change in catalytic behaviour with time-on-stream. Figure 8 shows that both the activity and the selectivity to acrylonitrile increase during the first 15 hours, after which the selectivity to acrylonitrile reaches 44 % at about 70 % of propane conversion. On the freshly charged catalyst, initially, mainly degradation products are formed.

421 Characterisation of the used sample with transmission electron microscopy combined with electron diffraction and EDX revealed the presence of a rutile phase with A1, Sb, V and W. This phase is present in form of polycrystalline aggregates consisting of very small crystallites, which are less than 10 nm in diameter, see Fig 9. Eight EDX analyses showed the average composition A10.1Sbo.8V0.7W0.404, which is quite different from the nominal metal composition. The polycrystalline aggregates were observed to be much less frequent in the freshly prepared sample. It appears that the activation behaviour is due to reduction in propane ammoxidation of W 6+ to form W 4+, which can accommodate in a structure of rutile-type. Considering the composition of the rutile with tungsten, it seems that it can be described as a solid solution between (V,A1)SbO4 and WO2, where two W 4+ replaces one Sb 5+ and one V3+/A13+.

Figure 9. Electron micrograph of polycrystalline (A1,Sb,V,W)204 (left), and the corresponding electron diffraction pattern with basic rutile tings (fight). Besides the rutile phase, XRD and electron diffraction of the flesh catalyst showed the presence of ff-Sb204 and WO3, while 8-A1203 was identified by electron diffraction only. The same phases were present in the used catalyst, though the tungsten oxide phase appeared less frequent and showed defects. The selectivity to propylene together with the selectivity and the yield to acrylonitfile are shown in Fig. 4 as a function of the conversion of propane. Compared with the Sb-V-O and A1Sb-V-O samples, the A1-Sb-V-W-O preparation is more selective to C3 useful products at high propane conversions (> 40 %). The selectivity to acrylonitrile reaches almost 50 % at 55 % propane conversion. The yield to acrylonitrile becomes almost 37 % at 80 % propane conversion, where the consumption of ammonia is complete. The results presented for the A1-Sb-V-W-O system show that in this system the active phase is formed in situ in propane ammoxidation. The improvement of the catalytic properties, compared with the less complex subsystems, can be considered due to the introduction of tungsten in the rutile, adjusting the structure and properties of the active vanadium ensemble. 4. C O N C L U S I O N S In the Sb-V-O system the structure that is active for propane ammoxidation is formed in the catalyst synthesis. It is comprised of --SbVO4 with a ruffle superstructure and a surface which is enriched in antimony. The formation of the active surface requires an excess of antimony in the synthesis (calcination), usually Sb:V - 2-5 [1-3,15], as compared with the equimolar ratio for forming Sbo.9Vo.904.

422 The active structure in the A1-Sb-V-O system is a bulk phase, which is directly formed in the catalyst synthesis. It is a trirutile-like phase All_xSbVxO4 (0 < x < 0.5), and the presence in the synthesis of an excess of aluminium is critical for its formation. Thus, the aluminium is not only a catalyst support in the form of ~-A1203, but it is also an element in the active phase. The Sb:V ratio in the trirutile usually falls in the range 2-5 (cf. Fig. 6). In the A1-Sb-V-W-O system a ruffle-type phase with the approximate composition A10.1Sb0.8V0.7W0.404 is active for propane ammoxidation. It is predominantly formed in situ during the ammoxidation. Compared with the rutiles in the other two systems, the A1-Sb-V-Woxide has a lower Sb:V ratio. Under the reaction conditions that were used in the present investigation the highest yields to acrylonitrile obtained were 11% for the Sb-V-oxide, 20 % for the A1-Sb-V-oxide and 37 % for the A1-Sb-V-W-oxide. Thus, considerable enhancement of the yield is achieved by adding aluminium and tungsten to the catalyst. The results can be rationalised in tenns of site isolation [4,16]. On active Sb-V-oxide the excess antimony at the surface creates appropriate isolation of the active vanadium centres. In the A1-Sb-V-oxide the aluminium effects isolation of vanadium not only at the surface, but also in the bulk of the active phase. Compared with these two systems, the introduction of tungsten in the rutile structure further tunes the active ensemble to a configuration which is more efficient for propane ammoxidation. Surrounding the vanadium sites with aluminium and tungsten produces a catalyst which not only gives less amount of C 1-2 products, but also less degradation of ammonia (see Fig. 4). REFERENCES

1. 2. 3.

4. 5. 6. 7. 8. 9. 10. 11. 12. 13. 14. 15. 16. 17. 18.

A.T. Guttmann, R.K. Grasselli and J.F. Brazdil, US Patents 4 746 641 (1988) and 4 788 317 (1988); assigned to The Standard Oil Company (Ohio). G. Centi, R.K. Grasselli, E. Patane and F. Trifir6, in G. Centi and F. Trifir6 (eds.), New Developments in Selective Oxidation, Studies in Surface Science and Catalysis, Vol. 55, Elsevier, Amsterdam, 1990, pp. 515-526. R. Nilsson, T. Lindblad, A. Andersson, C. Song and S. Hansen, in V. Cort6s Corber~in and S. Vic Bell6n (eds.), New Developments in Selective Oxidation II, Studies in Surface Science and Catalysis, Vol. 82, Elsevier, Amsterdam, 1994, pp. 293303. R.K. Grasselli and J.D. Burrington, in D.D. Eley, H. Pines and P.B. Weisz (eds.), Advances in Catalysis, Vol. 30, Academic Press, New York, 1981, pp. 133-163. J. Nilsson, A.R. Landa-C~inovas, S. Hansen and A. Andersson, J. Catal., 160 (1996) 244. R. Catani, G. Centi, F. Trifir6 and R.K. Grasselli, Ind. Eng. Chem. Res., 31 (1992) 107. T. Birchall and A.W. Sleight, Inorg. Chem., 15 (1976) 868. F.J. Berry, M.E. Brett and W.R. Patterson, J. Chem. Soc. Dalton Trans., (1983) 9. A. Landa-C~inovas, J. Nilsson, S. Hansen, K. StS.hl and A. Andersson, J. Solid State Chem., 116 (1995) 369. A.R. Landa-C~inovas, S. Hansen and K. Stgthl, Acta Crystallogr. B, in press. J. Nilsson, A. Landa-C~inovas, S. Hansen and A. Andersson, Catal. Today, 33 (1997) 97. G. Centi and P. Mazzoli, Catal. Today, 28 (1996) 351. S. Hansen, K. St~hl, R. Nilsson and A. Andersson, J. Solid State Chem., 102 (1993) 340. R.G. Teller, M.R. Antonio, J.F. Brazdil and R.K. Grasselli, J. Solid State Chem., 64 (1986) 249. R. Nilsson, T. Lindblad and A. Andersson, J. Catal., 148 (1994) 501. J.L. Callahan and R.K. Grasselli, AIChE J., 9 (1963) 755. K. Brandt, Ark. Kemi, 17 (1943) 1. S. Hansen, A. Landa-C~inovas, K. St~hl and J. Nilsson, Acta Crystallogr., A51 (1995) 514.

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

423

hlfluence of A n t i m o n y Content in the Iron A n t i m o n y Oxide Catalyst and Reaction Conditions on the ( A m m ) O x i d a t i o n of Propene and P r o p a n e Eric van Steen, Gunther Kuwert, Alvin Naidoo and Marco Williams Catalysis Research Unit, Department of Chemical Engineering, University of Cape Town, Private Bag, Rondebosch 7700, South Africa The influence of antimony content and of ammonia partial pressure on the selectivity of the (amm)oxidation of propene and propane with FeSbO4 can be understood in terms of the degree of reduction of the catalyst surface at steady state conditions. The higher the degree of reduction which can be caused either by a low antimony content or by a high ammonia partial pressure the higher the selectivity for the combustion/degradation products. 1. INTRODUCTION Valuable chemicals can principally be produced from paraffins if these unreactive compounds can be functionalized. This primarily requires the activation of the rather unreactive paraffinic C-H bond or C-C bond. Therefore, high temperatures are usually applied in the functionalization of paraffins. An interesting route is the partial oxidation of these compounds. This may yield olefins by oxidative dehydrogenation, oxygenates by oxygen insertion, or even nitrile compounds if a reactive nitrogen compound is added to the feed. The partial oxidation of olefins over a variety of mixed oxide catalysts is well known [ 1] and has been studied in great detail. For example, iron antimony oxide is known to be a selective catalyst for the partial oxidation of propene yielding acrolein [2-4] and of 1-butene yielding 1,3 butadiene and 2-butenal [5,6]. The oxidation of paraffins, like propane, with this type of catalysts yields only the combustion products CO and CO2 [7]. Propane can however be selectively oxidised over this type of catalysts in the presence of ammonia to yield acrylonitrile [8]. The active and selective phase in the iron antimony oxide catalyst for the selective partial oxidation of propene is FeSbO4 [3,8]. The selectivity to acrolein, however, improves dramatically if the antimony to iron ratio in this catalyst exceeds one [3 ]. It has been observed, that the surface of iron antimony oxide with Sb/Fe = 1 is enriched with antimony [8,10]. If the Sb/Fe ratio in the catalyst formulation is increased, then a higher enrichment of the surface with more antimony was observed [8]. It has been indicated for the analogue U-Sb system [9], that the surface antimony inhibits the reduction of the catalyst during the reaction. Different oxygen species are involved in the partial (amm)oxidation reaction [1]. The presence of e.g. 0 2. will lead to a nucleophilic addition to the adsorbed hydrocarbon species resulting in the formation of the selective partial (amm)oxidation products. The presence of electrophilic oxygen species like O2-and O lead to the formation of degradation and combustion products. All these oxygen species can be formed from gas phase oxygen. 0 2. can also act as lattice oxygen. Although nucleophilic oxygen species are formed in the re-oxidation process of the catalyst, the relative concentration of these species will be less on surface with a

424 higher degree of reduction. The degree of reduction of the catalyst surface during the reaction will therefore influence the relative composition of the pool of oxygen species at the surface. The presence of ammonia in the ammoxidation is expected to change the degree of reduction of the catalyst at steady state conditions [9], and hence a change in selectivity should be observed. In this paper we want to compare the influence of antimony and ammonia on the (amm)oxidation of propene and propane, which are expected to have opposite effects. We will explain the observed changes in terms of a change in the degree of reduction of the surface, and hence in terms of changed composition of the pool of oxygen species at the surface.

2. EXPERIMENTAL The iron antimonate catalysts with an antimony to iron ratio of 1:1 and 2:1 was prepared using the method described by Allen et al. [4]. Briefly, Sb203 was added to a solution of iron nitrate in its own crystal water at 60 ~ The temperature was subsequently raised to 80 ~ and an aqueous 25 % NH3-solution was added to raise the pH to 3. The precursor was then dried for 16 hours at 125 ~ in an oven. Finally, the catalyst precursor was calcined at 800 ~ for 8 hours. Two different methods were attempted to enrich the iron antimony catalyst with Sb/Fe = 1 with antimony. With both methods the catalyst was contacted with a Sb-containing solution (2 ml/gcatalyst), which was prepared by dissolving Sb203 in HC1. Method A was the impregnation of iron antimonate with antimony. For method B the antimony was precipitated by adjusting the pH to 8 by adding an aqueous 25 % NH3 solution. Subsequently, the catalyst was dried at 120 ~ for 4 hours and calcined at 800 ~ for 8 hours. For method A the impregnation and calcination steps were repeated once. The phases present in the catalysts were characterised using XRD. The bulk composition of the catalysts were determined using ICP and XRF. The ammoxidation was carried out in a fixed bed glass reactor [6]. The catalyst (0.5 g; dp < 0.1 mm) was diluted with sand (4 g; dp = 0.2 - 0.3 mm). In all experiments the inlet partial pressures of propene/propane and of oxygen were kept constant at respectively 15 kPa and 30 kPa. The inlet partial pressures of ammonia and nitrogen were adjusted to give a total of 90 kPa. The space velocity for the propene and propane (amm)oxidation was kept constant at 1.2 mmol hydrocarbon/(gcat min). Methane was added to the effluent as internal standard. The organic products in the effluent were analysed with a GC equipped with an FID using the offline ampoule sampling technique [ 11 ]. The combustion products CO and CO2 were analysed using an on-line IR-analyzer. HCN was trapped in a silver nitrate solution. Unconverted NH3 was trapped in a 0.1 M HCI solution. The amounts of HCN and unconverted NH3 were then determined by titration.

3. RESULTS AND DISCUSSION

3.1 Catalyst preparation and characterisation The XRD of the base iron antimony oxide catalyst with Sb/Fe = 1 showed only the presence of the FeSbOn-phase. The catalyst with Sb/Fe = 2 showed also the presence of c~-

425 Sb204. The two phases in this catalyst could not be visually identified and therefore not be separated easily. The calcined iron antimony oxide with Sb/Fe = 1 had a BET-surface area of 28.9 m2/g. Based on the surface area and an assumed diameter of Sb203 of 5.5 9 10a~ m the amount of antimony added to the catalyst either with method A or method B was calculated to be between 0.2 and 10 times a monolayer coverage. The final amount of antimony on the surface with method A will have been lower than intended, since Sb in a HCl-solution partly forms SbC13, which evaporates at 283 ~ The precipitation of the antimony from a HCl-solution (method B) led to the formation of the FeSbO4 and ot-Sb204 phase, which could be easily identified and separated. The phases were separated. The phase, which showed by XRD only the existence FeSbO4, was used in the catalytic studies. The preparation of the catalysts according to method B was more reproducible than the preparation according to method A. The catalysts used in the catalytic studies, which were prepared according to either method A or method B, showed by XRD solely the existence of FeSbO4. The antimony was either finely dispersed on the surface or very small Sb204 crystals were formed, which could not be detected by XRD. An attempt was made to measure the change in the bulk Sb/Fe ratios in these catalysts, but the observed changes were within the accuracy of the analyses. The results are therefore given based on the initial amount of Sb added to the base iron antimony oxide catalyst (Sb/Fe = 1).

3.2 Initial activity/selectivity in the ammoxidation of propene

Figure 1 shows the conversion of propene in the ammoxidation and the selectivity to the partial (amm)oxidation products, acrylonitrile plus acrolein, as a function of time on stream for the base iron antimony oxide catalyst (Sb/Fe =1) and for the impregnated catalyst with 0.27 g Sb added per g FeSbO4 (method A), which corresponds to a coverage with antimony of 5 100

30 Base CatTlyst

O

E 25

Impregnated catalys!

8O

.ooel 20

~60

9

tD

> 15

~ 10 o~

ml.m u

9

9

40

0

• m

5 !mpregnated ~

0 -1

0 1 log (Time on Stream/min)

Base Catalyst (Sb/Fe = 1)

~

......

> acrylonitrile

Ca Hs, ads +

>

......

~ acrolein

Oads

It can then be easily derived that the ratio of the rate of formation of acrylonitrile relative to the rate of formation of acrolein will be:

431 racr,/ionitrilc ~ [Oads] [NH3,~ds] racrolein [OHads] 2 If the concentration of oxygen species decreases due to the increase in the ammonia partial pressure, then a maximum is expected for acrylonitrile content, i.e. for the probability of NHinsertion instead of oxygen insertion in the ~-allylic complex. 3.5 Influence of t e m p e r a r e on the a m m o x i d a t i o n of p r o p e n e / p r o p a n e

The temperature influence on the ammoxidation of propene and propane was studied over FeSbO4 at a fixed ammonia partial pressure of 10 kPa (see Figure 7). The apparent activation energy for the formation of acrolein and acrylontrile from propane was almost twice the apparant activation energy for the formation of these compounds in the propene conversion (90 vs. 55 kJ/mol). This might reflect a different rate determining step in the propane conversion compared to the one in the propene conversion. In the latter case the formation of the allylic complex is most likely to be the rate determining step [ 13,14]. For the activation of propane a surface propyl group is a very likely initial intermediate [8], whose formation might be energetically less favoured than the second hydrogen abstraction and the third hydrogen abstraction yielding the ~-allylic adsorbed species. -2

Ea (kJ/mol) = -3 - 55 ~

~s

-4

E -5

-

100 Propene

~

9o

~60

-

e

+

' 1.3

1.4

n nnn~nn \9

40

"~ "~ "-6 ~ ~ 20 < 0

"~ -6 = -7

opane

80

1.5 1.6 1000/T, 1/K

1.7

Prop ' 300

ene~ ' ~ '

'

'

'

400 Temperature, ~

'

' 500

Figure 7 Influence of temperature on the formation of acrylonitrile plus acrolein in the ammoxidation of propene and propane (pNm, inlot = 10 kPa) over an iron antimony oxide catalyst (Sb/Fe = 1) left: Arrhenius' plot right: Selectivity of acrylonitrile plus acrolein The selectivity for the partial (amm)oxidation products decreased with increasing temperature. This indicates a higher activation energy for the formation of combustion/ degradation products than for the selective formation of the partial (amm)oxidation products. With propane a higher selectivity for acrylonitrile and acrolein was observed. In the propane (amm)oxidation the content of acrylonitrile increased with increasing temperature, which indicates that the insertion with oxygen becomes less probable with increasing temperature. This might be attributed to either the difference in the activation energy for the NH- and Oinsertion and/or a higher degree of reduction of the partial oxidation catalyst. In the propene

432 ammoxidation the acrylonitrile content showed a maximum, which can also be explained by an increase in the degree of reduction of the catalyst surface (vide supra). The decrease in the selectivity of the partial (amm)oxidation products acrolein plus acrylonitrile might therefore not only be attributed to the differences in activation energy, but can also be caused by the change in the degree of reduction. 4. CONCLUSIONS Iron antimony oxide catalysts enriched with antimony have been prepared. The impregnation of FeSbO4 is not very reproducible. The catalyst prepared according to method A showed in the propene partial oxidation results similar to iron antimony oxide catalysts with Sb/Fe = 2. In the preparation of the catalysts according to method B most of the antimony precipitated as a separate phase. The selectivity changes in the (amm)oxidation of propene and propane can be (partly) understood in terms of a change of the degree of reduction of the surface, which would cause a change in the distribution of nucleophilic and electrophilic oxygen species in the pool of surface oxygen species. Antimony inhibits the surface reduction and the catalysts with an antimony enriched surface therefore exhibit a higher selectivity with a lower activity. Ammonia inhibits the formation of partial (amm)oxidation products by competitve adsorption. It can also modify the degree of reduction of the surface and thereby shitting the selectivity of the reaction. To obtain a maximal selectivity to the partial (amm)oxidation products a molar ammonia/propene ratio of less than one should be applied. The ammoxidation of propene has a higher apparent activation energy than the ammoxidation of propene, but shows a higher selectivity. REFERENCES A. Bielanski and J. Haber, Oxygen in Catalysis, Marcel Dekker Inc., New York, 1991 [1] G.K. Boreskov, S.A. Venyaminov, V.A. Dzisko, D.V. Tarasova, V.M. Dindoin, N.N. [21 Sanobova, I.P. Olenkova, and L.M. Keteil, Kinet. Katal. 10 (1969) 1530 I. Aso, S. Furukawa, N. Yamazoe, and T. Seiyama, J. Catal. 64 (1980), 29 [31 M. Allen, R. Bettely, M. Bowker, and G. Hutchings, Catal. Today 9 (1991) 97 [4] G.I. Straguzzi, K.B. Bischoff, T.A. Koch, and G.C.A. Schuit, J. Catal. 104 (1987) 47 [5] E. van Steen, M. Schnobel, and C.T. O'Connor, in Heterogeneous Hydrocarbon [6] Oxidation (B.K. Warren, S.T. Oyama, eds.), ACS Symposium Series 638, ACS Washington DC, 1996, p. 276 M. Schnobel, PhD Thesis, University of Cape Town, 1997 [7] M. Bowker, C.R. Bicknell, and P. Kerwin, Appl. Catal. A: General 136 (1996) 205 [8] R.K. Grasselli and D.D. Suresh, J. Catal. 25 (1972), 273 [9] [10] S.R.G. Carraz/m, L. Cadus, Ph. Dieu, P. Ruiz, and B. Delmon, Catal. Today. 32 (1996), 311 [11] H. Schulz, W. Bohringer, C. Kohl, N. Rahman, and A. Well, DGMKForschungsbericht 320, DGMK Hamburg, 1984 [12] J.D. Burrington, C.T. Kartisek, and R.K. Grasselli, J. Catal. 81 (1983), 489 [13] J.D. Burrington, C.T. Kartisek, and R.K. Grasselli, J. Catal. 87 (1984), 363 [14] G.W. Keulks, and M.-Y. Lo, J. Phys. Chem. 90 (1986), 4768

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

433

Catalytic selective o x i d a t i o n of C2-C4 alkanes over reduced heteropolymolybdates Wen Li and Wataru Ueda Department of Environmental Chemistry and Engineering, Tokyo Institute of Technology, Nagatsuta-cho 4259, Midori-ku, Yokohama 226 JAPAN Catalytic oxidation of ethane, propane, and isobutane to corresponding unsaturated acids and acetic acid with molecular oxygen were carried out over reduced H3PMo12040 catalysts above 300~ A highly reduced state of H3PMo12040 was formed by the heat-treatment of the pyridinium salt under N2 flow up to 420~ This catalyst, denoted by H3PMo12040(Py), gave the highest conversion of alkanes and selectivities to the objective products and its reduced state was highly stable under the conditions of catalytic oxidation. The highly stable reduced state was composed of defect Keggin unit with oxygen vacancies and a secondary structure formed by the defect Keggin unit and remaining pyridine in the structure. Another reduced catalyst generated by the heat-treatment of (NH4)3PMo12040 was active for the propane and isobutane conversion but poorly selective to acrylic acid and methacrylic acid, because this catalyst was readily oxidized during the oxidation. Non-reduced H3PMo12040 catalysts showed poor activity for all of the reactions. A reaction mechanism is proposed over the reduced H3PMol2040(Py) catalyst. 1. INTRODUCTION Because of the global abundance of liquefied petroleum gas (LPG), interest in the potential use of ethane, propane, and butanes as sources of the corresponding alkenes or their derivatives is increasing [1]. In the last decade much progress has been made, particularly in the selective partial oxidation of light alkanes with molecular oxygen in gas phase [ 1, 2]. For economic reasons, molecular oxygen is usually used as the primary oxidant[3]. To promote both the conversion of reactants and the selectivity to partial oxidation products, many kinds of metal compounds are used to create catalytically active sites in different oxidation reaction processes [4]. The most well-known oxidation of lower alkanes is the selective oxidation of n-butane to maleic anhydride, which has been successfully demonstrated using crystalline V-P-O complex oxide catalysts [5] and the process has been commercialized. The selective conversions of methane to methanol, formaldehyde, and higher hydrocarbons (by oxidative coupling of methane [OCM]) are also widely investigated [6-8]. The oxidative dehydrogenation of ethane has also received attention [9,10]. Heteropolycompounds, having unique structures and the resulting acidic and redox properties, exhibit high oxidation abilities in the selective oxidations of alkanes [ 1, 2, 11, 12]. Especially, it should be noted an interesting fact that the heteropolycompounds containing molybdenum showed high catalytic activities in the partial oxidation of hydrocarbons to corresponding acids [10, 13-23]. The direct oxidation of light alkanes to corresponding acids has tremendous economic value. Although it seems very difficult, some related research results have been reported with the catalysts of heteropolycompounds containing molybdenum for the reactions of isobutane to methacrylic acid [3, 13, 19, 22, 23] and propane to acrylic acid [13, 17, 21]. There appears to be no published literature pertaining to the production of

434 acetic acid from ethane oxidation directly over heteropolycompound catalysts. Recently we found that highly reduced H3PMo12040 which was formed by the heattreatment of pyridinium salt can catalyze the propane oxidation to acrylic acid and acetic acid selectively [24, 25]. After activation in N2 flow at 420~ for 2hr, the catalyst of H3PMo 12040 (Py) shows reduced state of molybdenum and a new stable structure in which pyridine remains as the linkage of the secondary structure. The activated H3PMo12040 (Py) also gives catalytic activities in the partial oxidations of ethane and isobutane to acetic acid and methacrylic acid respectively. In this paper, we will report the oxidation results of C2-C4 alkanes and discuss the roles of reduced state and activation of molecular oxygen over this catalyst.

2. EXPERIMENTAL 2.1 Catalyst preparation 12-molybdophosphoric acid (I--I3PMol2040.xH20) was obtained from Nippon Inorganic Color & Chemical CO., LTD.. H3PMol2040.xH20 was dissolved in distilled water, then the solution was filtered, evaporated at 40-50~ with stirring until there was little crystals on the surface of the solution. The mixture was kept at 5~ for overnight to recrystallize, and the crystal was filtered and dried at 40~ for 8hr. By this way, H3PMo 120,0.10H20 was obtained. Pyridine-treated heteropolycompound, denoted by (Py), was prepared by precipitation method. Recrystallized H3PMo12040 was dissolved in distilled water with stirring at 40-50~ and desired amounts of pyridine was added slowly, then the solution containing precipitates was evaporated to dryness at 40~ The obtained solid was further dried in N2 flow at 120~ for 8hr. For the synthesis of (NH4)3PMo12040, desired amounts of (NH4)3MoTO24 and H3PO4 were dissolved in distilled water at about 40-50~ then HNO3 was added slowly with stirring at 50~ The addition of HNO3 resulted in a yellow precipitate, which was filtered, washed with distilled water, and dried at 40~ overnight. The obtained solid contains the desired composition and shows cubic structure as reported in literature [20]. Before catalytic reaction, H3PMol2040(Py) and (NH4)3PMo12040 were activated in N2 flow for 2h at 420~ and 450~ respectively.

2.2 Catalytic testing Alkane oxidations were carried out at an atmospheric pressure in a conventional flow system with a fixed bed Pyrex tubular reactor( qb 12mm). Catalyst(3g) was diluted by 2g sands to prevent the catalyst from overheating during reaction. The catalysts were heat-treated in a nitrogen flow at a selected temperature, then set to desired reaction temperatures. The feed compositions were controlled with mass flow controller (KOFLOC 3510). The total flow rate was 30 ml.min l. The feedstock and products were analyzed by on-line gas chromatography operating with two sequential columns, molecular sieve 13X lm at room temperature for separation of 02, N2 and CO, and Porapak Q 4m at 60"C-140~ for hydrocarbons and CO2. There was a ice cooling trap to collect other products at the exit of the reactor, and the collected products were analyzed quantitatively by another gas chromatography (TC-WAX 60m capillary).

2.3 Characterization TPD The experiments were carried out using a standard apparatus. Helium was used as a carrier gas. The thermal conductivity detector was employed to detect the changes of desorption. A dry ice trap was used to remove the water in the carrier gas. The weight of sample was 10mg. The sample was pretreated in He flow at room temperature for lh before the experiment. The heating rate was 10~ l. XPS A Shimadzu ESCA-750 electron spectrometer with an aluminum anode (1486.6eV) was used to obtained X-ray photoelectron spectra. All binding energies were referenced to gold (Au 4f7/2 line; 83.8eV) which was deposited on samples in vacuum. The activated sample in the flow system was outgassed in the preparation chamber of the spectrometer

435

under 10~ Torr for 15 minutes. 3. R E S U L T S AND DISCUSSION 3.1 Ethane Oxidation We examined the ethane oxidation over the mentioned catalysts under different reaction conditions at 280~176 About 1% conversion of ethane and 10% selectivity to acetic acid were obtained over the activated H3PMol2040(Py) at 3400C. Ethane oxidation did not occur over non-reduced H3PMo12040 and activated (NH4)3PMo12040 catalysts. There seems to be no literature about the direct oxidation of ethane to acetic acid over heteropolycompounds catalysts. Nevertheless, there is a limited amount of literature[ 10, 26-28] about direct oxidation of ethane to acetic acid over oxide catalysts at low temperature (200350~ It seems that vanadium and molybdenum are necessary to those catalysts, and the addition of water is useful to increase the production of acetic acid. Roy et al. [ 10] has proved that vanadium and molybdenum phosphates supported on TiO2-anatase were effective in the direct oxidation of ethane to acetic acid. Considering previous research results, it is suggested that other promoters, such as transition-metal oxides, are necessary to enhance the catalytic activity of the activated H3PMol2040(Py) in the direct oxidation of ethane to acetic acid. 3.2 Propane Oxidation The catalytic activities of H3PMo 12040, (NH4)3PMo 12040 and I-I3PMo 12040(Py) in the propane oxidation are shown in Figure 1. The non-reduced, acidic form of molybdophosphate catalyst, H3PMo12040, revealed an activity and yielded propene mainly. The activated H3PMol2040(Py) catalyst showed distinctly enhanced activities for the propane oxidation, and the catalyst gave a significantly different product distribution, where acrylic acid and acetic acid were main organic oxygenated products. Such drastic chaiage of the product distribution was observed even at a low conversion of propane under a short contact time. Obviously, this results from an intrinsic change of catalyst properties by the treatment with pyridine. The product distribution in activated H3PMo12040(Py) catalyst is different from that in the reported metal oxides catalyst or heteropolycompounds catalysts promoted by metal atoms, such as Bi, Fe, Ni, or Cs substitution for proton and V s+ substitution for Mo 6. in

Catalyst (precursor) [S.A.: m2/g ]

0

Conversion of Propane (%) 2 4 6 8 |

lO [] Acetic acid 9 Acrylic acid [] Propene

w

(NH4)3PMo12040 [ 15.21 H3PMol2040(Py) [12.8] H3PMo12040(Py)

(0.Sg) H3PMo12040 [3.8]

73% NgJ/F. .

.

[

!

0

10

.

.

. i

~ !

,t~ "

20 30 40 Selectivity(%)

Figure 1. Catalytic performance for the propane oxidation at 340~ Catalyst weight: 3g; total flow rate: 50ml/min; reactants" C3H8:O2:H20:N2=20.4:9.9:19.9:49.8 (mol%).

436 the direct oxidation of propane to acrylic acid [13, 17, 21, 29]. The former shows relatively higher ability to form organic oxygenates than the latter. The mentioned metal promoters seemed to be favorable to form propene, and oxidize it to CO and CO2 directly. Figure 1 also shows the catalytic performance of the activated (NH4)3PMo12040 catalyst which has a surface area similar to the H3PMo12040(Py)catalyst. Non-activated (NH4)3PMo 12040was completely inactive but became active after the heat-treatment under the N2 flow at 450~ The catalyst, however, gave a broad product distribution. The special ability of the activated H3PMo 12040(Py) in the partial oxidation of propane allows us to investigate its catalytic performance in detail. The conversion of propane and selectivities to acrylic acid and acetic acid as a function of time on stream at 3400C are shown in Figure 2. The main products were acrylic acid and acetic acid, and the remainder is carbon oxides. At the beginning of the reaction (t~ 40 ....

>

,,41.-I

O

(U

-~ 20 O0

0

' .9 - , ,,,,

0

.

I

10

I

I

20

I

I

30

i

I

4O

I

_ . _

I

50

Conversion, % F i g u r e 4. Variation of the selectivities to propylene (D); CO J ) and CO2 (V) in the ODH of propane on a VzOJT-A1203 catalyst (3.5 wt% V-atoms) at 500-500~ Experimental conditions in text.

Since a similar trend is observed on all the catalysts studied it can be concluded that propylene is a primary unstable products while CO and CO2 can be considered as primary and secondary products. However, oxygenated products other than COx were not observed. Figure 5 shows the variation of the catalytic activity for the oxidation of propane at 500~ with the reducibility of V-based catalysts, determined as the inverse of the temperature of maximum H2-consumption (1/Tm). A parallelism between the reducibility of catalyst and the catalytic activity for propane conversion similar to those observed on supported vanadium oxide catalysts [1, 10] can be proposed. On the other hand, the selectivity to propylene varies with the amount and type of the metal oxide added. Figure 6 shows the variation of the selectivity to propylene with the MeN atomic ratio, at a propane conversion of 10%. In all cases, the doped catalysts with a low MeN atomic ratio show a higher selectivity to propylene than the undoped sample, although K-containing catalysts, specially the sample with a KN atomic ratio of 0.7, are the most selective ones.

449

60

9 Bi.1 [3

K.0.7

O

K.0.1

I

P.1

o

A

Bi.0.2

o

V

Mo.0.2

-~ 20

+

V.AL

A

Mo.0.9

~7

P.0.5

40

>




~.~ o ~40C9

~9

~ ~20-

-

N2 (20h)

+ 20% H20

0

I

I

120

240

I

N2 (20h) I

360 480 t, min

I

I

600

720

Figure 4. Dependence of the conversion of n-butane (I I) and of the MA selectivity (A) on time on stream Feed: 1.46 % C4H10- air, T= 450 ~ ~ = 1.6 sec; after 320 min and 570 min t r e a t m e n t with butane containing feed: 20 h Nz (GHSV: 67 h a (NPT), T= 450~

3.2. T h e m o b i l i t y o f p h o s p h a t e a n i o n s in (VO)2P207 c a t a l y s t s In order to get more information on the factors influencing the mobility of phosphate anions and causing the loss of phosphorus in (VO)2P207 catalysts in situ-thermogravimetric experiments were performed. Firstly, the adsorption of water molecules on the surface of CVO)2P207 was studied. Selected results are depicted in the Figures 5 and 6. Figure 5 shows the water uptake of the sample versus t e m p e r a t u r e on changing stepwise the composition of the gas phase from dry to wet nitrogen and vice versa (see also Figure 7). With increasing temperature the amount of adsorbed water molecules decreased up to 300~ and then remained nearly constant at higher temperatures. Figure 6 shows the dependence of P:V ratios determined by ISS (uppermost layer) and by XPS (about 50% of the XPS signal intensity stems from first 6 layers) on the temperature of the water vapour treatment.

467

~0.6

~4.0

%.

55 55

~.3o

~0.4 r

=0.2 ~: 0

t

t

200

400

C~l. 0 600

T,~ Figure 5. Dependence of the mass of H2Oads on the temperature, CH2O: 20% in N2, m e a t . : 120mg

I

100

I

300 T,~

,

I

500

Figure 6. Dependence of the P/V ratios (XPS (A), ISS (E~i))on the temperature of the water vapour treatment, CH2O: 30% in N2

Both the ISS- and the XPS-P:V ratios decrease with increasing temperature, i.e., with decreasing amounts of adsorbed water molecules. This dependence illustrates again the driving force of the P enrichment, i.e., the gain in energy by hydratation of P-OH groups. After conventional butane oxidation (450~ dry feed) the ISS-P:V ratio of the surface layer in used catalysts exceeds 2.0 indicating both an excess of phosphorus and a marked presence of vanadium in the utmost layer. This is in contrast to the pyrophosphate termination model of Thompson and Ebner [6] from which primarily the presence of phosphorus on the surface follows. P:V ratios ranging from 1.3 up to 1.4 were measured by XPS in samples which were obtained after conventional butane oxidation. These ratios are higher than those of O k u h a r a et al. [21] (1.10 _+ 0.04) and Coulston et al. [3] (1.08) who also used reliable sensitivity factors of XPS obtained by calibration. The reason for this difference might be that both the precursor synthesis and the catalyst preparation of the authors differ from the methods applied in the present work. From the XPS and ISS results obtained the following conclusion may be drawn. If the P:V ratio in the surface layer is about 2:1 then the P:V ratio in the next subsurface layers should be of similar magnitude. This follows from an electron attenuation length in the sample 00 of about 8 layers roughly estimated according to [20] and the lattice data of (VO)2P207 and VO(PO3)2 [22,23]. This conclusion is s u p p o r t e d by the s t a t e m e n t of C o u l s t o n et al [3] t h a t at m o s t 25% of the X P S s i g n a l is expected to originate from the o u t e r 5A of the (VO)2P207 s a m p l e . As Shown in Figure 7, a total weight loss of 0.11% occurs during the alternate t r e a t m e n t with wet and dry nitrogen for 6h at 480~ This is confirmed by other adsorption studies in the temperature range from 100~ to 600~ which reveal weight losses in the same order of magnitude. In contrast, no weight loss was observed if dry nitrogen was dosed over a dry calcined fresh catalyst.

468 122.3 ~ .

wet

wet Figure 7. Dependence of the mass of the (VO)2P207 catalyst on the duration of the alternate t r e a t m e n t with wet (20% H20) and dry N2, T: 480~

~122 1 drY-drY" 121.9

I 0

~y I

120

y I

240 360 t, min

I 480

In order to elucidate the reasons of this effect the weight of (VO)2P207 samples was m e a s u r e d during the treatment with nitrogen streams containing different levels of water vapour. Typical results are presented in Figures 8 and 9. 0.4 0.3 0.2 0.1

0.5 ~Y h~

0

-0.5

o o0 oQ

-0.1 -O.2 -0.3

I 0

120

-1.5

-2

I '

60

-1

180

t, min Figure 8. Influence of water vapour on the mass of a (VO)2P207 catalyst CH2Oin N2 : 0 % (O) ,10 % (?~),20 % (A), 30 % (X), T = 400 ~

0

60 120 t , min

180

Figure 9. Influence of water vapour on the mass of a (VO)2P207 catalyst at different t e m p e r a t u r e s C H 2 0 : 30 % in N2, T = 300 ~ (':)), 400 oc (D), 550 oc (A)

Figure 8 shows that in the first 30 min the mass of (VO)2P207 rose with increasing concentration of water vapour in the feed due to the adsorption of H20. During treatment with 20% and 30% H20 in N2 a subsequent weight loss was observed which was more pronounced with the higher water content. Figure 9 exhibits that the weight loss strongly depends on the t e m p e r a t u r e s of the water vapour treatment. The weight of the sample remained constant at 300~ whereas a strong mass loss was observed at 550~ The chemical analysis showed the presence of phosphoric acid and, surprisingly, also traces of vanadium containing species in the condensate. This finding is explained by decomposition of vanadyloligophosphate structures on the catalyst surface which are assumed to be formed under the influence of adsorbed w a t e r molecules. It is assumed that this also occur during the technical process but on a much longer time scale. However, after some hundreds of hours time on s t r e a m of an industrial reactor the loss in phosphorus leads to a decline in

469 selectivity and to an enhancement in activity. The addition of volatile phosphorus containing compounds to the feed is required to retain the catalytic performance. 4. C O N C L U S I O N S As shown by the results presented the surface layer of an equilibrated (VO)2P20: catalyst consists of a VPO structure with a P:V ratio of about 2 and in the next subsurface layers also higher P:V ratios than one have to be assumed. To elucidate the structure of the active site of the (VO)2P207 catalyst first of all some characteristics of the known active site models are given: usually the active center is described in terms of the bulk structure sometimes terminated in pyrophosphate groups stressing the role of edge-shared VO6 double octahedra (e.g. [2,6,24]). However, the existence of these VO6 double octahedra on the surface plane parallel to (I00) has not been proven up to now for the equilibrated (VO)2P207 catalyst, the surface of which is enriched in phosphorus due to the formation of water in the butane oxidation. Furthermore, the dynamic character of the surface is not considered in the published models although Pepera et al. [25] have demonstrated a high mobility of surface oxygen and that only a few surface layers take part in the butane oxidation. Additionally, the special properties of the surface in contrast to the bulk of a solid catalyst must be considered. This surface must be seen as a two-dimensional phase with special s t r u c t u r a l and energetic properties (see e.g. [26]). Taking these ideas into account, in the following a model is proposed which is able to explain our findings as well as the results of other authors who investigated the catalytic properties of VO(PO3)2 [7,8]. The active site of an equilibrated (VO)2P207 catalyst is assumed to consist of a thin layer of si~gle VOw; octahedra linked to the (VO)2P207 bulk in a chain-like fashion. This chain s t r u c t u r e is thought to be similar to the one of VO(PO3)2 and the VO6 octahedra to be isolated from each other by oligo- and monophosphate anions resulting in a P:V ratio of the uppermost layer of about 2. This structure should comprise besides the surface layer only few subsurface layers. The valence state of v a n a d i u m in this layer is between +4 and +5 according to previous results [5] and both phosphate and oxygen ions are very mobile (see [25]). This model explains additionally the result t h a t the area-specific rates of the MA formation are similar on (VO)2P207 and VO(PO3)2 as shown by Morishige et al. [7], S a n a n e s et al. [8] and by us [9]. The remarkable resistance of the bulk to oxidation under the usual reaction conditions as shown by earlier TG m e a s u r e m e n t s [5] reveals the active surface to be confined to a few layers on an oxidation resistent (VO)2P207 core. Furthermore, applying this model it is also possible to explain the enhancement in MA selectivity of a fresh catalyst p r e p a r e d from aqueous medium in the conditioning procedure [5]: During this process the original VO6 double octahedra of the surface region are splitted by migratory acidic phosphate anions to corner-shared single octahedra chains linked to the bulk double octahedra chains. This leads to a structure with isolated vanadyl chains which act more selectively according to the well known

470 ,,principle-of-the-density-of-the-oxidizing-sites" [27]. The same result is obtained when the precursor VOHPO~-0.5H~O is calcined in the presence of water vapour which results in a selective catalyst without the need of any conditioning procedure [28]. REFERENCES

1. 2. 3. 4. 5.

G. Centi, Catal. Today, 16 (1993) 1 F. Cavani, F. TrifirS, Catalysis, 11 (1994) 246 G.W. Coulston, E.A. Thompson, N. Herron, J. Catal., 163 (1996) 122 D. Ye, A. Satsuma, T. Hattori, J. Murakami, Appl. Catal., 69 (1991) L1 B. Kubias, M. Meisel, G.-U. Wolf, U. Rodemerck, Stud. Surf. Sci. Catal., 82 (1994) 195 6. M. R. Thompson, J. R. Ebner, Stud. Surf. Sci. Catal., 72 (1992) 353 7. H. Morishige, J. Tamaki, N. Miura, N. Yamazoe, Chem. Lett., (1990) 1513 8. M. Sananes, G.J. Hutchings, J.C. Volta, J. Catal., 154 (1995) 253 9. F. Hannour, A. Martin, B. Kubias, B. Liicke, E. Bordes, P. Courtine, Catal. Today, submitted 10. E. W. Arnold, S. Sundaresan, Appl. Catal., 41 (1988) 225 11.J.R. Ebner, M.R. Thompson, Catal. Today, 16 (1993) 51 12.H. Papp, F. Richter, G.-U. Wolf, Th. G5tze, Surf. Interface Anal., in preparation 13.F.R. Landsberger, P.J. Bray, J. Chem. Phys., 53 (1970) 2757 14.T. Okuhara, T. Nakama, M. Misono, Chem. Lett.,(1990) 1941 15.D. Briggs, M.P. Seah, Practical Surface Analysis, John Wiley and Sons, Chichester 1983, 512 16.T.L. Barr, J. Phys. Chem., 82 (1978) 1801 17.J.Ph. Nogier, M. Delmar, Catal. Today, 20 (1994) 109 18.F. Garbassi, J.C.J. Bart, R. Tassinari, G. Vlaic, P. Lagarde, J. Catal. 98 (1986) 271 19.U. Rodemerck, B. Kubias, H.-W. Zanthoff, M. Baerns, Appl. Catal., in press 20.R.M. Friedman, Silicates Ind., 39 (1973) 247 21.T. Okuhara, M. Misono, Catal. Today, 16 (1993) 61 22.S. Linde, L. Gorbunova, Dokl. Akad. Nauk, SSSR, 245 (1979) 584 23.V.V. Krasnikov, S.A. Konstant, Inorganic Mat. (Russ.) 15 (1979) 2164 24.J.Ziolkowski, E. Bordes, P. Courtine, J. Mol. Catal., 84 (1993) 307 25.M.A. Pepera, J.C. Callahan, M.S. Desmond, E.C. Milberger, P.R. Blum, N.J. Bremer, J. Am. Chem. Soc., 107 (1985) 4883 26. J. Haber, 8th Internat. Congr. on Catal., Berlin, 2.-6.7.84, Proc. Vol. 1, 85 27.J.L.Callahan, R.K.Grasselli, AIChE J., 9 (1963) 755 28.B. Kubias, unpublished result This work was supported by the Deutsche Forschungsgemeinschaft (project No Ku 957/1-4) and by the Bundesministerium fiir Bildung, Wissenschaft, Forschung und Technologie of the FRG (project No 03D0024).

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

471

Partial oxidation of C5 hydrocarbons to phthalic and maleic anhydrides over suboxides of vanadia" Use of dicyclopentadiene as a probe molecule U.S. Ozkan*, G. Karakas, B.T. Schilf, and S. Ang Department of Chemical Engineering, The Ohio State University, Columbus, Ohio 43210, U.S.A.

Partially reduced vanadium oxide catalysts have been examined in the selective oxidation of pentane and pentene to phthalic and maleic anhydride. The anhydride selectivity has been shown to be a strong function of catalyst pre-reduction and reaction temperatures. Controlled-atmosphere, postreduction and post-reaction surface characterization experiments have shown the most selective catalyst surface to_be comprised of V6013, V409, and VO 2 species. In this phase of the research, dicyclopentadiene has been used as a probe molecule to elucidate the reaction network for the fQrmation of phthalic and maleic anhydrides. 1. INTRODUCTION There is growing interest in the partial oxidation of the C 5 fraction of the hydrocarbon stream from naphtha steam crackers since there is no real market for them at the present time. Furthermore, the partial oxidation of lower alkanes and alkenes continues to pose challenging problems for catalysis researchers. In the case of C 5 hydrocarbon oxidation to form phthalic anhydride, the challenge is even greater since the catalyst needs not only to insert oxygen selectively, but also promote the formation of C-C bonds in an oxidative medium. In recent years, several studies have been reported in the literature, focusing on C 5 oxidation using catalysts such as supported vanadia, VPO catalysts, and molybdates [1-11] . In this study, suboxides of vanadia catalysts were used in pentane, pentene, dicyclopentadiene and cyclopentane oxidation reactions. In the previous phase of the work [12,13], the role of alkali promoters on the catalyst selectivity was examined. The catalysts were reduced in situ at different temperatures and the effect of pre-reduction temperature was investigated. C~176176 characterization of pre-reduction, post-reduction, and post-reaction catalysts were performed using X-ray diffraction, X-ray photoelectron spectroscopy, laser Raman spectroscopy and temperatureprogrammed desorption experiments. The objectives of this study were to determine the activity and selectivity of different suboxides of vanadia in

472 maleic and phthalic anhydride formation reactions and to examine the effect of alkali promoters on the product distribution. In this article, we report the results from the transient response and temperature programmed desorption experiments performed using dicyclopentadiene and phthalic anhydride as probe molecules. 2.

EXPERIMENTAL

Crystalline V20 5 was prepared by the decomposition of a m m o n i u m metavanadate (NH4VO 3 from Aldrich) with oxygen using a t e m p e r a t u r e programmed heat treatment. V20 5 was promoted with alkali promoters in the form of carbonates. The suboxides were prepared by reducing the catalysts with hydrogen in a temperature range of 300-600 ~ The catalysts were characterized by krypton BET surface area measurement, X-ray diffraction, controlled-atmospher e Raman spectroscopy, controlled-atmosphere X-ray photoelectron spectroscopy, temperature programmed reduction (TPR), and temperature programmed desorption (TPD)using n-pentane, 1-pentene, and dicyclopentadiene as adsorbates. The steady-state partial oxidation reaction experiments were carried out in a quartz fixed-bed, flow reactor with 10 mm I.D. and a length of 50 mm. The analytical system consisted of two parts, Gas Chromatography (GC) and High Performance Liquid Chromatography (HPLC). GC was used to analyze on-line the reaction products that were in the gas phase. The GC (HP-5890)was equipped with a Chromosorb PAW 23% SP-1700 column (30 i~. 1/8") for FID, a Hayesep D column (15 ft. 1/8") and a Molecular Sieve 5A column (10 i~. 1/8") connected serially to TCD. Reverse-phase HPLC with a Spherisorb 5 ODS-2 (1 ft. 1/4")column was used off-line to analyze the products that were solid or liquid at room temperature. The details of the experimental procedures have been described elsewhere [12,13]. The transient response experiments outlined in this article were performed pulsing feed mixtures of dicyclopentadiene, nitrogen, and oxygen through the catalyst bed which was connected on-line to the mass spectrometer (M.S. Engine, Hewlett Packard). The feed mixture had a composition of 0.036% dicyclopentadiene (H.P. Gas Products Inc.), 0.280% oxygen, 71.650% nitrogen and balance helium. In oxygen-free runs, oxygen was replaced by helium. The pulse size was 1 cm 3. The major ions monitored through the selected ion mode of mass spectrometer operation were 44 for CO2, 32 for 02, 18 for H20 , 16 and 29 for CH4, 54 and 98 for maleic anhydride, 104 for phthalic anhydride and phthalic acid, 105 for phthalic anhydride, 66 for dicyclopentadiene, 68 for methyl-maleic anhydride, 34 for pentadiene, 46, 88, and 116 for maleic acid. Similar experiments were performed by pulsing mixtures of phthalic anhydride through the catalyst bed. In the TPD experiments, the catalyst was reduced in situ and was degassed in vacuum. The adsorption of dicyclopentadiene was done at room temperature for 2 hours. After 30 minutes of flushing the system, the temperature program was started with a heating rate of 15~ The effluents were monitored by the quadrupole mass spectrometer using the s a m e m/e ratios listed above. In the TPD experiments, m/e=28 was used to monitor CO since there was no nitrogen present in the system.

473 3.

RESULTS AND DISCUSSION

In our previous studies [13], we have found the product distribution and the anhydride yields to be a strong function of reduction t e m p e r a t u r e in 1-pentene oxidation reaction. Table 1 s u m m a r i z e s the yields of major reaction products over the pre-reduced vanadia catalysts at different reduction and reaction temperatures. Only the results from the unpromoted catalysts are included in this table. The yield, YA for any product A is defined as YA =100 X (moles of A in the product stream x 7A)/(moles of reactant in the feed) where 7A is the ratio of the number of C atoms in product A to the n u m b e r of C atoms in the reactant molecule. Table 1 Yield of major products in 1-pentene oxidation Reduction temp. (~ Reaction temp. (~ M.A.

P.A. CO CO2 Me-M.A.

450

400

600

350

375

400

350

375

400

350

375

400

0.2 0.3 2.5 3.5 0.3

0.8 0.5 4.5 8.6 3.0

5.7 7.7 34.2 37.1 7.3

4.3 0.5 6.6 9.2 10.6

6.3 0.9 12.5 14.1 12.8

5.8 1.1 16.6 19.7 10.2

0.3 0.1 2.9 3.6 1.3

0.2 0.1 2.9 4.8 0.8

0.3 0.1 4.1 6.1 0.8

M.A.- maleic anhydride, P.A.- phthalic anhydride, Me-M.A.- methyl maleic anhydride As seen in this table, the highest anhydride yield is obtained with a prereduction t e m p e r a t u r e of 400 ~ and at a reaction t e m p e r a t u r e of 400 ~ The other reaction products are acetone, acetaldehyde and formaldehyde. W h e n the catalysts pre-reduced at 400 ~ were characterized by controlledatmosphere X-ray photoelectron spectroscopy, the surface was found to be comprised of three different suboxides. The surface composition was found to be 64% V409, 5% V6013, and 31% VO 2. While our studies for identification of the active/selective phase continues, in a parallel effort, we have conducted some experiments to assess the role of dicyclopentadiene, one of the intermediates proposed in the literature [7]. The objective behind this phase of the study was to examine the routes that originate from dicyclopentadiene and result in phthalic anhydride v e r s u s maleic anhydride formation. The catalyst used in this set of experiments w a s v a n a d i u m pentoxide t h a t was pre-reduced in situ at 400 ~ with hydrogen.

474 Figure 1 shows signals when the feed mixture consisting of dicyclopentadiene, oxygen, and nitrogen was pulsed through the catalyst bed at room temperature. What is worth mentioning about this figure is the dicyclopentadiene signal which shows an initial spike, that coincides with the nitrogen signal and a second broader feature, which exhibits a tail. This provides a clear indication of the adsorption/desorption interaction between the catalyst surface and the dicyclopentadiene molecule. The time scale between the on-set of the signal and the end point of the tail was about 3.9 minutes.

m/e=66 o

i~ 3.9 min.

"o t..Q

~1


,

0.2

0.8 0.9 C6H~2 / O 2 Ratio

0.2

C 60lefins

0.25

I

.___5_'

(1) 03 _-. . ~ .

C6H6

0.05

/~'~6He

o: - - - - ~ , , 0.5 0.(5 0.7

, 0.8 0.w Cell2 / O 2 Ratio

, 1

~

,-----" 1.1 1.2

o[ 0.5

I~4H8 0.6

"=

,-~

0.7

= 0.8

27,H~ r 019

C8H12102Ratio

1

-

1.1

1.2

Figure 4. Conversions and selectivities of cyclohexane oxidation. Up to 70% selectivity to total olefins are observed with cyclohexene, butadiene, and ethylene as the major products. At high 02 and higher temperatures ethylene and butadiene dominate, while at lower 02 and lower temperature, cyclohexene dominates. Much more cyclohexene is obtained than benzene, and much more ethylene and butadiene are produced .than propylene. to -10% 02 breakthrough in the product. The 02 conversion can be increased by using pure 02 rather than air and by preheating (although both increase the occurrence of flames), although complete 02 conversion cannot be obtained in these experiments. We have no definitive explanations for O2breakthrough from larger alkanes, but we suspect it is due to larger coverages of carbon which inhibits 02 adsorption and allows its breakthrough.

497 The most interesting feature of cyclohexane oxidation is the types of olefins which are obtained. At low temperatures and higher C6H12/O2 ratios, we observe primarily cyclic C6 olefins, as shown in figure 4b. Figure 4c shows that the C6 olefins are primarily cyclohexene with much smaller amounts of cyclohexadiene and benzene. Thus, the singly dehydrogenated olefin dominates over the multiply dehydrogenated products, even though thermodynamic equilibrium predicts benzene over cyclohexene. At higher temperatures and lower C6H12/O2 ratios, more cracked products are observed. However, as shown in Figures 5b, 5c, and 5d, the dominant products are C2 and C4 rather than C3. Examination of the C4 products shows that the dominant species is butadiene, which is formed with four times the selectivity of all butylenes. The patterns of olefins from cyclohexane thus has a distinctive characteristic which shows that the products are either cyclohexene or ethylene plus butadiene, with very little propylene. Further, the ethylene is approximately equal to the butadiene. This suggests that these products are produced in the simple reactions: cyclohexane ---) C2H4 + butadiene or cyclohexane --->cyclohexene. It is interesting that cracking splits the cyclohexane into C2 and C4 fragments rather than into two C3 fragments. If the adsorbed species were singly bonded cyclohexyl, then [3 scission should produce a secondary alkyl, and a second 13 scission would produce propylene and leave adsorbed isopropyl which should yield a second propylene molecule. The formation of ethylene and butadiene seems to argue that the intermediate which leads to these products is a di-cy adsorbed cyclohexane fragment, from which two 13 scission events would liberate ethylene and leave adsorbed C4 which could desorb as butadiene. 5. C A T A L Y T I C C O M B U S T I O N Complete oxidation of alkanes to carbon dioxide and water can also be achieved in millisecond contact time catalytic reactors by operating in the fuel lean regime. Short contact time combustion has many potential applications[ 17] including power generation in gas turbines, incineration of waste containing air streams, and the development of radiant heat sources. We have demonstrated greater than 99.5% conversion of methane, ethane, propane, and butane to complete oxidation products over platinum coated ceramic monoliths at contact times as short as 2 ms at compositions ranging from 20 to 100% excess air [ 18,19]. These experiments showed that, as long as the catalyst was in the ignited state, very high fuel conversions (>99%) were realized. It was found that the catalyst extinction temperature for methane, ethane, and propane are approximately 1150, 1050, and 800 ~ respectively, for room temperature feeds in air at a residence time of 5 ms. Extinction with these fuels occurred at 5.5, 2.2, and 1.3% fuel in air for room temperature feeds with is much leaner than the 9.5, 5.7 and 4.0 for ~=1. We have investigated the use of this type of reactor as an incinerator and found that it is possible to achieve very high destruction efficiencies by adding methane to an air stream containing volatile organic compounds and passing the resulting gas stream over an ignited monolith catalyst. Methane addition is necessary in order to increase the fuel value of the stream to maintain a sufficiently high catalyst temperature so that extinction of the catalyst does not occur. Figure 5 shows a plot of the data obtained for the incineration of toluene containing air streams at varying residence times and toluene inlet concentrations. For these experiments, methane was added to air streams containing 500-2000 ppm of toluene until the methane concentration was 6.75% (~=0.7). This mixture was then passed over an ignited catalyst at residence times ranging from 3 to 10 ms.

498

1O0

2

g "

8c-

1.5

1

Fm s f

99.98

~ o

o_--. ! O" 1200

5 ms//f 10 ms

g

99.94 .(~ ~

0.5

5"

99.92 =

8 '

500

3 ms ~..._...~.-.---~ 5 ms ~---

7 ms 99.96 ,~ o~ ~ 1000

O

0

1400

,a

?lit

_

~ I

~

I

r

lOres ~

&&

I

II

800

? '

1000 1500 2000 Inlet Concentration Toluene (ppm)

99.9

600 I

500

I

/

t

1000 1500 2000 Inlet Concentration Toluene (ppm)

Figure 5. Toluene destruction by methane oxidation in air. For 50 to 2000 ppm toluene the toluene breakthrough is less than 0.5 ppm for a destruction of greater than 99.5%. The temperature of the catalyst is - 1100oc, and the velocity before the monolith is ~ 1 rn/sec for a residence time of several milliseconds. Figure 5a shows both the concentration of toluene at the reactor outlet as well as the overall conversion of toluene to carbon dioxide and water as a function of the inlet concentration of toluene and the reactor residence time. This plot shows that the outlet concentration of toluene remains constant at -0.3 ppm regardless of residence time or inlet composition. This results in calculated toluene conversions from 99.94 to 99.99%. Any scatter in the conversion data can be attributed to experimental error since no general trend of conversion with residence time is observed. Figure 5b shows how the catalyst temperature varies with residence time and toluene composition. There is a small dependence of temperature on the toluene composition as would be expected as the fuel value of the stream increases. The figure also shows a large increase in temperature with decreasing residence time. This occurs because at higher flow rates (shorter residence times), heat losses from the reactor to the surroundings become less important. The adiabatic temperature for the compositions examined in figure 5b is approximately 1400oc. It is also important to note that theaddition of toluene allows the catalyst to operate below the extinction temperature for pure methane. We have performed similar experiments using other compounds such as chlorobenzene, acetonitrile, and thiophene in simulated VOC streams and see similar high conversions (>99%) and no apparent catalyst deactivation due to C1, S, or N. The most important variable in these experiments is temperature. High conversions are realized for all cases when the catalyst is ignited, although sufficient fuel must be supplied to the catalyst to maintain the temperature above the extinction temperature. 6. C O N C L U S I O N S Catalytic partial oxidation at very short contact times is a promising route to new chemicals and to catalytic destruction of volatile organic compounds. Conversions of the limiting reaction are >99% at residence times less than 1 millisecond, and low concentrations of undesired products are observed. The results discussed in this paper are summarized in Table 1 below.

499 Table 1 Conversions and Selectivites in Millisecond Oxidation Reactions Reaction

Selectivity

Fuel Conversion

Yield

Syngas

> 0.9

> 0.9

-- 0.9

C2H 6 tO C2H 4

_~ 40 The reaction temperature 500~ tracks the iC4H10 conversion O 30 and both decrease at the higher 20 450Cj iC4H10:O2 ratios, further away = 1 0 from the flammability limits. At ..Q o these higher iC4H10:O2 ratios, =1 i , , I , . , I , , , I , , , I , , , I, 400 0 1 1.2 1.4 1.6 1.8 2 less cracking of the iC4H 10 occurs and the selectivity to iCeHlo / 02 iC4H8 increases. However, this Figure 3: Selectivity,iC4H10 conversion, and reaction temperature for selectivity to iC4H8 is somewhat the oxidative dehydrogenation of iC4HI0 over Pt/oc-A1203 in a Pd less than the typical selectivity in membrane reactor as pictured in Figure 2 as a function of the iC4H]0:O2 either dehydrogenation in a ratio. The reaction side contained 30% N2 dilution, had a total flow rate of 1 slpm, and was maintained at a pressure of 2 psig. The sweep side membrane reactor [ 1, 3, 4] or (N2) had a total flow rate of 4 slpm and was maintained at a pressure of oxidative dehydrogenation in a 1 psig. monolithic reactor [6]. Oxidative .

-

-

i

.

.

.

i

.

.

.

!

.

.

.

i

.

.

.

i

3.1

AND

.

505

dehydrogenation in the monolithic reactor occurs at a somewhat higher temperature (800~ and at contact times two orders of magnitude shorter. At the comparably long contact times used in the current membrane reactor module (10 -1 seconds), it is likely that iC4H8 is decomposing before leaving the reactor. Second generation reactor designs will allow for shorter residence times. Figure 3 shows the selectivity to H2 split between the fraction that remains on the reaction side and the fraction that is swept a w a y after p a s s i n g through the Pd membrane. For example, at an iC4H10:O2 ratio o~" 25 of 2.0, the total selectivity to H2 >,, 20 (on an H atom basis) is 42% (38 > ,~,,.,-.-''"~ F)o = 1.45 + 4) with 10% of the total H2 O 15 transferred to the sweep. . . . . .

I

. . . .

I

. . . .

I

. . . .

I

. . . .

v

i

co

F/O = 1.00

10

3.2 T r a n s i e n t B e h a v i o r of the C a t a l y s t / M e m b r a n e ~ 5 Figure 4 shows the changes in conversion and selectivity for 0 m 0 this reactor as a function of time 0 10 20 30 40 50 ~ 7o on-line for two iC4H10:O2 (F/O) ratios each with 10% N2 dilu60 C .I'" F/O- .oo tion. For simplicity, only the 0 "~ 50 iC4H 10 conversion, reaction t emperature, and iC 4H 8 > 40 cO selectivity are shown. High F/O = 1.45 0 30 conversion and low selectivity = 20 are observed for F/O = 1.00 while low conversion and high 10 .,Q selectivity are observed for F/O 0 0 = 1.45. In both cases, the con10 20 30 40 50 0~'850 version and selectivity both inL... crease with time on-line. The re.~600 action t e m p e r a t u r e is very :3 closely linked to the conversion ~550 and it increases with time as O. well. Although there is some E ~500 change in the selectivities to each E of the products during this o450 "warm-up" stage, the amount of O H2 removed on the sweep side ~400 .... I .... I .... ! .... I .... of the module does not indicate 0 10 20 30 40 50 any loss in permeability of the Time Online ('min) Pd membrane. After about 30 Figure 4: Selectivity,iC4Hlo conversion, and reaction temperature for minutes, there is little change in the oxidative dehydrogenation of iCnHlo over Pt/t~-A1203 in a Pd c o m p o s i t i o n of the product membrane reactor as a function of time on stream. The reaction side contained 10% N2 dilution, had a total flow rate of 1 slpm, and was stream or the reaction temperamaintained at a pressure of 2 psig. The sweep side (N2) had a total flow ture. All of the data presented in rate of 4 slpm and was maintained at a pressure of 1 psig. Data is the other figures corresponds to shown for iC4H10:O2 ratios of 1.0 and 1.45. E

, .

, ,

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.

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.

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506

data collected after 48 minutes on-line. 3.3

Hydrogen Removal Trials both with and without H2 removal were conducted under the same experimental conditions in the reactor shown in Figure 2. For trials without H2 removal, the Pd foil was backed by a solid copper disk which 25 blocked H2 removal from the re.i 20action side of the module. The , I> 0 Pd foil was present in the reactor 0 -~ 1 5 to account for any catalysis that may be occurring on the surface 0 r- 1 0 of the foil. As shown in Figure 0 >., 3, the for almost all composi:~ 5 tions, --10% of the total H2 pro..Q 0 duced diffuses through the Pd -0 membrane to the sweep side of 20 10 30 the reactor. For membrane reaco~" 70 tors, this is a very small amount t-- 6 0 o of H2 removal. However, resi.m -.,-.,-i,.. 50 dence times in more traditional O > membrane reactors are typically ,-- 4 0 o much longer than the residence 0 30times used in this study. c 20Figure 5 shows the effect of H2 removal on iC4H8 selectiv..Q 10 1.015 o ity, iC4H10 conversion, and re0 action t e m p e r a t u r e for two 20 10 30 iC4H10:O2 ratios, with varying 6-650 t._,. levels of N2 dilution. In all o cases, the selectivity to iC4H8 is t,,,,. = 600 greater at the higher iC4H10:O2 I... ratios and the iC4H10 conversion 0 -550 is greater at the lower iC4H 10:02 0 ratios. At all compositions, hyt-

5]

ro 500

.i

O

n - 9450

drogen removal leads to increased iC4H8 selectivity and increased iC4H I O conversion.

Therefore, by removing as little as 10% of the produced H2, this reactor design has shifted the Figure 5: Selectivity, conversion, and reaction temperature with and p r o d u c t d i s t r i b u t i o n t o w a r d without a Pd membrane for continuous H2 removal. The striped bars olefin formation. correspond to the data for the membrane reactor (H2 removal). The solid bars correspond to data collected when the membrane is replaced by an 3.4 D i l u t i o n of t h e Reacimpermeabledisk (without H2 removal). In all cases, the total flow rate t a n t s on the reaction side was 1 slpm at a pressure of 2 psig. Data is shown for iC4H10:O2 ratios of 1.0 and 1.5 at 10%, 20%, and 30% N2 dilution. F i g u r e 5 also shows the For the trials with the Pd membrane, the flow rate on the sweep side effect of N2 dilution on reaction was 4 slpm of N2 at a pressure of 1 psig. system. In all cases, increased 10

20 30 N i t r o g e n Dilution (%)

507 level of N2 dilution decreases both the iC4H10 conversion and the iC4H8 selectivity. This is largely a temperature effect since the presence of the diluent also reduces the reaction temperature. The differences with dilution are much less apparent when additional heat is added to the system to compare dilution levels at equivalent reaction temperatures. 4.

CONCLUSIONS

The kinetics of oxidative dehydrogenation in a membrane reactor are much more favorable than the kinetics of catalytic dehydrogenation (in the absence of oxygen). The goal of this research is to balance the reaction rate and the rate of H2 removal from the reactor by adjusting the residence time and availability of oxygen to create a highly efficient membrane reactor for the production of isobutylene. We have achieved over 65% isobutane conversion with 13% selectivity to isobutylene at compositions just outside the flammability limits, using a Pd foil, radial flow membrane reactor and a Pt/~-A1203 catalyst. Selectivities were low, mainly due to the decomposition of the product, which can be attributed to the relatively long contact times (I: -- 1 sec) of the reactor. We have also shown that, although the current reactor and membrane configuration only allowed for a 10% hydrogen removal, the removal of hydrogen substantially increased both isobutane conversion and isobutylene selectivity. REFERENCES

1. 2. 3. 4. 5. 6. 7. 8. 9.

T. Matsuda, I. Koike, N. Kubo, E. Kikuchi, Appl. Catal. 96 (1993) 3-13. S. Udomsak, R. Anthony, Ind. Eng. Chem. Res. 35 (1996) 47-53. J. Deng, J. Wu, Appl. Catal. A 109 (1994) 63-76. E. Gobina, R. Hughes, J. Membr. Sci. 90 (1994) 11-19. B.A. Raich, PhD Thesis Chemical Engineering, University of Delaware, Newark 1995. M. Huff, L. D. Schmidt, J. Catal. 155 (1995) 82-94. M. Huff, L. D. Schmidt, J. Phys. Chem. 97 (1993) 11815-11822. M. Huff, L. D. Schmidt, J. Catal. 149 (1994) 127-141. H. Armendariz, G. Aguilar-Rios, P. Salas, M. A. Valenzuela, I. Schifter, H. Arriola, N. Nava, Appl. Catal. A 92 (1992) 29-38.

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3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

509

Chemoselective catalytic oxidation of polyols with dioxygen on gold supported catalysts Laura Prati and Michele Rossi Dipartimento di Chimica Inorganica Metallorganica e Analitica, UniversitY. di Milano e Centro C.N.R., via Venezian 21, 1-20133 Milano, Italy

Supported gold catalysts in the presence of dioxygen have shown high selectivity toward polyols monooxygenation. In fact, commercially useful products such as glycolic and lactic acids from, respectively, ethane-l,2-diol and propane-l,2-diol, can be advantageously obtained. Studies on the influence of the support and preparation methods are reported along with comparison tests involving commercial Pd and Pt catalysts.

1. INTRODUCTION The oxidation of the alcohols to carbonylic or carboxylic derivatives is of great practical interest in the synthesis of commercially useful products, as is evident from the abundant literature that has appeared in recent years [ 1]. Particular interest has been shown in the catalytic oxidation of polyols with dioxygen using supported platinum-group metals as the catalysts. The most studied metals have been palladium and platinum which are, however, often affected by deactivation problems [2]. The introduction of cocatalysts such as bismuth or lead represents an enhancement in the use of these catalysts having the double effect of increasing catalytic activity and improving catalyst life [3]. One important application is in the field of carbohydrate transformations: the catalytic oxidation of D-glucose to D-gluconic acid represents an economically competitive route with respect to biochemical oxidation [4]. This new process is the result of extensive studies on the selective C 1-hydroxyl group oxidation in the presence of 0 2 using a Pt or Pd catalyst modified with cocatalysts [5]. In spite of the abundant literature dealing with carbohydrate transformations, there is relatively little on lower polyols [6]. In the case of glycerol particular attention has been paid to the influence of cocatalysts, like Bi, on the selectivity of Pd and Pt catalysts, which changes the hydroxyl group oxidation from primary to secondary producing dihydroxyacetone with 70 80% selectivity [7]. Industrial and academic investigations have dealt with the catalytic oxidation of C2 and C3 diols. In ethane-l,2-diol oxidation, a few studies indicate a limited application of platinum and palladium catalysts because of the overoxidation that normally occurs with C-C bond

510 cleavage, forming CO 2 via HCOOH. In fact, the patent literature reports mostly catalytic processes on protected subtrates such as methoxy glycol and polyethylene glycol [8]. The same type of limitation was encountered in the use of a copper-zinc catalyst in oxidizing ethane-1,2diol to glycol aldehyde where high selectivity (90%) was obtained only at low conversion (23%) [91 . Other catalytic systems have been employed to achieve high selectivity with more acceptable conversion. Thus, an Ag/Si-C catalyst was used for glyoxal production with 73% selectivity at 96% conversion [10] whereas, to the best of our knowledge, the most notable result in glycolic acid production from ethane-1,2-diol was claimed by using an Ir on carbon catalyst operating at 10 atm and 80~ (87% selectivity at 98% conversion) [ 11 ]. Also in the case of the C3-diol catalytic oxidation there is a lack of literature concerning the selective oxidation of the hydroxyl groups. According to a recent patent the selective oxidation of propane-1,2-diol has been claimed for either the primary or the secondary hydroxyl group, but no examples were reported for primary ones [6b]. Palladium on carbon has been reported to oxidize propane-l,2-diol in a non selective manner [12] whereas a patent claims 86% selectivity in propane-1,3-diol oxidation to 3-hydroxy-propanoic acid using a commercial Pd/C catalyst [6a]. This paper reports our studies on the selective oxidation of ethane-l,2-diol and propane1,2-diol using new gold-based catalysts. As an extension, commercial Pd and Pt catalysts were tested under similar reaction conditions to compare their activities and selectivities in the synthesis of important commercial products, such as glycolic acid and lactic acid.

2. EXPERIMENTAL

2.1 Reagents and apparatus Ethane-1,2-diol and propane-1,2-diol were of the highest purity from Fluka and were used without any further purification. NaOH was 99.9% pure from Merck and stored under nitrogen. Gaseous oxygen from SIAD was 99.99% pure. The gold metal was of the highest purity grade from Fluka. Commercial 5% Pd/C was supplied by Montecatini Tecnologie (MPT5 catalyst), 5% Pt/C by Engelhard. Activated carbon (5-100 kt) had a specific area of 1200 m2/g,; AI20 3 (type 507C 100-125 mesh), SiO 2 (220-440 mesh) and CeO 2 (>99%) were from Fluka. Reactions were carried out at the appropiate temperature in a thermostatted glass reactor (30 cm 3) provided with an electronically controlled magnetic stirrer connected to a large reservoir (5,000 cm 3) containing oxygen at 2 atm (310 KPa). The oxygen uptake was followed by the use of a mass flow controller connected to a PC through an A/D board, plotting a flow/time diagram.

2.2 Oxidation procedure

2.2.1 Oxidation of ethane-l,2-diol The substrate (0.50 g, 8.06 mmol), NaOH (0.96 g, 24 mmol) and the catalyst ( diol/metal = 1000) were mixed in distilled water (total volume 10 ml). The reactor was pressurized at 2 atm of 0 2 and thermostatted at the appropriate temperature. The mixture was stirred and the samples analyzed at various times by HPLC and 13C-nmr.

511

2. 2.2 Oxidation of propane-1, 2-diol The substrate (0.518 g, 6.8 retool), NaOH (0.82 g, 20.4 mmol) and the catalyst ( diol/M - 100 or 1000) were mixed in distilled water (total volume 10 ml). The reactor was pressurized at 2 atm of 0 2 and thermostatted at the appropriate temperature. The mixture was stirred and the samples analyzed at various times by HPLC and/or 13C_nmr.

2.3 Catalyst preparations 1% Ir/C was prepared as previously described [ 11 ] using IrC14 solution ( Ir 44.8 mg/ml). Gold catalysts were prepared using a HAuCI4 0.1M solution obtained by dissolving 1.97 g of metal gold in a minimum amount of a 3:1 (v/v) mixture of concentrated HC1 and HNO 3 and then diluted to 100 ml with distilled water. After reduction all the catalysts were filtered and then washed with hot water to obtain chloride ion free catalysts. The catalysts were used in wet form.

2.3.1 GoM on oxides 1% Au on Al203 and SiO2 were prepared by suspending the support (2 g) in water (3 ml) and then by adding the gold solution (1 ml). The slurry was evaporated to dryness and then calcined under H 2 at 250~ for 3h. 1% Au/otFe20 3 was prepared by the coprecipitation method as previously reported [ 14 ]. 1% Au/CeO 2 was prepared using the incipient wetness impregnation method (see below). 2.3. 2 GoM on carbon a) Impregnation method Carbon (2 g) was suspended in 10 ml of distilled water and while stirring the solution of gold (1 ml) was added. The mixture was allowed to stand at room temperature for lh then heated to 70~ and reduced by adding HCHO 37% (1.5 ml) dropwise. b) Incipient wetness impregnation The support (2 g) was impregnated with 1 ml of 0.1M HAuCI 4 diluted with distilled water to a volume equal to the pore volume of carbon. The suspension was mixed for 20 min then put in a hot solution of HCOONa ( 20 ml of water and 200 mg of formate). c) Precipitation method The solution of HAuC14 0.1M (1 ml) was diluted with distilled water (10 ml) and a saturated solution of Na2CO 3 was added until pH8 was reached. To the stirred solution carbon (2 g) was added. The slurry was allowed to stand for lh then heated to 70~ and reduced by adding HCHO 37% (1.5 ml) dropwise. Alternative reductive agents used were HCOONa (200 mg) or NaH2PO 2 (200 mg).

2.4 Analysis of products The products were identified by comparison with authentic samples. Quantitative analyses were performed by either HPLC or 13C-nmr analyses, using an internal standard (propionic acid). 2.4.1 HPLC analysis Analyses were performed on a Varian 9010 instrument equipped with a Varian 9050 UV detector (210 nm) using a Merck Lichrospher 100 RP18 (5~tm) column. An aqueous nBu4NHSO 4 5mM (1 ml/min) as the eluent was used for ethane-l,2-diol reaction products

512 and NaH2PO4/H3PO 4 0.15M (pH 2.47) ( 0.7 ml/min) for propane-l,2-diol reaction products. Samples of reaction mixture (0.1 ml) were diluted (10 ml) by using the eluent after adding the standard.

2. 4. 2 13C_NMR analysis Spectra were recorded in water on a Varian 200 MHz. Samples of reaction mixture were neutralized with HC1 12N before adding the standard.

3. RESULTS AND DISCUSSION Gold is reported to weakly interact with molecular oxygen, inhibiting subsurface oxygen diffusion and some studies (surface enhanced Raman spectroscopy, microgravimetric and temperature desorption techniques) indicate that molecular oxygen adsorption is preferred to the atomic one [13]. However most theoretical studies have been conducted on smooth gold surfaces or large gold particles. On the other hand it has been reported that the surface chemistry and reactivity change drammatically for supported gold particles [ 14]. The catalytic behavior of gold has recently been reviewed by Haruta who highlights the strong dependence of gold activity on the type of support, the preparation methods and the size particles during CO oxidation, hydrocarbon combustion and other selective partial oxidation [ 15]. With the aim of exploring the activation of 0 2 toward the oxidation of vicinal diols we tested different supported-gold catalysts under mild conditions. By working in neutral aqueous solution of ethane-1,2-diol up to 100~ and 2 atm of 0 2 in the presence of 1% Au supported on active carbon there was no oxidation, whereas in alkaline solution a smooth oxygen uptake at 50-90~ was observed. HPLC and 13C-NMR analyses of the reaction products showed quite good slectivity toward monooxygenation. In fact, operating with a substrate/metal ratio of 1000 (Table 1, entry 1) very high selectivity in glycolate production was reached (90%) at very high conversion (94%). The catalyst was recycled 10 times without loss of activity, there being a slight decrease in selectivity (2-3 %). The attractive result obtained with the gold catalyst prompted us to compare its performance with more convemional catalysts (Pd/C and Pt/C) and with others suggested by previous studies, such as Ir/C [ 11 ] and Cu/C [ 16 ]. Platinum and palladium on carbon were found to be more active but less selective than gold, being affected by overoxidation (entries 2 and 3) to C1 products (HCOOH). By reducing the temperature from 70 to 50~ the selectivity improved, rising to about 70% (entries 6 and 7). Copper on carbon produced relevant quantities of formate even at low conversion (entry 4) owing to the known activation toward C-C bond cleavage [ 16]. Under our reaction conditions the Ir on carbon catalyst, prepared according to the literature [ 11], was inactive (entry 5). Thus, due to the different experimental conditions no comparison between our and the previously reported ethane-1,2-diol oxidation with Ir/C [ 11] was possible.

513 Table 1 Oxidation of ethane-1,2-diol with various catalysts entry

catalyst

t(h)

T(~

GLA

OXA

FOR

(mol )

(mol%)

(tool%)

1%Au/C

3

70

85

5%Pt/C

2

70

50

6

5%Pd/C

2

70

57

6

I%Cu/C

6

70

5.4

l%Ir/C

5

70

5%Pt/C

2

5O

68

5%Pd/C

2

50

74

.

,

.

.

0.5

.

conv.%

selec.% GLA

5

94

90

10

100

50

3

100

57

16

54

10

95

71

100

74

,

.

10.5

Reaction conditions: ethane-l,2-diol/M = 1000; NaOH 2.6M; NaOH/ethane-l,2-diol = 3 GLA-glycolate; OXA=oxalate; FOR=formate. Na2CO 3 was also formed.

The influence of the support on the gold catalyzed oxidation of ethane-l,2-diol is very relevant. In fact a comparison of the results in Table 2 shows carbon (entry 1) to be peculiar in its high activity, with respect to A120 3, CeO 2, SiO 2 and ot-Fe20 3 (entries 2-5).

Table 2 Influence of support entry

r'

catalyst

t(h)

GLA

OXA

FOR

(tool%)

(~mol%)

(mol%)

conv.%

selec.% GLA

1

I%Au/C

3

85

5

94

90

2

1%Au/Al20~

6

36

23

70

51

3

1%Au/CeO 2

6

22

6

35

63

4

l%Au/SiO 2

6

-

5

1%Au/otFe203

6

17

29

50

34

Reaction conditions: ethane- 1,2-diol/M = 1000; NaOH 2.6M; NaOH/ethane- 1,2-diol = 3; T -- 70~ GLA=glycolate; OXA=oxalate; FOR=formate. Na2CO 3 was also formed. With a view to optimizing the activity and selectivity of the Au/C catalyst, a short screening on deposition and activation methods was carried out. Comparing, in Table 3, three different deposition methods, namely alkaline precipitation (entry 1), absorption from diluted solution of HAuC14 (entry 2) and incipient wetness impregnation (entry 3), followed by reduction of gold to metal, we observed that the first method performed the best. Although all the methods show

514 similar selectivity the precipitation one had the best activity, resulting in, for the given time (2 h), the highest conversion. On the contrary the effect of the different reductive agents (HCHO, HCOONa, NaH2PO2) was negligible. In any case, the activated catalyst had always to be dechlorinated to avoid the known poisoning effect of chloride ion [ 13]. The dependence of activity on the preparation methods observed for Au/C catalysts is presently under investigation. According to the literature data, larger gold particles (10 nm) can be expected with the adsorption method while the other two methods result in smaller particles (2-5 nm) [ 13]. Thus, a linear correlation between gold dispersion and activity seems to be ruled out. Table 3 Influence of preparation methods on the activity of 1% Au/C catalyst

GLA

preparation

(mol%)

1

a

75

-

4

2

b

58

-

2

3

c

64

-

5

entry

OXA

FOR

conv.%

selec.%

(mol%)(mol%)

GLA 93

81 65

,,

70

,

89 91

a. precipitation method b. absorption from diluted solution c. incipient wetness impregnation Reaction conditions: ethane- 1,2-diol/M = 1000; NaOH 2.6M; NaOH/ethane- 1,2-diol = 3; T = 70~ t = 2 h. GLA=glycolate; OXA=oxalate; FOR=formate.

As an extension of the ethane-1,2-diol oxidation, we investigated the catalytic oxidation of propane-1,2-diol. In this case the problem of chemoselectivity, arising from the presence of a primary and a secondary alcoholic function, is of great interest as both hydroxyacetone and lactic acid are products of synthetic importance (Fig. 1) Figure 1. Oxidation products of propane- 1,2-diol

~

OH

OH

[~

OH,-~COOH

0 hydroxyacetone

OH hctic acid

By considering the results reported in Table 4 we can outline that almost total selectivity has been shown by the I%Au/C catalyst toward lactic acid production under mild conditions (70~ 2 atm of 02) (entry 1). However this result Canbe achieved using a low substrate/ metal ratio (100) because at higher values (1000) the selectivity drops to 90% (entry 2).

515 By using palladium on carbon, from previous studies on C3-diol oxidation, poor selectivity can be expected in the case of propane-l,2-diol [12] and good selectivity in the case of propane-l,3-diol (86% of 3-hydroxy-propanoate) [6a]. In our experiments, however, oxidizing propane-l,2-diol we obtained good selectivity in lactic acid using either Pd/C or Pt/C (Table 4, entries 4 and 6). In these cases, higher substrate to metal ratio (1000) produced higher selectivity. Acetate was the main byproduct and a small amount of pyruvate was observed only with Pd/C catalyst. Table 4 Oxidation of propane-1,2-diol with various catalysts entry

catalyst

S/M

t(h)

ratio

LA

AC

PYR

(mol%)

(mol%)

(mol%)

conv.%

selec.% lactic a.

1

1%Au/C

100

1

98

-

-

100

98

2

1%Au/C

1000

2

76

8

-

84

90

3

5%Pd/C

100

0.5

68

16

2

86

79

4

5%Pd/C

1000

1

81

11

1

94

86

5

5%Pt/C

100

1

90

10

-

100

90

6

5%Pt/C

1000

2

5

-

83

94

i

78

Reaction conditions: NaOH 2.6M; NaOH/propane-l,2-diol = 3; T = 70~ LA=lactate; AC=acetate; PYR=pyruvate. Among the products hydroxyacetone was not detected.

4. CONCLUSIONS As has already been noted by Haruta [14] the tunable reactivity of gold catalysts by controlling the particle size, the type of support and the different preparation methods widens the potential of such catalysts in different fields of applications. The results presented in our studies point to the synthetic use of supported gold as the catalyst in oxidation reactions of industrial interest. In fact the selective synthesis of important products like lactic and glycolic acids could make the catalytic route competitive with others currently used.

REFERENCES 1. a) R.A.Sheldon, J.K.Kochi Metal Catalyzed Oxidations of Organic Compounds, Academic Press, New York, 1981 b) R.ASheldon Heterogeneous Catalysis and fine chemicals II M.Guisnet, J.Barrault, C.Bouchoule, D.Duprez, G.Perot, M.Maurel and C.Montassier Eds., Elseviers, Amsterdam, 1991, pages 33-54 c) S.T.Oyama, J.W.Hightower Eds. Catalytic Selective Oxidation A.C.S.,Symposium Series, Washington, 1993

516 d) M.Hudlicky Oxidations in Organic Chemistry A.C.S.Monograph 186, Washington, D.C., 1990 2. a) H.E.van Dam, A.P.G.Kieboom, H.van Bekkum Appl.Catal. 33 (1987) 361 b) P.J.M.Dukgraaf, H.A.M.Duisters, B.F.M.Kuster, K.van der Wiele J.Catal. 112(1988)337 3. T.Mallat, A.Baiker Catalysis Today 19 (1994) 247 and references cited therein 4. a) K.Deller, B.Despeyroux Chem.Ind (Dekker) 47 (1992) 261 b) M.Besson, F.Lahmer, P.Gallezot, P.Fuertes, G.Fleche J.Catal. 152 (1995) 116 c) M.Wenkin, R.TouiUaux, P.Ruiz, B.Delmon, M.Devillers Appl.Cat.A:General 148 (1996) 181 5. a) H.E.van Dam, A.P.G.Kieboom, H.van Bekkum Appl.Catal. 33 (1987) 361 b) P.J.M.Dukgraaf, H.A.M.Duisters, B.F.M.Kuster, K.van der Wiele J.Catal. 112(1988)337 c) C.Broennimann, Z.Bodnar, P, Hug, T.Mallat, A.Baiker J.Catal. 150 (1994) 199 d) C.Broenniman, T.Mallat, A.Baiker J.Chem.Soc.Chem.Commun. 1377 (1995) 6. a) A.Behr, A.Botulinski, F.J.Carduck, M.Schneider U.S.Patent 5,321,156 (1994) b)H.Kimura, K.Tsuto U.S.Patent 5,274,187 (1993) 7. a)H.Kimura, K.Tsuto, T.Wakisaka, Y.Kazumi, Y.Inaya Applied Cat.A:General 96(1993)217 b)R.Garcia, M.Besson, P.Gallezot Applied Cat.A:General 127 (1995) 165 c)H.Kimura, K.Tsuto U.S.Patent 5,274,187 (1993) 8. a) M.Nozue Jpn.KOKAI TOKKIO KOHO JP 63,211,251 (1988) b)M.B.Libman, V.F.Shvets, Yu.P.Suchov Khim.Prom.-st. 9 (1988) 520 c)M.Nozue Jpn.KOKAI TOKKIO KOHO JP 04,342,559 (1992) 9. T.Seto, M.Odagiri, M.Imanari Jpn.KOKAI TOKKIO KOHO JP 03,279,342 (1991) 10.P.GaUezot, S.Tretjak, Y.Christidis, G.Mattioda, A.Schouteeten J.Catal. 142 (1993) 729 11.T.Oku, Y.Onda, H.Tsuneki, Y.Sumino Jpn.KOKAI TOKKIO KOHO JP 07,112,953 (1995) 12.T.Tsujino, S.Ohigashi, S.Sugiyama, K.Kawashiro, H.Hayashi J.Mol.Cat. 71 (1992) 25 13.D.I.Kondarides, X.E.Verykios J.Catal. 158 (1996) 363 14.M.Hamta, N.Yamada, T.Kobayashi, S.Iijima J.Catal. 115 (1989) 301 15.M.Haruta Catalysis Today, inpress 16.M.Lanfranchi, L.Prati, M.Rossi, A:Tiripicchio J. Chem. Soc.Chem.Commun. 1698 (1993)

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama,A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

517

P r o m o t i n g e f f e c t s of b i s m u t h in c a r b o n - s u p p o r t e d b i m e t a l l i c P d - B i c a t a l y s t s for the selective o x i d a t i o n of glucose to g l u c o n i c acid M. Wenkin a, C. Renard a, p. Ruiz b, B. Delmon b and M. Devillers a, * Universit6 Catholique de Louvain, a Laboratoire de Chimie Inorganique et Analytique, place Louis Pasteur, 1 b Unit6 de Catalyse et de Chimie des Mat6riaux Divis6s, place Croix du Sud, 2 B-1348 Louvain-la-Neuve, Belgium

Experiments are carried out to improve the u n d e r s t a n d i n g of the behaviour of Bi-promoted P d / C catalysts during their use in the selective oxidation of glucose to gluconic acid by 02. Supported Bi(5 wt.%)-Pd(5 wt.%)/C catalysts are prepared by deposition from a suspension of several carboxylate precursors in heptane, followed by thermal degradation under N2 at 773 K. The catalysts are characterized by XRD, XPS and SEM-EDX. Because significant amounts of bismuth are leached from the catalysts under the reaction conditions, recycling experiments are performed to evaluate the influence of this process on the catalyst lifetime. Whereas the Bi losses are essentially restricted to the first few catalytic runs, the gluconic acid yield, normalized with respect to the catalyst mass, remains constant. Catalytic tests are also conducted in the presence of diethylenetriaminepentaacetate, which is a stronger chelating agent than the gluconate ions, to remove the major part of dissolved Bi from the solution. The behaviour of the bimetallic catalyst is also compared with that of a commercial trimetallic Pd(5 wt.%)-Pt(1 wt.%)-Bi(5 wt.%)/C catalyst.

1. INTRODUCTION Bismuth is known for displaying very attractive properties as promoting element in numerous heterogeneous oxidation catalysts and namely, in its association with noble metals like palladium and platinum for the selective oxidation of alcohols or aldehydes by molecular oxygen in aqueous solutions [17]. However, the actual origin of the promoting role of Bi and the question whether Bi-Pd alloys are present and do play, or not, a significant role in these catalysts is still under discussion. In addition, the fact that significant amounts of bismuth are leached from these catalysts during their functioning remains a critical point which impedes their broader use. In previous works [8,9], bimetallic Bi-Pd catalysts supported on activated carbon and characterized by various Bi/Pd molar ratios ((Pd+Bi)=10 wt.%) were prepared from the thermal degradation of Bi and Pd acetate-type precursors under nitrogen at 773 K. Because several binary Bi-Pd alloys were heavily suspected in the supported catalysts, three intermetallic compounds, Bi2Pd, BiPd and BiPd3 were also prepared from the same precursors, according to the same

518 procedure, upon thermal heating under nitrogen in appropriate conditions described elsewhere [9]. The catalytic performances of the supported and unsupported catalysts characterized as having the same overall composition were measured and compared for the selective oxidation of glucose to gluconic acid. At constant palladium weight, Bi2Pd was found to be the most active phase, whereas the BiPd3 alloy displayed no activity. For the carbon-supported catalysts, the highest performances were observed for Bi/Pd=l. When comparing the pure alloys with the supported catalysts, the highest absolute yields in gluconic acid were therefore found for different compositions, suggesting the multiphasic nature of the supported catalysts. Furthermore, bismuth was found systematically to dissolve in the reaction medium during the catalytic tests, the losses being significantly more extensive from the monometallic Bi/C than from the bimetallic PdBi/C catalysts. Bi2Pd is the phase that loses Bi to the largest extent, while the promoting element does not dissolve from the inactive phase BiPd3. Glucose and gluconate in solution were shown to be both responsible for Bi dissolution. However, we demonstrated that there was no simple relationship between the extent of Bi losses and the performances of the supported catalysts. The present work reports on further experiments devoted to the role played by dissolved bismuth in the catalytic reaction and its consequences on the deactivation process. Particular interest will be paid to the recycling capabilities of these catalysts. Because the complexing properties of gluconate ions were proved to influence the dissolution process and are therefore suspected to modulate the catalytic behaviour, experiments aimed at trapping the solubilized fraction of the promoting element were carried out by adding a stronger complexing agent than the gluconate ion (KD [Bi2(C6H1007)2(OH)] + = 10-10 [10]). Diethylenetriaminepentaacetic acid (H5dtpa), which is known as a strong chelating agent for Bi3+ (KD [Bi(dtpa)] 2- = 10-30 [11]) was therefore added at various amounts to the reaction mixture in the presence of the bimetallic catalyst. The behaviour of the Bi-Pd/C catalysts used in the present work was also compared with that of a well-established commercial trimetallic Pd(4 wt.%)Pt(1 wt.%)Bi(5 wt.%)/C from Degussa. Because the latter was shown to loose smaller amounts of bismuth than the bimetallic Pd-Bi/C catalysts tested so far [8], this comparison was aimed at finding and understanding the relationship between the catalytic behaviour and the extent of Bi losses. In addition, scanning electron microscopy coupled with energy-dispersive energy analysis (SEM-EDX) has been implemented as a complementary analytical tool for various purposes : (i) to estimate the mean particle size of the metallic particles and look at the eventual influence of the used precursors on these characteristics ; (ii) to investigate more deeply the composition and dispersion anomalies detected by XPS on certain catalysts; (iii) to find experimental evidence for bismuth redeposition on the catalyst surface after use.

519 2. EXPERIMENTAL 2.1. Preparation of the catalysts Carbon-supported (Activated carbon SXplus supplied by NORIT, SBET = 750 m2.g -1, particle size : 0.2-0.1mm) bimetallic and monometallic catalysts were prepared by deposition from a suspension of carboxylate particles in n-heptane chosen as inert organic solvent. Precursors used for the incorporation of the metals were either, palladium(II) acetate (ACROS) and bismuth(III) oxoacetate, BiO(O2CCH3) (synthesized as described elsewhere [8]), or diammine(pyrazine-2,3dicarboxylato-N,O)palladium(II) [12] and tris(monohydrogenopyrazine-2,3dicarboxylato)bismuth(III) (noted Bi(2,3-pzdcH) 3) [13]. Monometallic Pd(5 wt.%)/C and Bi(5 wt.%)/C catalysts : the selected carboxylate precursor of palladium or bismuth (see above) was dispersed in the presence of 7.2 g of the activated carbon in 250 ml n-heptane under ultrasonic stirring for 30 min. The hydrocarbon was evaporated very slowly at room temperature under vacuum and the precursor deposited on the support was subsequently decomposed upon heating under nitrogen at 773 K during 18h. Bimetallic Pd(5 wt.%)Bi(5 wt.%)/C catalyst : The n o n - d e g r a d e d monometallic P d / C catalyst was dispersed in n-heptane under ultrasonic stirring for 30 min in the presence of the bismuth carboxylate containing the same ligand as the precursor Pd compound. After slow evaporation of the hydrocarbon, the bimetallic catalyst was activated upon thermal heating under nitrogen at 773 K d u r i n g 18h. D e p e n d i n g on the p r e c u r s o r used, acetate (Ac) or pyrazinedicarboxylate (pzdc), these catalysts were noted (Ac.PdBi/C) and (pzdc.PdBi/C). The pure BiPd3 alloy was prepared from the acetate precursors according to the deposition procedure described above. The carboxylates were decomposed upon thermal heating under nitrogen at 1173 K during 18h. 2.2. Catalytic measurements 2.2.1. Reaction conditions and analysis of the reaction products The selective oxidation of D-glucose into gluconic acid was selected as catalytic test reaction. The reactor vessel and the experimental conditions were described in detail elsewhere [8]. The pH of the reaction mixture was kept at a constant value in the range 9.25-9.45 by adding a 20 wt.% aqueous solution of sodium hydroxide with an automatic titrator (Star Titrino 718) from METROHM. The base consumption was recorded as a function of time. The glucose solution (72 g glucose in 400 ml water) was heated in the reactor to 50~ Once the temperature was stabilized, the catalyst (mcat = 54-200 mg) was added to the solution and the oxidation reaction started by introducing oxygen (flow rate : 0.4 1.min -1) in the stirred (1000 rpm) slurry. Measurements performed with 50 to 100 mg catalyst at different stirring rates (in the range 10001800 rpm) confirmed the absence of diffusional limitations under these conditions. After 4 hours reaction, the oxygen inlet was turned off and the

520 catalyst was removed from the reaction mixture by filtration. The catalyst was washed with water, ethanol and ether and dried under vacuum at 30~ before being analyzed by XPS, XRD and SEM-EDX. For the recycling experiments, the catalysts were dried under vacuum at room temperature before being reused in a further catalytic test with 400 ml of a fresh 1 mol.1-1 glucose solution. The composition of the reaction mixture was determined by HPLC and 13C-NMR spectroscopy. The bismuth and palladium losses from the catalysts in the reaction mixture during the catalytic tests were determined by analyzing the collected filtrates by atomic absorption spectrometry. Analytical conditions were described elsewhere [8].

2.2.2. Expression of the catalytic results Because the 13C-NMR analyses showed that gluconic acid was the only carboxylic acid generated in the reaction medium, the yields in gluconic acid (YGLU, %) were calculated directly from the NaOH consumption. The main side product is fructose due to isomerisation in the presence of oxygen and appears at an extent between 2.6 and 4.6 % yield when YGLU is larger than 10 %.

2.3. Catalyst characterization techniques 2.3.1. X-ray diffractometry (XRD) Powder X-ray diffraction patterns were obtained with a SIEMENS D-5000 diffractometer using the Ks-radiation of a copper anode. The samples were analyzed after deposition on a quartz monocrystal sample-holder supplied by Siemens. The crystalline phases were identified by reference to the ASTM data files.

2.3.2. X-ray induced photoelectron spectroscopy (XPS) X-ray photoelectron spectroscopy was performed on a SSI-X-probe (SSX100/206) spectrometer from FISONS, using the AI-K~ radiation (E = 1486.6 eV). The energy scale was calibrated by taking the Au 4f7/2 binding energy at 84 eV. The C ls binding energy of contamination carbon fixed at 284.8 eV was used as internal standard value. The analysis of bismuth and palladium were based on the Bi 4f7/2 and Pd 3d5/2 photopeaks. The intensity ratios I(Bi4f7/2)/I(Bi4fs/2) and I(Pd3d5/2)/I(Pd3d3/2) were fixed at 1.33 and 1.5 respectively.

2.3.3. Electron microscopy Scanning electron microscopy was carried out on a LEICA stereoscope $260 equipped for energy dispersive X-ray analysis (EDAX 9100). Samples are deposited on a Cu-AI support from a slurry in acetone.

521

3. RESULTS 3.1. Characterization of the supported bimetallic catalysts The XRD characterization results of this type of catalysts have already been discussed in a previous paper [8]. The main observations consisted in the presence of metallic palladium in the monometallic P d / C catalyst, and the obtention of poorly resolved XRD spectra for the bimetallic catalysts, in which the formation of one of several binary BixPdy alloys was nevertheless suspected. Representative XPS results are listed in table 1 for fresh and used catalysts. Table 1 XPS atomic intensity ratios in fresh and used catalysts Catalyst (a)

P d / C (%)

Bi/Pd (b)

Bi/C (%) (b)

Other data

(Ac.Pd/C)f

0.69

-

-

(Ac.Pd/C)tl

0.65

-

-

(Ac.Pd/C)tl - 5g H5dtpa

0.65

-

-

(Ac.PdBi/C)f

0.72

1.10

0.80

(Ac.PdBi/C)tl

1.12

0.50

0.56

(Ac.PdBi/C)tl - 5g H5dtpa

0.77

0.39

0.30

(Ac. PdBi / C)t5

1.19

0.37

0.44

(Ac.PdBi/C)t13

0.90

0.37

0.33

(PdPtBi / C-Degussa)f

1.06

1.92

2.05

P t / C = 0.24 %

(PdPtBi/C-Degussa)tl

1.30

1.00

1.30

P t / C = 0.12 %

N / C = 0.4 %

N / C = 0.5 %

P t / C = 0.15 % (PdPtBi/C-Degussa)t5 1.02 0.90 0.92 (a) f: fresh catalyst tn: catalyst engaged in n consecutive tests (b) theoretical(bulk) values : Bi/Pd = 0.52 ; Pd/C (xl00) = 0.62 ; Bi/C (xl00) = 0.32 X P S : Bismuth and palladium appear in the metallic (Pd 0, Bi 0) and the oxidized form (Pd 2+, Bi3+). The binding energy values associated with the Bi 4f7/2 photopeak lie in the range 157.7-158.7 eV for Bi0 and 159.1-160.2 eV for Bi3+ and, for the Pd3d5/2 line, in the range 335.6-335.9 for Pd 0 and 337.3-338.0 eV for Pd 2+. The experimental B i / P d ratio in the fresh bimetallic catalyst is higher than the theoretical value calculated from the bulk composition of the catalyst, indicating a partial coverage of palladium by bismuth. This observation is in agreement with the sequential incorporation of Pd first, then Bi, during the preparation of this catalyst, but also with the propensity to segregation in Pd-Bi alloys. The experimental values for the atomic ratios P d / C and Bi/C in the fresh catalysts are higher than the theoretical values, indicating that a d e q u a t e d i s p e r s i o n is achieved for both metals on the surface. The B i / P d molar ratio in the used bimetallic catalysts decreases to reach the value of 0.5 after one run. This value

522 was already observed in previous experiments with bimetallic catalysts of several Bi-Pd compositions [9]. This value decreases slightly further up to about 0.4 after a series of 5 consecutive tests and remains constant afterwards. This decrease in the Bi/Pd but also in the Bi/C ratio after the catalytic tests is in line with the bismuth losses previously observed during the catalytic oxidation of D-glucose in Dgluconic acid. The observation of nitrogen compounds on the surface of the catalysts tested in presence of H5dtpa and the unusually intense lines in the Cls spectral region of those catalysts indicate that this complexing agent is adsorbed on the surface of the catalyst. Electron microscopy : The use of pyrazinedicarboxylate-type precursors was found to generate the largest metal particles with a maximum size of 20 ~tm, whereas the catalysts prepared from acetate precursors are so small that they are not detected by backscattered electron microscopy. EDX analysis of the (pzdc.PdBi/C) catalyst demonstrated the bimetallic nature of the metal particles. The SEM-EDX spectra of the used monometallic Ac.Pd/C catalyst tested in the presence of added dissolved bismuth, according to the procedure described elsewhere [8], gave evidence for bismuth deposition on the support and on Pd particles. 3.2. Catalytic results 3.2.1. Influence of dissolved bismuth Preliminary experiments described elsewhere allowed us to check the absence of any catalytic activity in the presence of the support alone, but also in the presence of the monometallic Bi/C catalyst [8]. The experiments performed in this work (fig. 1) indicate that the catalytic performances of the inactive phase BiPd3 are enhanced in the presence of a monometallic catalyst Ac.Bi(10 wt.% Bi)/C (after 4 h reaction, YGLU (50 mg BiPd3) = 2 %, to compare with YGLU (50 mg BiPd3 + 54 mg Ac.Bi/C) = 9.6 %). The same observation was made when two separate monometallic catalysts Ac.Pd/C and Ac.Bi/C were engaged simultaneously [8]. As 70 to 80 % of the initial bismuth content was found to dissolve from this monometallic catalyst in the reaction m e d i u m during the catalytic tests, the presence of this dissolved bismuth in solution could be responsible for the activity increase. Under conditions that need to be investigated further, the presence of bismuth in solution is therefore a sufficient condition to improve the catalytic activity of a monometallic catalyst P d / C , in disagreement with previous observations [8]. Complementary experiments in the presence of a chelating agent (Hsdtpa) were conducted to improve the understanding of this peculiar point. The gluconic acid yields obtained after 4h and the bismuth losses are listed in table 2. A decrease of the gluconic acid yield was observed when the mass of engaged H sdtpa in the catalytic operations was globally increased. This means that complexation with dtpa induces a decrease in activity because it removes complexation of dissolved Bi3+ by the gluconate ions generated by the catalytic

523 reaction. However, the observed loss of activity could also be assigned to a poisoning effect of the catalyst by the complexing agent itself. This is supported by the deactivation of a monometallic catalyst Ac.Pd/C in the presence of the same complexing agent (YGLU (Ac.Pd/C) = 11.9 %, to compare with YGLU (Ac.Pd/C + 5g H5dtpa) = 3.4 %) and also by the XPS results showing that nitrogen is present on the catalysts after these experiments. Further experiments should therefore be performed in the presence of an insoluble complexing agent able to trap the dissolved bismuth and to remove it physically from the solution, in order to avoid poisoning effects due to adsorption phenomena.

40

~Cy-0 30-

20-

10-

o

.~

,~

.................

,? .. .1. .1 "~

0

b

j

-

9............ ._

!

!

c _

Is Ii

d

_

I

I

I

!

1

2

3

4

time (h)

Figure 1 : Yield in gluconic acid vs reaction time in the presence of a) a mixture of two monometallic catalysts (Ac.Pd/C + Ac.Bi/C) b) a mixture of an inactive intermetallic phase and a monometallic catalyst (BiPd3 and Ac.Bi/C) c) the monometallic catalyst (Ac.Pd/C) and d) the intermetallic phase BiPd3 Table 2 Influence of the mass of H5dtpa engaged on the catalytic performances and Bi losses of the bimetallic catalyst Ac.PdBi/C m H5dtpa (mg)

YGLU (%)

Bi losses (%)

0

45.1

12

50

26.5

18

100

24.6

15

500

20.4

15

5000

10.9

23

524

3.2.2. Recycling Two series of tests were carried out to investigate the deactivation process of the bimetallic catalyst Ac.PdBi/C; 200 and 100 mg of the bimetallic catalyst were respectively engaged in 13 and 5 successive tests. Because the absolute amounts of catalyst engaged in consecutive tests was regularly decreasing, gluconic acid yields were normalized with respect to the catalyst mass. The absolute and normalized gluconic acid yields and the bismuth losses after each catalytic operation are listed in tables 3 and 4. When normalized with respect to the mass of engaged catalyst, the gluconic acid yield remains constant during the successive operations while significant Bi losses are observed only after the first one or two experiments. This is in agreement with the fact that constant Bi/Pd atomic ratios were observed by XPS on the surface of the bimetallic catalyst after 5 to 13 successive experiments (see Table 1). Table 3 Catalytic performances and Bi losses of the bimetallic catalyst Ac.PdBi/C engaged in 13 successive tests (180-200 mg catalyst - 50~ - 1000 rpm - pH 9.25-9.45 - 4h diffusional limitations) Test

YGLU(%)

YGLU/mcat (%.mg -1)

Bi losses (%)

1

89.4

0.45

23 + 3

2

79.7

0.40

97%). The reaction

551

may be catalytic (Run 12-Table 3) but the conversion is improved under stoichiometric conditions (Run 11-Table 3). Unexpectedly, the catalytic performances are much better than with epoxy-l,2-octane [39]. With transition metal-based systems, the reaction is slower. Using a solution of Co2(CO)8 in methanol, the isomerization of the epoxy-alcohol proceeds under mild conditions (35~ The reaction is catalytic but not complete (conversion c a . 50%). As for LiBr-HMPA (1:1), the keto-alcohol is obtained with good selectivity (Run 13). The conversion is lower than for epoxy- 1,2-octane. Table 3 Li-, Co- or Ni- based systems for catalysed ring opening reactions of 1,2-epoxy-alcohols, 3a, 3b and 3e; a two-step synthesis of keto-alcohols, 2a, 2b and 2c. Keto-alcohol fields Run

T (~

Time (h)

11

85

2

12

85

2

13

35

24

14

50

46

15

50

46

16

50

46

17

50

20

18

85

2

19

35

24

20

35

24

21

85

2

Precursor [M']" (mol.1-1)

Epoxyalcohol (mol.l1)

LiBr/HMPA 0.75 0.75 LiBr/HMPA 0.75 0.75 Co2(CO)8 0.05 NiBre(PPh3)2/Zn/PPh3 0.01 0.1 0.01 NiBr2(PPh3)JZn/PPh3 0.016 0.160.016 NiBr2(PPh3)2/Zn/TBA 0.016 0.160.016 NiBr2(PPh3)JZn/PPh3/ 0.016 0.16 0.016 TBA 0.016 LiBr/HMPA 0.75 0.75 Co2(CO)8 0.05 Co2(CO)8 0.05 LiBr/HMPA 0.75 0.75

3a (0.75) 3a (4.5) 3a (1.52) 3a (0.12) 3a (0.38) 3a (0.16) 3a (0.16)

3b (4.5) 31a (1.54) 3c (1.5) 3c (4.5)

Solvent (cm 3)

Conversion b

EpoxySelectivity c alcohol/[M'] (%)

Toluene 10 Toluene 10 MeOH 1.5 CH2C12 4 THF 3 THF 3 THF 3

100

1

2a (98)

85

5

2a (98)

50

15

2a (99)

37

-

43

10

16

2

24

2

la (97) 2a (3) 2a (96) la(4) 2a (96) la(2) 2a (96) la(3)

Toluene 10 MeOH 1.5 MeOH 1.5 Toluene 10

83

5

2b (99)

44

14

2b (98)

48

14

2c (98)

82

5

2e (97)

a [M]

----lithium, cobalt or nickel precursors; b0~ of substrate consumed; c 100 (mol 2/mol 3 converted); b'cproducts analysed by GC (OV-1701) and GC-MS coupling;TBA---tributylamine.

There is a dramatic solvent effect with the nickel-based system (Runs 14-15). A tetrahydrofuran (THF) solution of the nickel-based catalyst appears to be less active than the

552 dicobaltoctacarbonyl complex. NiBr2(PPh3)2/PPh3/Zn (1:1:10) (referred to as Ni/P/Zn) catalyses isomerization of 3a with a TON similar to those obtained by Miyashita et al. [21 ] for epoxy-propane isomerization. Good selectivity for the keto-alcohol, similar to that observed for the above systems is obtained. The same catalytic performances were observed for epoxy1,2-octane. On the other hand, with a non-coordinating solvent such as CH2C12, la is "regenerated". When PPh3 is replaced by tributylamine (TBA), the selectivity is unchanged but the reaction rate decreases and the catalyst deactivates quickly. NMR studies show that, in contrast to PPh3, TBA is not bound to the nickel centre. TBA is not convenient for building a stable active complex and, when used as adduct, it does not significantly modify the catalytic performances as it does with palladium catalysts in the Heck reaction [40]. Trialkyl phosphine is more promising and is under study in our group as well as novel heterogeneous systems. Epoxides 3b and 3c are selectively isomerized by LiBr/HMPA in a short time (Runs 18, 21).

2.5.2. EPR study of the nickel-based system In order to characterize the active site of the homogeneous nickel-based catalyst (oxidation state, composition of the coordination sphere), the reaction was monitored by EPR. Before addition of the epoxide The catalytic solution is obtained by mixing NiIIBr2(PPh3)2 (0.016 mol.r 1) with a reducing agent (Zn) in the presence of 1-1.6.10 -2 mol.1-] additional PPh3 which may also act as a reducing agent, as already described for palladium-based catalysts [40] or supported nickel complexes [41]. The best catalytic performances are obtained when the solution is prepared before adding the epoxide, suggesting that the active species is generated by the reduction of Ni II to a Ni 0 or Ni I complex. EPR is expected to be able to discriminate between these various oxidation states. Most Ni II complexes, especially low symmetry ones [42], are EPR-silent; only perfect octahedral Ni II complexes or Ni II ions with an isotropic environment exhibit an isotropic signal near g = 2.17 [43,44]. Ni 0 which is diamagnetic, is also EPR-silent except in metallic particles, which exhibit a broad ferromagnetic resonance signal [45]. EPR is however a suitable technique for characterizing Ni I. At room temperature, before addition of the epoxide, the catalytic solution exhibits an isotropic EPR signal A at giso = 2.21 (Figure 2a). However, a broad axial signal B located at g_L = 2.40 > gll = 2.07 > ge (Figure 2b) is observed when the EPR spectrum of the same solution is recorded at 77K. Both signals are due to the same complex because (i.e. the average g value expected from signal B). The initial giso = (2g• NiIIBr2(PPh3)2/THF solution presents a very broad EPR signal, which is ten to twenty times less intense than signals A and B. The g tensor values of signals A and B lie in the range expected for Ni I d 9 ions [46-48]. Our findings suggest that the Ni I complex is monomeric and rule out the formation of a ligandbridged bimetallic species [49] [the half-field region (g = 4) is EPR-silent; no broadening of the EPR signal related to spin-spin interaction was observed]. Many Ni I complex solutions show such an evolution of their EPR spectrum with temperature [50-52]. The relative order of the g tensor components of signal B suggests a dz2 ground state for the single electron, and consequently a distorted tetrahedral [53] or trigonal-bipyramidal symmetry [54,55] for the corresponding Ni I complex. Unfortunately, because of the poorly

553 resolved superhyperfine structure (more apparent on the parallel component), no doubt resulting from the interaction of the unpaired electron with 31p (I = 1/2), 79Br or 81Br ( I - 3/2) nuclei of the Ni I ligands, the numbers of each ligand cannot be deduced. Nevertheless, signal B is very similar to that reported by Gmznyh et al. [56] for NilCI[P(n-C4H9)3]3. We therefore propose to assign signal B to NiIBr[PPh3]3. The observed g value shift of signal B compared to that of NilBr[P(n-C4H9)3]3 agrees with the decrease of ligand field strength expected on replacing P(n-C4H9)3 by PPh3 and/or C1 by Br. In order to measure the amount of Ni II reduced to Ni I in the presence of Zn, the intensities of signals A and B were estimated using a copper-based standard [57]. The number of spins measured at 77K for a sample containing initially 6.6 10-3 mol Ni II is (4.7 + 1.0) 10-3. This means that (70 + 14)%, i.e. most of the initial Ni II, is reduced to NiI.

L ~ ~2H~ _

DPPH J 100G

g|so

~-

IDPPH [g//=.2 07

=2.21

Figure 2a

Figure 2b

These findings suggest that NiIBr(PPh3)3 is the catalytic centre and explain why additional equimolar PPh3 is required: it replaces one Br atom to form the active complex. After addition of the epoxide The conversion never reaches 100% whatever the reaction time, suggesting deactivation of the catalyst. Indeed, after addition of the epoxide, the intensity of signal B decreases with time. Similarly, the initial red-orange colour fades with time and vanishes at the end of substrate conversion. This colour may be attributed to the active complex. The catalytic site proposed above involves an oxophilic Ni I centre. Deactivation might arise from oxidation of this active complex and simultaneous epoxide deoxygenation yielding the small amount of allylic tertiary alcohol la observed (Table 3). Toluene or dichloromethane (Run 14) solutions of NilIBr2(PPh3)2 do not exhibit A or B

type EPR signals in the presence of Zn. We propose that if Ni II is reduced to Ni I, under such conditions, it cannot be stabilized. This could explain the low activity of these samples for epoxide isomerization and show once more the influence of the solvent on catalyst performance. Our findings suggest that NilBr(PPh3)3 is the active site, in contrast to Miyashita et al. [21 ] who proposed a Ni ~ active site.

554 2.5.3. Silica-supported Ni I complexes The catalytic properties of the homogeneous catalyst were compared with those of similar well defined supported Ni I complexes. In order to confirm the nickel oxidation state in the active entity and to study the influence of the composition of its coordination sphere [coordination number, nature of the trialkylphosphine (TAP) ligands] on the catalytic performances, we have used a set of silicasupported unsaturated Ni I complexes [58] with a given number of TAP ligands of different electronic and steric properties [59,60]. NiI(pPh3)2(O)2 (O = silica surface oxygen atom), whose coordination sphere may be similar to those of the homogeneous epoxide isomerization catalysts, led to the conversion of 15% of epoxy-l,2-octane to 2-octanone after a 24 h reaction at 50~ In contrast, NiI(PEt3)(O)2 and NiI(p(c-C6Hll)3)(O)2 showed no ability to isomerize epoxy-alcohol 3a. The immediate colour change observed after epoxide addition suggests the formation of Ni II by Ni I oxidation in the presence of epoxide. The latter catalysts are more oxophilic than NiI(pPh3)2(O)2 because of their high unsaturation and because PEt3 and P(c-C6Hll)3 are more electron-donating than PPh3. These findings are in agreement with the results of Miyashita et al. [21 ]: they found that the deoxygenation reaction becomes predominant when the c~-basicity of the trialkylphosphines bound to nickel is increased. In any case, these preliminary results obtained with NiI(pPh3)2(O)2 are promising; they are consistent with a Nikbased active species in the homogeneous catalytic system and encourage us to pursue this approach to improving catalyst performance.

3. CONCLUSION The advantage of these systems over traditional catalytic ketonization routes lies in the selectivities and conversions; the two-step reaction path is very convenient from a preparative point of view. Work is continuing to control these systems better, and to investigate their mechanisms and their synthetic utility. Novel and highly efficient ways of isomerizing epoxyalcohols by careful control of the oxophilic character of the active species are sought. ACKNOWLEDGMENTS We thank Dr. J.S. Lomas for constructive discussions and for correcting the manuscript. REFERENCES 1. A.B. Smith, P.A. Levenberg, P. J. Jerris, R. M. Scarborough and P. M. Wovkulich, J. Am. Chem. Soc., 103 (1981) 1501. 2. G. B0chi and H. W0est, J. Am. Chem. Soc., 100 (1978) 294. 3. W. Adam, M. Braun, A. Griesbeck, V. Lucchini, E. Staab and B. Will, J. Am. Chem. Soc., 111 (1989) 203. 4. A.L. Baumstark and P. C. Vasquez, J. Org. Chem., 53 (1988) 3437. 5. M.S. Newman and V. Lee, J. Org. Chem., 40 (1975) 381. 6. F. Derdar, J. Martin, C. Martin, J.-M. Br6geault and J. Mercier, J. Organometal. Chem., 338 (1988) C21. 7. R.E. Parker and N. S. Isaacs, Chem. Rev., 59 (1959) 737.

555

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10. 11. 12. 13. 14. 15. 16. 17. 18. 19. 20. 21. 22. 23. 24. 25. 26. 27. 28. 29. 30. 31. 32.

33. 34. 35. 36. 37. 38. 39.

R. J. Gritter, in Chemistry of Functional Groups, S. Patai Ed., Wiley, New York, 3 (1967) 400. A. S. Rao, S. K. Paknikak and J. G. Kirtane, Tetrahedron, 39 (1983) 2323. Organic Reactions, Wiley, New-York, 29 (1983) 345. V. N. Yandovskii, B. A. Ershov, Russ. Chem. Rev., 41 (1972) 403. B. Rickbom and R.M. Gerkin, J. Am. Chem. Soc., 90 (1968) 4193. ibid, 93 (1971) 1693. G. Magnbusson and S. Thoren, J. Org. Chem., 38 (1973) 1380. H. Alper, D. Des Roches, T. Durst and R. Legault, J. Org. Chem., 41 (1976) 3611. D. Milstein, O. Buchman, J. Blum, J. Org. Chem., 42 (1977) 2299. S. Kulasegaram and R. J. Kulawiec, J. Org. Chem., 59 (1994) 7195 and references therein. J. H. Kim and R. J. Kulawiec, J. Org. Chem., 61 (1996) 7656 and references therein. A. Cabrera, F. Mathe, Y. Castanet, A. Mortreux and F. Petit, J. Mol. Cat., 64 (1991) Lll. J. L. Eisenmann, J. Org. Chem., 27 (1962) 2706. J. Prandi, J. L. Namy, G. Menoret and H. B. Kagan, J. Organometal. Chem., 285 (1985) 449. A. Miyashita, T. Shimada, A. Sugawara and H.Nohira, Chem. Let., (1986) 1323. Japan Kokai Tokkyo Koho, JP 60023376, A2 850205, Showa to Mitsui Toatsu Chemicals, Inc., Japan ; CA 103 : 6207. W. Prandtl and L. Z. Hess, Z. Anorg. Chem., 82 (1913) 103. C. N. Caughlan, H. M. Smith and K. Watenpaugh, Inorg. Chem., 5 (1966) 2131. L. G. Hubert-Pfalzgraf, Chem. Rev., 90 (1990) 969. J.-M. Brrgeault, B. E1 Ali, J. Mercier, J. Martin and C. Martin, C.R. Acad. Sci. Paris, 307, Serie II (1988) 2011. J.-M. Bregeault, B. El Ali, J. Mercier, J. Martin and C. Martin, C.R. Acad. Sci. Paris, 309, Serie II (1989)459. A. Atlamsani, J.-M. Bregeault and M. Ziyad, J. Org. Chem., 58 (1993) 5663. W. A. Herrmann, R. W. Fischer and D. W. Marz, Angew. Chem. Int. Ed. Engl., 30 (1991) 1638. A. M. AI-Ajlouni and J. H. Espenson, J. Org. Chem., 61 (1996) 3969. J.-M. Bregeault et al., unpublished results. J.-M. Bregeault, R. Thouvenot, S. Zoughebi, L. Salles, A. Atlamsani, E. Duprey, C. Aubry, F. Robert and G. Chottard, in New Developments in Selective Oxidation II, V. Cortes Corberb.n and S. Vie Bellrn Eds., Elsevier Sci. Publ., Amsterdam (1994) 571. I. V. Kozhevnikov, Catal. Rev.-Sci. Eng., 37 (1995) 311. C. L. Hill and C. M. Prosser-McCartha, Coord. Chem. Rev., 143 (1995) 407. L. Salles, C. Aubry, R. Thouvenot, F. Robert, C. Doremieux-Morin, G. Chottard, H. Ledon, Y. Jeannin and J.-M. Bregeault, Inorg. Chem., 38 (1994) 871. A. Atlamsani, E. Pedraza, C. Potvin, E. Duprey, O. Mohammedi and J.-M. Bregeault, C.R. Acad. Sci. Paris, 317 (1993) 757. R. Grigg, R. Hayes and A. Sweeney, J. Chem. Soc., Chem. Comm., (1971) 1248. G. Adams, C. Bibby and R. Grigg, J. Chem. Soc., Chem. Comm., (1972) 491. F. Ziani-Derdar, Thesis, Univ. Pierre et Marie Curie, Paris, France (1988).

556 40. 41. 42. 43. 44. 45. 46. 47. 48. 49. 50. 51. 52. 53. 54. 55. 56. 57. 58. 59. 60.

C. Amatore, E. Carrd, A. Jutand, M. A. M'Barki and G. Meyer, Organometallics, 14 (1995) 5605. C. Lepetit, Thesis, Univ. Pierre et Marie Curie, Paris, France (1987). A. Bencini and D. Gatteschi, in Transition Metal Chemistry, G. A. Melson and B. N. Figgis Eds, Marcel Dekker, New York, 8 (1982) 1. J. E. Bolton and J. R. Wertz, Electron Spin Resonance, Me Graw-Hill, New York, (1972) 291. L. Bonneviot, M. Che, K. Dyrek, R. Sch611ner et G. Wen&, J. Phys. Chem., 90 (1986) 2379. L. Bonneviot, M. Che, D. Olivier, G. A. Martin and E. Freund, J. Phys. Chem., 90 (1986)2112. K. Nag, A. Chakravorty, Coord. Chem. Rev., 33 (1980) 87. C. Lepetit, M. Kermarec and D. Olivier, J. Mol. Catal., 51 (1989) 73. J. G. Van Ommen, J. G.M. Van Rens and P. J. Gellings, J. Mol. Catal., 13 (1981) 313. M. F. Rettig, E. A. Kirk and P. M. Maitlis, J. Organometal. Chem., 111 (1976) 113. M. J. Nilges, E. K. Barefield, R. L. Belford and P. H. Davis, J. Am. Chem. Sot., 99 (1977) 755. E. K. Barefield, D. A. Krost, D. S. Edwards, D. G. Van Derveer, R. L. Trytko and S. P. O' Rear, J. Am. Chem. Sot., 103 (1981) 6219. G. A. Bowmaker, P. D. W. Boyd and G. K. Campbell, Inorg. Chem., 22 (1983) 1208. reference 43, p. 324 M. Kermarec, D. Olivier, M. Richard, M. Che and F. Bozon-Verduraz, J. Phys. Chem., 86 (1982) 2818. R. Barbucci, A. Bertcini and D. Gatteschi, Inorg. Chem., 16 (1977) 2117. B. A. Gruznyh, B. B. Saraev, F. K. Schmidt, G. M. Warin and E. H. Siedyh, Koord. Khim., 9 (1983) 1400. K. Dyrek, A. Rokosz and A. Medej, Appl. Magn. Reson., 6 (1994) 309. F. X. Cai, C. Lepetit, M. Kermarec and D. Olivier, J. Mol. Catal., 43 (1987) 93. M. Kermarec, C. Lepetit, F. X. Cai and D. Olivier, J. Chem. Sot., Faraday Trans. I, 85 (1989) 1991. C. Lepetit, M. Kermarec and D. Olivier, J. Mol. Cat., 51 (1989) 95.

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

557

M e t a l - c a t a l y z e d oxidations with alkyl hydroperoxides: a c o m p a r i s o n b e t w e e n

tert-butyl

h y d r o p e r o x i d e and p i n a n e h y d r o p e r o x i d e

H.E.B. Lempers and R.A. Sheldon Laboratory of Organic Chemistry and Catalysis, Delft University of Technology, Julianalaan 136, 2628 BL, Delft, The Netherlands

The oxidizing capacities of the bulky pinane hydroperoxide (PHP) and

tert-butyl

hydroperoxide (TBHP) were compared in oxidations catalyzed by Mo, V, Se, Os and Ru. The general trend is that metals which react via an oxometal pathway show a very similar behaviour using these two hydroperoxides as oxidant, e.g. the selenium catalyzed allylic oxidation of olefins to the corresponding c~, [3-unsaturated alcohols. Reactions which involve a peroxometal pathway, e.g. the molybdenum catalyzed epoxidations of olefins, show a completely different behaviour using these two hydroperoxides, namely virtually no reaction is observed with the bulky PHP. We conclude that PHP is a suitable mechanistic probe for distinguishing between oxometal and peroxometal pathways in catalytic oxidation.

1. INTRODUCTION In homogeneous metal catalyzed oxidations with alkyl hydroperoxides two different reaction mechanisms can be distinguished: the peroxometal- and oxometal mechanism [1] (Scheme 1). Peroxometal mechanisms are favoured when the metal in its highest oxidation state is both a Lewis acid and a weak oxidant, e.g. Mo(VI), W(VI) or Ti(IV) [2,3]. Polar solvents, particularly alcohols greatly retard the reaction by competing with the hydroperoxide for coordination sites on the metal [4]. Substituent effects on the olefin are consistent with the active alkyl peroxometal species being electrophilic in nature [5]. Oxometal mechanisms are favoured with later and/or first row transition elements having a high oxidation potential, e.g. Os(VIII), Ru(VIII) and Cr(VI) [2]. Generally, polar solvents

558 have a much less negative effect than is observed for peroxometal catalyzed reactions. For example, osmium catalyzed dihydroxylations are camed out in tert-butanol as solvent [6].

-HX

M_O2 R

S

r_

MOR + SO

PEROXOMETAL PATHWAY MX + RO2H

~-

ROH

M--O

S

MX + SO

I

X OXOMETAL PATHWAY Scheme 1

A borderline case is V(V) which has an intermediate oxidation potential [3] and can react via both mechanisms depending on the substrate used. It catalyzes epoxidations with TBHP via the peroxometal mechanism and alcohol oxidations via the oxometal pathway [7]. In this present investigation we used pinane hydroperoxide (PHP) as a mechanistic probe to distinguish between peroxometal and oxometal pathways. In the peroxometal pathway the bulky pinane group is present in the active oxidant while in the oxometal pathway it is not. Hence, in the case of the peroxometal mechanism one might expect more steric constraints and consequently a slower reaction using the bulky PHP compared to the much less bulky TBHP [8]. In the case of oxometal mechanisms the difference between PHP and TBHP should be much smaller as the alkyl group is not present in the active oxidant. PHP is also of industrial interest [9]. Catalytic hydrogenation affords pinanol which is pyrolized to linalool. The latter is used in the fragrance industry and in the synthesis of vitamin E. Hence we were interested in utilizing the active oxygen in PHP to perform useful oxidations.

559

OH OH OOH TBHP

PHP 1

2

3

OH

HO

J J7

~ovi 6 ~

OH

2. RESULTS AND DISCUSSION 2.1. Molybdenum catalyzed epoxidation Previous studies indicated that the structure of the alkyl hydroperoxide in molybdenum catalyzed epoxidations has only a minor effect on the rate and selectivity [ 10]. Hence, we were initially surprised to observe that PHP failed to give the expected epoxidation of cyclohexene (1) and limonene (2) in the presence of a molybdenum catalyst (Table1). Epoxidation of limonene with TBHP as oxidant, in contrast, gave the epoxide of the more highly substituted double bond in 84% selectivity, consistent with nucleophilic attack of the olefin on the alkylperoxomolybdenum(VI) [3,5]. We tentatively concluded that this low reactivity of PHP is a result of steric hindrance in the putative alkylperoxomolybdenum(VI) intermediate. This prompted us to carry out a systematic investigation [8] of steric effects of the alkyl substituents in the alkyl hydroperoxide on the rate of molybdenum catalyzed epoxidations. Molybdenum-catalyzed epoxidations of cyclohexene under pseudo first-order reaction conditions showed the highest rate, of the investigated tertiary alkyl hydroperoxides (Figure 1), with TBHP.

560 Table 1. Molybdenum catalyzed a epoxidation with TBHP

Successive substitution

and PHP as oxidant

of the methyl groups in TBHP by higher

substrate

oxidant

conv. (%)

sel. (%)b

1

TBHP

89

96

reaction rate. Substitution

2

TBHP

81

84

of all three methyls by

1

PHP

0

2

PHP

4

alkyl groups resulted in a steadily decreasing

three ethyl groups 0

resulted in a 99% decrease in reaction

a Conditions: 10 mmol substrate, 10 mmol oxidant,

rate compared to TBHP

0.1 mmol Mo(CO)6, 1 g internal standard was heated

and a further increase

for 5 h at 80~ in 10 ml chlorobenzene

in the steric bulk of the

b selectivity to epoxide

alkyl group as in PHP resulted in a complete

loss of epoxidation activity. These results are rationalized on the basis of a peroxometal mechanism involving nucleophilic attack of the olefin on an alkylperoxomolybdenum(VI) intermediate [8]. Bulky substituents at the o~ position in the alkyl hydroperoxide seriously impede the approach of the olefin to the O - - O bond.

lOO O

,~==.(

9

~8 o 9 '--2.*"

'

~.,

0 9

I"1

(C3H7)(CH3)2COOH

. ~

(C2Hs)2(CH3)COOH

A

i

10

I

I

I

10

20

30

time (min)

Figure 1

(CH3)3COOH

(C2H5)3COO H

561 2.2. Vanadium catalyzed oxidations

Similar results were observed with vanadium catalyzed epoxidation of cyclohexene (1) and limonene (2) with TBHP and PHP (Table 2). Epoxidations with TBHP gave reasonable conversions and Table 2. Vanadium catalyzed a epoxidation with TBHP

selectivities while

and PHP as oxidant

PHP gave virtually no reaction, consistent with

substrate

oxidant

conv. (%)

selec. (%)b

a peroxometal mechanism [ 11, 12].

1

TBHP

66

34

2

TBHP

60

70

catalyzed epoxidation of

1

PHP

6

7

cyclohexenol (3) and

2

PHP

10

0

carveol (4) gave high

3

TBHP

89

98

substrate conversions

4

TB HP

100

80 c

and high epoxide

3

PHP

61

100

selectivities both with

4

PHP

100

85 c

PHP and TBHP [13].

5

TBHP

100

> 95

These results can be

6

TBHP

99

> 95

rationalized by assuming

7

TBHP

99

> 95

that efficient coordination

5

PHP

0

6

PHP

57

> 95

of the allylic alcohol to

7

PHP

88

> 95

the alkylperoxovana-

In contrast, the vanadium

of the hydroxyl group

dium(V) forces the a Conditions: 10 mmol substrate, 10 mmol oxidant,

oxidant and substrate

0.1 mmol VO(acac)2 and 1 g internal standard were heated

into close proximity, thus

for 5h at 80~

facilitating intramolecu-

b selectivity to o~,13-epoxy alcohol.

lar oxygen transfer to the

c the rest is o~,13-unsaturated ketone.

double bond. The electron rich allylic

alcohol (7) exhibited high conversions with both hydroperoxides. When the electron density of the substrate (6) was decreased we observed no effect when TBHP was used, while with PHP a decrease in conversion was observed from, 88% to 57%. A further decrease in electron density of the double bond, as in allyl alcohol (5) resulted in a complete loss of activity with PHP while with TBHP no effect was observed. Hence, the observed differences between

562 cyclohexenol epoxidation with TBHP and PHP are a reflection of the intermediate reactivity of this olefin. The double bond of carveol, on the other hand, is more substituted and shows no difference in reactivity with these two alkyl hydroperoxides. 2.3 Selenium catalyzed oxidation SeO2 displays a unique mode of interaction with olefins, involving an initial ene reaction followed by a [2,3] sigmatropic rearrangement to a Se(II) species [14] (Scheme 2). H ene

[2,31

H O • e

+ TBHP ~ - TBA

O

H

+ SeO 2

Scheme 2 The resulting Se(II) is reoxidized by the TBHP. Hence, selenium-catalyzed allylic oxidations with RO2H involve an oxometal pathway and, assuming that this is the rate-limiting step, the rate should not significantly be influenced by the structure of RO2H. This is indeed the case: using 0.05 equivalent of SeO2 and 1.5 equivalent of oxidant at room temperature smooth allylic oxidation was observed with both TBHP and PHP (Table 3). The reactivity order for Table 3. Selenium catalyzed a allylic oxidation

selenium catalyzed oxida-

substrate

tion of allylic C - - H

oxidant

conv. (%)

sel. (%)b

groups is CH2>CH3>CH 8

TBHP

100

96

[ 15, 16]. This order of

9

TBHP

68

92

reactivity was observed

8

PHP

100

99

in the selenium-catalyzed

9

PHP

91

95

allylic oxidation of ~-pinene (8) and

a Conditions: 10 mmol substrate, 15 mmol oxidant, 0.2 mmol SeO2 and 1 g internal standard were stirred

2-carene (9) which both

for 24 h at room temperature.

dation. We studied the

b selectivity to the cx,13-unsaturated alcohol

selenium-catalyzed

showed mainly CH 2 oxi-

allylic oxidation of 13-pinene with TBHP and PHP in more detail, using an olefin/alkyl hydroperoxide ratio of 10. Under these conditions we observed zero order kinetics rather than the expected pseudo-first

563 order reaction. The selenium catalyzed allylic oxidation of 13-pinene gave similar reaction rates with PHP (4.8 910 -4 mol-1-1 .min -1 ) [ 17] and TBHP (5.0.10-4 mol.1-1 .min -1 ) which is consistent with the idea that the reoxidation of Se(II) by RO2H is not the rate-limiting step. Synthetic applications of the SeO2 allylic oxidation have revealed that the reaction gives predominantly E-allylic alcohol products [ 18]. These results parallel the geometric preference shown by [3,3] sigmatropic rearrangements of the Cope and Claisen variety and apparently arise from the most favourable transition state geometry being a pseudo chair cyclohexane with the largest groups in the equatorial position [ 19, 20]. This similar geometric selectivity in the SeO2 oxidation of olefins indicates an analogous steric effect might be operating in the [2,3] pseudo cyclopentane transition state [21]. This preference for the E-allylic alcohol product for selenium catalyzed allylic oxidation was also investigated with PHP as oxidant. Allylic oxidation of a mixture of cis/trans dupical (10) yielded with both TBHP and PHP more than 95% of the E allylic alcohol (Figure 2), further demonstrating that selenium catalyzed oxidations are independent of the oxidant used.

~

O

I0

O

fO SeO 2

TBHP/PHP

OH

~21

10

13 C (ppm) substrate product

1

2

3

115.3 114.6 120.1

147.4 146.7 151.5

35.t3 37.8 76.13

Figure 2 2.4 Ruthenium and osmium catalyzed oxidation

Two other examples of metal-catalyzed oxidations with RO2H which involve an oxometal pathway are the OsO4 catalyzed dihydroxylation of olefins [22] and the RuC13 catalyzed oxidation of alcohols [23]. OsO4 catalyzed dihydroxylation is believed to involve a 2 + 2 cycloaddition of oxoosmium(VIII) to the olefinic double bond followed by

564 rearrangement to a cyclic osmate(VI) ester (Scheme 3). Reaction of the latter with RO2H/H 2 0 affords the diol product with regeneration of OsO4.

t)

+ OsO4

o os (v

O3Os/O~

o 2.

NO

+.

o.o HO

-t- OsO4 +

ROH

Scheme 3

This reoxidation of the relatively substitution inert osmium(VI) ester to a substitution labile osmium (VIII) ester is the rate-limiting step [22]. This step would be expected to be slower with the more sterically demanding PHP than with TBHP. This is indeed what we observed: OsO4-catalyzed dihydroxylation of cyclohexene (1, Table 4) gave a faster reaction with TBHP than with PHP, the final conversion being reached in 6h and 24h respectively. Table 4. Ruthenium a and Osmium b catalyzed oxidations with TBHP and PHP as oxidant substrate

catalyst

oxidant

conv. (%)

sel. (%)

4

RuC13

TBHP

57

72

4

RuC13

PHP

33

75

1

OsO4

TBHP

100

78

1

OsO4

PHP

73

65

a Conditions: 10 mmol substrate, 10 mmol oxidant, 0.05 mmol RuC13 and 1 g internal standard were stirred for 24 h at room temperature. b Conditions: 10 mmol olefin, 16 mmol oxidant, 20 ml tert-butanol, 0.75 m120% aqueous tetraethyl ammonium hydroxide and 0.005 mmol OsO4 were stirred for 24 h at 0~ RuC13-catalyzed alcohol oxidation of carveol (4, Table 4) similarly showed a higher rate with TBHP than with PHP, suggesting that the rate-limiting step in ruthenium catalyzed oxidation of alcohols may involve reaction of a ruthenium alkoxide with RO2H, resulting in formation of the carbonyl compound with simultaneous reoxidation of the ruthenium (Scheme 4 ).

565

/OH

n+

Ru--~- O

N

+/ck H

/ 0-

/c\ H

Rn+OH RO 2H

O.

\

X

~-

H

Ru/n+

n+

Ru~O

+

H20

"OH \ + /C=O

Scheme 4 3. Conclusions

In conclusion, we believe that a comparison of the oxidizing capacities of PHP and TBHP in metal-catalyzed oxidations can provide valuable mechanistic insights. When the reaction involves rate-limiting oxygen transfer from a peroxometal species to the substrate, e.g. in Moand V-catalyzed epoxidations the bulky PHP is not reactive. The steric constraints are less of a problem, however, when coordination of the substrate, e.g. in V-catalyzed epoxidation of allylic alcohols, provides for an intramolecular oxygen transfer. When the reaction involves reaction of an oxometal species with the substrate as the rate-limiting step little or no difference is observed. Intermediate reactivities may be observed when reoxidation of the catalyst by RO2H to the active oxometal species, is the rate-limiting step. Hence, we conclude that PHP is a suitable probe for distinguishing between alternative mechanistic pathways in catalytic oxidations with RO2H. Acknowledgement: We wish to thank the Netherlands Institute for Research on Catalysis (NIOK) for financial support and Quest International for supplying us with pinane hydroperoxide. 4. References

1

R.A. Sheldon and J. Dakka, Catal. Today, 19 (1994) 215.

2

R.A. Sheldon and J.K. Kochi, Metal-catalyzed Oxidations of Organic Compounds, Academic press, New York, 1981.

3

R.A. Sheldon, J.A. van Doom, C.W.A. Schram and A.J. de Jong, J. Catal., 31 (1973) 438.

566 R.A. Sheldon and J.A. van Doom, J. Catal., 31 (1973) 427. M.N. Sheng and J. G. Zajacek, J. Org. Chem., 35 (1970) 1839. K.B. Sharpless and K. Akashi, J. Am. Chem. Soc., 98 (1976) 1986. R.B. Dehnel and G.H. Witham, J. Chem. Soc., Perkin Trans. I, 4 (1979) 935. H.E.B. Lempers, M.J. van Crey and R.A. Sheldon, Recl. Trav. chim. Pays-Bas, 115 (1996) 542. Ullman's Encyclopedia of Industrial Chemistry, 4th ed., VCH Weinheim (Germany), 20 (1981) 199. 10

R.A. Sheldon, Applied Homogeneous Catalysis, Vol. 1 (W.A. Herrmann and B. Cornils, Eds.) VCH Weinheim (Germany) pp. 411-423.

11

H. Mimoun, M. Mignard, P. Brechot and L.Saussine, J. Am. Chem. Soc., 108 (1986) 3711.

12 13

A.O. Chong and K.B. Sharpless, J. Org. Chem., 42 (1977) 1587. H.E.B. Lempers, A. Ripoll~s i Garcia and R.A. Sheldon, to be submitted to J. Am. Chem. Soc.

14

M.A. Warpehoski, B. Chabaud and K.B. Sharpless, J. Org. Chem., 47 (1982) 2897.

15

A. Guillemonat, Ann. Chim., 11 (1939) 143.

16

U.T. Bhalerao and H. Rapoport, J. Am. Chem. Soc., 93 (1971) 4835.

17

Unpublished results, H.E.B. Lempers and R.A. Sheldon.

18

J.J. Plattner, U.T. Bhalerao and H. Rapoport, J. Am. Chem. Soc., 91 (1969) 4933.

19

D.J. Faulkner and M.R. Peterson, Tetrahedron Lett., (1969) 3243.

20

C.L. Perrin and D.J. Faulkner, Tetrahedron Lett., (1969) 2783.

21

D. Arigoni, A. Vasella, K.B. Sharpless and H.P. Jensen, J. Am. Chem. Soc., 95 (1973) 7917.

22

K.B. Sharpless, A.Y. Teranishi and J-E. B~ickvall, J. Am. Chem. Soc., 99 (1977) 3120.

23

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24

E. Erdik and D.S. Matteson, J. Org. Chem., 54 (1989) 2742.

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

567

On the Way to Redox-Molecular Sieves as Multifunctional Solid Catalysts for the One-Step Conversion of Olefins to Aldehydes or Ketones M. van Klaveren and R. A. Sheldon Delft University of Technology, Laboratory of Organic Chemistry and Catalysis, Julianalaan 136, 2628 BL Delft, The Netherlands Ti-A1-Beta was used as a bifunctional catalyst in the one-step conversion of 2,3-dimethyl-2butene to 3,3-dimethyl-2-butanone. The Ti-AI-Beta catalyzed epoxidation of the olefin with aqueous H202 to the corresponding diol 2,3-dimethyl-2,3-butanediol using different solvents was studied. Good yields and high H202 selectivities were obtained when using polar aprotic solvents, such as dioxane or diglyme. Furthermore, the rearrangement of 2,3-dimethyl-2,3epoxybutane to 3,3-dimethyl-2-butanone was studied. Results obtained from these separate reaction steps were used in the one-step conversion.

1. Introduction Since the pioneering work of Enichem workers on titanium silicalite (TS-1)[1] there has been great interest in the use of redox-molecular sieves as oxidation catalysts for the synthesis of fine chemicals. I2'31 TS-1 is a highly efficient catalyst for the oxidation of a large number of substrates such as amines, m alcohols, tSl paraffins t6~ or olefins I71 with H202 (30 wt% aqueous) under mild reaction conditions. The high reactivity of TS-1 as an epoxidation catalyst has been attributed to the greater hydrophobic character of the internal surface compared to other molecular sieves, allowing a more facile diffusion of the olefin into the pore channels. R--CHO

H202 R--CH=CH2

~

O / \ R--CH--CH2

[Ti(IV)]

R'OH ~ H20

OR' OH I I R--CH--CH2 (R ' = H or alkyl)

R--CH2CHO + R--C(O)CH3 Scheme 1

568 However, the use of TS-1 is limited due to its rather small pore size, i.e. 5.6 x 5.3 A, [7'8] SO that only linear olefins are efficiently epoxidized. For this reason, a large pore molecular sieve, i.e. Ti-A1-Beta, was developed by Camblor et al. tg'~~ Ti-AI-Beta was much more active than TS-1 for the epoxidation of branched olefins, due to its larger pore size (7.6 x 6.4 ,~).[11,12] However, owing to the Bronsted acidity (due to the presence of framework AI) of Ti-A1-Beta, the main products were those resulting from acid-catalyzed ring-opening reactions, i.e. diol and diol monoethers, or rearrangement products (Scheme 1). t12'~3'~4~Sato et al. showed that these ring-opening reactions are suppressed by neutralization of the Bronsted acid sites by ionexchange with Li+ or Na+.t~4j In contrast to TS-1, the presence of A1 and a large amount of SiOH groups give the pore channels ofTi-A1-Beta a more hydrophilic character. Interestingly, in the case of TS-1 a large solvent effect was found during the epoxidation of propylene, i.e. protic polar solvents such as MeOH largely enhanced the reactivity. ~15'16jIt was suggested that this is due to formation of a hydroperoxo species I, in which the protic alcohol ROH (R = alkyl) coordinates to the Ti site and stabilizes the complex via additional H bonding (Scheme 2). [17]

SiO

O

SiO

.O_--H

\/

O~O~H -

SiO~Ti~ SiO

SiO/'

/ R

I

--_~'N? - H

II Scheme 2

An analogous solvent effect was observed in the Ti-A1-Beta catalyzed epoxidation of 1hexene with H202, with an even higher reactivity being observed in MeCN than in MeOH. rlsl This enhanced reactivity was explained by the formation of Ti-peroxo species II in which water, instead of an alcohol, coordinates to the Ti site (Scheme 2). Furthermore, MeCN is able to coordinate to the Bronsted acidic sites of the zeolite, thus preventing ring-opening reactions &the epoxide to glycol and glycol derivatives (Scheme 1). tlsl Recently, advantage was taken of the fact that several Ti-molecular sieves contain both Lewis as well as Bronsted acidic properties. For example, Corma et al. showed that Ti-MCM41 and Ti-A1-Beta are able to catalyze the epoxidation of linalool in the presence of TBHP followed by ring-closure to cyclic hydroxy ethers in a one-pot reaction, t191 Neri et al. found that TS-1 catalyzes the epoxidation of styrene, with aqueous H202 in MeOH, with in-situ rearrangement to the corresponding aldehyde in 75% yield within 4h. [2~ Based on the above mentioned literature examples, we were especially interested in the use of bifunctional catalysts based on redox-molecular sieves, such as Ti-A1-Beta, 111'12] TAPSO [2H

569 or Ti-MCM-41, [22'23] in the one-step conversion of an olefin to the corresponding aldehyde or ketone. Such processes have industrial potential based on the simplicity of a one-pot procedure, its salt-flee nature and the facile recovery and recycling of the zeolite. In this way several pharmaceutical, flavor and fragrance intermediates can, in principle, be prepared from cheap raw materials. In this paper we report on the use of Ti-Al-Beta as a bifunctional catalyst in the one-step conversion of an olefin to an aldehyde or ketone. As a model system the Ti-AlBeta catalyzed oxidation of 2,3-dimethyl-2-butene 1 with H202 (30 wt% aqueous) to 2,3dimethyl-2,3-epoxybutane 2 (or 2,3-dimethyl-2,3-butanediol, pinacol, 3), followed by subsequent rearrangement to 3,3-dimethyl-2-butanone 4 (pinacolone) was studied (Scheme 3). 0 ~_~

[Ti(IV)]u202 ~

HO

OH

~--~---~

1

0 -H20~

2

3

[[

4

Scheme 3 Initially, the separate reaction steps were investigated, i.e. the epoxidation of olefin 1 via epoxide 2 to diol 3 (Section 2.1.), as well as the rearrangement of epoxide 2 to ketone 4 (Section 2.2.). Results obtained from these investigations were used in the one-step conversion of the olefin 1 to the ketone 4 (Section 2.3.). Research is currently in progress to broaden the scope to other substrates, such as styrene, ~-methylstyrene or ~-pinene as well as the use of other redox-molecular sieves such as TAPSO t21] or Ti-MCM-4 1.t22'23] 2. R e s u l t s and D i s c u s s i o n 2.1. Epoxidation of olefin to diol

The Ti-AI-Beta catalyzed epoxidation of 2,3-dimethyl-2-butene 1 with H 2 0 2 (30 wt% aqueous; molar ratio 1 : H 2 0 2 - 2.4 : 1), via the epoxide 2 to pinacol 3 (Scheme 3), using different types of solvents, i.e. methyl-tert-butylether (MTBE), acetone, t-BuOH, i-PrOH, diethyleneglycoldimethylether (diglyme) and 1,4-dioxane was studied in detail (Table 1). It is important to note that no epoxide 2 was detected by GC during the reaction, indicating that 2 is converted immediately to pinacol 3 over the strong Bronsted acidic sites of Ti-A1Beta. Furthermore, the results presented in Figure 1 and Table 1 show that the nature of the solvent has a large influence on the reactivity of the epoxidation, resulting in different yields of pinacol 3 aider 24h. These solvent differences can be explained by competitive sorption of both solvent and reagents in the zeolite pore channels. MTBE and acetone gave a very low yield of pinacol 3 after 24h of 8% and 21%, respectively. The low reactivity in the case of MTBE can be explained by mass transfer problems due to a liquid two-phase system. In the case of the alcoholic solvents t-BuOH and i-PrOH an increased reactivity was found, which resulted in a

570 better yield of 3 atter 24h of 48% and 66%, respectively. This higher reactivity when using protic polar solvents, which is in agreement with literature reports, tlsl is explained by preferential formation of complex I as the catalytic active species (Scheme 2). However, we found a rather low H202 selectivity in the case of both solvents. Table 1. Epoxidation of 2,3-dimethyl-2-butene 1 to pinacol 3. a) Entry

Solvent

H__2_O_Q2 conv. (%)b) (at t = 6h)

Yield of 3 (%)c) (at t = 6h)

H202 select, d) (at t = 6h)

TON c) (at t = 6h)

1.

M T B E f'g)

-

6

-

42

2. 3. 4. 5. 6.

acetone g) t-BuOH i-PrOH diglyme dioxane

18 44 40 19 25

10 13 28 19 25

56 30 70 100 100

68 90 197 133 172

a) 2,3-dimethyl-2-butene 1 (3.24 g; 38.5 mmol), Ti-AI-Beta (50 mg; 0.02276 mmol Ti), H202 (30 wt% aqueous; 1.8 g; 15.88 mmol), internal standard 1,3,5-tri-tert-butylbenzene, 65 ~ 20 mL of solvent, one-pot procedure, b) Determined by iodometric titration, c) Determined by capillary GC; based on added amount of H202. d) Yield of 3/H202 -conv. x 100. e) mmol of 3/mmol Ti. f) two-phase system; H202 conversion not determined, g) 50 ~ reaction temperature. lOO o

90

;~

80

+

dioxane +

diglyme

---tr-t_BuOH

~ acetone

= MTBE

"-- i-PrOH

70 60 50 40 30 20 10 I

I

10

20

Figure 1

!

Time (h)

571 Interestingly, diglyme and even better dioxane, resulted in a reasonable yield after 24h of 49% and 62%, respectively, as well as a 100% H202 selectivity, i.e. no other products were identified by GC beside the desired product. No reaction took place under analogous experimental conditions without catalyst. From the results presented above can be concluded that the Ti-A1-Beta catalyzed epoxidation of olefin 1 to diol 3 can best be performed in a polar aprotic solvent, such as dioxane or diglyme, resulting in both a good yield of 3 as well as a high H202 selectivity. 2.2. Rearrangement of epoxide to ketone The direct rearrangement of 2,3-dimethyl-2,3-epoxybutane 2 to pinacolone 4 using several catalysts, i.e. TS-1 (1.13 wt% Ti; Si/Ti = 75.2), t241Al-free Ti-Beta (1.7 wt% Ti; Si/Ti = 59), I251 Beta (Si/Al = 12.5), [261 amorphous silica-alumina and Ti-AI-Beta (2.18 wt% Ti; Si/Ti = 44.3; Si/Al = 72.1) ll~ in benzene (50 mL) at 45 ~ was investigated in detail (Scheme 3). The results are presented in Table 2. Table 2. Rearrangement of epoxide 2 to pinacolone 4. a) Entry 1. 2. 3. 4. 5.

Catalyst

Time (h) b)

Conv of 2 (%)c)

Yield of 4 (%)~

TS-1 Ti-Beta Beta Silica-Alumina Ti-Al-Beta

0-24 0-24 4 4 4

0 0 100 100 100

0 0 49 a) 52 a) 39 a)

a) Epoxide 2 (5 mmol), catalyst (80 mg; 0.0364 mmol Ti; 16 wt%), benzene (50 mL), internal standard 1,3,5-tri-tert-butylbenzene, 45 ~ one-pot procedure, b) Time at which a sample is taken from the reaction mixture, c) Determined by capillary GC. d) Side products (identified by GC-MS) were pinacol 3 and allylic alcohol CH2=C(Me)C(Me)2OH. With TS-1 (entry 1) and N-free Ti-Beta (entry 2) no reaction was observed, even after 24h. However, Beta (entry 3), amorphous silica-alumina (entry 4) as well as Ti-A1-Beta (entry 5) gave a moderate yield (39-52%) of 4. Besides pinacolone 4, pinacol 3 and allylic alcohol CHz=C(Me)C(Me)2OH were also formed. Ongoing studies show that the allylic alcohol can be further dehydrated to the corresponding diene in the presence of the acidic sites of a zeolite. [2vl Because the above mentioned experiments gave pinacol 3 in addition to pinacolone 4, attention was focused on the rearrangement of 3 to 4 over the various catalysts. In order to promote the rearrangement, water was removed from the reaction medium using a Dean-Stark apparatus. The results are presented in Table 3. TS-1 (entry 1) and Al-free Ti-Beta (entry 2) again gave no product formation within 24h. Beta gave after 24h a 90% conversion of 3 and a 61% yield of 4 (entry 3). Amorphous silica-alumina (entry 4) resulted in a rather poor conversion of 3 (26%) and yield of 4 (5%). Interestingly, Ti-AI-Beta gave after 4h a 100%

572 conversion of 3 and a 62% yield of 4, indicating that the Bronsted acidic properties of Ti-AlBeta are sufficient to catalyze the rearrangement of pinacol to pinacolone (entry 5). Table 3. Rearrangement of pinacol 3 to pinacolone 4. a~ Catalyst

Time (h) b)

TS-1 Ti-Beta Beta Silica-Alumina Ti-Al-Beta

0-24 0-24 24 24 4

Entry 1. 2. 3. 4. 5.

Conv. of 3 (%)c~ Yield of 4 (%)c~ 0 0 90 26 100

0 0 61 5 d)

62 d>

a) Pinacol 3 (5 mmol), catalyst (80 mg; 0.0364 mmol Ti; 16 wt%), benzene (50 mL), 80 ~ internal standard 1,3,5-tri-tert-butyl-benzene, Dean-Stark apparatus, b) Time at which a sample is taken from the reaction mixture, c) Determined by capillary GC. d) Side product (identified by GC-MS) was allylic alcohol CH2=C(Me)C(Me)2OH. From Table 2 (entry 5) and Table 3 (entry 5) it is apparent that the Ti-Al-Beta catalyzed pinacol rearrangement results in a higher yield of pinacolone 4 than the direct rearrangement starting from the epoxide. Consequently, in order to promote the rearrangement of epoxide 2 to pinacolone 4 via pinacol 3, 2 equivalents of water (relative to the amount of epoxide) were added at the start of the reaction (Scheme 4; conditions A). Furthermore, after formation of the diol a Dean-Stark apparatus was used to remove the excess water (conditions B). O

Y
0.05 bar) a sharp decrease in CH3OH conversion and in the production of H2 and CO2 is observed. In the Po2 range between 0.1 and 0.2 bar no significant change was observed. Interestingly, if the 02 partial pressure is then decreased it was found that the catalyst does not recover the initial conversion and H2 and CO2 production. This irreversible decrease in activity suggests that at 488 K and Po2 higher than 0.05 bar the catalyst oxidizes, and this might have been responsible for the large decrease in methanol conversion and H2 and CO2 production. The unreduced catalyst, i.e., copper oxide (CuO/ZnO), showed similar results with very low methanol conversion ( -15.2 (D LL. -15.4 0 -15.6

t"

J 9 -15.8 -16.0 -16.2 20

10

30

40

Time / sec Figure 2. First order plot for reduction of OFervTDCSPP (from lxl 0 -6 mol dm "3FemTDCSPP) by 5x10 5 mol dm 3 Acid Orange 12, pH 6.93; la = 0.05 tool dm 3 ; 30 ~ C 3.0 2.5 ,~ ,',2,

2.0 1.5

J

A 1.0 .5

0.0

X

J

0.0

1.0

2.0

3.0

4.0

5.0

[Acid Orange 12] / 10.5 mol dm 3 Figure 3. Dependence of pseudo-first order rate constant for reduction of OFervTDCSPP by Acid Orange 12 on [Dye], pH 6.93; ~t = 0.05 mol dm'3; 30 ~ C 2.4. The effect of substituents, on the phenyl group of Acid Orange 12, on the second order rate constants of dye oxidation Linear kobsvs dye concentration plots, equivalent to Figure 2, and hence k 2 values were also obtained for the seven derivatives of Acid Orange 12 with substituents on the m e t a or p a r a position of the phenyl ring (Table 1). A Hammett analysis of the rate data reveals a good correlation of the log k2 values against ~ with a p value of-1.66 (R = 0,981) (Figure 4), whereas

658 the corresponding plot using o gives a poorer correlation (p = -2.06 with R = 0.936). Unlike previous studies on the oxidation of phenols with cationic oxoiron(IV) and oxomanganese(IV) porphyfins, ~5'16with the dye oxidations there is no significant correlation with o', suggesting that any radical character in the latter transition state is small and dominated by the polar substituent effect.

2.0 1"5t 1.0"

,~ 0.5, 0.0. -0.5.

4-COMe

-1.0. -1.5

'

0

I

-05

'

'

0'0

,

0'5

10

o+

Figure 4. Correlation of log k 2 v s . G+ for the oxidation of 1-phenylazo-2-naphthol-6-sulfonate dyes by OFervTDCSPP 2.5. Kinetic isotope effect studies on the oxidation of l-(4-methylphenylaz~)-2-naphthol-6sulfonate in deuterated buffer The rate of oxidation of 1-(4-methylphenylazo)-2-naphthol-6-sulfonate by 1 was followed as described above using D20 in place of water.When the trace amounts of HzO, arising from the buffer salts and other proton sources, had been taken into account the substrate was calculated to be >99% in the deuterated form. Under these conditions the pD value of the reaction mixture (obtained by adding 0.38 to the value obtained with the pH meter)26 was 7.50. This increase from the value of 6.93 obtained in 1-I20 arises from an increase in pKa of the buffeting salts in DzO) 8 The pI~ of the dye (10.71 in water at ionic strength 0.05 mol dm3) also increases (calculated value 11.14 in I)20) 27 this compensates for the change from pH 6.93 to pD 7.50 and as a result there is a negligible effect on the degee of ionisation of the dye. The reaction between I and the deuterated dye was shown to be first order in both the dye and the oxoiron(IV) species with a second order rate constant, after correcting for the selfreaction of 1, of 1829 + 49 dm 3 mol "~s~. This leads to a kinetic isotope effect, kn/kD, of 1.58. This value is comparable to but larger than that observed for H atom abstraction from 4fluorophenol (kH/k, 1.32) by oxoiron(IV) tetra(2-N-methylpyridyl)porphyrin at pH 7.7.15

659

2.6. The stoichiometry of the oxidation of Acid Orange 12 The stoichiometry of the oxidations was examined with the most reactive of the dyes used in the present study, 1-(4-methoxyphenylazo)-2-naphthol-6-sulfonate, to minimise problems from the slow self-reaction of 1. This revealed that the reaction requires four oxoiron(IV) species to oxidise each molecule of dye. The consequence of this result is that the true k2 value for the reduction of I by an azonaphthol dye in this study is a quarter of the value shown in Table 1. This correction does not, however, affect the Hammett correlations and p values. 2.7. The mechanism of azonaphthol dye oxidation by oxoiron(lV) porphyrins Our recent studies on phenol oxidation by oxoiron(IV) and oxomanganese(IV) porphyrins in aqueous solution (pH 7.6) have shown that the initial step in these reactions is H atom abstraction to generate a phenoxyl radical (Scheme 3).ts'~6The kinetic isotope effect measured in the present study indicates that H atom abstraction also occurs in the rate determining step of the oxidation of azonaphthol dyes by 1. However, although the structures of azonaphthol dyes are normally shown as azo compounds, in aqueous solution they are in a rapid dynamic equilibrium with their hydrazone tautomers; the latter isomer being the dominant species (Scheme 4)fl This complicates kinetic studies on the dyes, since the substrate is in effect a mixture of two compounds. Consequently one or both the tautomers may be the active form of the substrate providing the H atom for the oxidant. It is important to note that, irrespective of which tautomer is the reactive substrate, H atom abstraction leads to a common azonaphthoxyl radical intermediate and subsequent reactions of this species should be independent of the initial tautomerism (Scheme 5).

ArOH

+

O~:e !

-~ArO"

+

HO

Scheme 3

xxN

Azo-tautomer.

------

SO3Na

N,, N

Hydrazone-tautomer.SO3Na Scheme 4

The position of the azo-hydrazone equilibrium has been studied using a range of spectroscopic methods which show that it is sensitive to solvent polarity, solvent H-bonding and

660 to substituents on the phenyl ring: electron-withdrawing groups favour the hydrazone and elcctron-donors the azo tautomer. 29 A further difference between simple phenols and the azonaphthol dyes used in this study is the strong internal H-bonding that occurs in both of the dye tautomers. This results in a large downfield shitt of the 1H NMR signal associated with this proton and an increase in pK a(Table 1). Examination of the measured second order rate constants for the dye oxidations shows that the reaction is significantly favoured by electron-donating substituents: the 4-methoxylated compound (9) reacts 274 times faster than the 4-nitro analogue (3). Furthermore, the excellent linear Hammer correlation of log(second order rate constant) with o ~ and the p value of-1.66 conf'trm this conclusion and show that all the substrates are oxidised by the same mechanism. However, the size of the p value is larger than that obtained from the earlier oxidation of phenols (p-1.10) 15and is larger than expected for a mechanism involving H atom abstraction in the rate determining step. A possible explanation is that the substituents have mo effects on the reaction which reinforce each other: they influence H atom abstraction by the electrophilic oxidant, as noted previously this is favoured by electron-donation,~5'16and they shift the position of the azohydmzone equilibrium. Berrie et al. 3~ carded out an NMR study on the tautomeric equilibria of an analogous series of unsulfonated dyes in chloroform and measured the percentage tautomer distribution for each dye. Their data show that the azo tautomer is favoured by electron donation and vice v e r s a , furthermore, the tautomer distributions can be used to calculate equilibrium constants and to examine the influence of substituents on the equilibrium. The data from eight dyes (omitting the 4-cyano compound) give a Hammett p value of-0.46 (vs 6+, R = 0.930). In agreement with this study in chloroform, UV-Vis spectroscopy shows that in aqueous solution (pH 6.93) the hydrazone is the major tautomer for both 3 and 9 and that the former, with an electronwithdrawing 4-nitro substituent, is almost entirely in this form (strong absorbance at 486 nm) whereas the latter, with an electron-donating 4-methoxy group, is a more equal mixture of tautomers (absorbances at 410 nm and 502 nm). It follows that if the azo tautomer is the reactive compound, then electron-donation by the substituent will increase the measured rate of reaction by increasing the proportion of this species in the reaction mixture. The p value for H atom abstraction (PH.~) is then given by equation 2, where Pobsis the measured p value and Ptaut is the p value for the azo dye tautomerism. PH.abs = Dobs- Ptaut

2

From the p value obtained in this study and from the value calculated from the work of Berrie et al., PH.absis -1.20. Although this p value from equation 2 is close to that obtained from our earlier study of phenol oxidation by oxoiron(IV)tetra(2-N-methylpyridyl)porphyrin (p-1.10), it is important to note that P~ut was obtained from a study in chloroform whereas the present investigations used aqueous buffer (pH 6.93). Spadaro and Renganathan, s from their studies on the peroxidase-catalysed oxidation of azo dyes with hydrogen peroxide, have proposed a mechanism involving the initial formation of the azonaphthoxyl radical. Subsequent one-electron oxidation and reaction with water result in the cleavage of the azophenyl group from the 1-position of the naphthol ring and loss of dye colour. Based on this mechanism and the 4:1 stoichiometry of the oxidation, we suggest the mechanism

661 in Scheme 5 to account for the oxidation of azonaphthol dyes by oxoiron(IV) porphyrins.

"0

oFivp

0

2>N= N X

X

SO 3-

SO 3-

SO 3OFeWP H20

~

O ~---N:NH

X

+ O

+ O

12OFeWp

~

H20

SO 3-

SO 3-

X

Scheme 5 ACKNOWLEDGEMENTS We thank Unilever Research and the University of York for fmancial assistance and the EPSRC Mass Spectrometry Service, Unversity of Wales, Swansea for measuring the electrospray mass spectrum of iron(Ill) tetra(2,6-dichloro- 3-sulfonatophenyl)porphyrin. REFERENCES

1. 2. 3. 4. 5.

K.E. Hammel and P. J. Tardone, Biochemistry, 27 (1988) 6563. K. Valli and M. H. Gold, J. Bacteriol., 173 (I 991) 345. A. Paszczynski and R. L. Crawford, Biochem. Biophys. Res. Commun., 2178 (1991) 1056. J. Spadaro, M. H. Gold and Renganathan, Appl. Environ. Microbiol. 58 (1992) 2397. A. Paszczynski, M. B. Pasti- Grigsby, S. Goszczynski, R. L. Crawford and D. L. Crawford, Appl. Environ. Microbiol. 58 (1992) 3598. 6. M.B. Pasti-Grigsby, A. Paszczynski, S. Gozczynski, D. L. Crawford and R. L. Crawford, Appl. Environ. MicrobioL.58 (1992) 3605.

662 7. 8. 9. 10. 11. 12. 13. 14. 15. 16. 17. 18 19. 20. 21. 22. 23. 24, 25. 26. 27. 28. 29. 30.

C. Muralikrishna and V. Renganthan, Biochem. Biophys. Res. Commun. 197 (1993) 798. J.T. Spadaro and V. Renganathan, Arch. Bioehem. Biophys. 312 (1994) 301. M. Stibrova, B. Asfaw, E. Frei and H. H. Schmeiser, Collect. Czech. Chem. Commun. 61 (1996) 962. R. Labeque and L. J. Marnett, J. Am. Chem. Soc., 111 (1989) 6621. T.G. Traylor J. P. Ciccone, J. Am. Chem. Soc., 111 (1989) 8413. J.R. Lindsay Smith and R. J. Lower J. Chem. Soc. Perkin Trans. 2, (1991) 31. T.C. Bruice, Acc. Chem. Res. 24 (1991) 243. T.G.Traylor, S. Tsuchiya, Y.-S. Byun and C. Kim, J. Am. Chem. Soc. 115 (1993) 2775. N. Colclough and J. R. Lindsay Smith, J. Chem. Soc. Perkin Trans. 2, (1994) 1139. N.W.J. Kamp and J. R. Lindsay Smith, J. Mol. Catal. A: Chem., 113 (1996) 131. G. Hodges and J. R. Lindsay Smith, unpublished observations. R.A. Robinson and R. H. Stokes, in Electolyte Solutions, Butterworths, London, 1965. S.J. Bell, P. R. Cooke, P. Inchley, D. R. Leanord, J. R. Lindsay Smith and A. Robbins, J. Chem. Soc. Perkin Trans. 2, (1991) 549. T.C. Bruice, Acc. Chem. Res., 24 (1991) 243. P.S. Traylor, D. Dolphin and T. G. Traylor, J. Chem. Soc. Chem. Commun., (1984) 279. B. Meunier, Chem. Rev., 92 (1992) 1411. N.B. Chapman and J. Shorter in Correlation Analysis in Chemistry: Recent Advances, ed. O. Exner, Plenum Press,New York, 1978, p.439. Y. Okamoto and H. C. Brown, J. Am. Chem. Soc., 80 (1958) 4979. D.R. Arnold in Sustituent Effcts in Radical Chemistry, eds. H. G. Viehe, Z. Janonsek and R.Merenyi, Reidel, Dordrecht, 1986, p 171. T.H. Fife and T. C. Bruice, J. Phys. Chem., 65 (1961) 1079. N. Isaacs, Phycal Organic Chemistry, Longman, Belfast, 1987. P. Ball and C. H. Nicholls, Dyes Pigments, 3 (1982) 5. Y. Onari, Bull Chem. Soc. Jpn. 58 (1985) 2526. A. H. Berrie, P. Hampson, S. W. Longworth and A. Mathias, J. Chem. Soc. Section B, (1968) 1308.

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

663

VOC's ABATEMENT: PHOTOCATALYTIC OXIDATION OF T O L U E N E I N V A P O U R P H A S E O N A N A T A S E TiO~ C A T A L Y S T V. Augugliaro a, S. Coluccia b, V. Loddo ~, L. Marchese b, G. Martra b, L. Palmisano ~, M. Pantaleone ~ and M. Schiavello a aDipartimento di Ingegneria Chimica dei Processi e dei Materiali, Universith degli Studi di Palermo, Viale delle Scienze, 90128 Palermo, Italy bDipartimento di Chimica I.F.M., Universit/l degli Studi di Torino, Via P. Giuria 7, 10125 Torino, Italy

Photocatalytic oxidation of toluene has been carried out in a gas-solid regime by using polycrystalline anatase TiO2. A fixed bed continuous photoreactor of cylindrical shape was used for performing the photoreactivity runs; the Pyrex glass photoreactor was irradiated by a medium pressure Hg lamp. Air containing toluene and water vapours in various molar ratios was fed to the photoreactor. Toluene was mainly photooxldised to benzaldehyde although benzene, benzyl alcohol and traces of benzoic acid and phenol were also detected. The presence of oxygen was essential for the occurrence of the photoreaction while water played an important role in the mamtainance of the catalytic activity. The results obtained in a preliminary Fourier transform infrared (FT-IR) investigation indicate that toluene is weakly stabilised on the TiO2 particles by hydrogenbonding between the aromatic ring and surface hydroxyl groups.

1. I N T R O D U C T I O N Volatile organic compounds, VOC's, are an important class of mr pollutants usually found in the atmosphere of all urban and industrial areas. Toluene is one of these compounds and, due to its noxious nature, several strategies have been identified in order to reduce its presence in indoor and industrial emissions. Among the methods effective to oxidize toluene, heterogeneous photocatalysis is one of the most attractive, due to the mild conditions under which this process is usually carried out. Photocatalysis has been largely used for the photooxidation of many organic molecules in the hquid-solid regime [1-5], but a few papers report photoreactions in the gas-solid regime [6-10]. The photocatalytic oxidation of toluene in the presence of water was performed m the gas-solid regime by Ibusuki and Takeuchi [6] at room temperature by using TiO2. They found that the presence of water was beneficial in order to achieve the almost complete photo-omdation of toluene to CO2 and HzO, in fact,

664 only very small amounts of benzaldehyde, which is the main product of toluene partial oxidation, were detected. The present paper reports the results of the toluene photooxidation reaction using polycrystalline anatase TiO2 as catalyst. The photoreactivity runs were carried out in a continuous photoreactor fed by a mixture of air, toluene, and water in various molar ratios and irradiated in the near-LW region. The influence of toluene concentration, gas flow rate, and water presence on the photoprocess performance was investigated. A preliminary investigation of the interaction between toluene and the catalyst surface was carried out by Fourier transform infrared (FT-IR) spectroscopy.

2. E X P E R I M E N T A L The set-up of the experimental apparatus is reported in Figure 1. The reactivity experiments were carried out in a flow apparatus using a Pyrex cylindrical reactor whose dimensions were: internal diameter, 1 cm; external diameter, 1.2 cm; and height, 30 cm. A porous frit at the bottom of the cylinder was used to support the fixed bed of powder and to distribute the inlet gaseous mixture. The reactor was vertically positioned inside a thermostatted chamber and was irradiated through a circular window made on a wall of the chamber and covered by a Pyrex glass sheet. An aluminum parabolic reflector was located behind the photoreactor in order to mcrease the illumination. The radiation source was a 400 W medium pressure Hg lamp (Polymer GN ZS, Helios Italquartz) put at approximate 30 cm from the reactor. The radiation power impinging on the photoreactor was measured by usmg a radiometer UVX Digital and its mean value was 5 mWcm 2. The catalyst was polycrystalline TiO2 (Merck, anatase, BET surface area: 10 m2gl); the powder was classified by sieving and the fraction with particle size in the 45-90 ~m range was used. The amount of catalyst was 8 g, corresponding to a fixed bed height of ca. 10 cm. The reactant mixture was generated by bubbling air through saturators containing water and toluene at room temperature. For some runs benzoic acid was used instead of toluene. The gas flows were then mixed and fed to the photoreactor. The temperature of the reactor was always 413 K. The gas flow rates were m the 0.17-10 cm3s-1 range and the toluene molar fraction ranged from 4.0x10 4 to 1.3x10 2. The water molar fraction was held constant to 2.5x10 2. The catalyst was irradiated only when steady state conditions were achieved in the system, i.e. after about 24 h from the beginning of the photoreactor feeding. The runs lasted 170, 350 or 470 h. The gas leaving the photoreactor was periodically analyzed by a gas chromatograph (Varian, Vista 6000), equipped with FID detector. A 0.1% AT-1000 on Carbograph column (2 m x 2 ram) and a Porapak QS column (2 m x 2 ram) were used. For some experiments the gas exiting from the photoreactor was continuously bubbled in liquid acetomtrile. The resulting solution was analyzed by high pressure liquid chromatography (HPLC) (Varian 9050) in order to detect

665 compounds produced in small quantities. At the end of each run, the catalyst was held for 24 h in twice distilled water or acetonitrile in order to dissolve products adsorbed on the surface; the resulting solutions were analyzed by HPLC.

(b)

(c) ~1

(a)

_i

(h)

ill

(e)

vent

(i) GC ._1

(mt

(1)

I~176I

(o) (n)

Figure 1. Photoreactivity set-up. (a) Air cylinder, (b) control valves, (c) bottle with water, (d) bottle with toluene, (e) switch valves, (f) thermostatted chamber, (g) parabolic reflector, (h) cylindrical photoreactor, (i) lamp, (1) power supply, (m) water filter, (n) gas chromatograph, (o) bubbling bottle containing acetonitrile. Thick line was electrically warmed in order to avoid product condensation.

For selected runs the gas at the outlet of the photoreactor was bubbled in a saturated aqueous solution of barium hydroxide in order to trap CO2 as BaCO3. The IR spectra were obtained with a Bruker IFS 48 spectrometer. The catalyst powders, as self-supporting pellets, were placed m an infrared cell allowing adsorption-desorption experiments to be carried out in situ. Prior to the adsorption of toluene, the cell was evacuated (1.0xl0 -~ Torr) at room temperature.

3. R E S U L T S AND D I S C U S S I O N Blank reactivity tests were performed at the same experimental conditions used for the photoreactivity experiments but in the absence of catalyst, oxygen or light. Other runs were carried out by using COz or Nz instead of 02. No reactivity

666 was observed in all these cases so that it may be concluded that 02, catalyst, and irradiation areneeded for the occurrence of the photoprocess. The photoreactivity results showed that the reactor reaches steady state conditions after a long period of time (ca. 70 h) from the beginning of the irradiation. At steady state conditions the mare photooxidation product was benzaldehyde but also benzyl alcohol and traces of benzoic acid and phenol were detected at all the experimental conditions used. During the transient period, benzene together with CO2 were also produced. No significant evidence of CO2 production was observed at steady state conditions. These results obviously indicate that at the experimental conditions used the photoprocess does not give rise to a complete degradation of toluene. In Figure 2 the experimental data of benzaldehyde steady state production rate are reported as a function of the gas flow rate. An increase in the reaction rate occurred above a flow rate of 1.7 cm3s1, and remained constant at higher values. This indicates that external mass transfer limitations occur for flow rates less than 1.7 cmas ' . Because of this all the runs carried out for investigating the influence of toluene concentration were performed at a flow rate of 2.5 cm3s -1.

CD

A v

6

4

~~

3

.~~

2

"~

1

o 0

5

10

F l o w r a t e [cm~s -1] Figure 2. Benzaldehyde steady state production rate versus the gas flow rate. Toluene molar fraction: 4.0x10-4. Catalyst amount: 8 g. Photon flux: 5 mWcm 2.

The inlet toluene concentration greatly affected the benzaldehyde production rate. The runs performed at higher molar fraction of toluene (4.0x10 4, 7,0x10 4, and 1.3x10 -2) showed an increasing benzaldehyde production rate (4.1x10 a, 6.2x10 "a, and 3.4"x10 2 ~molsl).

667 The occurrence of catalyst deactivation and of the role played by water, was investigated by performing very long reactivity runs at the flow rate of 0.42 cm3s 1. Deactivation is not affected by the presence of mass transfer limitations, while the use of a low flow rate allows variations of outlet gas composition to be detected more accurately. Figure 3 reports the fractional conversion of toluene to benzaldehyde versus irradiation time for two long runs carried out in the presence and in the absence of water vapour. For the run in the presence of water, a maximum conversion of 0.19 (corresponding to an

O O

0 r cD

0.1

A

0

AA A m

(D ,.Q

A O

m

A

0.01 0.1

1

10

100

1000

T i m e [h] Figure 3. Toluene fractional conversion to benzaldehyde versus irradiation time for runs carried out in the presence of water vapour (m) and in the absence of water vapour (A). Flow rate: 0.42 cmasl; toluene molar fraction" 1.3x10 -~. Catalyst amount: 8 g. Photon flux: 5mWcm -2.

oxidation rate of 5.6x10 -2 ~mols -1) was achieved after 2 h, while a steady state conversion of 0.08 was achieved after 70 h of irradiation (corresponding to an oxidation rate of 2.4x10 -2 ~mols-i). No decrease of the photoreactivity was observed after 350 h. A different behaviour can be observed for the run carried out in the absence of water. Indeed, a maximum conversion of 0.1 (corresponding to an oxidation rate of 2.9x10 -2 ~mols -1) was reached after ca. 12 h. After that time, the photo-reactivity continuously decreased down to negligible values. In Figure 4 the fractional conversion of toluene to benzene is reported versus irradiation time for the same runs reported in Fig. 3. The presence of water was beneficial for benzene production, but benzene virtually disappeared after 3-4 h of irradiation, independent of the presence of water.

668 In order to determine if the catalyst deactivation was irreversible, a run was performed where first the wet reagent mixture was fed to the irradiated photoreactor and time was allowed for the achievement of a constant photoreactivity level. Then the dry reagent mixture was fed for a prolonged time; and finally the wet reagent mixture was again fed. In Figure 5 the results obtained in this run are reported as toluene fractional conversion to benzaldehyde versus irradiation time. In the absence of water a sharp decrease of toluene conversion occur from 0.08 to 0.04 after ca. 6 h and thereafter from 0.04 to 0.01 after ca. 180 h. When water vapour was again added to the reaction mixture, the photoreactivity initially increased but then slowly decreased until a

0.005 o o o0

-mmq

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mm A i

0

1

m

m A

m I

2

nI

|

ILl 3

n I

n I

4

Time [h] Figure 4. Toluene fractional conversion to benzene versus irradiation time for the runs reported in Figure 3.

constant value of conversion was achieved. For this run benzene was produced only by the fresh catalyst in the first hours of irradiation; the partially spent catalyst did not produce benzene when water vapour was again present. The results reported in Figure 5 indicate that water is an essential reagent for the photoprocess. The highest activity is shown by the fresh catalyst working m the presence of water vapour; m the absence of water the catalyst progressively deactivates. The catalyst deactivation seems to be a partially reversible process; the restoration of water m the reaction mixture allows a partial recovery of the activity whose final level, however, is less than that of the fresh catalyst.

669 0.1 O O r~

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-

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O r

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~.

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150

250

350

450

Time [hi Figure 5. Effect of water on the photoreactivity.

Figure 6 reports FT-IR spectra obtained at various experimental conditions. The admission of toluene onto the catalyst causes the depletion of the band originally present in the background spectrum at 3670 cm -1 (curve a), due to a stretching mode of surface hydroxyl groups [11]. This band is transformed to a complex and much broader absorption in the 3600-3450 cm -I range (curve b). On the basis of such behaviour it can be concluded that surface OH groups are revolved in the adsorp.tion of toluene onto the catalyst, the resulting organic molecules are probably stabilized on the surface by hydrogen bonding between the aromatic ring and the hydroxyl groups [12]. IR bands assignable to adsorbed toluene appear in spectrum b, namely in the 3200-2800 cm 1 (CH stretchings) and 2000-1300 cm 1 (summation bands, ring stretchings). By simply outgassing at room temperature these bands disappear (spectra not reported) and the original spectral band of the surface hydroxyls is progressively fully restored (Figure 6, reset). Such reversibility suggests a weak character of the observed interaction between toluene and the catalyst. By taking into account all the results above reported, the foUowmg reaction mechanism for the production of benzaldehyde can be suggested. Under photoexcitation of the semiconductor oxide with band gap irradiation, electron-hole (e-h) pairs are photogenerated: TiO2 + hv -~ TiO2 + e- + h §

(1)

The pairs, once separated, can induce chemical redox transformations with the species adsorbed on the surface, subject to thermodynamic constraints. It is generally assumed that surface hydroxyls act as hole traps by producing OH" radicals:

670

2.0 1.3 1.2

!.5 "~ e~

1.0 0.9

l.O

0.8

r,~

0.7


0 2

--~ 02

+

H202

H202(.~) +O2-(.~) --> OH- +

OH" + 02.

(7)

671 The radical species are very reactive and may attack toluene molecules according to the following reactions: OH"

+ C6H~CH3(.~)

-+ H20 + C6H5CH2"

C6HsCH2" + 02(,d~) --> C6HsCH200" C6HsCH2OO"

+ e

~

C~H~CHO + O H .

(8) (9) (10)

The formation of benzyl alcohol may occur by the following reaction: C6HsCH2" + OH"

~

C6H~CH2OH.

(11)

The small amount of benzyl alcohol found as a product is understandable since two radicals are involved in reaction (11). Concerning the simultaneous appearance of benzene and C02, the experimental results suggest that the breakage of the bond between the CH3 group and the ct carbon of toluene does not occur. In this last case, in fact, CH4 and/or CH30H molecules should be produced by the reaction between CH3 radicals and OH groups present on the catalyst surface but neither CH4 or CH3OH were ever detected in our experiments. On this basis the occurrence of the following reaction steps may be suggested: C6HsCHO(.~) + OH" C6H~CO" + 0 2 ( ~ ) ~ C6H5C000"

-~ C6H~CO" + H20 C6H5C000"

+ C6HsCHO(a~)-~

C6H5C0" + C6HsCOOOH(a~)

(12) (13) (14)

C6HsCOOOH(.~) + C6HsCHO(.~)---> 2 C6H~COOH(.~)

(15)

C6H~COOH(~) --> C6H6 + C02.

(16)

According to reactions (12)-(16), CO2 results from the oxidation of toluene to benzoic acid whose traces were found in our experiments and its subsequent photodecarboxylation. It must be reported that some runs carried out by using benzoic acid at the same experimental conditions used for toluene photooxidation, afforded benzene and CO2 in large amounts. It is well known [7, 8] that ethanoic acid is easily photodecarboxylated in a gas-solid regime in the presence of irradiated polycrystallme semiconductor oxides. The small amounts of phenol are probably due to an attack of benzene by OH radicals (see eqn.(16)). The progressive deactivation of the catalyst in the absence of water could be due to some surface dehydroxylation and/or to the formation of stable intermediate species which can not evolve in the absence of water so that they remain strongly adsorbed on the catalyst surface. It is worth reporting that the

672 catalyst was found of ochre colour at the end of the run and this colour completely disappeared by irradiating the catalyst in liquid water. The presence on the catalyst surface of sites with different oxidation strength can explain the role played by water and the deactivation-activation pattern exhibited by the catalyst. The partial restoration of activity when water was restored to the reacting ambient could be due to the rehydroxylation of surface sites. The finding that benzene is produced only in the first hours of irradiation while the partially spent catalyst never produced benzene, may be justified by considering that the fresh catalyst has higher oxidant properties due to the presence of highly oxidant hydroxyl groups. Benzaldehyde, instead, can be obtained both on these sites and on less oxidant sites. The experimental results suggest that the more oxidizmg sites react irreversibly under irradiation. By concluding, the photo-oxidation of toluene mainly to benzeldehyde and benzene (as a transient product) was proved to occur in mild conditions in gassolid regime. The presence of sites with different oxidant properties is proposed and a preliminary FTIR mvestigation indicates a weak interaction between toluene molecules and TiO2 surface. ACKNOWLEDGEMENT The authors wish to thank the 'TIinistero delrUmversith e della Ricerca Scientifica e Tecnologica" (Rome) for financially supporting this work.

REFERENCES 1. M. Schiavello (ed.), Photocatalysis and Environment. Trends and Applications, Kluwer, Dordrecht, 1988. 2. V. Augugliaro, L, Palmisano, A. Sclafani, C. Mmero and E. Pelizzetti, Toxicol. Envir. Chem., 16 (1988) 89. 3. E. Pelizzetti and N. Serpone (eds.), Photocatalysis. Fundamentals and Applications, Wiley, New York, 1989. 4. E. Pelizzetti and M. Schiavello (eds.), Photochemical Conversion and Storage of Solar Energy, Kluwer, Dordrecht, 1991. 5. D. F, Ollis and H. A1-Ekabi (eds.), Photocatalytic Purification and Treatment of Water and Air, Elsevier, Amsterdam, 1993. 6. T. Ibusuki and K. Takeuchi, Atmos. Environ., 20 (1986) 1711. 7. L. Palmisano, M. Schiavello, A. Sclafani, S. Coluccia and L. Marchese, New J. Chem., 12 (1988) 847. 8. T. Matsuura and M. Anpo (eds.), Photochemistry on Solid Surfaces, Elsevier, Amsterdam, 1989. 9. M.L. Sauer and D. F. Ollis, J. Catal., 158 (1996) 570. 10. A. J. Malta Vidal, J. Soria, V. Augugliaro, V. Loddo, Chem. Biochem. Eng. Quart., 1997 (in press). 11. C. Morterra, J. Chem. Soc. Faraday Trans. I, 84 (1988) 1617. 12. E. A. Paukshits and E. N. Yurchenko, Russ. Chem. Rev., 52 (1983)242. 13. R. I. Bickley, G. Munuera and F. S. Stone, J. Catal., 31 (1973) 398.

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

673

Oxidation Processes on Stoichiometric and Nonstoichiometric Hydroxyapatites H. Hayashi', H. Kanai b, Y. Matsumura c, S. Sugiyama', and J.B. Moffat d "Department of Chemical Science and Technology, Faculty of Engineering, The University of Tokushima, Minamijosanjima, Tokushima 770, Japan ~Faculty of Life Science, Kyoto Prefectural University, Simogamo, Sakyo-ku, Kyoto 606, Japan COsaka National Research Institute, AIST, Midorigaoka, Ikeda, Osaka 563, Japan dDepartment of Chemistry and the Guelph-Waterloo Center for Graduate Work in Chemistry, University of Waterloo, Waterloo, Ontario, Canada N2L 3G 1

1. INTRODUCTION Solids whose crystallographic structures are stable at high temperatures, permit ionic exchange and changes in their elementary composition while remaining structurally invariant are of considerable interest for utilization as and fundamental studies of heterogeneous catalysts. Solids with these properties but additionally with catalytic functionalities which are altered by one or more of the aforementioned changes are particularly useful in understanding the nature of catalytically active sites. Calcium hydroxyapatite [Ca~o(PO4)6(OH)2] (Fig. 1) is an inorganic solid which exhibits these characteristics [ 1]. It does not readily lose hydroxyl groups and the lattice is believed to be stable 9. I

.,

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10

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8

0.6

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0.4

4

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0-0

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g, r

0 10

20

30

40

50

I /mA

Fig. 5 Current-votage and current-power (reaction conditions are the same as Fig.4)

3.3 Effect of Temperature Figure 6 shows the effect of temperature on C~ selectivity and CH, conversion at constant feed composition and total flow rate. It can be seen that the C2-selectivity initially increases with a rise in temperature and then appears to reach a plateau and for

689

T>750 ~ it starts decreasing. The C2-selectivity shows maxim at 750 ~ There is the optimum temperature for the synthesis of C2 hydrocarbons. It also indicates that the total oxidation reaction is much easier to occur under high temperatures. The CH4 conversion shows a increase with increasing temperature. Figure 7 shows the effect of temperature on the current generated. It can be seen that the current generated increases with temperature. This is due to the increased ionic conductivity of the YSZ at the higher temperatures. 100

5

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50 660

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,

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680

700

720

740

760

780

0 800

T/C

Fig.6 C2 selectivity and CH4 conversion as a function of temperature( R=0.1 ohm, anode: Q=100ml/min, Pch4=20.3 kPa, cathode: Q=250ml/min, Po2=20.3kPa) 100 4,

Exp.

80
r

o

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0

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0 Methane Conversion

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9 Oxygen Conversion Temperature I I I J 0 60 70 80 90 100

Catalyst Loading (mg)

v

25-

0 CO Selectivity H2 Selectivity

O~ 40

I

50

I

60

I

70

I

80

I

90

Catalyst Loading (mg)

Figure 2. Effect of catalyst loading completely cover the permeable section of the membrane tube. The minimum amount of catalyst that can be used in the membrane reactor is 50 mg because this amount is required to completely surround the permeable portion of the membrane. Conversion is highest at loadings which are just over this minimum loading. At the higher loadings, methane conversion and hydrogen selectivity both decrease, which is an indication of oxygen reacting with the hydrogen rather than methane. It is clear that oxygen is the limiting factor in methane conversion. Increasing the catalyst loading exposes only a small fraction of the additional catalyst to the methane/oxygen mixture because the volume of catalyst is higher than the permeable section of the membrane tube. Simulation results in a fixed bed reactor (not shown) predicts that the reaction occurs only in a small region at the bed entrance. In the case of the membrane reactor, the active region is located near the wall of the membrane tube. Accordingly, increasing the catalyst loading does not significantly increase the active region as shown by the experimental observations. The catalyst temperature, like methane conversion and CO selectivity, does not vary too much in these experiments. It remains constant near 515~ throughout regardless of the catalyst loading. At the higher loadings, methane conversion decreases so it would be expected that the temperature would decrease. The results show that the temperature does not change at the point at which it is measured. However, it is expected that the temperature gradient in the reactor is affected by the increasing amount of catalyst. The membrane reactor yields slightly lower conversions and selectivities than the fixed bed reactor. Nonetheless, the safer reaction environment that the membrane reactor provides, permits us to study the reaction at conditions that could be hazardous in a fixed bed reactor. In the membrane reactor, it is possible to lower the methane/oxygen feed ratio without the potential for an explosion. Conversions and selectivities obtained when varying the methane/oxygen feed ratio are shown in Figure 3. In these experiments 60 mg of catalyst

100

696

02

are used, and the methane feed rate is held constant at 500 cc/min while varying the flow rate.

900 100 ~

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600 ~ "r, o.75'

75-

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0

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400 = ~50

5025-

It 1 .~3

Methane Conversion Oxygen Conversion Temperature ~

2.~

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3.

Methane/Oxygen Feed Ratio

200

25 0 I

1

1.5

H2 Selectivity I

2

I

2.5

I

3

3.6

Methane/Oxygen Feed Ratio

Figure 3. Effect of methane/oxygen feed ratio Although oxygen conversion is complete at 100%, methane conversion, syngas selectivity and temperature are affected by changes in the methane/oxygen feed ratio. Methane conversion varies from a high value of 64% at a ratio of 1/1 to a low value of 44% at a ratio of 3/1. Selectivities vary from 22 to 82% for hydrogen and 67% to 90% for CO. Both CO and hydrogen selectivities are significantly lower at the lower feed ratios, particularly when the ratio is less than 2/1. This is an indication that the complete combustion reaction begins to occur. This is corroborated by the large amount of water condensed in the traps and the larger amount of CO2 detected in product gas analysis. This effect can be explained by the fact that at the lower feed ratios, more oxygen is available in the reactor to oxidize CO and H2 to form CO2 and water. At higher feed ratios, the concentration of oxygen is low and hence methane reacts to form CO instead of CO2. It should be noted however that the CO and hydrogen selectivities remain nearly constant around 90 and 80% ,respectively, for values of CH4/O2 above 2/1. The decrease in temperature with increasing methane/oxygen ratio also is consistent with the role of the oxidation of the products as the reason for the lower selectivities. At feed ratios less than 2/1, the steady state temperature rises above 700~ but at ratios equal to or greater than 2/1 the temperature is in the 500~ range. The higher temperature can be accounted for by the reaction of hydrogen and oxygen to form water and the occurrence of the complete combustion reaction, which has a higher heat of reaction than the partial oxidation reaction. In addition, as methane conversion decreases, the heat generated decreases and the temperature decreases.

697 Although 100% conversion of oxygen is achieved at low feed ratios, the methane conversion levels off at about 65%. While the residence time is higher at the low C H j O 2 ratio, we expected higher methane conversion since oxygen is no longer the limiting reagent in this range of feed ratios. Additional oxygen should result in more methane being converted, but instead it only results in more of the products being converted to CO2 and water. At higher feed ratios, the methane conversion begins to decrease because the amount of oxygen in the reactor decreases. In addition, at higher feed ratios the temperature is lower, further decreasing methane conversion. At a 2/1 feed ratio, which is the stoichiometric ratio of the direct partial oxidation reaction, methane conversion and both CO and hydrogen selectivities are maximized. The effect of varying the total feed rate was studied using 60 mg of catalyst and a methane/oxygen feed ratio of 2/1. The total feed rate varies from 300 to 1200 cc/min corresponding to residence times of about 3 to 0.8 ms respectively. The results ( Figure 4) show that, contrary to expectations, upon doubling the flow rate from 300 to 600 cc/min, methane conversion and H2 and CO selectivities remain nearly constant at approximately 65%, 70%, and 80% respectively. In addition, as in the previous results, oxygen conversion remains at 100%. Methane conversion, which is nearly constant at 65% at feed rates of 600 cc/min and below, falls to only 35% at 1080 cc/min. The CO and H2

IOE

9 W

A

9

A W

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9 9

9 9

750

100

!

-.I

o~75

75

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,m

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0

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250"-"

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0 MethaneConversion I OxygenConversion 0 Temperature

0 200 4()0

I

600

25

CO Selectivity H2 Selectivity

0

800 1000 1200

0

i

200

Total Feed Rate (cc/min)

400

I

600

I

I

800 1000 1200

Total Feed Rate (cc/min)

Figure 4. Effect of Total Feed Rate selectivities remain nearly constant throughout the whole range of flow rates, even when methane conversion decreases. The temperature increases at the higher feed rates, even when a decrease in methane conversion occurs. The temperature is nearly constant, around 515~ at the lower feed rates, but increases to 680~ at 1080 cc/min. This is because, although methane conversion is decreasing, overall syngas production is increasing due to the larger total flow rate. Thus, the temperature increase results from an increase in heat generation due to higher rate of syngas production. It should be noted that higher feed rates were attempted, but in such

698 cases the temperature rose out of control due to the large amount of heat evolved from an increased amount of syngas production. The effect of flow rate observed in a fixed bed reactor is somewhat different than what is achieved in the membrane reactor. In the fixed bed reactor, methane conversion remains nearly constant at higher flow rates, but in the membrane reactor, the conversion decreases significantly. This is probably due to the fact that in the fixed bed reactor, the reaction occurs in a narrow region at the top of the bed, but in the membrane reactor, the reaction occurs close to the membrane wall. In the membrane reactor, as the total flow rate increases, the amount of oxygen near the outside reactor wall does not increase, but the amount of methane throughout the bed does increase. In addition, there is a greater area for mixing in the fixed bed reactor, which becomes more important to maintaining high conversions at high flow rates. It is this greater reaction area which also results in lower hydrogen selectivities in the fixed bed. In that case, there is a greater probability that hydrogen will react with oxygen to form water. A detailed elementary step model similar to that proposed by Hickman and Schmidt [7] was developed to interpret the fixed bed results. The model assumes that methane adsorbs dissociatively forming adsorbed carbon and adsorbed hydrogen. Oxygen adsorbs dissociatively and the reaction of adsorbed carbon and adsorbed oxygen yields CO. Recombination of adsorbed hydrogen yields dihydrogen which desorbs into the gas phase. Further oxidation of CO yields CO2 and the reaction of adsorbed hydrogen and adsorbed oxygen yields adsorbed hydroxyls which recombine or react with adsorbed hydrogen to yield adsorbed water. Using this model we were able to reproduce the results reported by Schmidt and Hickman using their reaction parameters. After accounting for the differences between the monolith and fixed bed reactors we attempted to fit the model to our fixed bed results using the same reaction parameters as Hickman and Schmidt [7], we found a significant discrepancy between the model predictions and our experimental fixed bed results. The discrepancy derived mainly from the fact that the reaction temperatures in our experiment are significantly lower than in the those reported by Schmidt group using a monolith reactor. The model predicted that under our reaction conditions the surface will be covered mainly by CO and under predicted the methane conversion (30% instead of 70%) and CO selectivity (20% instead of 90%). Upon conducting a sensitivity analysis we found that the model prediction fitted our results fairly well when the CO desorption energy was lowered from the value of 31.6 kcal/mol used by Hickman and Schmidt to 25.5 kcal/mol. Hence we conclude that in our case the TiO2 support increases the rate of CO desorption. One of the main results of the model is that it showed that the reaction occurred in a narrow region near the entrance of the bed, which explains the lack of sensitivity of the results to the changes in loading, flow rate and methane/oxygen ratio. The fixed bed reactor model is one dimensional and thus its solution is not so involved. In the case of the membrane reactor it is necessary to use a two dimensional model to account for the radial flow of oxygen at the wall. Work is underway to solve this model, however, the results of the one dimensional model are qualitatively useful to interpret the membrane reactor results. After studying the effect of catalyst loading, methane/oxygen feed ratio, and total feed rate, the effect of operating the reactor at temperatures above the autothermal steady state temperatures was studied. Because only 65% methane conversion was achieved in the membrane reactor, an attempt was made to increase the conversion by raising the reaction temperature using external heating. In these experiments, 500 cc/min of methane and 250 cc/min of oxygen were reacted over 60 mg of catalyst. As the temperature is increased to 650~ the oxygen conversion remains at 100%, methane conversion remains around 60% and the CO and hydrogen selectivities remain around 90 and 75% respectively. Due to 100% oxygen conversion, there is not enough oxygen throughout the catalyst bed to react with the methane and no further methane conversion is achieved. Oxygen is consumed near the membrane wall and does not reach the outer, external wall of the reactor, resulting in

699 unconverted methane. An alternative method of increasing the methane conversion is for the unreacted methane to participate in an additional reaction. A potential candidate is the reaction between methane and CO2 (dry reforming). So a specified amount of CO was added to the methane feed. It was found that the methane conversion and syngas selectivity decreased with the addition of CO2. Methane conversion decreased from 64% to 41% at a carbon dioxide feed rate of 145 cc/min, and CO and hydrogen selectivities decreased from 90% and 82% to 37% and 62% under these conditions. The decrease occurs because the added gas dilutes the methane concentration in the feed. This smaller CH4 concentration allows for a larger amount of CO2 to be produced because the local methane/oxygen concentration is relatively low. In addition, the reaction temperature remained around 510~ which is below the temperature at which the dry reforming reaction occurs. The fixed bed reactor can achieve CO and hydrogen selectivities of 93% with methane conversions around 70%. In the membrane reactor the highest methane conversion is about 65% with CO and hydrogen selectivities of 90% and 82% respectively. One of the reasons for this difference lies in the different concentration profiles in the membrane and fixed bed reactors. In the fixed bed, methane and oxygen are mixed throughout the reactor at the top of the catalyst bed, but in the membrane reactor, the oxygen concentration is much greater near the membrane tube than towards the outer wall of the reactor. The concentration profile determines the rate of reaction and thus the rate of heat evolution and consequently the temperature distribution in the bed, which is also an important factor in determining conversion. Further increases in methane conversion were attained by using an additional bed downstream from the membrane bed. In addition, the reactor temperature was increased so that the second bed operates at temperatures higher than the autothermal operation. This allows for the dry reforming reaction to occur in the second bed thus increasing the conversion of the methane not consumed in the first bed. In this case the highest methane conversion was about 90% with CO and H2 selectivities of about 90% when the external temperature is 700~ Similar results can be attained without heating if a third feed of 02 is added between the membrane bed and just before the second fixed bed. In this case, the temperature increase is realized by the partial oxidation reaction with no major loss of selectivity. At this point only a few catalysts have been studied in the membrane reactor. XRD and XPS analyses have been conducted mainly for the 3% Rh/TiO2 catalyst. XRD results show that after impregnation and calcination the TiO2 phase is mainly anatase. Some small peaks corresponding to metallic Rh are also detected. The XRD pattern for the catalyst after reaction shows that the TiO2 phase has changed from mainly anatase to rutile. It should be noted that the catalyst did not deactivate during the duration of these experiments. Furthermore, the results were reproducible when tested several times using the same batch of catalysts. So the phase change of the support does not seem to affect the activity of the Rh catalyst. XPS results indicate that the Rh 3d peak is located at 309.1 eV, while the reported binding energy for metallic Rh is 307.2 eV. After reaction, there is a significant change in binding energy to 306.7 eV. These results indicate that there is a change in the oxidation state of Rh during reaction. The reported binding energies of Rh203 and RhC1 are 308.8 and 310.1 eV respectively. The specific state of Rh cannot be determined by these ex-situ experiments and work is underway to establish the state of the catalyst surface after reaction. The changes occurring during reaction do not appear to significantly affect the catalyst activity and selectivity, at least during the duration of these experiments. BET surface areas for the calcined but unreacted 3%Rh~iO2 was 117.5 m2/g whereas the value for a 0.3%Rh/TiO2 catalyst was 43.7 m2/g. These results, obtained in a Quantachrome apparatus, were confirmed by several measurements. It is not clear, however, why the surface area of the catalyst with higher Rh loading is so high.

700 Hydrogen chemisorption results, obtained by the pulse method, indicated that the dispersion of Rh on the calcined 3%RhffiO2 catalyst was 56.8 % which corresponded to an average crystallite size of 1.9 nm. It should be noted that this value is rather low, probably because some of the surface was in an oxidized state as shown by the XPS results. In conclusion, the catalytic partial oxidation of methane on a RhffiO2 catalyst has been demonstrated to occur at high conversions and selectivities under fast flow conditions when a membrane is used to separate the methane and oxygen feeds. This configuration allows operation of the reactor at low methane/oxygen ratios and high feed rates while eliminating the possibility of a flame flashback leading to an explosion. The low methane/oxygen ratio and the high flow rates are the key factors to attain autothermal behavior. The RhffiO2 catalyst exhibited a low ignition temperature and did not exhibit deactivation during the duration of these experiments. REFERENCES

1. Prettre, M. Eichner, C. and M. Perrin, Trans. Fraday Soc., 43 (1946) 335. 2. Gavalas, G. R.Phichticul, C. and G. E. Voecks, J. Catal., 88 (1984) 54. 3. Blanks, R. E., Witrig, T. S.and D. A. Peterson, Chem. Eng. Sci., 45 (1990) 2407. 4. Vermeiren, W. J. M., Blomsma, E. and P. A. Jacobs, Catal. Today, 13 (1992) 427. 5. Ashcroft, A. T. Cheetham, A. K, Foord, J. S., Green, M. L. H., Grey, C. P., Murrell, A.J. and P. D. F. Vernon, Nature, 344 (1990) 319. 6. Hickman, D.A. and L.D. Schmidt, J. Catal 138 (1992) 267. 7. Hickman, L.D. and L.D,. Schimidt, AICHE J. 39 (1993) 1164. 8. Torniamen,P.M., Chu,X., and L.D. Schmidt, J. Catal, 146 (1994) 1. 9. Santamaria, J.M., Miro, E.E., and E.E. Wolf. 10. Shiraha, T., M.S. Thesis, Unversity of Notre Dame (1995).

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 1997 Elsevier Science B.V.

701

Sustainable N i / B a T i O 3 Catalysts for Partial O x i d a t i o n of M e t h a n e to Synthesis Gas R. Shiozaki, a A. G. Andersen, b T. Hayakawa, S. Hamakawa, K. Suzuki, M. Shimizu and K. Takehira* National Institute of Materials and Chemical Research, Tsukuba Research Center, AIST, Higashi 1-1, Tsukuba, Ibaraki 305, Japan. aChemical Research Department, Institute of Research and Innovation, Takada 1201, Kashiwa, Chiba 277, Japan. bNycomed Imaging, P.o.box 4220 Torshov, 0401 Oslo, Norway. Ni/BaTiO3 catalyst has been prepared by solid phase crystallization (SPC) method and used successfully for partial oxidation of CH4 into synthesis gas at 800~

The catalyst was

further tested for 75 hrs with no observable degradation and negligible coke formation. The

SPC method bestows the catalyst with high Ni dispersion over the perovskite as well as strong metal-support interaction between Ni and the perovskite, resulting in both high activity and high sustainability against coke formation.

I. I N T R O D U C T I O N Recently, intensive studies have been carded out on the catalytic partial oxidation of CH4 to synthesis gas [1-3]. This process has advantages over the conventional steam reforming of CH4 to make synthesis gas, as the latter process is highly endothermic and produces synthesis gas having a H2/CO ratio > 3. The partial oxidation of CH4, expected to afford synthesis gas having H2/CO ratio of about 2, makes methanol synthesis an ideal follow-up process. In the partial oxidation of CH4 to synthesis gas, coke formation over the catalyst frequently takes place, resulting in catalyst deactivation. Claridge et al. [4] observed that the relative rate of coke formation follows the order Ni>Pd>> Rh, Ru, Pt, Ir. Nickel catalysts are highly effective for partial oxidation of CH4 to synthesis gas, but they are unsatisfactory with respect to coke formation [5]. From the industrial view point, Ni is preferable as the active species compared to expensive precious metal such as Rh, Pd or Ru. High dispersion of metal species over catalyst may reduce coke formation [6]. Nickel-supported catalysts are conventionally

702 prepared by wet impregnation of different supports. This method is not fully reproducible and may give rise to some unhomogeneity in the distribution of the metal on the surface. Therefore, a new method of catalyst preparation able to produce homogeneous distribution of nickel is proposed, i.e., use of the precursors containing homogeneously distributed nickel inside in the structure, followed by further calcination and reduction, may result in the formation of well dispersed and stable metal particles. This method gives a higher reproducibility and easier characterization in comarison to the analogous samples prepared by wet impregnation. Here, we named this method "solid phase crystallization (SPC) method". We have reported sustainability and high activity of Ni/(Ca, Sr)TiO3 in the oxidation of CH4 to synthesis gas [7-10], where the Ni catalyst supported on (Ca,Sr)TiO3 perovskite was prepared by the SPC method. However, the precursor was not homogeneous and contained two types of nickel species; one was dissolved in the Ti site in (Ca, Sr)TiO3 and another was separated as NiO from the perovskite structure. Nickel dissolved in the perovskite structure was reduced to well dispersed and stable metal species, while NiO was reduced to large size of Ni metal particles after the reduction. The well dispersed and stable nickel particles seem to be important for operation at low coke formation over the catalyst during the reaction. BaTiO3 can contain higher amount of nickel in the Ti site compared to CaTiO3 and give rise to more completely dispersed metal on the surface during the SPC preparation. Here, we report the preparation and thermal evolution of new perovskite containing well dispersed and stable nickel on the surface, resulting in the sustainable activity for the partial oxidation of methane.

2. EXPERIMENTAL

Catalysts of the composition BaTi 1-xNixO3-5, 0'~_x99.95%), ethane (BOC, C.P. grade), hydrogen (BOC, C.P. grade), argon (BOC, C.P. grade), oxygen (BOC, C.P. grade) and carbon dioxide (BOC, C.P. grade) were used as received without further purification.

2.2. Apparatus The apparatus used in this work was a modified version of the commercial Labcon microreactor described previously [31 ]. Briefly, the reactor was built using 1/8" and 1/16" o.d. 316 stainless steel tubing and 316 stainless steel Swagelok fittings throughout. The catalyst sample was placed between two quartz wool plugs in the centre of a 4 mm i.d. silica tube, and inserted into a vertical Severn Science tube furnace. This was heated to the required reaction temperature controlled from a Eurotherm 905 temperature controller. For safety reasons, in experiments carried out at elevated pressures the silica tube was placed inside a steel tube. Inlet gas flow rates were controlled using Brooks 5850TR mass flow controllers, and the water flow was controlled using a Bronkhorst LIQUI-FLOW liquid flow controller with a helium over pressure and operated from a Bronkhorst HighTech EPA2 control module. The exit gas stream from the reactor passed through a Tescom two stage back-pressure regulator to allow elevated pressure experiments to be carried out. All the pipework was heated to prevent condensation of the products.

2.3. Product Analysis Product analysis was carried out using a Hewlett-Packard 5890II gas chromatograph, fitted with both a thermal conductivity detector, and a methanator/fiame ionisation detector. Separation of the products was achieved using a 3m Porapak Q packed column, with argon carrier gas. Reference data and pure component injections were used to identify the major peaks, and response factors for the products and reactants were determined and taken into account in the calculation of the conversion and product distribution. In all cases stoichiometric gas mixtures were used and carbon balances were better than 97%. Conversions and yields were calculated as follows: Conversion, C[CHa]or C[CO2] = (% conversion of CH4 or CO2 into all products). Carbon monoxide yield, Y[CO] = (CO in products)/((CO2 + CH4) in reagents) x 100.

3. RESULTS AND DISCUSSION

3.1. Synthesis of High Surface Area Metal Carbides Before the transition metal carbides could be assessed as catalysts it was necessary to prepare samples of high surface area. Boudart and his co-workers revealed that by using a CHaffI-I2 mixture and temperature programmed reduction (TPR) it was possible to prepare high surface area binary carbides. Further work by Claridge et al. showed that by substituting C2H6

714 for CH4 in the gas mixture during TPR carbidation, a further increase in surface area could be achieved. Table 1 shows that with this technique a wide variety of group V and VI transition metal oxides, with BET surface areas 140 hours). Furthermore, no traces of M002 or WO2 could be seen in the XRD patterns of the post-reaction samples, and no carbide phase changes were observed.

Table 2 Results for the dry reforming of methane over group V and VI transition metal carbides. (GHSV = 2.87 x 103 h -1, CH4]CO2 = 1) Catalyst T/K P / bar C[CH4]/% C[CO2]/% Y[CO]/% H2/CO NbCx 1223 8.0 t 67.6 77.3 72.4 0.82 1373 8.0 83.7 96.3 90.0 1.33 TaCx 1223" 8.0 t 54.7 61.5 58.1 0.67 1123 1.0 t 92.4 92.5 92.5 0.93 I]-Mo2C 1223 1.0 t 98.8 95.9 95.9 0.92 1123 8.3 62.5 75.9 69.5 0.78 1223 8.3 83.3 89.5 86.5 0.88 1123 1.0 t 92.0 93.1 92.6 0.94 c~-WC 1123 8.3 62.7 75.4 68.6 0.79 * catalyst deactivates; * initial result could not be obtained.

1400 1200

A1

>~ 800 600

9 MoO2 " 13-Mo2C

,, 9

Ai

9

9

(c) _.zz____3

*

400

.

All

9 i AI

. .* . . .

9

*

9

9

II

200

20

3'o

4'0

5'0

6'0

7'o

Bragg Angle (20) Figure 2. XRD patterns of (a) [~-Mo2C as prepared by CH4 TPR, (b) [~-Mo2C post-dry reforming (8 bar, 1123 K), and (c) ~-Mo2C post-dry reforming (1 bar, 1123 K).

717

The primary cause for deactivation in current reforming catalysts, such as supported nickel, is the formation of carbon on the catalyst during reaction [20]. However, when postreaction samples of ~-MozC and a-WC were examined by HRTEM, no observable carbon deposition had occurred on the catalyst surface during the reaction. In addition, activity studies demonstrated that Mo2C had a methane dry reforming activity similar to an active supported noble metal catalyst, namely 5% Ir/Al203 [27]. Table 2 also shows that the group V transition metal carbides, NbCx and TaCx, are less active than the carbides of the group VI metals. At 1223 K, both NbCx and TaCx deactivated rapidly, even at elevated pressure, while catalyst stability and high reactant conversions were achieved at 1373 K with NbCx, where autothermal gas-phase reactions are likely to play a significant role in the reaction and carbon deposition was observed, as is demonstrated by the high H2/CO ratio. The reason for the low activity of these materials is the relative ease of their conversion back to the oxide, i.e. the rate of carbidation over these materials is slower than the rate of oxidation. Since MozC and WC were found to be active for the dry reforming of methane, they were also tested as catalysts for the partial oxidation of methane with air and methane steam reforming. As before, the product distributions were close to those predicted by thermodynamic equilibrium calculations, and deactivation occurred at atmospheric pressure due to oxidation of the carbide to inactive MoO2; in fact, with 02 (air) as the oxidant the catalyst deactivation was too fast for the initial activity to be measured. However, Table 3 shows that at elevated pressures the catalysts were highly active for both reactions, while no catalyst deactivation was observed for the duration of the experiments (>72 h), even with the strongest oxidant, i.e. oxygen. This is surprising, since the high surface area carbides are extremely reactive towards oxygen, even at room temperature, but is probably due to the relative rates of carbidation/oxidation during the reaction. At atmospheric pressure the oxidation is fastest, while the carbidation is favoured by elevated pressures.

Table 3 Catalytic results for the partial oxidation and steam reforming of methane using ~-WC. (GHSV = 5.2 x 103 h ~ (with air); 2.8 x 103 h -1 (with H20)) Catalyst CH4/ P/ T/ C[CH4]/% Y[CO]/% oxidant bar K ~-MozC 2:1 (air) 4.0 1173 94.7 89.6 2:1 (air) 8.7 1073 73.5 59.2 2:1 (air) 8.7 1173 88.5 81.3 a-WC 2:1 (air) 8.7 1173 87.5 78.2 ~-Mo2C 1:1 (H20) 1.0* 1223 91.5 90.1 1:1 (H20) 8.3 1223 81.8 77.7 8.3 1223 81.9 77.8 a-WC 1:1 (H20) catalyst deactivates

~-MozC and H2/ CO 2.01 2.07 2.01 2.04 3.08 3.07 3.06

718

3.3. Dry Reforming of Methane Using Ternary Metal Carbides as Catalysts Following the development of synthetic routes to produce single phase group V and VI ternary carbides, and the observation that the carbidation temperature of the ternaries was lower than the group V binary carbides, it was postulated that the synergistic combination of the group V and VI transition metals may result in the synthesis of highly active and stable carbide catalysts for methane dry reforming.

Table 4 Results for methane dry reforming over the ternary metal carbides (C2H6 TPR; T = 1223 K, p = 8 bar, GHSV = 5 x 103 h-1) Catalyst precursor C [ C H 4 ]/ % C[CO2]/ % Y[CO] / % MoO3 83.3 89.5 86.5 Mo2Ta2Oll 86.8 89.8 88.3 Mo3Nb2Ol4 82.0 89.9 86.0 Mo3Nb14044 47.7 64.3 56.0 W9Nb8047 73.3 92.1 82.7 W3Nb14044 26.1 22.8 24.4 WNb12033 22.4 25.5 24.0

H2/CO 0.88 0.84 0.90 0.60 0.83 0.78 0.63

4OOO

3O0O

(-. :3

0

23OO

o

i

i

I

~

(c~

1000

J '

II0

'

I

20

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i

3o 20

,

i

40

,

i

50

'

i

m

,

i

7o

Figure 3. XRD patterns for (a) MoETaECx post-catalysis (1223 K, 8 bar), (b) Mo2Ta2Oll as prepared, and (c) Ta2Os.

At first glance, the results in Table 4 indicate that some of the ternary metal carbides, namely those synthesised from Mo2Ta2Oll, Mo3Nb2O14, and W9Nb8047, are highly active for

719 methane dry reforming, with reactant conversions and yields very similar to those obtained with [~-MozC. Further, the catalysts appeared to be stable during lifetime investigations of more than 50 hours (not shown). However, a comparison of the pre- and post-reaction XRD patterns revealed that all the single phase ternary carbides are phase separating during the reaction. In fact, as was the case with the binary metal carbides, the group V metal carbides are oxidised, leading to the formation of the group V metal oxide (Nb205 or Ta2Os), while the group VI metal carbide is stable under these conditions (8 bar). Figure 3 shows that the postreaction XRD pattern of the carbide of MozTazOll has a strong resemblance to the pattern for pure Ta2Os; no evidence for any tungsten oxides was observed, and it is probable that the tungsten is present in the sample in the form of tungsten carbide, which is characterised by weak and broad peaks, and thus cannot be seen. Similar XRD patterns were also obtained for the other ternary carbides, and these observations were also supported by TEM and energy dispersive X-ray analysis (not shown). Therefore, it would appear that the stability of some of the ternary carbides is actually due to the presence of stable Mo2C or WC. This is further confirmed by the observation that only those carbides with a high ratio of group VI to group V metal are active for dry reforming, while in the other ternaries, such as WNblzCx, the amount of tungsten is small and the catalyst surface mainly consists of the metal oxide; therefore, the number of catalytically active carbide species at the surface is very small.

4. CONCLUSIONS It has been demonstrated that the use of ethane in the temperature programmed synthesis of transition metal carbides results in the formation of materials with higher surface areas than with methane TPR. In addition, the conversion process with ethane appears to proceed in a topotactic fashion. Of all the materials synthesised, only molybdenum and tungsten carbide were active and selective for the stoichiometric dry reforming, partial oxidation and steam reforming of methane to synthesis gas. These materials deactivated at atmospheric pressure, but were very stable when elevated pressures were employed, and no carbon deposition was observed on the catalysts. Niobium and tantalum carbide were inactive for the dry reforming of methane, due to their rapid deactivation by oxidation. Some of the ternary carbides were found to be active and stable for dry reforming, but the catalyst activity could be attributed to the presence of molybdenum or tungsten carbide. In fact, the group V metal appeared to play no role in the processes contributing to the catalytic reforming of methane to synthesis gas. The observations presented in this paper are potentially important, since these materials offer the chemical industry cheap and abundant, non-coking reforming catalysts, where previously the choice was between nickel, which promotes carbon formation, and expensive noble metals.

REFERENCES 1. J.M. Muller and F. G. Gault, Bull. Soc. Chim. Fran., 2 (1970) 416. 2. R.B. Levy and M. Boudart, Science, 181 (1973) 547. 3. L. Volpe and M. Boudart, J. Solid State Chem., 59 (1985) 348. 4. J.S. Lee, S. T. Oyama and M. Boudart, J. Catal., 106 (1987) 125. 5. D. Zeng and M. J. Hampden-Smith, Chem. Mater., 4 (1992) 968.

720 6. 7. 8. 9. 10. 11.

12.

13. 14. 15. 16. 17. 18. 19. 20. 21. 22. 23. 24. 25. 26. 27. 28. 29. 30. 31. 32. 33.

M.J. Ledoux and C. Pham-Huu, Catal. Today, 15 (1992) 263. J. Lemaitre, B. Vidick and B. Delmon, J. Catal., 99 (1986) 415. L. Leclercq, M. Provost, H. Pastor, J. Grimblot, A. M. Hardy, L. Gengembre and G. Leclercq, J. Catal., 117 (1989) 371. F. Sherif and W. Vreugdenhil, in "The Chemistry of Transition Metal Carbides and Nitrides", S. T. Oyama (ed.), p.414, Blackie Academic & Professional, Glasgow, 1996. G.B. Raupp and W. N. Delgass, J. Catal., 53 (1979) 361. C. Pham-Huu, M. J. Ledoux and J. Guille, J. Catal., 143 (1993) 249. M. J. Ledoux, C. Pham-Huu, A. P. E. York, E. A. Blekkan, P. Delporte and P. Del Gallo, in "The Chemistry of Transition Metal Carbides and Nitrides", S. T. Oyama (ed.), p.373, B lackie Academic & Professional, Glasgow, 1996. F. H, Ribeiro, R. A. Dalla Betta, M. Boudart, J. Baumgartner and E. Iglesia, J. Catal., 130 (1991) 86. F. H. Ribeiro, R. A. Dalla Betta, M. Boudart and E. Iglesia, J. Catal., 130 (1991) 498. J. C. Schlatter, S. T. Oyama, J. E. Metcalfe III and J. M. Lambert Jr., Ind. Eng. Chem. Res., 27 (1988) 1648. J. S. Lee and M. Boudart, Appl. Catal., 19 (1983) 207. L. Volpe and M. Boudart, J. Phys. Chem., 90 (1986) 4874. L. Volpe and M. Boudart, J. Solid State Chem., 59 (1985) 332. J. T. Wrobleski and M. Boudart, Catal. Today, 15 (1992) 349. J. R. Rostrup-Nielsen, in "Catalysis, Science and Technology", J. R. Anderson and M. Boudart (eds.), vol. 5, p. 1, Springer, Berlin, 1984. S. T. Sie, M. M. G. Senden and H. M. H. Wechem, Catal. Today, 8 (1991) 371. S. C. Tsang, J. B. Claridge and M. L. H. Green, Catal. Today, 23 (1995) 3. J. R. Rostrup-Nielsen, J. Catal., 85 (1984) 31. J. R. Rostrup-Nielsen, J. Catal., 31 (1973) 173. A. T. Ashcroft, A. K. Cheetham, M. L. H. Green and P. D. F. Vernon, Nature, 352 (1991) 225. J. B. Claridge, M. L. H. Green, S. C. Tsang, A. P. E. York, A. T. Ashcroft and P. D. Battle, Catal. Lett., 22 (1993) 299. A. P. E. York, J. B. Claridge, A. J. Brungs, S. C. Tsang and M. L. H. Green, J. Chem. Soc., Chem. Commun., (1997) 39. J. B. Claridge, A. J. Brungs and M. L. H. Green, Chem. Mat., submitted. M.W. Viccary and R. J. D. Tilley, J. Solid State Chem., 104 (1993) 131. T. Ekstrom, Acta Chem. Scand., 25 (1971) 2591. J. B. Claridge, M. L. H. Green, S. C. Tsang and A. P. E. York, Appl. Catal., 89 (1992) 103. E. K. Storms, The Refractory Carbides, Academic Press: New York, 1967, vol. 2. S. T. Oyama, R. Kapoor, H. T. Oyama, D. J. Hofmann and E. Matijevic, J. Mater. Res., 8 (1993) 1450.

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

721

Partial o x i d a t i o n of m e t h a n e to s y n t h e s i s gas using L n C o O 3 p e r o v s k i t e s as catalyst precursors R. Lag&, G. Bini b, M.A. Pefia" and J.L.G. Fierro" 'llnstituto de Catfilisis y Petroleoqu/mica, CSIC, Campus UAM, Cantoblanco, 28049 Madrid, Spain [FAX +34 1 585 4760; E-mail: [email protected]] bDipartimento de Chimica lndustriale e Materiali, University of Bologna, Vie. Risorgimento 4, 40136 Bologna, Italy

In this work Co supported catalysts have been prepared by the reduction of the perovskites LnCoO~ (Ln=La, Pr, Nd, Sm and Gd) to produce Co"/Ln20~. Detailed TPR and XRD studies showed that the perovskite NdCoO3 is reduced in two steps, first to NdCoO2s and further to Co~ and in both stages it was demonstrated that the reoxidation with Oe is capable to recover the perovskite structure. TPO experiments with reduced Ln-Co-O (Ln = La, Nd, Sm and Gd) catalysts indicated that reoxidation takes place in two steps: first oxidation of the supported Co" to the spinel Co~O4 (Co:+Co3+204) and further the oxidation of the Co :+ to Co 3+ with a simultaneous solid state reaction with Ln20~ regenerating the perovskite structure. It was observed that the temperature for the second oxidation step is strongly dependent on the nature of the lanthanide increasing in the following order La > Nd > Sm > Gd. This trend seems to be determined by the thermodynamic stability of the parent perovskite. These catalysts (Co"/Ln:O3) have been tested for the partial oxidation of methane to synthesis gas showing remarkable differences in activity [1]. The system Gd-Co--O showed exceptionally better performance for CO and H2 production whereas the activity for the other catalysts decreased in the following order Sm-Co-O > > Nd-Co-O > Pr-Co-O. The catalyst La-Co-O was active for methane combustion and only traces of CO and H: were observed. It was found that these differences are determined by the lanthanide which plays a fundamental role on the stability of the catalyst.

1. INTRODUCTION Perovskites of the type LnMO 3 (Ln = lanthanide and M - transition metal) may offer interesting features as precursors for supported metal catalysts. For example, careful reduction can be carried out in order to produce a finely dispersed transition metal over the sesquioxides Ln:O~. Also, the flexibility in the perovskite composition allows the preparation of compounds of the type LnM'~_,,M"xO~ (M' and M" = different transition metals) or Ln~_,,A,,MO~ (A = for example an alkaline or alkaline earth metal), which show the unique possibility as precursor of producing well dispersed bimetallic catalyst or doped metal catalyst. The research on CH4 conversion has recently focused on the direct oxidation to synthesis gas:

722

CH4 + 1 / 2 0 2 - - >

CO +

2H2

( A N~ ,, 2,~8~ = - 35.6 kJ/mol)

(1)

The major advantages of this route over the steam reforming are the H:/CO ratio of ca. 2 suitable for downstream processes and the exothermicity of the reaction which eliminates the need for a fuel gas [2]. Many catalysts for the partial oxidation of methane to synthesis gas, consisting of supported metals such as Ni, Co and Fe and noble metals Pd, Ir, Rh, Ru, Pt, etc have been described in the literature [2-9]. Hayakawa et al. [ 11] studied the perovskite Ca~_,,Sr,,TiO3 mixed with nickel oxide for methane oxidation at 1028 K, both before and after pre-treatment with methane. Before methane pretreatment the catalyst produced mainly C2 hydrocarbons and CO2 but no syngas, whereas after methane pre-treatment at 1048 K for 1 h, the mixture became very selective to synthesis gas giving 70.9% CH4 conversion with 94% selectivity to CO and H 2. Based on XRD analyses it was proposed that the methane pre-treatment produced Ni" supported on the perovskite which was responsible for the synthesis gas formation. Similar results were obtained for cobalt and iron (8). Slagten and Olsbye [10] studied the systems La-M-O (M = Co, Ni, Rh and Cr) for the partial oxidation of methane to syngas. They observed very high activity for the system La-Rh-O whereas the catalyst La-Co-O (which was a mixture of LaCoO3, La203 and Co304) produced mainly CO:. If the catalyst La-Co-O was kept at 1073 K after 30 h reaction the activity changed to give mainly CO which they assigned to the in situ reduction of cobalt. In this work we have studied the solid state reactions taking place with supported cobalt metal catalysts prepared by the reduction of LnCoO 3 precursors under the experimental conditions of the partial oxidation of methane.

2. EXPERIMENTAL All the perovskites were prepared by the method of the amorphous citrate percursor [12] and were present as pure phases with no contaminants of LnzO3 or cobalt oxides according to the powder XRD analyses. The catalytic tests were carried out in a quartz fix bed microreactor with 50 mg of catalyst and a 2:1"4 mixture of CH4:O:'Ar at 166.6 ml/min. All the catalysts were reduced in situ with a H2 (33%):Ar mixture at 1023 K for 3.5 h prior the reaction. Two thermocouples were used, one outside the reactor to control the furnace temperature and the other one inside in contact with the catalyst to measure the bed temperature. X-ray photoelectron spectroscopy (XPS) analyses were carried out in a ESCALAB 200R spectrometer equipped with a MgKa 120 W X-ray source (hv = 1253.6 eV). After reaction the catalysts were all quenched to room temperature under argon and immediately drenched in isooctane. Despite this careful procedure to isolate the catalyst it was not possible to completely avoid oxidation of superficial metallic cobalt during the transfering of the sample from the reactor to the XPS spectrometer. Powder XRD were obtained in a Seifert 3000P diffractometer. Temperature-programmed reduction (TPR) and oxidation (TPO) analyses were carried out in a Micromeritics 2900 apparatus. For the Co" area determination, H2 chemisorption measurements were performed in a conventional volumetric adsorption apparatus. A 200 mg sample of the perovskite was reduced at 1023 K for 3.5 h under a stream of 33% H2/N2 and then outgassed at 773 K overnight at pressures lower than 10 -4 Torr. The measurements were carried out at 315 K and the reproducibility was within 10%.

723 3. RESULTS 3.1. Reduction of the Cobalt Containing Perovskites TPR profiles of the cobalt containing perovskites are displayed in Fig. I. All the perovskites showed similar reduction profiles consisting of two sets of peaks at approximately 633 and 833 K. It can be observed that especially for LaCoO3 and PrCoO 3 the second reduction peak shifts to higher temperatures. For example, the reduction peak for SmCoO3 at c a . 785 K is observed at 844 K for LaCoO3. In all cases the hydrogen consumption for the first reduction step (peak at c a . 633 K) was always approximately half of the hydrogen consumption obtained for the second reduction step (peak at 833 K). Careful thermogravimetric reduction experiments (not shown here) demonstrated that the first step is a 1 electron reduction process whereas the second step is a 2 electrons reduction process. To identify the reduced species formed in the TPR steps in Fig. 1, XRD studies of NdCoO 3 A

-k 9

&

Q

__

_a =

Gd-Co-O

C /

C

v

~qoU]

~

,~

IC C

t ?'i" 9

At'?

,_ 5,,L_.Y.__.2_J,._..,' -b._.,,,.._...._..,,._~

=

=

C

c

Sm-Co-O B

B

Nd-Co-O Pr-Co-O La-Co-O I

400

A tO ~,,.

I

600

,

I

~()0

T e m p e r a t u r e (K)

1()()()

,

7o

6tl

50

40

30

21)

2 0 (")

Figure 1. TPR profiles of LnCoO3 perovskites. Figure 2. XRD patterns for NdCoO3: (a), as prepared" (b), reduced in H: at 623 K for 15 min; (c), after reduction (623 K, 15 min) and heated under He at 1123 K for 1 h (sintering); (d), reduced in H2 at 778 K for 15 min" (e), after reduction (778 K, 15 min) and sintering; (f), after reduction (778 K, 15 rain), sintering at 1123 K and reoxidation under O2 flow at 973 K. A = N d C o O 3 " B = N d : O 3 ( c u b ) ; C = N d : O 3 ( h e x ) " D = C o O ; E = Co~O4" a n d F = C o ~.

724 subjected to different reduction treatments were carried out (Fig. 2). The XRD pattern displayed in Fig. 2a corresponds to NdCoO~ cubic perovskite structure. Upon reduction at 623 K under H2 flow for 15 min the diffraction peaks are essencially at the same position but with a very strong broadening and in some peaks a splitting is observed (Fig. 2b). This indicates that the perovskite structure is basically the same however distorted due to the creation of anion vacancies. These results suggest that the stoichiometry of the phase formed in the first step 1 electron reduction can be represented as NdCoO2.~. The reoxidation of this sample at 973 K under oxygen flow can easily regenerate the original perovskite structure (XRD not shown here) demonstrating the reversibility of this process. On the other hand, the phase NdCoOz.5 is not stable and at high temperatures (1123 K for 1 h) under He phases of CoO and Nd203 hexagonal are formed (Fig. 2c). When the perovskite was treated with hydrogen at 776 K for 15 min, a 3 electron reduction was observed and XRD analysis showed the presence of Nd203 hexagonal and very weak and broad peaks which could be due to metallic coba!t (Fig. 2d). Upon heating at 1123 K in helium the highly dispersed cobalt metal sinters and shows much stronger lines in the XRD pattern (Fig. 2e). It was also observed the presence of small amounts of CoO, probably formed by the reoxidation of cobalt by water or oxygen still present in the sample during sintering. The reoxidation of the reduced sample in Fig. 2d (before sintering) also reproduced the perovskite original structurewith the same XRD pattern (not shown here). However, after sintering (Fig. 2e) the reoxidation produces a mixture of the phases NdCoO3, Co304 and Nd203 (Fig. 2f), showing that the process is not completely reversible. It was observed that if the sample in Fig. 2f was kept under oxygen flow at temperatures of 1123 K for 2 h a slow solid state reaction takes place regenerating the perovskite NdCoO~ structure. Therefore, the two main reduction steps of NdCoO 3 can be written as" 1 e/mol"

2NdCoO~ + H: ~

3 e/mol"

2 NdCoO 3 + 3H: ~

2NdCoOes

+ H:O

Nd203 + 2Co" + 3H20

(2) (3)

These reactions are reversible and as long as the metallic cobalt remains well dispersed the perovskite structure can be easily regenerated by oxidation. These results are in agreement with the work of Crespin and Hall [ 13], who observed basically the same reduction steps for LaCoO 3. The cobalt metal area of the reduced perovskites was determined by hydrogen chemisorption experiments. The results are shown in Table 1. The chemisorption measurements revealed that the cobalt metallic surface area was similar for all the perovskites. This is supported by the Co/Ln surface ratio (Table 1) obtained by XPS which also suggests similar metallic dispersion. The XPS analyses of the reduced perovskites showed the presence of Co" (778.6 eV) but also a doublet at approximately 780.5 and 796.2 eV which correspond to Co 2p3/2 and Co 2p~/: peaks respectively, for the Co -~+ ion. Shake-up satellite lines with 4.7 eV over the Co 3+ lines were also detected indicating the presence of Co :+ [12]. These oxidised species of cobalt are probably formed by air oxidation during the transference of the reduced sample from the reactor to the XPS spectrometer. Also, Marcos et al. [15] have shown that the reduction of the perovskite LaCoO 3 produced a La203 oxide covered by hydroxyl groups which upon heating and evacuation in the XPS pretreatment chamber partly reoxidises the cobalt crystallites. Powder XRD analysis of the reduced LnCoO~ perovskites showed only the presence of the sesquioxides La20~ (hex), NdzO~ (hex), Sm203 (cub) and Gd203 (cub). The apparent absence of reflections for the metallic cobalt indicates a high metallic dispersion with Co" particles smaller

725

Table 1 Characterization of the reduced perovskites Precursor

BET area (m:/g)"

Co" area (atom/g) ~'

LaCoO3 NdCoO3 SmCoO3 GdCoO 3

8.6 (5.4) 5.6 (4.5) 5.1 4.0 (5.0)

3.77- 10 I'~ 3.98-10 ~'~ 3.69.10 ~'~ 3.21.10 j''

XPS Co/La ratio (reduced cat.y 0.71 0.50 0.58 0.75

XRD (reduced cat.) La203 (hexagonal) Nd203 (hexagonal) Sm203 (cubic) Gd:O3 (cubic)

a BET areas for reduced catalysts are in parenthesis; b Catalysts reduced at 1023 K under H2 (33%) in Ar; c For Co/La ratio measurement all Co signals (reduced, oxidised forms and shakeup lines) were included. than 2 nm. Considering the approximate value of 3.10 t'~ atoms/m: for a close packing arrangement [ 16], the number of surface cobalt metallic atoms obtained for the reduced catalysts is similar to the number of surface cobalt ions on the original perovskite surface (between 2.10 ~~ to 5.10 ~~ at/g for GdCoO 3 with BET surface area of 4.0 m:/g and LaCoO 3 8.6 m2/g, respectively). Therefore, apparently after reduction of the perovskites great part of the cobalt metal is located in the bulk of the material.

3.2. Reoxidation of Co~ TPO experiments were performed for the reduced catalysts in order to investigate the stability of these systems towards gas phase oxygen (Fig. 3). The perovskites were completely reduced in a TPR experiment with a stream of 10% H: in Ar heating from room temperature to 923 K at a rate of 10 K/min. The temperature was then rapidly decreased to avoid sintering of the cobalt metal and TPO analysis performed. TPO profiles showed that the oxidation takes place in two steps. Thc first one at near 473 K (peak I) has an oxygen consumption which suggests that the cobalt is oxidised to two Co ~+ and one Co :+. Therefore, this oxidation probably corresponds to the formation of the cobalt oxide spinel according to the reaction" 3Co" + 202

~

Co304

(4)

The second step (peak II) takes place at a much higher temperature and the 02 consumption suggests the following process:

Co304 --I- 1/4 02 + 3 Ln203 ~

LnCoO 3

(5)

In this step apparently the Co :+ is oxidised to Co 3+ simultaneously with a solid state reaction to regenerate the perovskite structure. The occurrence of these reactions are also supported by a simple experiments involving sequences of TPO and TPR measurements, shown in Fig. 4A and 4B. It can be observed that the TPR profile, obtained with the sample reduced in the first TPR (Fig. 4B, profile a), and reoxidised in the first TPO up to 1123 K (Fig. 4A, profile a) is very similar to the TPR of the fresh sample (Fig. 4A, profile a) suggesting the presence of a SmCoO3 perovskite structure (Fig. 4B, profiles a and b). On the other hand, if the reduced sample (after

726

Nd-Co-O --

c

400

600

800

1000

12(

Temperature (K) Figure 3. TPO profiles for the reduced systems La-Co-O, Nd-Co-O, Sm-Co-O and Gd-Co-O. TPR) is reoxidised in a TPO experiment which is interrupted at 623 K and the sample heated up to 1123 K under He, the TPR analysis showed only a broad peak at ca. 623 K (Fig. 4B, profile c) which is similar to the reduction of the spinel Co304 reported in the literature [17]. It is interesting to note in Fig. 3 that the temperature for the first oxidation step is similar for all the catalysts whereas the temperature of the second step is strongly dependent on the nature of the lanthanide. For La, oxidation occurs at 939 K whereas for Gd a much higher temperature is necessary (1109 K). Also the area and the shape of the second peak varied significantly for the different lanthanides. Very sharp peaks were obtained for Ga and Sm whereas a broad and

02

:11

A

B

gl,

'2 //-"~

r

eq

/ /

0 inl c riu p ted C

I

400

~

1

~

60t)

I

800

T e m p e r a t u r e (K)

~

I

1000

,

f

-

,

120{}

J

I

600

750 T e m p e r a t u r e (K)

Figure 4. Sequences of TPO (A) and TPR (B) experiments with the system Sm-Co-O.

900

727 weak peak was observed for La. The Apcak i/Apcak II area ratio of the first and the second peak from the TPO profiles shown in Fig. 3 are 12.5; 10.0; 9.0 and 7.5 for La-Co-O, Nd-Co-O, Sm-Co-O and Gd-Co-O, respectively. The systems La-Co-O and Nd-Co-O can be observed to deviate significantly from the expected area ratio which is 8.0 considering the stoichiometry of Eqs. [4] and [5]. This suggests that part of the cobalt which in the Gd-Co-O and Sm-Co-O systems is only oxidised at temperatures higher than 973 K, is being oxidised at lower temperatures for La-Co-O and Nd-Co-O. In fact, Crespin et al. [23] showed that if the reduced LaCoO 3 was kept at a temperature as low as 673 K under oxygen the Co" was completely reoxidised regenerating the perovskite structure. 3.2. Influence of the ianthanide on the reduction-oxidation of the system Ln-Co-O As observed by TPR, the nature of the Ln affects the reducibility of Co in the perovskites LnCoO 3. The Goldschmidt tolerance factor t - (r~,+ ro)/[~2 (r(.o+ro) ] obtained for the structures of LaCoO3, PrCoO3, NdCoO3, SmCoO3 and GdCoO3 were 0.899, 0.885, 0.878, 0.867 and 0.857, respectively. These tolerance factors indicate that considering solely geometric factors lanthanum, the largest ion in the series, forms the most stable perovskite structure. This trend is reflected in the TPR results where the perovskite LaCoO3, the most stable structure, is reduced at the higher temperatures, 844 K (Fig. 5). Likewise, the TPO experiments showed that reoxidation of cobalt to form the perovskite structure is more favorable for larger lanthanides. Figure 5 shows a good correlation between the oxidation temperature obtained from the TPO profiles with the Goldschmidt's tolerance factor. Katsura et al. [ 18] studied the thermodynamics between 1473 and 1673 K of the oxidation of iron to the rare earth perovskites according to the reaction:

Fe(s) + 1/2 Ln203 + 3 / 4 0 : (g) ~

(6)

LnFeO3 (s)

They showed that reoxidation to form the perovskite structure is favored in the order La > Nd > Sm > Gd with standard Gibbs energies o f - 2 8 8 . 0 , -274.6, -267.9 and -263.7 kJ/mol, respectively. k;i o

0.90

[]

Pr c]

d

Nd El

.~ 0.88 t_ ..... SITI cl 0.86

Gd

o 780

i

I

800 Reduction

i

I

i

820 Temperature

TPR (K)

I

I

I

840

950

1000 Oxidation

i

Temperature

I

~

1050

I

1100

T P O (K)

Figure 5. Goldschmidt's tolerance factor t versus (a) reduction and (b) oxidation temperatures obtained by TPR and TPO experiments for the perovskites LnCoO~.

728

3.3. Catalytic Testing Among the cobalt containing perovskites GdCoO3, SmCoO~, NdCoO~, PrCoO3 and LaCoO3 tested as catalyst precursors for the partial oxidation of methane the Gd-Co-O system showed exceptionally better performance for synthesis gas formation (Figs. 6A-6C). At 1009 K a steadystate methane conversion of 73% with selectivities of 79 and 81% for CO and H2, respectively, is observed for the catalyst Gd-Co-O. The catalysts Sm-Co-O and Nd-Co-O, of lower activity, show similar steady-state methane conversions in the temperature range studied. On the other hand, the H2 and CO selectivities are much higher over Sm-Co-O. The catalyst La-Co-O is active for the methane combustion and only traces of H2 and CO were observed under the reaction conditions used (Fig. 6A-6C). It is interesting to observe that although the reduced perovskites possess similar cobalt metallic areas, as revealed by H2 80 A

~ ~ - - - - [] . - - - - ' ~ []

' ~ o

60

...__~0

411 211-

~

~

0

91)!I

800 100

75

~

B

o_i ~ o ~ O ~

1000

Gd-Co-O

.______-----o-~ - - - [ ]

~

1100

Sm-Co-O

~

~.. ":

51)

-:'

25

,._co_o

-

o ~

____+__________T_o~

(} 800

Pr-Co-O La-Co-O

900

1000

I 1O0

1000

1100

I00

c ~"

75

~

_

~

o

"= 50 25 800

900

T e m p e r a t u r e (K)

Figure 6. (A), Methane conversion and (B), H2 and (C), CO selectivities in the presence of LnCo-O systems prereduced under a 30% HJAr flow at 1023 K for 3.5 h.

729

Table 2 XRD and XPS analyses of the catalyst after reaction" Precursor

Co/Ln XPS ratio b

LaCoO3 NdCoO.~ SmCoO~ GdCoO~

0.7 0.2 0.3 0.5

XRD LaCoO3 and La203 c NdCoO~ and Nd203 (hexagonal) SmzO3 (cubic) Gd203 (cubic)

'~ After reaction reaction at 1023 K for 19 h; b For Co/Ln ratio measurement all Co signals (reduced, oxidised forms and shake-up lines were included" c traces chemisorption and XPS data, they showed strikingly different catalytic properties. XRD and XPS characterization data of the catalysts after reaction are summarized in Table 2. The XPS Co/Ln ratios suggest that the Co dispersion over the catalyst surface after reaction follows the order LaCo-O > Gd-Co-O > Sm-Co-O > Nd-Co-O. XRD analyses of the used catalysts Gd-Co-O and Sm-Co-O showed similar patterns to the reduced catalysts with very strong and sharp peaks for the sesquioxides Gd203 and Sm203. On the other hand, the XRD analysis of the La-Co-O catalyst after reaction at 1023 K for 19 h clearly showed the formation of the perovskite LaCoO 3. Therefore, it is not surprising that the only reaction products observed were water and carbon dioxide. This agrees with previous works on this perovskite and other forms of cobalt oxide which have been shown to be active catalysts for methane combustion and also for CO and H: oxidation [19]. The high Co/Ln surface ratio determined by XPS for the used catalyst is expected for a perovskite like surface. Slagten and Olsbye [ 10] studied the perovskite LaCoO~ (containing some impurities of La20~ and Co~O4) for the partial oxidation of methane to syngas and observed the production of mainly CO,,. If the catalyst was kept at 1073 K after 30 h on-stream the activity changed to give mainly CO which they assigned to the in situ reduction of cobalt. The XRD for Nd-Co-O after reaction revealed the presence of the phases Nd20~ and also the perovskite NdCoO3. For all used catalysts no clear evidence for the presence of simple cobalt oxides such as CoO, Co,O~ and Co304 could be found by XRD.

4. CONCLUSION This work suggests that the high activity and selectivity of the catalysts Gd-Co-O and Sm-Co-O for the partial oxidation of methane to synthesis gas is due to the stability of the cobalt in its reduced state over the sesquioxides Gd:O3 and Sm20~. In the case of La-Co-O and Nd-Co-O reoxidation of cobalt to the original perovskite structure causes loss of activity and selectivity. TPO experiments with reduced Ln-Co-O (Ln = La, Nd, Sm and Gd) catalysts indicated that reoxidation takes place in two steps: first oxidation of the supported Co ~ to the spinel Co304 (Co2+Co3+204) and further the oxidation of the Co `'+ to Co ~+ with a simultaneous solid state reaction with Ln20~ regenerating the perovskite structure. It was observed that the temperature for the second oxidation step is strongly dependent on the nature of the lanthanide increasing in

730 the following order La > Nd > Sm > Gd. This trend seems to be determined by the thermodynamic stability of the parent perovskite.

Acknowledgements This work was supported by CICYT, Spain (Contract MAT95-0894). One of the authors (R.M.L.) is grateful to the Ministerio de Educaci6n y Ciencia, Spain, for the award of postdoctoral fellowship.

REFERENCES .

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6. 7.

10. 11. 12. 13. 14. 15. 16. 17. 18. 19.

R.M. Lago, G. Bini, M.A. Pefia and J.L.G. Fierro, J. Catal., (1997), in press. B.M. Tindall and M.A. Crews, Hydroc. Proc., 11 (1995) 75. S.T. Tsang, J.B. Claridge and M.L.H. Green, Catal. Today, 23 (1995) 3. A.T. Ashcroft, A.K. Cheetham, J.S. Food, P.D.F. Vernon and M.L.H. Green, Nature, 344 (1990) 319. D.A. Hickman and L.D. Schmidt, Science, 259 (1993) 343. D. Dissanayake, M.P. Rosynek and J.H. Lunsford, J. Phys. Chem., 97 (1993) 3644. D. Dissanayake, M.P. Rosynek, K.C.C. Kharas and J.H. Lunsford, J. Catal., 132 (1991) 117. F. Looji, J.C. Giezen, E.R. Stobbe and J.W. Geus, Catal. Today, 21 (1994) 495. T. Hayakawa, A.G. Andersen, M. Shimizu, K. Suzuki and K. Takehira, Catal. Lett., 22 (1993) 307. A. Slagten and U. Olbsbye, Appl. Catal., 110 (1994) 99. T. Hayakawa, A.G. Andersen, M. Shimizu, K. Suzuki and K. Takehira, Catal. Today, 24 (1995) 237. J.M.D. Tascon, S. Mendioroz and L.G. Tejuca, Z. Phys. Chem. NF, 124 (1981) 109. M. Crespin and W.K. Hall, J. Catal., 69 (1981) 359. T.J. Chuang, C.R. Brundle and D.W. Rice, Surf. Sci., 60 (1976) 286. J.M. Marcos, R.H. Buitrago and E.A. Lombardo, J. Catal., 105 (1987) 95. G.A. Somorjai, "Introduction to Surface Chemistry and Catalysis", Wiley Interscience Pub., New York, 1994. J.G. Choi, Catal. Lett., 35 (1995) 291. T. Katsura, K. Kitayama, T. Sugihara and M. Kimizura, Bull. Chem. Soc. Jpn., 48 (1975) 1809. M. Futai, C. Yonghua and L. Louhui, React. Kinet. Catal. Lett., 31 (1986)47.

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

731

Performance o f catalytic properties o f reagent catalyst in the processes such as methane oxidative coupling and h y d r o g e n production by methane conversion M.I. Levinbuk a, N.Y. Usachev b, M.L. Pavlov c, A.U. Loginovd, L.V. Surkova ~ E.M. Savin ~, V.K. Smirnov e, I.V. Ivkova~ aGubkin State Academy of Oil and Gas. 117917, Leninsky prospect 65, Moscow, Russia blnstitute of Organic Chemistry, 117913 Leninsky prospect 47, Moscow, Russia ~

Catalyst Plant, 453210, Ishimbai, Russia

dMoscow State University, 117234, Moscow, Russia eCatachem Company Ltd., 129832, Giliarovskogo 3 l, Moscow, Russia

1. INTRODUCTION More than 400 articles on methane oxidative conversion have been published for the past decade [1]. In order to realize this process in industry a new method based on a welldeveloped production technology has to be created. As it was described in the literature, metals with variable valency can be utilized as catalysts to convert methane into higher hydrocarbons [2]. Introduction of variable valency metals into cracking catalysts initiates methane partial oxidation reactions based on catalytic cracking technology. The catalyst used in this process circulates between the reactor and regenerator thus providing the partial oxidation of the feed in the first vessel while further oxidation of the reduced metal oxide occurs in the second one. We call it as a "reagent catalyst" since it functions as the carrier of the reagent (oxygen in this case) [3]. The feasibility of catalytic cracking technology utilization for the processes of production of gCz-hydrocarbons and pure hydrogen by methane oxidative catalytic conversion are considered in this article. US oil refinery plants will probably suffer from the shortage of hydrogen because the "reformulated gasoline" program demands to decrease the reforming capacity.

2. E X P E R I M E N T A L

The investigated catalysts were prepared using TCC (Thermofor Catalytic Cracking) catalysts as the base. Starting oxides MnO, Mn304, MnzO3 and MnO2 (at 10 wt% of Mn• per dry catalyst) were introduced by a special technique into a gel-forming solution. All the TCC catalysts contained 5% of rare-earth Y zeolite (with a molar ratio S I O 2 / A 1 2 0 3 = 6 obtained by direct synthesis). Nickel incorporating pentasil zeolites (3 - 8 wt % of NiO) were

732 prepared by impregnation of zeolite samples having different molar ratios SIO2/A1203 with a nickel-containing salt solution. Tested samples of the catalysts were exposed to various pretreatments: at air temperatures up to 1073 K without steam and at 1023 K with 100 % steam. Methane coupling reaction in the absence of oxidants in feed was carried out in an impulse microreactor; the charge was 0.1 g of the catalyst; one pulse of methane was 1.17 ml. Methane coupling reaction over TCC catalysts involving magnesium oxides with different starting oxidation levels was studied in a plug-flow reactor (with a volume of 1 ml) at atmospheric pressure, methane feed volume rate of 1000 h -1 (without oxygen) and reaction time of 1 min. The reaction was investigated within a temperature range of 973 - 1073 K. Methane catalytic conversion into carbon and hydrogen was examined over nickelcontaining pentasil zeolites in a vacuum-circulation laboratory unit [4] at a catalyst/feed mass ratio of 5.0 and within a temperature range of 743 - 843 K. The Mn +2 ion intensity variation in TCC catalysts in oxidation - reduction (reactionregeneration) cycle was observed by EPR method on a JeoI-JES-3BS-Q spectrometer (9 GHz) at 77 K and 295 K. Changes in the state of Ni +2 applied on pentasil in oxidation-reduction cycle were detected by oxygen absorption technique at a temperature of 77 K.

3. RESULTS AND DISCUSSION The utilization of metals with variable valency in catalysts in an oxidation-reduction cycle (separately in the reactor and regenerator when used in the catalytic cracking process) allows us to use the following reaction equation so as to depict the methane partial oxidation model: (p+oQCH4 + MeOm --> [3C2H4 + (o~- [~)C02 + pCH4 + MeOm_n

(1)

From the molar ratio of metal oxide to feed for the model reaction (1) a mathematical dependence can be obtained: y = Z * AS/(13/~t) * (L * Mf)/M m

(2)

where y = Gt / (p + or) - methane conversion, Z - catalyst/feed mass ratio, L - the metal oxide catalyst content, Mr, Mm - the molecular weights of feed and metal oxide, AS - variation in the valency of metal oxide. The verity of the equation (2) was confirmed by an experiment on methane conversion over TCC catalyst containing 10 wt% of Mn203 in an impulse microreactor at 1073 K and different catalyst/feed ratios (Table 1). From Table 1 it follows that the selectivity of the yield of C2-hydrocarbons and complete oxidation products depends antibately on the catalyst/feed ratio. In the range of the studied values of the catalyst/feed ratios the minimal yield of the complete oxidation products corresponds to only 2.5 wt% methane conversion.

733 Table 1 Conversion and selectivity of methane coupling vs. pulse number during the reduction of the sample Pulse Number 1

1. Methane conversion, wt% 5.0 120 2. Calculated values of catalyst/feed mass ratio 3. Productyields, wt% C2H6 20 C2H4 10 ZC 2 30 CO 25 CO2 45 ECO + CO2 70

2

3

4

5

4.0 60

3.5 40

3.0 30

2.5 24

25 14 39 22 39 61

31 16 47 16 37 53

38 19 57 10 33 43

45 22 67 5 28 33

Table 2 Conversion and selectivity of metal coupling vs. the type of initial Mn• a TCC catalyst Sample .No Starting Conversion, Product Yield (wt%) MnxOy (wt%) 1 2 3 4

MnO

Mn203 Mn304 Mn02

3.8 4.4 5.6 6.0

introduced into

C2H 6

C2H 4

CO + CO 2

40.2 40.5 41.6 39.8

35.6 38.0 40.4 41.6

24.2 21.5 18.0 18.6

The investigation of the influence of the oxidation level of starting Mn• oxides introduced into TCC catalysts on the yield of methane conversion products (in the absence of oxidants in the feed) was performed in a plug flow microreactor at 973 K (Table 2). From Table 2 it follows that the methane conversion and ethylene/ethane ratio increase, and the yield of the complete oxidation products decreases with the growth of the oxidation level of the starting oxides Mn• The examination of TCC catalyst samples by EPR-method before and after the performance of the methane conversion reaction (reaction/regeneration cycle) disclosed two major signal types (g-factor = 2.00), which do not change their parameters with registration temperature changes (Table 3). As is evident from Table 3, the distinctions between the EPR signals (AH) are connected with only the redox properties of the starting MnxOy oxides introduced into the TCC catalysts. Changes in the integral density of the EPR signals (Mn 2§ ion concentration) in oxidation-reduction cycle point at MnxOy valency variation in the TCC catalyst. Hence the Mn• oxide is the carrier of the reagent (oxygen in the present case) in the reaction-regeneration cycle, which provides controlled selective

734 oxidation of methane into higher hydrocarbons. However, industrial realization of such a process is not promising due to low methane conversion. It seems more interesting to utilize the discovered mechanisms of methane conversion into higher hydrocarbons for catalytic cracking of methane into carbon and hydrogen. Table 3 Parameters of EPR spectra before and after methane conversion reaction vs. the type of initial Mn• v introduced into TCC catalysts Sample .No Initial MnxOy Oxidized samples Reduced samples AH (G) EPR signal AH (G) EPR signal intensity intensity (Mn+2* 102~ MnO 215 6.4 360 46 Mn203 207 20 370 75 Mn304 203 33 437 51 MnO2 200 34 380 130

(Mn+2*lO2~

2.5 748K (Methane)

2 o

. ~ 15

O

773K (Methane)

--

798K (Methane) 823K (Methane)

Aq Fo ,

"

9

~

~

-- ---. 823K (Hydrogen)

~

~

-"

1

,-

~

o

0"

.

-

-

~

-- t - . 798K (Hydrogen)

"'" "'" ::t: ..... ~~ " - - -

-- A--. 773K (Hydrogen)

~ 9

0.5

_

,.o

9149 o'~

-- o - . 748K (Hydrogen)

0

10

20

30

Experiment time (min)

Figure 1 Kinetic characteristics of methane conversion and hydrogen yield at different temperatures of methane decomposition reaction on zeolite with 8 wt% Ni.

Kinetic characteristics of methane conversion and hydrogen yield on nickel-containing pentasil zeolite with a molar SIO2/A1203ratio of 210 are represented in Fig. 1. The distinctive feature of the methane conversion into C and O is the absence of any feed oxidation products in the gas phase. Carbon obviously forms a chemical compound with NiO on the catalyst surface (NiC•

735 The methane conversion rate increases by a factor of 4 or 5 with dcclilic in tile alumillulll content of pentasil zeolites (Fig.2). The highest methane collvelsioll (75 wl%) ,,vas obtained at a reaction te~nperature o1"843 K and a catalyst/feed ratio o1" 10. l-ligh-temperature treatment of nickel-incorporating pct~tasil zcolitcs by Ilyd~t~gcl~ deactivates them in methane conversion into carbon and hydrogen; in this case tl~e quantity o1 the absorbed oxygen in a sample decreases by 0.35 mmol/g at 77 K as compared with the oxidized catalyst.

0.8 0.7 .,lt-

9- 0.6 ---....

o

E

0.5

Pentasil (Ni 8 wt%) treated by Oxygen

0.4

Pentasil (Ni 8 wt%) treated by llydrogcn

I-.,

= o

.,..~ k,i,

to >.

- a - I n i t i a l pcntasil treated by Oxygen

= 0.3

o r

.= 0.2 to

0.1

1 0.75

i m

i m

1 . 7 5 2.75

I m

3.75

.

I

4.75

Aluminum content of pentasil framework (wt%) Figure 2. Methane conversion rate vs. the aluminum content of zeolite and kind o1' its preliminary treatment.

4. C O N C L U S I O N

The introduction of Ni-containing pentasii zeolites into TCC and FCC (Fluid Catalytic Cracking) catalyst matrices makcs it possible to use (without considerable reconstruction)the technology of the production of these catalysts and the reactor-regenerator vessels of crackling units for tile process of hydrogen generation from natural gas. For example, a regenerator with a coke burning capacity of 1400 kg/h possesses a calculated hydrogen yield o1"460 kg/l~ (the vacuum gas-oil capacity of the initial cracking unit is 50000 kg/h). A TCC unit with a vacuum

736 gas oil capacity of 300000 t/year can be partially reconstructed for the process of hydrogen obtaining from natural gas with 4000 t/year output of the target product. For comparison, 4500 to 5000 t of hydrogen a year can be produced at a reforming unit of 300000 t/year capacity.

REFERENCES

1. O.V. Krylov, Catalysis Today, 18 (1993) 209. 2. G.E Keller and M. Bhasin, J.Catal., 73 (1982) 9. 3. M.I. Levinbuk and V.M. Melnikov, 211 th National Meeting, American Chemical Society, New Orleans, (1996) 410. 4. U.V. Shumovski, React. Kinet. Calal. Lett., 21 (1983) 3.

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

T h e Effect of t h e P b O L o a d i n g in Methane over PbO/SiO2 Catalysts.

the

737

Oxidative

Coupling

of

H. J. Lugo, N. Teran, L. Villasmil, G. Castillo and D. M. Finol Centro de Superficies y Cat~lisis, Facultad de Ingenieria, Universidad del Zulia, Apartado 15251, Maracaibo 4003A, Venezuela

The effects of the PbO loading and the C H 4 / O 2 ratio on the oxidative coupling of methane (OCM) over PbO/SiO2 catalysts were studied. Special emphasis was made in the interpretation of the product distribution in the gas phase and its relationship with the nature of the catalyst surface. At CH4/O2 > 2 ratio, the catalytic behavior of the 2 and 6 % PbO/SiO2 catalysts was very similar. The activity for the deep oxidation of CH4 to CO2 and the activity for the formation of C2 hydrocarbons were almost the same in both catalysts. At CH4/O2 = 1 ratio, the behavior of the 2 and 6 % PbO/SiO2 catalysts differed. There were a lower activity for the oxidation of CH4 to CO2 and a higher activity for the formation of C2 hydrocarbons in the 2 % PbO/SiO2 catalyst than in the 6 % PbO/SiO2 catalyst. At any C H 4 / O 2 ratio, the 10 % PbO/SiO2 catalyst was very different from the 2 and 6 % PbO/SiO2 catalysts. It had a very low activity for the oxidation of CH4 to CO2 and was selective for the generation of C2 hydrocarbons. By doubling the a m o u n t of catalyst, the amount of reacted methane doubled, while the selectivity remained almost constant. Temperature programmed reduction experiments showed almost the same behavior in all the catalysts. However, the 10 % PbO/SiO2 catalyst showed clearly reducible species at about 670 K, which were practically absent in the 2 and 6 % PbO/SiO2 catalysts. Probably, these species were responsible for the low activity for the oxidation of CH4 to CO2 in the 10 % PbO/SiO2 catalyst.

1. INTRODUCTION

The last decade has witnessed great efforts by scientists from many countries to convert methane to value added products. The pioneering work of Keller and Bhasin [1] stimulated great interest in the oxidative coupling of methane. Redox-types oxides constitute a category of catalysts extensively studied for this purpose. Within this category, lead oxide and supported lead oxide

738 have been reported to be suitable for the OCM reaction [2-7]. These studies have demonstrated that supported lead oxide catalysts are very active and that their catalytic behavior depends strongly on the lead oxide loading; however, a clear u n d e r s t a n d i n g has not been reached yet. With respect to the mechanism, Bytin and Baerns [8] distinguished two adsorption steps for m e t h a n e on lead oxide: (a) dissociative adsorption with the s u b s e q u e n t recombination of the adsorbed fragments to yield ethane, and (b} adsorption as methylcarbonium species on acid sites, which then undergo attack by surface O"- ions, yielding methoxide species, which then undergo deep oxidation. In this study, we tried to explain the effect of the PbO loading and the CH4/O2 ratio on the oxidative coupling of m e t h a n e (OCM) over PbO/SiO2 catalysts. Special emphasis was made in the interpretation of the product distribution in the gas phase and its relationship with the n a t u r e of the catalyst surface. The methodology of the investigation involves incorporation of variable a m o u n t s of PbO to the SiO2 to examine the product distribution for both low and high lead oxide loadings. Temperature-programmed reduction (TPR) reveals the surface changes of PbO.

2. EXPERIMENTAL 2.1. Catalyst preparation Pb(NO3)2 (99.101%) and SiO2 (Davisil, grade 646, SB~=245 m2/g) were p u r c h a s e d from Riedel and Fisher, respectively. Before preparation of the catalysts, the silica was calcined at 1200 K for 4 h. PbO/SiO2 catalysts were prepared by impregnating a m o r p h o u s SiO2 (60-80 mesh) with aqueous solutions of Pb(NO3)2 of appropriate concentration to yield Pb0 loadings of 2, 6 and 10 wt. %. Excess water was removed in a rotary evaporator. The catalysts were then dried at 393 K in an oven for 12 h and subsequently calcined at 1073 K for 4 h. 2.2. Reaction s y s t e m Methane conversion was performed in a conventional ruxed-bed continuous flow reactor operated u n d e r atmospheric pressure. The reactor consisted of a quartz, U type tube of 9 m m internal diameter. The a m o u n t of catalyst used for a test r u n was about 0.25 g, which was held in place by quartz wool plugs. The reactor was placed in an electric furnace with approximately 20 cm of the quartz-filled tube serving as a preheater. Before the reaction the catalysts were pretreated in an oxygen flow at 1048 K for 1 h. The reactant mixture of CH4 and 5% 02 in He was adjusted to meet several CH4/O2 ratios and a total flow rate of about 1.2 dm3/h, keeping constant the oxygen

739 partial p r e s s u r e (4.5 kPa). The coupling reaction w a s carried o u t at 1048 K for at least three h o u r s to establish steady state conditions. Catalyst deactivation w a s not observed over this time period. The r e a c t a n t s a n d p r o d u c t s were analyzed with a n o n - s t r e a m gas c h r o m a t o g r a p h equipped with a TCD. Two c o l u m n s , a C h r o m o s o r b 102 (3 m), a n d a Molecular Sieve 5A (2.5 m) were employed in the analyses. Care w a s t a k e n to avoid c o n d e n s a t i o n of the p r o d u c t s at the outlet of the reactor. The conversion a n d selectivities were calculated from the a m o u n t s of reaction p r o d u c t s formed (carbon a t o m basis) as d e t e r m i n e d by the GC analysis. The error in c a r b o n balance w a s found to be below 5% in all cases. Total conversion of r e a c t a n t (XT) a n d selectivity to p r o d u c t i (Si) are defined as;

XT --'--

Si =

moles of reactant transformed X 100 moles of reactant in the feed

C atoms of i 9moles formed of i x 100 moles of CH4 transformed

2.3. Experimental

techniques

The surface a r e a s of the catalysts were m e a s u r e d by the conventional BET nitrogen a d s o r p t i o n method. Values of 26, 7 a n d 4 m 2 / g were o b t a i n e d for the 2, 6 a n d 10 % PbO/SiO2 catalysts, respectively. T e m p e r a t u r e p r o g r a m m e d reduction e x p e r i m e n t s were performed u s i n g an a p p a r a t u s described by Robertson et al. [9]. The reduction w a s carried o u t with a purified h y d r o g e n - a r g o n mixture (10 vol.% hydrogen) at a h e a t i n g rate = 10 K min -I u p to 1048 K. The TPR reactor w a s charged with 0.25 g of c a l c i n e d (fresh) catalyst. Before the reduction, the catalyst w a s oxidized in a 02 flow to 1048 K for 1 h, a n d t h e n cooled down to 300 K in Ar. This reduction-oxidation cycle w a s repeated several times.

3. RESULTS AND DISCUSSION

Table 1 s h o w s the catalytic properties for the oxidative coupling of m e t h a n e over 2, 6 a n d 10 % PbO/SiO2 catalysts at different CH4/O2 ratios at 1048 K, u s i n g 250 mg of catalyst. It w a s also r u n u s i n g 500 mg of 10 % PbO/SiO2 catalyst at a CH4/O2 = 2 ratio. As it is well k n o w n on OCM catalysts, the CH4 conversion decreases a n d the C2 selectivity increases w h e n the CH4/O2 ratio increases. For the 2 a n d 6 % PbO/SiO2 catalysts,

740 the 02 conversion is less t h a n 100 % only for a CH4/O2 = 1 ratio, while for the 10% PbO/SiO2 catalyst, it is always well below 100 %. Besides, by doubling the a m o u n t of the 10% PbO/SiO2 catalyst, the conversion of m e t h a n e almost doubles, while the selectivity remains nearly constant. However, Table 1 does not provide sufficient information a b o u t the behavior of these catalysts. Therefore it is necessary to look further into the feed and p ro du c t gas composition.

Table 1 C a t a l ~ i c Properties of PbO/SiO2 catalysts PbO load P(CH4) Conv CH4 Conv 02 /P{O2)

(%1

(O/o)

(o/o)

C2H6

Selectivity

(%)

C2H4

CO

CO=

2 (250 mg)

1 2 5

39.3 29.0 15.6

65.4 100 100

7.9 9.9 13.3

11.1 15.8 25.0

5.9 4.8 4.2

75.1 69.5 57.5

6 (250 mg)

1 2 5

48.8 30.4 16.4

86.4 100 100

5.8 13.2 14.3

7.8 17.5 25.4

2.3 4.2 5.3

84.1 65.1 55.0

10 (250 mg)

1 2 5

8.2 6.4 3.6

6.5 7.4 12.0

46.7 53.9 60.7

19.1 22.5 19.7

8.7 4.7 3.6

25.5 19.0 16.0

10 (500 mg)

2

11.3

17.3

44.1

25.6

5.2

25.1

Reaction conditions: F = 1.2 din3/h, P(O2)ffi 4.5 kPa, Ptot~ffi 101 kPa, T = 1048 K, Inert = helium.

Figure 1 shows the c o n s u m p t i o n of CH4 and 02, and the a m o u n t of the different p r o d u c t s formed for several CH4/O2 ratios over a 6 % PbO/SiO2 c a t a l y s t . For a CH4/O2 = 2 ratio, the c o n s u m p t i o n of CH4 reaches a value n e a r or equal to the m a x i m u m a m o u n t allowed by the a m o u n t of oxygen available for the reaction. At this point all of the available oxygen is c o n s u m e d (100% O2 conversionl. For a CH4/O2 = 1 ratio, the c o n s u m p t i o n of CH4 does not reach the m a x i m u m allowed value (86.4 % 02 conversion). This is because with a decreased a m o u n t of m et hane, and a c o n s t a n t

741 a m o u n t of oxygen, competition for the active sites o c c u r s where the m e t h a n e is disfavored, so t h a t the m e t h a n e c o n s u m p t i o n decreases. This competition between m e t h a n e a n d oxygen h a s been noted in the literature [10]. For a CH4/O2 --- 5 ratio, the excess m e t h a n e displaces the oxygen on the active sites. It i n c r e a s e s the m e t h a n e c o n s u m p t i o n until there is no more oxygen available (100% 02 conversion) to regenerate the active sites, resulting in c o n s t a n t a m o u n t of c o n s u m e d m e t h a n e at the steady state. In this situation, the n u m b e r of active sites utilized can be less or equal to the total n u m b e r of active sites. Figure 1 also shows t h a t w h e n m e t h a n e is in excess with respect to oxygen, t h a t is, at higher CH4/O2 ratios, the oxidative coupling of m e t h a n e is favored over total c o m b u s t i o n . It is evidenced by the increasing p r o d u c t i o n of C2 h y d r o c a r b o n s a n d the decreasing a m o u n t of CO2. At these conditions, the m e t h a n e m u s t occupy a higher proportion of the surface, limiting the a m o u n t of oxygen t h a t h a s access to it. This situation favors a methyl radical p r o d u c t i o n a n d their coupling in the gas p h a s e a n d disfavors the oxidation of CH4 to CO2 on the surface. When m e t h a n e is in deficit, t h a t is, CH4/O2 < 1 ratio, the excess surface oxygen restricts the formation of m e t h y l radicals. Rather, m e t h a n e is deeply oxidized, favoring the p r o d u c t i o n of CO2.

o

3

D

2,5

uJ O am

--o--02

o

-r -o-- C02

n,,o O(9 1,5

a

o----o.-._...._

~

~-. C2H4 (*)

I--

n,,l~ w > =-0,5 z

o0

C2H6 (*) --

0 0

1

2

3

4

5

CO

6

CH4/O2 RATIO

Figure i. Effect of the CH4/O2 ratio on the gas p h a s e composition in a 6 % PbO/SiO2 catalyst. (*) Equivalent ~mol of CH4 = C atoms of i product. #Jmol formed of i product

742 Figure 2 shows the consumption of CH+ and 02, and the a m o u n t of several p r o d u c t s versus the CH4/O2 ratio for a 2 % PbO/SiO2 catalyst. The behavior of this catalyst is the same as the 6 % PbO/SiO2 catalyst for CH4/O2 > 2 ratio. The activity for the deep oxidation of CH+ to CO2 and the activity for the formation of C2 hydrocarbons are almost the same in both catalysts.

uJ

3

o

2,5

o

2

.--0--02 --r 9

CH4

--o--- CO2 x

Z

o

(3

C2H4 (*)

--b-- C2H6 (*)

0,s

--

0 0

1

2

3

4

5

CO

6

CH4/O2 RATIO

Figure 2. Effect of the CH4/O2 ratio on the gas p h a s e composition in a 2 % PbO/SiO2 catalyst. (*) Equivalent pmol of CH, = C atoms of i product, pmol formed of i product For a CH+/O2 = 1 ratio, the behavior in both catalysts differs slightly. There are a lower activity for the oxidation of CH4 to CO2 and a higher activity for the formation of C2 hydrocarbons in the 2 % PbO/SiO2 catalyst t h a n in the 6 % PbO/SiO2 catalyst. For this ratio, the adsorption of oxygen competes with that of methane, and results in a lower m e t h a n e c o n s u m p t i o n t h a n in the case of a CH4/O2 = 2 ratio. The lower consumption of m e t h a n e in the 2 % t h a n in the 6 % PbO/SiO2 catalyst at CH4/O2 = 1 ratio, corresponds to a decrease in the conversion of m e t h a n e to CO2 in the 2 % PbO/SiO2 catalyst, and an increase in the conversion of m e t h a n e to C~ hydrocarbons. The decrease in the production of CO2 a n d the increase in the production of C2 h y d r o c a r b o n s would be related to a lower n u m b e r of active sites for the deep oxidation of m e t h a n e a n d to a higher n u m b e r of active sites for the formation of methyl radicals in the 2 % PbO/Si02 catalyst. Figure 3 shows the consumption of CH4 and 02, and the a m o u n t of p r o d u c t s versus the CH4/02 ratio for the 10 % PbO/SiO2 catalyst. It is

743 observed that the oxygen consumption is very low, so there is excess of oxygen in the gas phase at any time, and oxygen suppy is not limiting the rate. The consumption of methane increases with the a m o u n t of methane in the gas phase. This consumption is strongly related to the production of C2 hydrocarbons and to the active sites able to generate methyl radicals, while the low production of CO2 suggests a shortage of active sites for the oxidation of methane to CO2.

0,5 o LU o :::) am O r,r n,'o 0(.9

uJ~

0,4

---o-- CH4 ---0.-02

0,3

C2H6 (*) 7. C2H4 (*)

0,2

IuJ :~. > 0,1 z O o

-o--CO2 #, CO

0 0

1

2

3

4

5

6

CH4/O2 RATIO

Figure 3. Effect of the CH4/O2 ratio on the gas phase composition in a 10% PbO/SiO2 catalyst. (*) Equivalent #Jmol of CH, = C atoms of i product, pmol formed of i product Evidently the I0 % PbO/SiO2 catalyst has a different nature than the two lower loading s a m p l e s , in the sense that it has something that inhibits the deep oxidation of methane, increasing the generation of methyl radicals and so the production of C2 hydrocarbons. Figure 4(a} compares the production of ethane, ethylene and carbon monoxide from the various catalysts. Figure 4(b) compares the generation of carbon dioxide from these catalysts, and shows that only the 2 % and 6 % PbO/SiO2 catalysts have high production of CO2. Clearly the 10 % PbO/SiO2 catalyst has undergone a modification in its structure, which allows lower formation of CO2.

744

,2

a uJ

o

1,2

......

1

a

o ~ . 0,8 n~

--o--- 6 % PbO

0.6

---a-- 10% PbO

~.~ +

o.j

0,4

0 o "0,2

+"0,2

0

0,8

- . - o - 2% PbO

0,4

"1r

DU)

|

0

1 2 3 4 5 6 CH4102 RATIO

0

,

|

1 2 3 4 5 CH4/O2 RATIO

(b)

(a)

Figure 4. Effect of the PbO loading in the OCM reaction over PbO/SiO2 catalysts. (*) Equivalent lJmol of CH4 = C atoms of i product, lJmol formed of i product

Ill 1 I'--

2 O E ia,.

PbO 10%

=E

6%

z

o

2%

O N

:z: 0

400

I

|

I

I

600

800

1000

1200

TEMPERATURE (K)

Figure 5. TPR profiles for PbO/SiO2 catalysts.

|

6

745 Figure 5 compares the results of the TPR experiments and shows that all the catalysts have similar behavior. However, the 10 % PbO/SiO2 catalyst shows clearly a reduction feature at about 670 K, which is absent in the 2 and 6 % PbO/SiO2 catalysts. The presence of some reducible species may be responsible for the very low activity for the deep oxidation of CH4 to CO2 in the 10 % PbO/SiO2 catalyst.

4. CONCLUSIONS

At CH4/O2 -> 2 ratio, the behavior of the 2 and 6 % PbO/SiO2 catalysts is very similar. The activity for the deep oxidation of CH4 to CO2 and the activity for the formation of methyl radicals ( C2H6, C2H4 and CO) are almost the same in both catalysts. At CH4/O2 = 1 ratio, the behavior of the 2 and 6 % PbO/SiO2 catalysts differs slightly. There is a lower activity for the oxidation of CH4 to CO2 and a higher activity for the formation of C2 hydrocarbons in the 2 % PbO/SiO2 catalyst than in the 6 % PbO/SiO2 catalyst. At any CH4/O2 ratio, the 10 % PbO/SiO2 catalyst has a very low activity for the deep oxidation of CH4 to CO2 and its catalytic activity is for the generation of methyl radicals. By doubling the a m o u n t of catalyst, the a m o u n t of reacted methane doubled, while the selectivity remained almost constant. Temperature programmed reduction experiments show almost the same behavior in all the catalysts. However, the 10 % PbO/SiO2 catalyst shows a reducible species at about 670 K, which is absent in the 2 and 6 % PbO/SiO2 catalysts. Probably, these species are responsible of the very low activity for the oxidation of CH4 to CO2 in the 10 % PbO/SiO2 catalyst.

REFERENCES

I. G.E. Keller and M.M. Bhasin, J. Catal., 73 (1982) 9. 2. W.Hinsen, W. Bytyn and M. Baerns, in Proceedings 8 th International Congress on Catalysis, Berlin, 1984, Vol. 3, Verlag Chemie, Weinheim, 1984, p.581. 3. K. Asami, S. Hashimoto, T. Shikada, K. Fujimoto and H. Tominaga, Ind. Eng. Chem. Res., 26 (1987) 7. 4. G. Wendt, C.D. Meinecke and W. Schmitz, Appl. Catal., 45 (1988) 209. 5. A. Machocki, A. Denis, T. Boroniecki and J. Barcicki, Appl. Catal., 72 ( 1991) 283. 6. S.E. Park and J. S. Chang, Appl. Catal. A 85 (1992] 117,

746 7. R. Mariscal, J. Soria, M. A. Pefia and J. L. G. Fierro, Appl. Catal., A: General 111 (1994) 79. 8. W. Bytyn and M. Baems, Appl. Catal., 28 (1986) 199. 9. S. D. Robertson, B.D. McNicol, J.H. De Bass, S. C. Kloel and J. W. Jenkins, J. Catal., 37 (1975) 424. 10. G.A. Martin and C. Mirodatos, Fuel Processing Technology, 42 (1995) 179.

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

747

Catalytic Combustion of Ethane over High Surface Area Lnl.xKxMnO3 (Ln La, Nd) Perovskites: The effect of Potassium Substitution. =

Y. Ng Lee a*, F. Sapifiab, E. Martinezb, J. V. Folgado b and V. Cort6s Corberfina+. a Instituto de Cat~lisis y Petroleoquimica, CSIC, Campus U.A.M. Cantoblanco, 28049 Madrid,

Spain. b Instituto de Ciencia de los Materiales de la Universitat de Valencia, Dr. Moliner 50, 46100 Burjassot (Valencia), Spain.

Monophasic Lnl.xKxMnO3 (Ln = La, Nd) perovskites with high surface areas (8-27 m2/g) were synthesised at mild conditions by the ~eeze-drying method, and were found to be active for the catalytic combustion of ethane at low temperatures (5 73 to 648 K). As a general trend, the substitution of the rare earth cation by potassium decreased the intrinsic activity, reduced the reaction order in oxygen and, for the more substituted samples (x>0.10), it increased the selectivity to ethene. It was found that the rare earth cation also influenced the catalytic activity of the substituted perovskites. These effects were analysed in terms of structural modifications induced by the introduction of potassium in the perovskites.

1. INTRODUCTION The advantages of the catalytic (flameless) combustion of light hydrocarbons are numerous [1-3]. In the presence of active catalysts complete combustion may be achieved at temperatures several hundred degrees lower than in the flame (thermal noncatalytic) combustion [3]. This low temperature is the basic advantage of catalytic combustion. For example, noxious nitrogen oxides formation is almost eliminated [3] and energy savings are achieved. Supported platinum or palladium catalysts have so far been used almost exclusively because of their excellent activity [2-3]. However, their high price, tendency to be poisoned, and volatility in some applications, have estimulated the search for suitable substitutes. In 1971 Ll"bby [4] proposed LaCoO3 as a potential auto exhaust catalyst, starting a wide interest in the catalytic properties of perovskites. However, their specific application as catalysts in flameless combustion has been investigated only since 1980 [6-15], mainly in Japan. Nakamura et aL [6] began these studies reporting the catalytic oxidation of propane, methane and CO on several LaBO3 (B = Co, Mn, Ni, Fe), and lanthanum substituted (Ca, Sr, Ba, Ce) perovskites. The order in catalytic activity was found to be Co > Mn > Ni > Fe > Cr, On leave from: Universidad Central de Venezuela, Facultad de Ciencias, A.P. 47586, Los Chaguaramos, 1041 A, Venezuela. + Correspondingauthor. E-mail: [email protected] Carar

748 and Sr-"doping" was the most effective in enhancing the activity of these compounds. Seiyama et al. [7] reported the activity for the catalytic oxidation of propylene of several perovskite oxides (ABO3) and their corresponding B oxides (BO~), showing that their activity was determined mainly by the nature of the metal in B positions. The most active catalysts were Co and Mn perovskites. Arai et aL [8] investigated the catalytic activity of various transition-metal perovskites for methane combustion. In every case, carbon monoxide or partial oxidation products were scarcely detected during oxidation. The activities of LaCoO3, LaMnO3 and LaFeO3 were quite close to that ofa Pt/ahxmina catalyst. An interesting feature of rare-earth perovskites is the possl~oility to vary the unit cell dimensions by controlling the nature of the A ion, and thereby the covalence of the B-O bond. A thorough study of the role of the A- and B-site ions on the catalytic properties of ABO3 perovskites for propane and methanol oxidation has been reported by Nitadori et aL [9]. They concluded that the influence of the rare-earth ions in the A-site on the oxidation properties of these compounds was secondary. Zhang et aL [10] studied the oxygen sorption and catalytic properties for methane and n-butane combustion, of Lal.xSrxCo~.yFeyO3. Whil~ the catalytic activity for n-butane oxidation was affected both by the transition-metal substitution and rareearth substitution, the catalytic activity for methane oxidation was only influenced by the rareearth ion substitution. Marti and Balker [ 16] have found that the activity of ACoO3 oxides for methane oxidation was, with exception of the PrCoO3 perovskite, only slightly influenced by the A-site cations. Perovskite-like compounds derived from laflVlnO3 by partial substitution of La3+ by divalent ions (Sr, Ca, Ba, Ce) are well known. However, the effect of partial substitution of the A site by alkaline ions on the catalytic activity in combustion reactions has been very little studied. Voorhoeve et aL [ 17] reported for the first time that the substitution of La 3+ by alkali ions in LaMnO3 increases its catalytic activity for NO reduction. Apart from their intrinsic activity, the performance of perovskite catalysts depends on their particle morphology, and, in particular on their specific surface area (SSA) [5, 12, 18]. To obtain homogeneous perovskite ~mples, the conventional ceramic route requires repetitive grinding and heating procedures as well as high temperatures, which results in large grain sizes and low SSA (< 2 m2/g), unsuitable for their application as catalysts. The use of alternative synthetic pathways produces high purity, homogeneous powders, and requires low temperatures to obtain the phases, thus leading to products with smaller particles and high SSA [19]. Besides these advantages, these alternative methods play an essential role in the preparation of potassium containing materials, due to the volat'flity of potassium oxide at high temperatures. In the alternative method involving freeze-drying all the cations are mixed at the atomic scale in the solution and this facilitates the incorporation of potassium to the perovskite lattice since it occurs at the initial stage of the preparative procedure [20]. Freeze-drying has proved to be a powerful and versatile technique for mild preparation of complex oxides, such as high temperature superconducting material and other perovskites (lanthanide, nickelates and cobaltates) [21, 22]. We report in this work the low temperature synthesis and catalytic properties of two series ofLal.xKxMnO3 (Ln = La, Nd; 0 _<x _ 0.15 when calcined at much higher temperatures (1273 K) than those used for catalyst preparation. 3.2. Surface areas

BET specific surfaces areas of the catalysts are given in Table 1. The lower values for the Ndl.xKxMnO3 series are due to the higher calcination temperature used (973 K), needed to obtain monophasic samples. In general, these values are relatively high. Thus, SSA values obtained for the series Lal.xK~MnO3are similar for those reported for La~.xSrxMnO3 prepared by the sol-gel method (17.5-23.5 m2/g) [24] and LaMno.sCt~503 (27 m2/g) prepared by t~eeze-drying [25], but much higher than that ofLao 95K0.osM O3 prepared by the ceramic

NKx

.......

z

20

3'o

4'0

s'o

2e (degrees)

6'o

20

3'o

4'0

20 (degrees)

Figure 1. XRD patterns of LKx (left) and NKx (right) catalysts.

s'o

6'o

751 Table 1 BET Specific Surface Area# ofLn~.xK~MnO3 (Ln = La, Nd). Catalysts x Specific Surface Area (m2/g) Lal.xKxMnO3 Ndl_xKxMnO3 0 19.7 19.0 0.05 24.3 14.8 0.10 26.5 10.0 0.15 24.4 9.6 0.20 21.7 8.0 0.25 20.8 10.6 From N2 adsorption at 77 K. Sm= 0.162 nm2. method (0.62 m2/g) [17]. The SSA of NdMnO3 (NK0.0) is also higher than that of NdCoO3 (12.3 m2/g) obtained by fleeze-drying of amorphous nitrate precursors [21], which was found to be the highest reported in the literature for the perovskite.

3.3. Scanning Electron Microscopy (Fidd Emission) The morphology evolution has been followed by Scanning Electron Microscopy Field Emission (SEM-FE). Micrographs shown in Fig. 2 correspond to two representative samples: sample Lao.ss~.tsMnO3 prepared at 873 K (catalyst LK0.15) and sample Ndo.ssKo.~sMnO3 prepared at 973 K (catalyst NK0.15). No segregation of secondary phases can be observed in the micrographs. No variation of morphology and particle size due to composition was observed within the homologous series for this temperature.

Fig.

2.

SEM-FE

micrographs

of

samples

LK0.15

(A)

and

NK0.15

(B).

752 SEM-FE images show the existence of a noticeable difference in the dism'bution of particle size between the LKx and NKx samples. The lanthanum sample (LK0.15, monophasic by XRD), is composed of homogeneous nanometer particles, with a narrow particle size distEoution (30-50 nm), while the Nd sample shows a broader distn'bution (20-100 nm). Nevertheless, the average particle size of both series was similar, about 40 nm In both La and Nd materials the particles were rounded, and non-faceted. Only when ~mples were treated at high temperature (1173 K) did they exls an increase in particle size, accompanied by the development of crystalline facets. The particle size range of La samples remained narrower than that of Nd samples (50-200 nm~ 60-300 nm~ respectively). 3.4. Catalytic Activity In the conditions used, all the catalysts were active for the combustion of ethane, producing conversions ~om 1-5 % at 573 K up to 40 % at 648 IL The only detected products were CO2, H20 and small amounts of ethene, that is consistent with a triangular reaction scheme involving parallel formation of CO2 and ethene and the consecutive combustion of ethene. In separate measurements of CO oxidation on these catalysts, made in a FTIR cell [26], CO was transformed completely into CO2 at temperatures as low as 573 K. This would explain the absence of CO among the reaction products of ethane oxidation, carried out here at higher temperatures. Fig. 3 shows the variation of areal rates of ethane oxidation over the LKx and NKx catalysts with temperature. In all cases, ethane conversion increased with increasing

x I

t

E450 /

O=! 0".I

INKx[

x 0 0.05 O 0.00

[] o.ool

0.10

(3//

0.15

;" ".~ o 2o

0 575

600 625 650 675 Temperature (K)

0.20

5"/5 600 6~ 6~ Temperature (K)

675

Figure 3. Intrinsic rates of ethane oxidation on Lai.xKxMnO3 (leit) and Ndl.xK~MnO3 (right) catalysts as a function of temperature and degree of substitution (x). Feed composition: 4% C2I~, 12% 02, He balance; W/F: 38 g. h/mol C2I-I6.

753 2O

INI 15

/

-I-

ID

.r. ,t

o

-o 0

0

10

20

30

C2H6 Conversion (%)

40

0

/

10

I

20

i

I

30

C2H 6 Conversion

40

(%)

Figure 4. Variation of ethene selectivity with ethane conversion on Lal.xKxMnO3 (left) and Ndl.xKxMnO3 (right) catalysts. Experimental conditions and symbols as in Fig. 3. temperature, but in both series the intrinsic activity decreased with increase in potassium substitution (x), with the only exception of catalyst NK0.05, where the areal rate was higher than that of the unsubstituted NdMnO3. Within experimental error, the calculated values of the apparent activation energy for CO2 formation were the same, 18.4_+0.2 kcal/mol, for all the LKx catalysts, while they were lower, 15.8_+0.9 kcal/mol, for the NKx catalysts. Another interesting effect of potassium substitution can be seen in Fig. 4 which shows the evolution of ethene selectivity as a function of ethane conversion in both series. In general, as potassium substitution increased, the ethene selectivity for a given conversion also increased. For the unsubstituted samples (x = 0.0) and for x = 0.05, CzI-I4selectivity remained constant c a . 2%, in the whole temperature range. However, in the case of catalysts with higher substitution degrees (x > 0.10), this selectivity tends to increase with increasing x and also with increasing conversion, with the later effect much more marked in the Nd-containing samples. This indicates that the presence of potassium favours either the formation and/or the desorption of ethene formed via oxidehydrogenation (OXD) of ethane. The effect of the cation in the A-position can be more clearly seen in Fig. 5, in which the areal rates of ethane transformation as a function of x are compared. The activity of the unsubstituted samples was very similar, being slightly higher for the La catalyst. By contrast, the ethane combustion rate was higher on the substituted Nd~.xKxMnO3 perovskites as compared with the homologous La~.~K~MnO3 samples with the same substitution level. This reveals that the nature of the rare-earth cation also has a si~ificant effect in the catalytic performance, especially in the presence of potassium The effect of oxygen partial pressure (Po2) on the areal rates of ethane transformation at 648 K is shown in Figure 6. In the samples without potassium or with low substitution, the activity increased as Po2 increased. However, in the samples with x = 0.20=0.25, only little changes in the intrinsic rate of ethane conversion could be observed when Po2 increased. Reaction orders

754

600

- - [ 3 - - LKx

N"

- - O - - NKx 450

r| O

E

300

ID

150 < 0

I

I

0.00

I

I

I

I

0 . 0 5 0 . 1 0 0 . 1 5 0 . 2 0 0.25 Substitution x

Figure 5. Effect ofpotassium on the areal rate. Temperature: 648 K, Po2" 1.1 kPa.

E '7" e"

jo O

45O

=L v

m

3OO V

150

+__._+_____+_ I

3

6

~

V

+

;

O2 Pressure 0oal

12

3

~

~

12

0 2 Pressure (kPa)

Figure 6. Ethane conversion areal rates at 648 K as a fimction of Po2. W/F: 38 g . h/mol C2I-I6.. Symbols as in Fig. 3. in oxygen for catalysts with x = 0-0.20 are summarised in Table 2. It can be seen that, within experimental error, the reaction order in oxygen decreased monotonically with increasing K substitution, while being practically equal for the homologous LKx and NKx samples with the same x. This may be attn~buted either to a change in the mob'dity of the oxygen species, or to a change in oxygen species that participate in the reaction, specifically, from molecularly

755 Table 2. Dependence of ethane oxidation rate on Llll.xKxlVInO3 ( L n = La, Nd). catalysts on the partial pressure of oxygen x Reaction order n Lal.xKxMnO3 Ndl.xKxMnO3 0 0.44 + 0.11 0.42 + 0.02 0.05 0.38 + 0.06 0.31 + 0.02 0.10 0.35 + 0.07 0.25 + 0.05 0.15 0.23 + 0.03 0.17 + 0.03 0.20 0.15 + 0.04 0.29 + 0.16 a

Oxidation rate, r = k (Pc2n6) ~ (Po2)"

adsorbed to lattice oxygen [8, 12]. Neutron di~action studies of LaMnO3.15 [27] are consistent with a structure consisting of a compact packing of oxygen ions with cationic vacancies at both La and Mn positions. The substitution of one trivalent rare earth ion by K+ will reduce these vacancies, that become filled for x = 0.15, thus reducing the mobility of oxygen species.. On the other hand, the promoting effect of potassium doping in selective oxidation reactions is well known. So, addition to metal oxide catalysts, such as V2Os/TiO2 and MoO3/TiO2 increases the selectivity to oxidehydrogenation (OXD) of propane [28]. It has been shown that potassium addition brings about a decrease in acidity, lowers surface potential (work function) and hinders the formation of electrophilic O species, which are the responsible for total combustion. Thus, the reduction of total conversion as well as the increase in the selectivity to OXD products when the potassium content increases could be interpreted as due to the modifications of these properties induced by the presence ofpotassiunt

CONCLUSIONS The freeze-drying of acetates solutions is a suitable method for preparing high SSA monophasic Lnl.xKxMnO3 (In = La, Nd) perovskites. The presence of potassium strongly affects the catalytic properties, probably by modifying the nature and concentration of surface species on the catalyst surface. The increase of K content would decrease the ability of the perovskite to activate 02 and to store active oxygen species, and would reduce the electrophilicity of these species, thus increasing the C2I-I6selectivity.

Acknowledgements The authors thank Dr. Rochel M. Lago for useful discussions. Y. Ng Lee was supported by a Fellowship from CDCH de la UCV, Venezuela. Financial support from the Spanish CICYT under project MAT96-0688-C02-02 is acknowledged.

756 REFERENCES

[ 1] G. E. Voecks, 3rd Workshop on Catalytic Combustion (1977). [2] R. Prasad, L. A. Kennedy and E. Ruckenstein, Catal. Rev.-Sci. Eng., 26 (1984) 1. [3] D. L. Trimm~ Appl.Catak, 7 (1984) 249. [4] W. F. L~by, Science, 171 (1971) 499. [5] N. Yamazoe and Y. Teraoka, Catal. Today, 8 (1990) 175. [6] T. Nakamura, M. Misono, T. Uchijima and Y. Yoneda, Nippon Kagaku Kaishi (1980) 679. [7] T. Seiyama, N. Yamazoe and K. Eguchi, Ind. Eng. Chem_, Prod. Res. Dev., 24 (1985) 19. [8] I-1.Arai, T. Yamada, IC Eguchi and T. Seiyama, AppL Catal., 26 (1986) 265. [9] T. Nitadori, T. Ichiki and M. Misono, Bull Chem_ Soc. Japan, 61 (1988) 621. [10] H. M. Zhang, Y. Shimizu, Y. Teraoka, N. M ~ a and N. Yamazoe, J. Catal., 121 (1990) 432. [11] Z. Kaiji, L. Jian, and B. Yingli, Catal. Lett., 1 (1988) 299. [ 12] L. G. Tejuca, J. L. G. Fierro and J. M. D. Tasc6n, in: Advances in Catalysis, Vol. 36, eds. D. D. Eley, H. Pines and P. B. Weisz (Academic Press, New York, 1989) p. 237. [13] J. G. McCarty and H. Wise, Catal. Today, 8 (1990) 213. [14] B. de Collongue, E. Garbowsky and H. Primet, J. Chem_ Sot., Faraday Trans., 87 (1991) 2493. [ 15] C. B. Alcock and J. J. Carberry, Solid State Ionics, 50 (1992) 197. [16] P. E. Marti and A. Baiker, Catal. Lea., 26 (1994) 71. [ 17] R. J. H. Voorhoeve, J. P. Remeika, L. E. Trimble, A. S. Cooper, F. J. Disalvo and P. K. CraHagher, J. Solid State Chem_, 14 (1975), 395. [18] R. J. I-1. Voorhoeve, D. W. Johnson Jr., J. P. Remeika and P. K. Gallagher, Science, 195 (1977) 4281. [19] J. Kirchnerova, D. Klvana, J. Vaillancourt and J. Chaouki, Catal. Lett., 21 (1993) 77. [20] Y. Ng Lee, F. Sapifia, E. Martinez, J. V. Folgado, R. Ibafiez, D. Beltrfin, F. Lloret and A. Segura, J. Mater. Chem_, submitted. [21] A. Gonzhlez, E. Martinez Tamayo, A. Beltrhn Porter, V. Cortes Corberhn, CataL Today, 33 (1997) 361. [22] V. Primo, F. Sapifia, M. J. Sanchis, R. lbafiez, A. Beltrhn, D. Beltrhn, Solid State Ionics, 63 (1993) 872. [23] A. Wold and R. J. Amott, J. Phys. Chem_ Solids, 9 (1959) 176. [24] I-I. Taguchi, D. Matsuda, M. Nagao, K. Tanihata and Y. Miyamoto, J. Ant Ceram. Sot., 75 (1992) 201. [25] D. W. Johnson, P. K. Gallagher, F. Schrey and W. W. Rhodes, Ant Ceram. Soc. Bull., 55 (1976) 520. [26] Y. Ng Lee, F. Sapifia, E. Martlnez, J. V. Folgado, V. Cort6s Corberfin, in preparation [27] J. A_ N. van Roosmalen, E. I-1. P. Cordfunke, R. B. Helmholdt and H.W. Zandbergen, J. Solid State Chem_, 110 (1994) 100. [28] B. Grzybowska, P. Mekss, R. Grabowski, K. Wcislo, Y. Barbaux, and L. Gengembre, in "New Developments in Selective Oxidation IF' (V. Cortes Corberfin and S. Vie Bell6n, Eds.). Elsevier, Am~erdam 1994, Studies in Surf. Sci. Catal., 82 (1994) 151.

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 1997 Elsevier Science B.V.

757

Effect o f R e d o x T r e a t m e n t o n M e t h a n e O x i d a t i o n o v e r B i n a r y C a t a l y s t . Yu. P. Tulenin, M. Yu. Sinev, V. V. Savkin, and V. N. Korchak Semenov Institute of Chemical Physics, Russian Academy of Sciences 4 Kosygin street, Moscow 117334, Russia The analysis of critical phenomena, such as hysteresis and self-oscillations, gives valuable information about the intrinsic mechanism of catalytic reactions [1,2]. Recently we have observed a synergistic behavior and kinetic oscillations during methane oxidation in a binary catalytic bed containing oxide and metal components [3]. Whereas the oxide component (10% Nd/MgO) itself is very efficient as a catalyst for oxidative coupling of methane (OCM) to higher hydrocarbons, in the presence of an inactive low-surface area metal filament (Ni-based alloy) a sharp increase in the rate of reaction accompanied by a selectivity shift towards CO and H2 takes place and the oscillatory behavior arises. In the present work the following aspects of these phenomena have been studied: (i) the effect of the state of the surface on the kinetic behavior; (ii) the localization of the oscillatory process. In this work the oxidative transformations of methane were studied with a catalyst system that combines an oxide and a metal component. The presence of both components gave rise to complex oscillation phenomena. The influence of pretreatment and reaction conditions over a wide range of parameters (temperature, total pressure, and oxygen concentration) on the oscillatory process was studied. The possible role of mass transfer and the balance of heat in the reactor were analyzed, and a model for the role of the components in the binary catalyst system is suggested. 1. EXPERIMENTAL The scheme of the experimental setup used for the studies of methane oxidation is given in Fig.l. The oxide catalyst (10% Nd/MgO, specific surface area 15 m2g-1, obtained by impregnation of magnesium hydroxide with neodymium nitrate solution, drying and final calcination at 850~ was placed into a quartz reactor (inner diameter 3 mm). The heating device ensured a constant temperature along the whole length of the catalyst bed (-- 1 cm) and a sharp decrease of temperature outside it, as indicated in Fig. 1. Concentrations of components of the reaction mixture were measured by online gas chromatography (GC). The initial CH4- 02 mixtures containing 5 - 15 vol. % oxygen were prepared in a cylinder and supplied to the reactor. Total pressure

This study was carried out under the financial support of the Russian Foundation for Basic Research (research grant No. 96-03-32440)

758

~

3

5

11

10 2 ///.// /7//,"

1.1

8

T/\

reactor axis

9

>

Figure 1. Experimental setup and temperature profile along the axis o f reactor 1 - quartz reactor; 2 - oxide catalyst bed; 3 - cylinder with gas mixture; 4,5 pressure/flow adjusting valves; 6 - thermocouple (metal filament); 7 temperature recorder; 8 - GC line; 9 - cartier gas inlet line; 10 - GC-sampling valve; 11 - exhaust line

Table 1. Methane and methane/ethane mixtures oxidation over oxide, metal, and combined catalysts (650oC, moxide = 15 mg, W = 9 ml/min, P = 20 kPa) Catalyst

Mixture

0 2 Conversion,

CH4:O2:C2H6

Concentration, vol.%

%

C2H6

CO

C02

H2

Oxide

90: 10:0

37

0.25

1.1

1.2

1.3

Oxide+metal*

90: 10:0

82-88

0.035

3.2-4.3

3.6

6.3-8.3

Metal

90: 10:0

-~0

.

Oxide+metal*

88.5:10:1.5

Metal

88.5:10 : 1.5

* - temperature oscillates near 650oC

89.5-93 -4)

.

.

1.2

3.2-4.1

1.5

-

. 3.5 3.8-5.7 -

-

759 in the reactor ranged from 20 to 100 kPa, and reaction temperature was varied from 500 to 750~ The details of preparation and catalytic performance of the mixed MgNd oxides, as well as the experimental procedure are described elsewhere [5]. A chromel-alumel thermocouple (diameter 0.3 mm, sheathed in a quartz cover or bare) was placed co-axially into the reactor filled with oxide catalyst making it possible to detect temperature oscillations accompanying concentration oscillations. This thermocouple in bare form also acts as the metal component in the oxide-metal binary system. The surface area of the thermocouple filament is ~ 7.5x10 -5 m -2. The effect of the state of the surface on the kinetic behavior was studied using various feeder including the methane-oxygen mixture alternating with inert (He), oxidizing (02), and reducing (HE) gases. 2. RESULTS 2.1 Regularities of Oscillatory Process When the metal filament (thermocouple) is placed inside the quartz cover, methane oxidation proceeds in a steady-state regime with high selectivity to C2 hydrocarbons (Table 1). However, if the cover is removed, i.e. the reaction mixture is in direct contact with the metal surface in the hot zone of the reactor, a sharp increase of conversion accompanied by a dramatic change of selectivity from O C M products to CO and H2 takes place. If the oxide component is removed from the reactor, no conversion of reactants is observed indicating a very low activity of the metal filament in methane oxidation. No reaction occurs also when ethane is added to the methane-oxygen mixture. However, if both oxide and metal components are present in the reactor, some part of ethane undergoes conversion causing additional consumption of oxygen (Table 1). Oxidation of methane in the presence of such a binary oxide-metal catalyst proceeds in an oscillatory regime, and both temperature and concentration oscillations take place. Oscillations arise at the temperature at which the rate of reaction over the oxide component becomes noticeable (-500~ As temperature increases, the oscillation amplitude passes through a maximum. The oscillatory behavior disappears when complete conversion of oxygen is reached. In other words, the range of temperatures in which the oscillations are observed covers the range of oxygen conversions from -- 0 to -~ 100%. Variations of total pressure and of oxygen concentrations in the initial gas mixture change significantly the parameters of oscillation. At reduced pressures in oxygen-rich mixtures complex, but regular temperature oscillations are observed. An example of such a behavior is given in Fig. 2: a simple cyclic temperature oscillation observed at low oxygen levels (5 %) changes to a more complex repetitive pattern if oxygen concentration in the initial mixture is increased up to 15 %. At a total pressure 40 kPa and "bare" temperature Tn~n = 630-640~ the number of peaks in one group increases from l to 7 when oxygen concentration is increased from 5 to 15 %. However, the intervals between maxima in one group remain nearly constant, whereas the magnitudes of the main maxima are very sensitive to oxygen concentration (Fig. 3).

760

T,~

Conc. 02 ,o o ,,

720 -

700 -

680

660

640

-

-

JJL

500

1.000

time, sec.

Figure 2. Variations of the shape of oscillations vs. oxygen concentration

A Tmax, ~

A tl,

sec.

1500

(Ptot =

40 kPa)

t~T, ~

~ t, sec.

50

80 -

450

60

200

40

150

30 - 300

40

100

20 150

10

20

O

I I I I I I I 1 11

0

5

I I I I I I I I I I I I I I I I I I I I I I I I

10

0

15

02 Conc., % Figure 3. Magnitudes of the main maxima and intervals between peaks in one group (At1) vs. oxygen concentration (40 kPa, Tmi~= 630~

0

50

IIIIIIIIIlllllllllllllllllllllllllllllllllllllllllllllllll

0

20

40

60

80

100

P, kPa Figure 4. Oscillation amplitude and period vs. total pressure (10 % 02, Train = 640~

761

T

10 min. in He flow

) L time ~/1 ~ 0 rain. in air flow

| l

n

hydrogen flow

time

T CH4 +02

~H4+02

1

time

time

Figure 5. Effect of reduction / oxidation on the oscillatory behavior (Ptot = 100 kPa, 13 % Train= 550~

02,

their amplitude (AT - Tmax - Tmm) increases at increasing total pressure (Fig. 4). However, the range of variation in the parameters of the oscillations in these experiments was not as wide as that caused by the variations of oxygen concentration at constant total pressure. Since the variations of oxygen pressure have such a strong effect on the oscillatory behavior, some additional information about the mechanism of the observed phenomena can be derived from experiments with preliminary reduction and oxidation treatments of the catalyst(s).

2.2 Effect of Redox Treatment of Catalyst(s) Effects on oscillatory behavior of the treatment of the binary oxide-metal catalyst bed in different gases are presented in Fig.5. If the binary catalyst was treated in inert gas, the sharp increase of temperature begins immediately after the reactants are supplied to the reactor, and then the process proceeds in the regular oscillatory manner, despite a phase of oscillation in which the flow of reactants was switched to the inert gas flow. The oxidative treatment leads to an initial disordered set of oscillations at increased activity. After some period of time the regular oscillations are restored. The reductive treatment resulted in a low-rate steady-state regime. The oscillatory behavior can be restored only by oxidative treatment.

762 3. D I S C U S S I O N 3.1 Localization of Oscillatory Process Synergistic effects and oscillatory behavior are observed only when oxide and metal components are present in the reactor simultaneously, i.e. there is a cooperative effect. Their roles can be better understood, if the space where the oscillatory process is localized is determined. This can be done by an analysis of heat transfer in the reactor. The balance of heat in the reactor may be described as follows:

(1)

W •(AXi AHi) = (dAT/dt) Z(cj mj) + k AT where W AT k q mj AXi AHi

- total flow rate; = (T - Tmm), Tmm - temperature in the reactor in the absence o f reaction or in a low-activity phase of oscillations; heat dissipation coefficient; heat capacities of the substances under heating; masses of the substances under heating; - fraction of reactants converted into different products; - AHf (CH4) - ~AHt (prod.), AHf - enthalpies of formation of methane and of different products of its oxidation via the following reactions: -

-

-

- AHi, kJ/mol a. CH4 + 2 02

--> CO2 -[- 2 H20

801.5

b. CH4 + 02

=> CO2 "[- 2 H2

318.3

c. CH4 + 1.5 0 2

~--> CO 4" 2 H20

518.7

d. CH4 + 0.5 02

=> CO + 2 H2

35.5

e. CH4 + 0.25 02 => 0.5 C2H6 + 0.5 H20

469.0

The value of heat dissipation coefficient k may be estimated from the steady-state data (at dAT/dt = 0), for example from the experiment in which a mixture of reactants is switched to a non-reactive gas flow (Fig.5). Estimates of heat transfer in the reactor based on the data from Table 2 give k -- 6.7x10 -4 J K -1 s-1. Taking into account the additional heat produced in the peak of oscillations, their amplitude and the rate of temperature rise, we estimated the value of Z(q mj) for the substance which undergoes heating as -~ 2xl 0 .3 J K -~. In our experiments presented in Table 2 the masses of oxide and metal components were 0.023 and 0.005 g respectively, giving the (q mj) values equal to 2.15x10 -2 and 2.35x10 -3 J K -l respectively. This estimate indicates that the additional heat evolution in the peak of the oscillatory process is not enough to heat the mass of the oxide component, and the stages of the overall reaction responsible for such a non-steady state behavior and for the formation of final products are likely localized on the metal surface.

763 Table 2. Temperatures and concentrations of products in different phases of oscillation and in the non-reactive mixture (moxide = 23 mg, W = 20 ml/min, P = 100 kPa) Mixture CH4:O2:N2

Phase

T, ~

90: 1 0 : 0 90: 1 0 : 0 90 : 0 : 1 0

minimum maximum no reaction

655 675 585

Concentration, vol.% C2H6 CO CO2 H2 --0 -0 .

3.0 3.4 . .

0.35 0.45 .

6.3 6.5

Effects of the treatment of the binary oxide-metal catalyst in oxidizing and reducing gases on the oscillatory behavior provide further evidence for this conclusion. As is shown elsewhere [5], the treatment of oxide O C M catalyst in oxygen and hydrogen leads to a sharp increase and decrease in their activity, respectively. However, this effect is of short duration: after few minutes the rate of reaction undergoes a relaxation to a steady-state level which is the same in all cases, i.e., if the treatment in hydrogen has an irreversible effect on oscillatory behavior, this can be explained by its influence on the metal component. 3.2 Mechanism of Synergy Since the metal filament is inert in both methane and ethane activation, but active in the binary catalyst, this effect is likely due to reactions involving some intermediates. In the absence of the metal filament, the oxide component is a very efficient catalyst for the O C M process, which is well-known to proceed via the formation and recombination of free methyl radicals [6]:

[O1

+ CH4

=> [OH] + CH3

CH3

+ CH3 (+ M) => C2H6 (+ M)

(2) (3)

where [O] and [OH] represent active sites on the oxide catalyst surface in oxidized and reduced states, and M represents a third body. Kinetic simulations based on the model described elsewhere [7] show that, if the probability of repeated collisions with the oxide catalyst surface is high, up to 90% of CH3 radicals formed by reaction (2) can undergo the reverse transformation into methane [OH] + CH3

=>

[O]

+ CH4

(4)

The apparent conversion rate measured as the rate of the formation of the final products is much lower than the rate of reaction (2) due to competition by reaction (4). A complete analysis of mass transfer should include consideration of diffusion (in the pores of the oxide catalyst and in the space between the grains) and of the accompanying reactions of radicals with different species in the gas phase and with surface active sites. This work is presently in progress, and here only a brief discussion is providedThe diffusion coefficient for methyl radicals at the conditions of our

764 experiments (500-600~ and 20-100 kPa) is 0.5-5 cm 2 s-l. If the maximum diffusion path is equal to the radius of reactor, the diffusion time (5x10 -2 - 5x10 -3 S.) is substantially less than the residence time in the reactor (-0.1 s.). This estimate indicates that CH3 radicals escaping from the grains of the oxide catalyst can reach the metal surface and undergo secondary transformations, i.e. additional processes of CH3 removal can occur that decrease the fraction of radicals transformed back into methane and increase the apparent rate of conversion. Since O C M products do not form over the binary catalyst, this suggests that CH3 radical capture by the metal surface is highly probable. As is mentioned above, the parameters of the oscillatory process are very sensitive both to the variations in oxygen pressure and to the pretreatment of the catalyst. In general, the increasing degree of oxidation causes a larger oscillation amplitude. On the contrary, after the preliminary treatment in reducing gas (hydrogen) the process proceeds in a steady-state regime and no major cyclic changes in activity are observed. Since the oscillatory process is localized on the metal component, it is likely that the effects of reduction and oxidation treatments are caused by the differences in the degree of oxidation of the metal surface. According to existing notions [1], rate oscillations in oxidative catalytic reactions may be caused by the existence of some "buffers", or "reservoirs" of oxygen which are able to supply it to the zone of reaction. In particular, a kinetic model based on such ideas is in satisfactory agreement with the experimental data on self-oscillations observed during CO oxidation on a Pt single crystal [8]. In the case of our experiments, the strong effect of oxygen concentration on oscillatory behavior and the sharp increase of activity after the treatment in oxygen is likely due to the variations of surface concentration of reactive oxygen species participating in transformations of methyl radicals:

I01

+ CHa

=> CO

+ 3/2H2

(5)

w h e r e / O / - oxygen species able to oxidize adsorbed CH3 radicals on the metal surface. The data presented above indicate that, although the metal component is not able to activate saturated hydrocarbon molecules, it is very active as a trap for free radicals generated by the oxide catalyst. As a result, reaction (5) competing with (4) leads to the apparent increase in activity of the binary catalyst, and the selectivity of the overall process is determined by the competition between reactions (3) and (5). On the other hand, the treatment in hydrogen flow causes exhaustion of this oxygen "buffer", which cannot be restored in the presence of both reactants in the reaction mixture due to the high reducing activity of methyl radicals. The positive effect of the treatment in helium is likely due to the removal of some species which are present in the reaction mixture and compete with CH3 radicals for the surface active sites. Their removal from the surface in the inert gas flow leads to the immediate restoration of activity when the reaction mixture is switched back. The combination of reactions (2) and (5) may be considered as a scheme for direct methane oxidation to synthesis gas (CO + H2). Similar reactions may determine the high efficiency of mixed catalysts containing Ni and rare-earth oxides for the partial oxidation of methane to synthesis gas [9]. This mechanism does not require a preliminary total oxidation of methane followed by its reforming with CO2 and/or water which was considered as the main route for synthesis gas formation [10,11]

765 CH4

+ 2 02

=> CO2

CH4

+ CO2

=> 2 CO + 2 H2

CH4

+ H20

=> CO2 + 3 H2

+

2 H20

(6) (7) (8)

and, consequently, has no thermodynamical limitations resulting from the positive values of AG in reactions (7) and (8). CONCLUSIONS 1. The cooperative effects observed during methane oxidation over a binary oxidemetal system are due to the formation of active intermediates (free methyl radicals) over the oxide component, their escape from the grains of oxide, and transformation into the final products (including CO and HE) over the metal component, which proceeds in a non-steady-state oscillatory regime. 2. The strong effects of the variations of oxygen concentration and of the preliminary reduction and oxidation treatments of the catalyst indicate that active oxygen species formed on the metal component participate in the transformations of methyl radicals into the final products. REFERENCES 1. V. I. Bykov, Modeling of Critical Phenomena in Chemical Kinetics, Nauka, Moscow, 1988. 2. M.M.Slin'ko and N. I. Eaeger, Oscillating Heterogeneous Catalytic Systems. Studies in Surface Science and Catalysis, v.86, Elsevier, 1994. 3. Yu. P. Tulenin, M. Yu. Sinev, and V. N. Korchak, I l th Int. Congress on Catalysis, June 30 - July 5, 1996, Baltimore, ML, USA, Programme and Book of Abstracts, P-275. 4. D.G. Filkova, L.A. Petrov, M.Yu. Sinev, and Yu.P. Tyulenin, Catal. Lett., 13 (1992) 323. 5. M. Yu. Sinev, G. A. Vorob'eva, and V. N. Korchak, Russ. Kinetics and Catalysis, 27 (1986) I 164. 6. D.J. DriscoU and J. H. Lunsford, J. Phys. Chem., 89 (1985) 4415. 7. M. Yu. Sinev, Catal. Today, 24 (1995) 389. 8. A. L. Vishnevskii, V. I. Elokhin, and M. L. Kutsovskaya, React. Kinet. Catal.Lett., 51 (1993) 211. 9. V.R. Choudhary, V. N. Rane, A. M. Rajput, Catal. Today, 22 (1993) 289. 10. D. Dissanayake, M. P. Rosynek, K. C. C. Kharash, and J. H. Lunsford, J. Catalysis, 132 (199 l) I 17. I I. O. V. Krylov, Russ. Chem. Rev., 61 (1992) 2040.

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3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

767

C a t a l y t i c Combustion of Methane: Activation and Characterization of Pd/AI203 Maria GraCa Carneiro da Rochaa and Roger Fretyb "Chemical Engineering Department, Polytechnic School, Federal University of Bahia, Brazil bLACE, Laboratoire des Applications de la Chimie/t l'Environnement, CNRS, France

Pd/ml203 catalysts for methane oxidation were observed to show an increase in activity after treatment under a reaction mixture. In order to better understand this phenomenon, 2% Pd supported on 6 and a-alumina were pretreated under 1-12or the reaction mixture (CI-h, 02, N2, CO2,1-120) at 500, 700 and 900~ then tested at 320~ (OJCI-I4=2/1) and I atm. The catalysts were characterized by temperature programmed reduction, X-ray diffraction, transmission electron microscopy and X-ray photoelectron spectroscopy. It was observed that the activity was related to the presence of both Pd/PdO species in the bulk and at the surface of the catalyst. In the active state the PdO in the biggest particles appeared more reactive towards hydrogen than in smaller particles, suggesting a better mobility of the existing oxygen species. The relative proportion between PdO and Pd and the increase of the extent of their interface is probably the origin of the activity of the catalysts. 1. Introduction Methane oxidation at mild or low temperatures can be catalyzed by platinum group metals. Palladium is one of the most efficient metals (1) and has been studied over many supports (26). This particular metal, when supported on alumina, begins to show an increase in its activity between 350 and 420~ At these conditions a general increase in the active species particle size is observed. Primet and Briot (7,8) defined two states for the Pd/AI203 supported catalyst: a state I, obtained after simple reduction and a state II after the catalyst had reacted at 600~ for 14 h under O2/CI-h=4/1. State II was more active than state I and showed a lower binding energy of oxygen with palladium. However, the state of the active phase was not clear. The differences in activity, also observed by others, have also been related to the formation/decomposition of PdO (9), to the oxygen adsorbed on metallic Pd (2), to the modification of Pd surface species (3), and to the reconstruction ofPdO crystallites (4, 10). One of the hypotheses for the activation of the Pd catalysts was the establishment of an epitaxy between the metal and the support (8, 11). In the present work, palladium catalysts supported on a well-crystallized alumina, the aalumina, and a transition alumina were studied. The state of the catalysts was examined after a reproducible standard test procedure and the catalysts samples were characterized by different physicochemical techniques.

2. Experimental Pd catalysts supported on aluminas (Puralox transition alumina from Condea and SPH 512 a-alumina from Rhone~Poulenc) were prepared by wet impregnation with PdCI2 to give approximately 2 wt.% Pd, and were then dried overnight in air at 120~ treated in N2 at 550 oC, reduced in H2 at 600 ~C, cooled in N2 to room temperature and stored in a dry box. The catalysts supported on ~-alumina and on a-alumina will be referred to DA and AA respectively. DA-SR means a simply reduced catalyst, DA-500 a catalyst treated under reaction mixture at 500~ C and so on.

768 The powdered catalysts were evaluated in an integral continuous fixed bed reactor (1 g cat, total gas flow rate 6 l/h), with on line gas analysis by gas chromatograhy using both thermal conductivity detector and flame ionization detector. A standardized procedure was adopted to get reproducible samples for characterization. The stored catalysts were pretreated "in situ" in H2 at 400 ~C and heated under N2 flow to 500, 700 or 900 ~C. At those temperatures the catalysts were treated under reaction mixture (molar ratio O2/CI-h=2/1, N2 balance ) for 3 h and cooled in N2 to 320~ to be isothermally tested for 4 hours consecutively at 320, 340, 350 and 360~ (standard reaction). After reaction, the system was cooled to room temperature in N2 and the catalysts were stored in a dry box after air exposure. The catalysts were characterized in the fresh state and at the end of the standard reaction. Temperature programmed reduction was performed in a dynamic system with TCD and quadrupole mass spectrometer analysis (Balzers QM1420) with a 1% H2-Ar, total flow rate of 1.1 l/h and a heating rate of 10 ~C/min. TPR analysis was performed in three steps between -78 and 500~ in order to determine differences in reactivity of the oxygen present in the samples. This analysis allowed also the quantification of the total amount of oxygen present in the catalyst after the standard reaction and to determine the behavior of palladium hydrides. At the end of the TPR at 500 ~C, the catalyst was cooled under Ar to room temperature and oxygen was introduced. This oxygen was then titrated with hydrogen and the Pd dispersion was determined from the amount of irreversible consumption of hydrogen. Powder X-ray diffTaction patterns were recorded using Cu Ktx (1.5418 A) radiation on a Philips 1050/81 vertical goniometer, fired with a diffracted beam graphite monochromator. X-ray photoelectron experiments were carried out in a Escalab 200R (Fisons Instruments) using monochromatic AI Kot radiation and operating at 10.9 Tort base pressure. Transmission electron microscopy analysis was performed in a JEOL 100CX either by direct observation or after extractive replication of samples (support dissolution). 3. R e s u l t s

The Pd coment determined by atomic absorption was 1.77 % wt for the DA catalyst and 2.13 % wt for the AA. Table 1 Initial activities and Pd dispersion Catalyst

Dispersion(%)

Initial activities 320~ (molCHa/h.molPd)

A c t ~ tr~t/Actn~

DA-SR

26.8

0.9

DA-500

20.7

2.6

3

DA-700

5.0

19.6

22

DA-900

2.4

37.3

42

AA-SR

16.6

1.1

-

AA-500

7.9

1.3

1

AA-700

3.5

15.3

13

AA-900

-

8.7

8

769 Table 1 summarizes the results of activity tests at 320~ and the values of dispersion, after reaction. The dispersion of the fresh catalysts was 29% for DA and 23% for AA. Despite a decrease in the dispersion of the active phase after the high temperature pretreatment in the reaction mixture, there was an increase in activity with the DA catalyst showing a higher degree of activation (DA-900/DA-SR=42). Examination of the values of initial and final activity of the AA catalyst also indicates unstable behavior for the AA-SR and AA-500 that could be called an extensive induction period. These results indicate that the support has a strong influence on the catalyst stability (Figs. 1 and 2).

~oo

Conversion(%)

.

Conversion(%) 100 - - - -

~

....

8o~

*'-~

i

80

~ ....

80 . 99

"-~."

60

i:::S

.~.~

..,.-

40i

%

40

~-~ ..:

........ .~

20

..... ., ............

.."

6

8

0

0

2

4

~176

,~

20

,4~.~. e * ' t t'

I

i

I

10

12

14

0:~

16

0

'

2

Time (h) Figure

1

~

Conversion x time for DA

4

6

8

i

i

I

10

12

14

16

Time (h)

series:(-)DA-SR, (+)DA-500, (~)DA-700,

Figure 2: Conversion x time for AA series:(-)AA-SR, (+)AA-500, (~-)AA-700,

(~)DA-900.

(D)AA-900

Figure 3 shows a typical TPR analysis profile between -78 and 500~ Three different temperature domains were revealed in the hydrogen consumption profile when contacting the catalysts with the reducing mixture: a first peak ~ - 7 8 ) at -78~ a second peak when the temperature rose from -78~ to room temperature (HT25) and the fourth one, only present for the AA catalysts, at temperatures higher than 120 o C (I-IT120). In the case of HT-78 two peaks were obtained, their difference giving the amount of irreversibly consumed hydrogen at these conditions. The hydrogen consumption features HT-78 and HT25 are both assigned to reaction with oxygen and to Pd hydride formation. The third peak observed between 75 and 120~ during the heating period, corresponds to an increase in the gas phase hydrogen concentration, associated with the decomposition of the previously formed hydrides. Finally, a fourth peak is observed, only for the AA catalysts, indicating a new hydrogen consumption feature at temperatures close to 120 ~C, that is likely due to the reduction of particular PdOx species. At higher temperature, a small hydrogen consumption is observed and is ascribed to the reduction to methane of the COx species adsorbed on the catalyst surface. The oxygen content of the catalysts in their fresh state and at the end of the standard reaction were estimated, after correction for the hydrogen consumption due to Pd hydride formation. It was thus possible to calculate the average oxidation degree of the catalysts and the number of

770 oxygen layers refered to the exposed metallic surface area determined by hydrogen titration. These results are summarized in Table 2.

Hydrogen Pressure (arbitrary units) HT(-78) ...........................

500 ~

................ . T ( t a m b i

0

.......................................................................

200

400

...............

600

Time (arbitrary unit) Figure 3" Typical TPR profile

Table 2: Hydrogen balance and Catalyst HT-78 HT25 . . . . Fresh DA 0.04 1.1 DA-SR 0.01 2.5 DA-500 0.06 2.4 DA-700 0.2 1.2 DA-900 0.3 0.9 Fresh AA 0.02 0.7 AA-SR 0.03 0.4 AA-500 0 0.4 AA-700 0.05 0.9 AA-900 0.09 0.5

oxygen estimation during TPR (H atoms/Pd atoms) Hyd X HT120 HT (1) O/Pd PdO(%) O . . . . . layers 0.15 1.0 0.9 0.5 50 1.1 0.2 2.3 0.8 1.1 100 3 0.4 2.1 0.6 1.0 100 3 0.4 1.0 0.2 0.5 54 6 0.4 0.7 0.06 0.4 36 11 0.1 0.4 0.2 0.7 0.2 22 0.6 0.01 1.1 0.7 0.5 0.5 54 2 0.03 1.2 0.8 0.2 0.6 60 6 0.2 0.8 0.1 0.1 0.5 46 9 0.3 0.9 0.2 0.01 1 -

(1) Hydrogentitration For the AA and DA catalysts the "HT-78" peak increases with the severity of the treatment in the reaction mixture. The "HT25" decreases for DA catalysts but increases for AA catalysts, indicating that the reactivity of the oxygen of the active phase is dependent on the support. The amount of hydride formed is lower for the AA catalysts because of the existence of a quantity of PdO which is not reducible in the temperature range where hydrides can be formed. These results indicate that the low initial activity of DA-SR and AA- SR and AA-500 catalysts does not have the same origin. The fresh catalysts possess a total amount of oxygen which is close to that of an oxygen monolayer on the Pd surface. This monolayer is likely formed upon contact with air. The CH4 oxidation reaction changes this figure. Difference in the ratio O/Pd is observed. If it is assumed that PdO is the stable oxide in the catalysts, all the available Pd is oxidized in the catalysts DA-SR and 500. It appears that a fully oxidized Pd is not an active state for the

771

catalyst (Table 1). On the other hand, in the highly active catalysts, as are those after treatment at 700 and 900 ~C, the Pd phase is partially in a metallic state and partly an oxidic state. The AA-SR and A-500 catalysts showed a high activation which can be related to the presence of a metal-oxide mixture. In all cases, the metallic character increases after treatment in the reaction mixture at 900 oC. The preceding results indicate that the small particles (D=15-25%) generated by the low temperature treatment do not display a good initial catalytic activity. On the contrary, the large particles obtained alter the high temperature pretreatment (D=2-5%), consist of a mixture of a reduced and an oxidized phase, and exhibit higher specific activities. Furthermore, the amount of hydrogen consumed for hydride formation relative to all the hydrogen consumed increases with the pretreatment temperature. In agreement with the literature, hydrides are more easily formed in big particles (12). XRD diagrams (Fig. 4) show the evolution of the structure in the DA series catalysts. The only peaks that appears in the DA-SR catalyst and DA-500 are due to PdO. The metallic Pd content increases for the catalysts pretreated at 700 and 900 ~C under the reactant mixture. Calculation of the crystallites sizes indicates that the PdO crystallites are smaller than the metallic ones, suggesting a breakup of the latter upon oxidation (Table 3). The more active catalysts are composed of both oxidized and reduced phases, in agreemem with the results of the TPR experiments. The relative amounts of each phase was calculated alter standardization (Table 3). The results are consistent with the complete accounting of Pd in the more active catalysts. In the lower activity catalysts, there is a fraction of palladium that is not visible by XRD, probably because of the small size of the particles, below the XRD detection limit. Pd [11t1 PdO[lOI]

oo-

oC

'

' I 30

"'

i

1 35

'

1~ 40

"

1 2/)"

Figure 4: XRD diagram for the DA series catalyst: a) pure A1203, b) DA-SR, c) DA-500, d) DA-700, e) DA-900. Table 3 Particle size calculated by XRD by the _S.cherrer equation and from H T Catalyst ....DRX (/~) . Phase (_%) . . . . . VdO(10!),, P d ( l l l ) ,PdO Vd Total SR 42 70 70 500 56 70 70 700 78 130 56 36 92 900 89 180 44 56 100

HT(/~) 40 60 220 560

772 As shown in Table 3, the titration values also give an idea of the size of the metallic particles. Unfortunately, the technique does not allow to distinguish between what was PdO and Pd before the reduction. Another AA catalyst with 2.7 wt % Pd content showed the presence of both phases by XRD, after simple reduction and after treatment at 900 ~C, the later with a higher metallic content. In general, the spectra do not show significant crystal modifications, neither for the active phase, nor for the support. TEM micrographs (Figs.5-7) show the sintering and evolution of the Pd particle morphology in going from the reduced sample to the catalysts pretreated under severe conditions of reaction. The results agree with those of Chert and Ruckenstein for oxidized Pd (13). The DA-SR catalyst are composed of regular, spherical small particles (20-50A). The DA-700 catalyst has a heterogeneous texture with small spherical particles (20-80 A) decorating larger irregular particles (>280 A). On DA-900 the small particles disappear and a particle size between 210 A and 700 A is observed. Both, large particles decorated by small ones and smooth particles, are observed. Many particles seem to be polycrystalline. The AA series shows the same particle evolution with a slightly different size distribution. The fresh catalyst has particles of size between 20 and 300 ~ most of them around 20 K The AA-SR has basically the same size distribution but large particles begin to appear (250-500 A). In the AA-700 catalyst, particles with size lower than 40/~ disappear and sintering is observed with particles up to 900 A. Finally in the AA-900 only big particles are observed. Micrographs show the size and morphology evolution of the DA and AA series of catalysts (Fig. 5 and 6 ). The morphology changes in these catalysts with the high temperature pretreatment is evident: initial spherical particles change to irregular ones decorated by small ones. Perforated particles are also observed (Fig. 7).

Figure 5: TEM micrographs for DA series catalysts: (a)DA-700, (b)DA-900.

773

Figure 6: TEM micrographs for AA series catalysts: (a)AA-fresh, (b)AA-SR

80 nm Figure 7 TEM micrographs details of DA and AA catalysts: (a) DA-700, (b) AA-SR.

....

!ili!ii:i(ii!il

774 XPS results are presented in Table 4 . The decrease in Pd/AI ratios is consistent with the catalyst sintering. Higher values for the AA catalysts are related to the low specific area of the support compared to that of the DA catalysts. The CI/AI ratio is reported because the literature suggests that dechlorination could be related to catalyst activation during methane oxidation (3, 10, 14). This could be true for the DA catalyst, but not for the AA catalyst which shows a lower activity for a lower CI/Pd ratio. Table 4: Binding enerBies for Pd 3d 5/z and atomic ratio by XPS Catalyst BE (eV) Atomic ratio Pd/Al Cl/Al DA-SR 336.0 337.9 0.007 0.009 DA-900 334.6 336.5 0.002 0.002 AA-SR 335.7 336.9 0.044 0.01 AA-900 334.3 336.1 0.011 0.004 After treatment at 900~ both catalysts series present similar BE values (334.6 and 334.3 eV) which are assigned to metallic Pd, even though these values are lower than that reported in the literature, 335.2 eV (15). Such negative deviations have been observed in the literature and related to a partial reduction of the support (15). The 336.5 and 336.1 eV values may be assigned to PdO, whose reported B.E. is 336.5 eV (16). These results suggest that the active phase contains both metallic Pd and PdO. The DA-SR presents BE values of 336.0 and 337.9 eV close to those reported by Otto et al. (16) assigned to two species of PdO: the first one, normal, and the second one, with BE of 338.3 eV possibly belonging to small clusters ofPd § in strong interaction with the support or PdO2. The AA-SR presents a mixture of oxidized and reduced palladium, different from the species obtained after the reactive high temperature pretreatment. 4. Discussion

An increase in the activity of alumina supported Pd catalysts under conditions of methane complete oxidation was once again observed in the present work. Even catalysts which were treated under the reaction mixture at 900 ~C, an elevated temperature even for a combustion reaction, were activated. The samples can only be activated under reaction conditions and at high temperatures. Furthermore, the increase in activity was accompanied by an increase in the particle size. However, the catalytic activity could not be correlated directly with the particle size as already observed by others (4, 5). The absence of a correlation is not surprising since it was observed that the active catalyst contained both metallic and oxidized palladium. By quantification of the oxygen content it was possible to estimate the 'thickness' of the oxide layer, related to the exposed surface of Pd (whose size was estimated after complete reduction of the sample). It was observed that the more active catalysts had a larger number of oxygen "layers" which suggested the participation of bulk Pd in the reaction. Therefore, for methane oxidation, the Pd/AI203 catalytic properties seem to be controlled by a redox mechanism, where the oxidation and reduction rates of the active phase depend on the size and stability of the Pd and PdO species. The small PdO particles seemed to be poorly reactive, due to their higher stability or to a greater interaction with the support. In fact, analysis of TPO and TPD data revealed that the small particles of PdO decompose at

775 temperatures 40 to 50 ~C higher than those present on the particles after reactive treatment at 700 and 900 ~C. These results, associated with the heat of adsorption of oxygen on different sizeA Pd particles (17) or to hydrogen reactivity on PdO (7,18) show that oxygen mobility is higher in the larger particles than in the smaller ones. In summary, for the complete oxidation of methane, Pd/A1203 catalysts with small particles do not constitute the active phase. A metallic phase alone is not effective. On the metallic phase, methane undergoes dissociative adsorption, probably favored by a high electronic density as observed by XPS. The high electronic density of Pd could be due to a partial reduction of the support, to the presence of carbon donor species around the particles, or to adsorbed oxygen species on PdO localized very close to the metallic state. QMS analysis during TPD of the used catalysts indicated the evolution of CO2 from the sample at the temperature of PdO decomposition. The presence of carbon filaments after exposure to CI-h at high temperature was also observed by SEM. It is known that PdO decomposes around 850~ so the treatment at 900~ under the reaction mixture might be expected to form only metallic Pd. However the DA-900 catalyst, was more active than DA-SR and DA-500 and at the end of reaction a mixture of Pd-PdO was observed. This means that the reaction itself allowed a partial oxidation of the catalyst. Hence, it can be concluded that catalysts which are fully reduced or fully oxidized are not active (DA-S1L AA-SR and AA-900). The determination of the active sites remains a challenging subject because we are not able to quantify directly the active fraction of the catalyst constituted by PdO. The overall metallic surface area is not directly related to the active sites working during the reaction. The TEM observations in this study are similar to those obtained by Chen and Ruckenstein (13), although in a different range of temperature. In this case the atmosphere and its interaction with palladium was more complex. An initial strong sintering was observed, and, subsequently in some particular sites there was an apparent oxidation which was responsible for the formation of cavities. In a following step, the particles started to fragment and their shape became irregular (cauliflower shaped particles). Further differential oxidation was observed which seemed to depend on certain active sites present on the catalyst. At 700 ~C, a fraction of the Pd could reduce and apparently this led to the agglomeration of particles, while at the same time a spreading of particles occurred due to their oxidation. This is probably the reason for the decorated cauliflower shaped particles; these particles are likely metallic inside with small PdO particles outside. At 900~ PdO is not stable and the particles are more crystalline, even if some decorated particles still exist. The disappearance of those small particles was also observed after reduction of the used catalyst (17) and was the possible cause for the decrease of activity, even if some particular sites were still oxidized. EDX analysis of some particles indicated the presence of both Pd and PdO in the very same particle in different proportions depending on the position. In this particular system, well crystallized particles are not active, but particles rich in defects, capable of inducing modifications in the Pd state are. In this situation, the catalyst activity is probably related to the existence of a Pd-PdO interface, not to an active phasesupport interface From the differences observed by TPR analysis, the species present in the fresh catalyst underwent changes after reaction. All results indicated that the active catalyst contained a PdPdO mixture, coexisting mainly in the same particles, which was formed during the reaction. These phases were of a size and morphology that allowed oxygen mobility. The small

776 particles alone, initially in complete oxidized or reduced state, were not active, due probably to a strong interaction with the cartier. Depending on the support, the activation can proceed in a different way. It is probable that during the reaction there is oxidation and reduction of the active phase, with the catalytic reaction occurring at the oxide-metal interface which is continuously undergoing reconstruction. Acknowledgements: M. G. C. da Rocha acknowledges the financial support of CNPq, Conselho Nacional de Desenvolvimento Cientilico e Tecnol6gico of Brazil and the contributions of M. Bran, P. D61ich6re in the XPS experiments, G. Bergeret in XRD and C. Leclercq, F. Beauchesne in the electron microscopy analysis. References 1. R. B. Anderson, K. Stein, J. J. Feenan and L. J. E. Hofer, Ind. Eng.Chem., 53 (1961) 809. 2. C. F. Curtis and B. M. Willatt, J. Catal., 83 (1983) 267. 3. R~ F. Hicks, H. Qi, M. L. Young. and R. G. Lee, J. Catal., 122 (1990) 295. 4. T. R. Baldwin and R. Butch, Appl. Catal. 66 (1990) 359. 5. F. H. Ribeiro, M. Chow and R. A. Dalla Betta, J. Catal., 146 (1994) 537. 6. J. G. McCarty, Catal. Lett., 26 (1995) 283. 7. P. Bdot and M. Primet, Appl. Catal., 68 (1991) 301. 8. P. Bdot, Thesis, Universit6 Claude Bernard, Lyon I, France, 1991. 9. R. J. Farrauto, M. C. Hobson, T. Kennelly and E. M. Waterman, Appl. Catal. A: Gen., 81 (1992) 227. 10. T. R. Baldwin and R. Butch, Appl. Catal., 66 (1990) 33 7. 11. P. Briot, P. G~ezot, C. Lederc and M. Pdmet, Microsc. Microanal. Micr0strut., 1 (1990) 149. 12. M. Boudart and H. S. Hwang, J. Catal., 39 (1975) 44. 13. J. J. Chert and E. Ruckenstein, J. Phys. Chem., 85 (198 l) 1606. 14. D. O. Simone, T. Kennelly, N. L. Bmngard and R. J. Farrauto, Appl. Catal., 70 (1991) 87. 15. T. H. Fleisch, R. F. Hicks and A .T. Bell, J. Catal., 87 (1984) 398. 16. K. Otto, L. P. Haagk and J. E. de Vales, Appl. Catal. B: Environ., 1 (1992) 1. 17. P. Chou and M. A. Vannice, J. Catal., 105 (1987) 342. 18. M. G. Almeida, F. Beauchesne, M. Pfimet and R. Frety, Book of Abstracts, vol 2, Europacat, p 917, Montpellier, France, 1993.

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

777

Activity o f manganese dioxides towards V O C total oxidation in relation with their crystallographic characteristics C. LAHOUSSE, A. BERNIER, E. GAIGNEAUX, P. RUIZ, P. GRANGE, B. DELMON Unit6 de Catalyse et Chimie des Mat6riaux Divis6s, Universit6 catholique de Louvain, 2/17 P1. Croix du Sud, B 1348 Louvain-la-Neuve, Belgium This paper describes the catalytic activity of various forms of manganese dioxides towards volatile organic compounds deep oxidation. Important differences in activity are evidenced for very closely related structures ; the most active sample is a high surface area nsutite. A parallel is drawn between the findings in the literature of battery applications and the catalytic activity results. The superior activity of nsutite is attributed to a different oxygen coordination and to the clustering of cationic vacancies in the bulk which improves electronic and protonic conductivity.

1. INTRODUCTION Whatever the application, mild or deep oxidation, a limited number of parameters are usually thought to influence the catalytic properties of oxides. These include the nature of the metal cation(s), their oxidation state, their first coordination sphere and the corresponding density of each type of site at the surface. In this work, the catalytic activity towards Volatile Organic Compound (VOC) deep oxidation of 3 crystallographic forms of manganese dioxides: pyrolusite (13), nsutite (~,)and ramsdellite was studied. These 3 forms, respectively named in this article: py-MnO2, nsu-MnO2 and rams-MnO2, are very closely related in structure [ 1]. In each structure the Mn IV cation is in an octahedral environment with each oxygen comer of the MnO6 octahedron shared by 3 octahedra, the only difference being in the arrangement of the O coordination which can be either planar or tetragonal truncated. So, all of the above characteristics namely cation nature, oxidation state of manganese and first coordination sphere are identical in these 3 phases. Moreover, a SEM study did not detect any anisotropy and preferential exposure of specific faces which could give rise to a different catalytic behavior. The aim of this paper is first to report on the catalytic activity of these very closely related structures of MnO2, then to show that characteristics like oxygen coordination, which are not usually considered for explaining the catalytic behaviour of oxides in catalysis, do affect their catalytic properties. One sample of py-MnO2 and rams-MnO2 and two samples of nsu-MnO2 of different specific surface areas were used. 2. EXPERIMENTAL

2.1. Catalysts origin The catalysts were provided by SEDEMA-SADACEM, Belgium. Table 1 gathers the

---

Table 1 : Comparison of the characteristics of the different forms of MnOz

Catalysts studied y-Mnoz (1 00) y-Mnoz ( 100) y-Mnoz ( 100) y-Mnoz ( 100) Y-MnOz (35) Y-MnOz (35) Y-MnOz (35) Y-MnOz (35) (3-Mn02 P-MnOz P-MnOz P-MnOz

MnOzrams MnOzrams MnOzrams MnOzrams

data p Crystalline form nsutite nsutite nsutite nsutite nsutite nsutite nsutite nsutite pyrolusite pyrolusite pyrolusite pyrolusite ramsdellite ramsdellite ramsdellite ramsdellite

EDEk )article

surf.

mVg ratio ize/l.tn 1.95 35 100 1.95 1.95 1.95 1.95 1.95 1.95 1.95 1.96 1.96 1.96 1.96 1.95 1.95 1.95 1.95

35 35 35 37 37 37 37 34 34 34 34 20 20 20 20

100 100 100 35 35 35 35 50 50 50 50 18 18 18 18

voc

7

Ea Hex Ea Hex Ea Hex Ea Hex Ea Hex Ea Hex Ea Hex Ea Hex

----

Conversion % (W=O. Ig ; [VOC]=2SO feed composition 100 120 140 150 160 21 60 88 100 25 Ea+Hex 0 0 0 0 51 Ea+Hex 16 39 48 77 Ea+Hex+HzO 10 4 4 5 5 Ea+Hex+HzO 7 84 7 6 24 48 Ea+Hex 1 5 0 0 0 Ea+Hex 10 32 47 20 Ea+Hex+H20 12 7 9 10 9 7 Ea+Hex+H,O 41 7 10 21 2 Ea+Hex 0 0 0 1 0 Ea+Hex 20 23 32 12 Ea+Hex+HzO 10 7 3 7 7 Ea+Hex+HzO 8 7 40 13 20 2 Ea+Hex 1 1 0 0 0 Ea+Hex 40 9 16 24 Ea+Hex+H20 10 7 6 5 5 Ea+Hex+HzO 6

10 d m i n air)

TI --- - - -180 200 220 240 --- -- 260 - 280 O(

93

100

100

50

80

98

100

87

98

100

57

84

98

7

18

38

90

100

100

70 99 20 95 0 63 6 90 10 61 4

99 60 100 22 98 8

61

100

40 95 16

--- ---- --43

64

81

90

779 characteristics of these products. Thanks to careful preparation by SEDEMA, each of these samples also presents a reasonable surface area together with a "good" crystallinity. According to the data provided by SEDEMA, the amount of CO2, S, alkaline, alkaline-earth and transition metal cation content is extremely small on the nsu-MnO2 samples ; 0,3 % CO2 and 0,02% S can be found in the rams-MnO2 sample. In addition, the XPS analysis shows a trace amount of S (SO4) on the surface of the py-MnO2 sample.

2.2. Catalyst characterization and testing XPS enables the determination of the oxidation state of the Mn ions. The most precise evaluation rests on the measurement of the binding energy difference between the Mn 3s main peak and its shake-up satellite [2-4].The XPS analysis of the samples was performed on a SSI X-probe (SSX 100/206) spectrometer of FISONS equipped with a monochromatised microfocused AI I ~ X-ray source (1486.6 eV). The angle between the sample surface and the electron detection axis was 55 ~ The analyser pass energy was set at 50 eV and the analyzed area was 1.4 mm 2. At these conditions, the energy resolution determined by the Au 4f7/2 full width at half maximum (FWHM) of gold was 1.1 eV. The XPS results are reported in Table 1. The BET surface area of the samples was measured by nitrogen physisorption at 77 K (196~ using a Micromeretics ASAP 2000 apparatus (Table 1). The samples were examined using a Hitachi S-70 scanning electron microscope (accelerating voltage : 15 kV). XRD patterns of the samples were obtained on a Siemens D5000 apparatus working with the copper K~ line. Ethylacetate (Ea) and n-hexane (Hex) were chosen as representative VOC's reactants [5]. 250 ppm of the two molecules in air were used with a contact time of 60 kg s/m3 (which corresponds to a space velocity of 72000 h-1 (NTP) with nsu-MnO2). For experiments with water vapor, the stream was saturated at 25~ with 20000 ppm of water. The contact time, the concentration, the nature of the VOCS and the presence of water vapour are representative of the conditions of VOC removal in printing industries [5]. Before measuring the conversion as a function of temperature, the catalysts were stabilized by letting the reaction proceed for 16 h at 150~ The conversion was measured as a function of time during this period, at which end it reached a constant value. Then, the temperature was decreased to 100~ and increased by steps of 20 or 10~ until complete combustion was obtained. The conversion was then again measured at 150~ to verify that the conversion at the end of the test was the same as after the preliminary stabilisation period. Table 1 presents the conversions for the same catalyst weight, and thus allows a comparison of the specific conversions (conversion per gram of catalyst). The intrinsic conversions (conversion per square meter of surface area) will be compared in the figures presented in this paper. 3. RESULTS The conversion of the two reactants on the different catalysts in the presence or in the absence of extra water vapour is given in Table 1. The existence of a competition for adsorption between the different VOCs has been demonstrated in an earlier communication [6,7]. As far as Ea and Hex are concerned, the presence of Hex in the flow has no influence on Ea conversion, but Ea, strongly inhibits Hex conversion as long as some of it remains present. Nevertheless, for a same catalyst weight, the conversions measured with nsu-MnO2 are about five time higher than those measured with the other forms of manganese dioxide (see Table 1). Figure 1 provides a comparison between the different catalysts corrected from the variation of surface area. The differences in activity appear when both specific conversion

780

and conversion corrected for surface area are considered. According to the Ea and Hex conversion, the MnO2's efficiency for the VOC removal can be ranked as follows: py-MnO2< rams-MnO2< nsu-MnO2 of 35 m2/g< nsu-MnO2 of 100 m2/g. A mechanical mixture of pyMnO2 and rams-MnO2 shows an activity which corresponds well to the average value between those of rams-MnO2 and py-MnO2. Figure 1: Comparison of the intrinsic activity (towards Ea)of different form of MnO2 30 25

, 2o ~ 5 ~10 5 0 py-MnO 2 + rams-MnO 2 35 m2/g

rams-MnO 2 18 m2/g

nsu-MnO 2 35 mVg

nsu-MnO 2 100 m2/g

The effect of water will be described more thoroughly in a future paper [8]. As shown by Table 1 data, water affects the activity of manganese dioxides. It also dramatically reduces the time needed to obtain a stable conversion. The two effects can be explained by a hydration of the oxide surface. As far as the specific activity of the different catalysts is concerned, the order of catalyst activity is not modified, but the differences are very much attenuated. Figure 2 presents the activity of the different MnO2 samples in the presence of water vapor corrected for the variation in surface area. The order of activity is different from Figure 2: Comparison of the intrinsic activity (towards Ea) of different form of MnO2, in presence of water 25 20 [] 100"C, [] 120"C [] 140"C

~9 lO

9150"C i

I! 160"C, ~

5 0 py-MnO 2 50 m 2/ g

nsu-MnO 2 35 m 2/ g

nsu-MnO 2 100 m 2/ g

rams-MnO 2 18 m 2/ g

781 that attained at the "dry" conditions. Indeed, the conversion measured with 7-MnO2 of 35 m2/g is in between those measured with py-MnO2 and rams-MnO2. The activity of the nsu-MnO2 of 100 m2/g is still above any of the others. 4. DISCUSSION Considering that the metal cation (s), their oxidation state, their first coordination sphere and the crystallite morphology are identical, the differences in activity (figs. 1 and 2) between these different forms of MnO2 are unexpected. The impurity level of all these samples is very low. Miscellaneous data collected throughout our work suggests that the wellknow acid-base modifiers (e.g." K, C1), as well as probably sulphur containing compounds, have negligible effect on activity. The differences of activity between the different samples cannot therefore be attributed to the effect of impurities. Furthermore, the O/Mn ratio in the rams-MnO2 and the nsu-MnO2 samples are identical" this implies that they present exactly the same oxidation state (O/Mn=l.95). The py-MnO2 sample initially contains a slightly more oxidised Mn (O/Mn=l.96). But XPS shows that the surface oxidation state becomes identical to that of the other forms under the reaction conditions (XPS Mn 3s shake up satellite = 4,9 +_ 0.1 eV). The variation of activity reported above cannot be linked to differences in the Mn mean oxidation state. In addition, as mentioned in the introduction, Py-MnO2, nsu-MnO2 and rams-MnO2 are so closely related structurally that the first coordination sphere of the Mn TMions is the same in these 3 phases. A preferential exposure of a given type of site could in principle explain the differences in catalytic behavior. But SEM did not reveal any such differences. Py-MnO2 was constituted of particles in the form of "columns" with an average size of 180xlS0x500 nm; rams-MnO2 particles were more like thick "platelets" of 200x5 nm; nsu-MnO2 particles did not exhibit angles ("pebble form") and the particles had a diameter of 160 nm. The elemental crystallites of the nsu-MnO2 samples were smaller than the apparatus resolution. However, SEM micrographs seemed to indicate that the smallest observable particle was constituted of smaller elements in the case of the 100 m2/g nsu-MnOz than in the 35 m:/g sample. But, none of these differences corresponded to important changes in the development of different type of faces. In summary, the parameters which are usually considered to affect the performances of oxidation catalysts are almost identical. Accordingly, our results point to a new important effect. The only meaningful difference known between these phases concern the crystalline structure and the oxide texture. This is discussed herafler. As far as texture is concerned, the particles size (Table 1) and the crystallite size shown by SEM pictures are comparable for py-MnO2, rams-MnO2 and the 35 m2/g nsu-MnO2. Considering surface area, the exposed surface of rams-MnO2 (18 m2/g), py-MnO2 (50 m2/g) and of the sample of nsu-MnO2 (35 m2/g) could be considered as reasonably similar. But the nsu-MnO2 with 100 m2/g could be considered as substantially different. Its high surface area could be explained by or linked to a higher density of point defects on this oxide surface. This conclusion is consistent with the SEM study results, where the elementary crystallites appeared smaller for this sample. As for crystalline structure, we already mentioned that Mn coordination is the same in each of the samples. But there is a difference in the arrangement of the MnO6 strings. 3D views provided in the figs 3,4,5 attempt to reproduce the ones found in literature [ 1]. In py-MnO2, MnO6 octahedra form single strings linked to each other by shared O octahedron summits. Rams-MnO2 has a chain structure similar to that of py-MnO2. But in rams-MnO2 the chains are sorts of ribbons constituted of 2 strings, which involves the sharing of edges between two adjacent chains. Nsu-MnO2 is a structural intergrowth of py-MnO2 and rams-MnO2 in which

782 layers of single strings and ribbons alternate in a random fashion. The shift from rams-MnO2 to py-MnO2 is actually linked to a change in oxygen coordination. Normal oxygen bulk coordination number is identical, namely 3 for rams-MnO2 and py-MnO2, but the form of the coordination sphere differs. Indeed, in py-MnO2, the 3 Mn linked to one oxygen are placed in a single plane, whereas in rams-MnO=, 2 types of oxygen coordination are encountered in equal amounts: the py-MnO2 planar type (Opl) and a tetrahedral truncated type (Ott) (see fig. 6 and 7). The Opl type is the O linking two ribbons. The Ott occurs at the junction between the two strings constituting the ribbons. The nsutite being a random intergrowth, it presents also the two types of oxygen ions but in variable amounts 0 < Ott/(Ott+Opl) 0.67. The constituents of the prepared sample wcrc found to bc also (VO)2P2OT, VO(PO3)2 and some amorphous compounds. At this takes place, an the cations wcrc considered by authors [12] to bc bonded in vanaclyl groups V - O and phosphorus atoms form Bromtcd acidic center each. It has bccn found an increased concentration of phosphorus over the surface as compared to the bulk and the hishcr phosphorus content m the sample the 8town bulk concentration of the reduced vanadium ions were observed. As for V-P containin8 samples prepared by impregnation method, one can expect predominated mamtmnin8 of the supported compounds structure over the surface which would be very similar to those synthesized in bulk samples. On the other hand, it is well known that amon8 oxide supports SiO2 is the weakest dispersant of vanadium oxide over the surface [13] and always there is some V2Os phase-like formations despite the very low surface occupation de8rcc [11,14,15]. Quite limited information on the mixed vanadium-phosphorus catalysts supported on silica arc represented m open scientific literature. For cxamplc, in [4] it was proposed the method to prepare well dispersed VPO composition over SiO2 but, unfortunately, the authors did not ttun out successfully to determine the chemical composition of the surface phase formed (supposcly VOPO4 to8cthcr with VOx). In [3] r 4 as a major phase of the similar samples was d e s c ~ d and r toscthcr with umdenfilicd compound were found to be constituents of the surface layer in [6]. In the present paper, impregnated IPV samples on acrosil, containin8 5, 10 and 15 wt.% vanadium at fF of about 0.6 were prepared. All the XRD patterns for the samples arc proved to be very similar and contain the most characteristic reflections of [~-VOPO4, cc-VOPO4 and (VO)2P~O7 phases a8~in.~t the considerable halo-effect. By means of scannin8 electron microscopy it has been established (Figure 1), that there are crystal resions over the surface of IVP imprc8natcd samples, winch resemble the structure of ~-VOPO4 phase. These regions size arc defined by supported compounds amount and the more active components supported the lar8cr domains formed. Microprobe analysis in 10 sites on the surface showed the ratio V/P m these domains was close to that observed for bulk VOPO4 phase (see Table 1). It can be asstmacd that thermal treatment of the catalysts as prepared causes the phosphorus migration over the surface to produce I~-VOPO4 phase and the rest phosphorus and vanadium form the prcerystal compounds or remain unbonded toscthcr. The microprobe analysis of the surface between the mentioned above VP oxide domains confirmed this asstmaption (Table 1). The

790

~,~

Figure 1. Microphotographs of the impregnated and graBcd VPO/SiO2 catalysts surface: a -bulk p-VOPO4, b - IPV-3 (15 wt.% V), c - IPV-4 (5 wt.% V), d - PVD-5 (4 wt. % V), c - PVD-3 (2.72 wt.9~ V).

791 binding cn~gy of V 2p-clcctrom lacing 517.2 cV (B.E. P 2Pv4~ 133.7, O ls 532.4, Si 2p = 102.6 r can br indicative of a presence of ions on the samples surface d ~ to its lower value as compared to that for ~-VOP04 bulk phase (518.6 r It is conceivable that these vanadium ions do not r into the composition of crystal formations and they arc part of prcc~stal structures. Tablc 1 El cnts ratio on the surface of imprcgnatcd and graftcd VPO/SiO~_ samples a I ~ microprobe analysis Preparation Concenmethod tration of vanadium, wt.% Iml~csnahon Grafting

[3-VOPO4

Bulk

Ratio V:P:Si on the surface

P/V

ratio

15.0 5.0 4.0 2.7 1.8 -

MPA

1.2 1.2 1.4 1.5 I.I 1.0

XPS

C~taIs

Support

1:1.10:0.52 1:1.05:0.70 I: 1.02:0_90 I:I.II:0.96 1:1.02:1.08

1:1.24:3.02 1:1.16:5.96 1:1.34:6.04 I: 1.38:6.45 I:I. I0:7.32

I:I.03:0

-

1:1.36:4.07

1:1.32:5.25 -

I:I.08:0

Table 2 Properties of impregnated IPV and bulk VPO catalysts m the reactions of benzene and n - b ~ c oxidation*

Catalyst

IPV- 1 IPV-2 IPV-3 I3-VOPO4 tt-VOPO4 V20s

(VO) 2P207

Vanadium on support, wt.%

SSA,

5 10 15 -

120.5 71.4 25.3 2.0 -

-

4.5

-

m2/g

4.2

Benzene oxidahon

Tm~, *C 450 420 460 550 500 -

410

Sin, % 65 68 58 60 56 0

55

n-Butane oxidation Sst~, 17 20 16 17 35 I

54

* - Tmx - tcml)~aturc of the maximmn of malcic anhydride yield; selectivity to malcic anhydride

Sst~

792 A correlation between bulk and supported vanadium phosphorus catalysts properties (Table 2) shows much higher activity of the latter samples as related to former ones. The most active catalyst is IPV-2 where VI'O microerystals over the surface are of an average size. The f m ~ e r growth of these formations size leads to decrease of the catalytic activity and selectivity (see sample IPV-3 in Table 2). From the catalysts efficiency comparison it can be concluded that no mdividtml phase tested is responsible for the supported samples catalytic performance. However, when the crystal formations on the surface arc large c n o u ~ the properties of the samples draw near the bulk compounds activity and selectivity. 3,2, Properties of gralted VPO catalysts, At simultaneous graltin8 of vanadium and phosphorus compounds followed by hydrolysis of the surface grOUl~ (samples PVD-4, -5, -6) the components can be fixed on the surface accordin8 to the reac~om:

H20

>Si-OH + EOC1,--> >Si-O-EOC12 + HCI . . . . > >Si-O-EO(OH)2,

)Si-OH

+ EOCI~ -->

)si-o\

2

)si-5

si_o/Roci + 2I-ICa .... >

"

si-o

o(oI-I),

~

Si-OH )Si-Ox Si-OH + EOC1,--> ) S i - O - - ) E = O + 3HC1, Si-OH ~Si-O/

where E = V or P. Table 3 Catalytic properties of the grafted PVD samples m benzene and n-butane oxidation*. Catalyst

Benzene oxidation Vanadium fp = content, wt.% P/(P+V)

Xuc, ~; PVD-6 PVD-5 PVD-3 PVD-4

2.78 4.00 2.72 2.87

0.50 0.58 0.60 0.67

90.0 89.6 89.0 90.5

S~, ~ 60.0 62.0 72.0 74.6

n-Butane oxidation

S~, % 19 20 24 22

W.10v 9.4 8.8 9.1 9.3

* X H C - hydrocarbon conversion, SMA - selectivity to malcic anhydride, W. 10~ rate of n-butane oxidation, mole/s.atom V

793 From the Table 3 it can be seen that the total m o u n t of grafted compounds exceed that of OH 8roulm on aerosfl surface. It can be explained by additional centers folvning dl~ to either partial dissociation of the surface siloxanc groups or the mixed (V-O-P) structmca formin8 during the grafting. At the consecutive g t a f l i ~ of the phosphorus and vanadium compounds the latter can bc fixed to not only fzcc OH grOUl~ but also to nascent P-OH groups to produce mixed P-O-V structures. Due to close catalytic properties of the samples prepared by simultaneous and consecutive grafting one can conclude that the obtained in both cases P-O-V structmes arc very similar despite the order of grafting. An investigation of their surface by microscopy supports this idea (scc Figure 1). Over the surface of the grafted samples island crystal structures, similar to those for bulk 13-VOPO4 and impregnated catalysts, arc observe& The mcasurcd by XPS clcctrom bmdin8 cncrgics for all the clcmcnts of PVD-5 sample dillcr ncgli81bly from those found for impregnated catalyst. At the vanadium concentration reduced to 0.8 wt.% (samples series SVP) such crystal structmcs arc not found_ In the last case, an influence of phosphorus atomic fraction fp at the constant concentration of vanadium over the surface on the prepared samples p r o p o s e s were studied (Table 4). Table 4. The suffacc state for SVP catalysts and thc sclcctivity towards malcic anhydridv at benzene and ~ - b ~ a n c o~daho~ Catalyst

fp

Degree of Content V~ wt.9~ SiOH $roulm covcra$c before after rcaction rcaction

Oxidation* bcn2~nc, SMA, %

n-butane SuA,

SVP-0 SVP-3 SVP-4 SVP-1 SVP-8 SVP-7 SVP-9

0 0.27 0.49 0.62 0.71 0.82 0.89

0.33 0.48 0.66 0.85 1.21 1.72 2.77

0.09 0.09 0.12 0.15 0.20 0.38 0.55

0.11 0.13 0.43 0.69 0.63 0.45 0.55

36 57 52 0 0

4 6 7 33 36 18 19

E 205 202 170 120 126 182 176

* - SMA - selectivity to malcic anhydridv, E - activation cncrsy of n-butane oxidation, kl/molc

An introduction of phosphorus to vanadium-containing aerosfl leads to significant changes m its catalytic properties. In so doing, an increase of the phosphorus concentration (up to certain value) provokes growth m both activity and sclcctivity of the obtained samples at hydroca~ons oxidation (scc Table 3 and 4). Low occupation dcsrce as compared to OH g r o u ~ concentration favours

794 vanadium fixation sepazately from phosphorus or/and little amount of mixed PO-V structures formins. At the essential excess of phosphorus over the optimum amount the catalytic performance changes for the worse. An activation cncrsy for n-butane oxidation rcachcs its minimum value over thc most active samplc and rises asain with increase of the phosphorus concentration. F_~SRspectra (Figure 2) recorded for the samples with relatively low phosphorus concentration shows high rcsolulion structure attributed to V4t ions at their low concentration and an ~ c c of cxchansc interaction between them. With increase of the phosphorus concentration m the samples V4t ions amount growth and cxchansc interaction between them or in V4~-V~+ pairs appeaxs. Further increase (over optimum) phosphorus content leads to rcappcaxancc of high resolution structure of the spectra.

/

Figure 2. ESR spectra of the grafted V-P/SiO2 catalysts: a - SVI'-4 (fp = 0.49), b - SVP-1 (fp = 0.61), c - SVP-9 (fF = 0.89), d - SVP-8 (fp = 0.71).

795 Compann8 V4+ ions concentration in the samples with cor~cspondm8 specific reaction rates and selectivity towards malcic anhydride one can notice their symbatic growth (except samples having fp > 0.7). T h o u ~ V4+ ions concentration arc found to play a decisive role m selective oxidation, it is important but not sufficient requ~ement for the VPO/SiO2 catalysts to be selective (Table 4). As for the participation of intermediate O~- and O species m selective run of the hydrocarbon oxidation, there arc some doubt on such mechanism. The mason for this conclusion can bc the observed increase of the proccascs selectivity with the decrease of 0 7" and O- concentration [16] and reaclmtg the maximum selectivity at surface oxygen radicals content of about zero. So, it should be assumed mechamsm for the hydrocatbom selective oxidation revolving oxysen from surface grou~ V=O. The same conclusion has been drawn for the bulk V2Os-P205 catalysts in butene-1 [7] and b e ~ e oxidation [17]. The presented results on grafted VPO composition allows to assume that at the low phosphorus content oxide clusters of V20~ type predominate over the surface. As for the samples containing only supported vanadium, it should be mentioned that the catalyst prepared using aerosil preheated at 250 ~ and having 0.78 mmol/8 OH groups displays catalytic properties very similar to those of the bulk V205. At the same time, another preparation using aerosil preheated at 800 ~ (0.25 mmol/g OH groulm) catalizcs the benzene selective oxidation to produce about 19 % tool. yield of M& The best catalytic performance was found to be the property of grafted/ impregnated systems having fp = 0.65-0.72 for all the tested reactions. When fp > 0.7 the activity and selectivity sharply decreased due to isolation of vanadium ions and fmlhcr covering them with phosphorus-oxide groups. Study of an the grafted samples by EXAFS [18] showed the presence of peaks attiibutcd to V-O, P-O, V-O and O-O bonds length m p - v o P 0 4 phase. Moreover, some of the lines (0.400, 0.517, 0.650, 0.695 nm) were indicative for the presence of several ordered layers of this phase. Conclusively, it should bc emphasized that at the vanadium and phosphorus g r ~ on the surface of acrosil mainly p-VOPO 4 phase flasmcnts arc folmm8 for the V-P/SiO~. efficient catalysts. At the same time grafted (VO)2P207 phase foxtmng was failed, probably due to its space structmc which cannot bc created on the surface. REFERENCES 1. Vanadyl pyrophosphatr catalysts (Edit. O.Ccnti), Catal.Today, 16 (1993) 1. 2. V.A Zazhigalov, YtLP. Zajtscv, V.M. Bclousov, B. Parlitz, W. Hankc and O. Ohlmann, React. Kinet. Catal. Lett., 32 (1986) 209. 3. KE. Birkcland, H.H. Kung, S.R. Barl, O.W. Coulston, R. Harlow and P.L. L ~ , Hctcrog.Hydrocarbons Oxidation, 21 lth Na~onal Mo~t. Amcr. Chem. Soc., New Orleans, 41 (1996) 197. 4. RA. Overbeek, A.RC.L PekelhaxinS, KJ. van Dillen and ff.W. Geus, Appl. Catal. A, 135 (1996) 231. 5. R & Ovexbeek, P.A Wamnsa, M.J.D. Crombas, L.M. Vmser, AJ. van DBlen and J.W. Gem, Appl. Catal. A, 135 (1996) 209.

796 6. M. Maxtmcz-Laxa, L. Momno-Rr R Pozas-Torm0, A Jmaencz-Lopez, S. Bruque, P. Ruiz and O.Pocelet, Can. L Chem., 70 (1992) 5. 7. M. Nakamura, K. Kawai and Y.Ftowara, L Catal., 34 (1974) 345. 8. RamstcRcr and M. Batrrns, J. Catal., 109 (1988) 303. 9. P.S. Kuo and B.L. Yan8, L Catal., 117 (1989) 301. 10.V.A_ Zazl~alov, J. H ~ , J. Stoch, h~I. Pyatmtskaya, O.A_ Komashko and V.M Br Appl. Catal. A, 96 (1993) 135. 11.V.& Zazt~alov, V.M. Br R Kozaowski and YuLP.Zaitscv, Tcorct. and Expenm Chem., 23 (1987) 650. 12.A. Satsuma, K Hatton, A_ Furuta, A_ Miyamoto, T. Hattori and Y. Murakami, I. Phys. Chem., 92 (1988) 2275. 13.M. Niwa, Y. Matsuoka and Y. Murakami, I. Phys. Chem., 93 (1989) 3660. 14.Z. Roozeboom, M.C. Mittelmeijer-I-Iazaleser, J. Medema, V.H.L de Beer and P.J. Gellings, I. Phys. Chem., 84 (1980) 2783. 15.M. Dercwinsld, I. Haber, R Kozlowski, W.A_ Zazhigalow, J.P. Zajcev, I.W. Bachcrikowa and W.M. Bclousow, Bull. Polish Acad., Sci. Chem., 39 (1991) 403. 16.R. Frickc, H.-G. lcrschkcwitz, O. Lischkc and O.Ohlmatm, Z. anor8, a118. Chem., 448 (1979) 23. 17.& Satsuma, F. Okada, & Hatton, & Miyamoto, T. Hattori and Y. Murakami, Appl. Catal., 72 (1991) 295. 18.V.A_ Zazl~alov, V.M. Belousov, I.V. Bacherikova and Yu.P. Zaitsev, Teoret. and Exl~rim. Chom., 27 (1991) 370.

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

797

Effects of cesium doping on the kinetics and mechanism of the n-butane oxidative dehydrogenation over nickel molybdate catalysts L.M. Madeira* and M.F. Portela GRECAT. Instituto Superior T6cnico - Technical University of Lisbon Av. Rovisco Pais, 1096 Lisboa Codex 0~ortugal)

The kinetics of the ODH of n-butane has been investigated for unpromoted and cesium

promoted tx-Ni]~oO4catalysts. The reaction rates of dehydrogenation products as functions of the butane and oxygen partial pressures are described by a kinetic model based on the Mars and van Krevelen mechanism. The effects of Cs on the kinetic parameters can be interpreted on the basis of recently published results concerning the properties of those catalyst3. 1. INTRODUCTION Nickel molybdates proved to be promising catalysts for n-butane oxidative dehydrogenation (ODH) [ 1,2]. A detailed study about the effects of alkali metal promoters (Li, Na, K and Cs) showed that selectivity for dehydrogenation products (C4's) can be increased by suitable modification of the nickel molybdate catalysts with those promoters, specially with cesium [3]. In this way, the effects of doping them with increasing Cs loadings (1, 3 and 6 %) were carefully investigated [4,5] and a maximum of selectivity for dehydrogenation products was found with 3 % promoter loading. Several kinetic studies have been carried out with lower alkanes but they were usually performed with vanadium and phosphorus containing catalysts. According with Mamedov and Cort6s Corber/m [6], the ODH of lower alkanes on V205 - based catalysts appears to occur by redox mechanism but it is evident that no clear interpretation regarding the alkane activation can be derived from the literature. Other studies with the V-P-O system showed that the oxidation of butane to maleic anhydride occurs by redox reaction kinetically consistent with the Mars and van Krevelen model [7,8]. For this system, the participation of lattice oxygen was demonstrated [9] and when oxygen is consumed in the reaction, the vacancy is filled by gaseous oxygen. Centi et al. [ 10] also found that a redox process is involved in the ODH of nbutane on a V-P mixed oxide. In what concerns to nickel molybdate catalysts, a kinetic study showed that the propane ODH could proceed by redox mechanism [11 ]. For improving our knowledge of the factors determining the kinetics and mechanism of nbutane ODH, and due to the small amount of available information for the Ni-Mo-O system, a systematic kinetic study with unpromoted and 3% Cs promoted NiMoO4 was undertaken. "L.M. Madeira thanks PRAXIS XXI Program from JNICT (Junta Nacional de Investiga~o Cientifica e Tecnol6gica) for financial sup~rt.

798 2. EXPERIMENTAL

2.1. Catalysts preparation and characterization The ct-NiMoO4 catalyst was prepared by eoprecipitation [2] and afterwards doped by wet impregnation with a solution of cesium nitrate The impregnated sample was filtered, dried and finally calcined in air for 2 h at 550 ~ The catalysts were carefully characterized by several techniques such as BET, ICP (inductively coupled plasma spectroscopy), AA (atomic absorption), HTXRD, FTIR, XPS, CO2-TPD, TPR and electric conductivity. Experimental details and results can be found elsewhere [3-5,12].

2.2. Oxygen temperature programmed desorption experiments To investigate the interaction of oxygen with the catalyst surface, several TPD experiments were undertaken. A pretreatment was first performed by heating the sample in He flow up to the desired temperature but avoiding the transition to the metastable 13-phase, which is already present at 595 ~ [13]. After 02 adsorption at different temperatures the sample was cooled to room temperature, purged in He and finally 02 was desorbed by heating (10 ~ in a He flow (1 ml/s) with on line gas analysis performed with the TC detector of a gas chromatograph.

2.3. Catalytic tests Catalytic runs were performed in a fixed bed, continuous flow tubular quartz reactor with a coaxially centered thermocouple. Reactants and products (1-butene, trans-2-butene, cis-2butene, butadiene and carbon oxides) were analyzed with an on line gas chromatograph with two columns [3,4]. The catalyst was mixed with inert quartz (180-254 ttm) in a volume ratio of 1:2 catalyst to quartz. The catalyst charge was chosen to yield differential conversions, in order to minimize the effects of products and secondary reactions. With unpromoted NiMoO4 the catalyst charge was 0.150 g and the contact time 4.45 g.h/(mol butane). With the cesium promoted sample W = 0.300 g and W/F = 5.80 g.h/(mol butane). The feed was a mixture of nbutane, 02 and N2 and for both catalysts, the influence of the reactants partial pressure was studied. The total pressure in the reactor was 1.10 bar and the temperature range studied was 500 - 560 ~ Blank runs proved that under the experimental conditions used the homogeneous reactions can be neglected. The stability of the NiMoO4 activity was previously checked for 50 h and a good stability in butane conversion and products distribution was found [ 14]. 3. RESULTS

3.1. Catalysts characterization Characterization data evidenced that the prepared NiMoO4 is stoichiometric and that Cs is deposited only on the catalyst surface (atomic ratio Cs/Mo = 0.03) not affecting the molybdate structure. However, Cs doping causes a decrease of the catalyst surface area: SBET(NiMoO4) = 44.1 m2/g and SBET(3% Cs-NiMoO4) = 28.7 m2/g. Moreover, the promoted sample exhibits a higher surface basicity, electrical conductivity and also a larger resistance to reduction [4,5,12].

3.2. Oxygen temperature programmed desorption experiments Adsorbed oxygen on a stoiehiometric surface of a oxidized oxide is readily identifiable because it is usually desorbed at temperatures lower than the sublimation temperature of

799 surface lattice oxygen [9]. One method to detect the presence of adsorbed oxygen is by temperature programmed desorption (TPD) after exposing the oxide to 02. While MoO3 does not adsorb oxygen, NiO exhibits an O2-TPD profile with several peaks, corresponding to different forms of adsorbed oxygen, either molecular or atomic [15]. However, we have not found in the literature any reference with respect to nickel molybdates. In this way, several O2-TPD experiments were performed with unpromoted and promoted NiMoO4 with different experimental conditions such as the kind of pretreatment, the way of 02 adsorption or even changing the adsorption temperature because it is well known that for certain species the adsorption is activated [9]. Nevertheless, no adsorption of 02 was detected. Only a slight increase in the TCD response at temperatures higher than 600 ~ was observed which can be due to the sublimation of lattice oxygen. 3.3. Catalytic tests To investigate some aspects of the kinetics of the reaction as well the effects of Cs doping, a systematic study over the unpromoted and 3% Cs doped NiMoO4 catalysts was undertaken for n-butane ODH by changing the temperature and the reactants partial pressures (Pi). However with the unpromoted NiMoO4 it was not possible to operate with butane pressures higher than 0. l0 bar, when Po2 was fixed at 0.05 bar, due to catalyst reduction and/or coke formation. Therefore, some tests were performed with Po2 = 0.15 bar but again, catalyst reduction and/or coke formation was observed at butane pressures higher than 0.15 bar. With the Cs promoted sample we worked with Pb~t~ up to 0.25 bar (with Po2 = 0.05 bar) without significant reduction of the catalyst or coke formation. The influence of the reactants partial pressure in the dehydrogenation products formation rates (based on surface area) are shown in Figure 1 for NiMoO4 and in Figure 2 for 3% Cs - NiMoO4. It is noteworthy that a zero order in O2 is apparent for all C4 products with both catalysts. Moreover, a positive order (smaller than one) in butane is visible with NiMoO4 while an almost linearity is found with the Cs doped catalyst, specially with 2-butenes. In what concerns to products distribution, the effect of temperature is similar with both catalysts and corresponds to an increase in butane conversion: the butenes selectivity decreases as the butadiene selectivity increases and the carbon oxides formation also increases, specially CO. These effects are more pronounced with the Cs doped sample. Butane partial pressure does not affect the products distribution with NiMoO4 but increases the C4's selectivity (specially butenes) decreasing mainly the CO2 formation with the 3% Cs doped catalyst. The effects of increasing Po2 are the same for both catalysts. A decrease of C4"s selectivities and an increase of COx formation at low pressures is mainly observed. It is noteworthy that the main effects of Cs doping in the selectivities are increase in 1-butene selectivity and decrease in CO formation. 4. DISCUSSION It is well known that in oxidation reactions oxygen acts by adsorption on the oxide catalyst as O, Os etc. or incorporation as lattice O 2 species. The solid is oxidized in this step, and the electrons received by adsorbed oxygen could come from reduced surface cations or anionic vacancies with trapped electrons. If oxygen is incorporated as lattice oxygen ion, the sites for adsorption of oxygen and for oxygen attack in the catalytic reaction may be different, and migration of oxide ions in the solid between the two sites would occur [9].

800

~ - -

L ~ L

1.0E~4

0.05 bar

x _~

x

pc~= O05 bet

1.2D04

":

A

pc~=Q15bar

,= ~0E~5

~ S,~~ 8 4

0,0E~ 0,00 7,~E-05

0~1~C10

o,05

,

~

*

t

,

~

,

,

ao5

I

,

~

0,0E~

a~o

P t ~ , (bar)

&0E-0~

F~u~, = (I05 mr

E

~

am

0,10 0,15 F~ (t~)

aoo ~0E-~

P=-0.05bar

ao5

a~o

P ~ , , (bar)

a~5

am

Pcz - 0.15 I:=r

x

o

o

,,,,,,

....

(~oo 7,~-,.05 E ,'-: 5,0E,.05

o,o5

~1o q15 Pc= ~a0

.._

0,~,

q2o

pbuw== gO5 i:=r

x=_..__..___~~

: : : : I : : : : : ~ : o,oo o.o5 O.lO 8m,,(t=0

~0~05

P o ~ " 0.05 bar

9,0E~5 -

~

~

0,05

1,6E-04 --'-o 8,0E-05

0,0E+E)

~ A ~

I

0,05

.

I

t

0,10 0,15 Pc~ ~a~

: : ; : ; : : : ; ; : : 0,00 0,05 0,10

0,0E-Kg) 0,00

P = m (ti)

..-.

.

Pc~ = O 15 bar

~0E~4-

P~-O.05bar

x

~0~04 T

/

9

c~ 1,(E-04 "O

I

0,0EH~,

0,20

Pmana" G05 bm

o

=-.

0,0E*(I),

0,20

pha~=0.05ber

0,00

~0E~04

0,10 0,15 Pc~ 0~a0

0,00

"O i,. ::

~5E~04-

qo5

O,lO

I---r-=---1----~

=.. O,0E+(]0

0,00

I

0,05

I

0,10 Paz~

I

0,15

I

0,20

0,0EN:}0/

0,00

po==0.05ber

6,QE-04

E

E

J=..

J=_ (~0E-t{]~

0,00

~

I

t

I

I

I

m

i

~

I

o,o5 qlO Pt,am (bar)

t

I

o,2o

Po~-(1151:ar

t

I

0,(IS+{]0

o,oo

t

0,05 0,10 0,15 Ra=m ~a~

x

I

0,05

I

I

A - 540 ~

x - 560 ~

I

qlO o,15 o,20 Pu=~ Oa~r)

Figure 1. Influence of reactants partial pressure on the dehydrogenation products formation rates over NiMoO4 (e - 500 ~ 9 520 ~ cq. (7) with the data presented in Figure 3.

I

0,21)

po~=Q15b ~

E ~ 4,0E-04

....

o,15

Pt,,-,,(bar)

+

::;~: : : ; : '. 0,05 0,10 Raam~

.=

o= ZOE.O4

o.2o

x

r~~6'0E~5 ~ 3,~--05 ~ -

=.. 0,00

o.o5 O.lO o . 1 5 P=m (t~

E

x ,,E~ ~ 40,E 0-5,

~OE+~

0,0E+G) o.oo

Full lines obtained from

801

8,0E-05

Pbutane

T

/

N

x

x

~ 4,0E-05

Po2 = 0.05 bar

= 0.05 bar .

~"

___JK

o E

m

9

2,0E-04

1,0E-04

T=...

:

1

0,0E+00 0,00

! i 0,05 0,10 Po2 (bar)

4,0E-05 -1-

P~tane = 0.05 bar

2,0E-05

~

I

9

~,

:

I

5,0E-05

~~

t I 0,05 0,10 Po2 (bar)

T

Pbut=~ = 0.05 bar

t,

x

x

1,0E-04

T

[

0,0E+00 0,00

I 0,15

2,0E-04 x

9

I 0,05 0,10 P02 (bar)

x

=j9

,

T

"S

.

.-. 1,0E-04

! ~

t

t 0,0E+00 l 0,00

I

0,05 0,10 Po2 (bar)

X

9

9

9

0,15

0,20

0,25

0,20

0,25

0,20

0,25

0,20

0,25

Po2 = 0.05 bar

T

0,0E+00 0,00

0,05

3,0E-04

Po2 = 0.05 bar

0,10 0,15 Pbut=~ (bar)

g to 1,0E-04

.~

O,OE§

i

0,15

1,0E-03 T x

9

9

9

I 0,05

0,00

&

_ -

t 0,10 P 0 2 (bar)

0,05

0,10

0,15

Pbun~ (bar)

Pbut=~ = 0.05 bar x

0,10

Pb.t~,, (bar)

~ ~,o~-0,

9

0,00

0,05

~

K"

_=

,,-,, 2,0E'--04 ~

0,15

x

O,OF+I~

0,25

9

Pbutane= 0.05 bar

~

E 5,0E-05 o9

0,20

Po2 = 0.05 bar

T

, O,IXI

0,15

.,~~,OE-O~ = l

~

0,0E+O0 i

0,10

~ 1,0E-04 J

9

i

0,05

Pbutane (bar)

.... 1,5E-04

_=

0,0E+00 0,00

0,0E+00 0,00

I

0,15

X

9

9

9

4 0,15

i|

~'~ Z

Po2 = 0.1)5 bar

o

E 5,0l::-04

~"

" d

I i

0,0E+O0 0,00

0,05

0,10 0,15 P~t=~ (bar)

Figure 2. Influence of reactants partial pressure on the dehydrogenation products formation rates over 3% Cs - NiMoO4 (symbols and legend as in Figure 1). Approaching from a different point of view, Bielanski and Haber analyzed the question of selectivity by considering the types of oxygen available [ 16]. They propose that the reactivity of oxygen depends on whether it is electrophilic or nucleophilic. Electrophilie oxygen species,

802 such as adsorbed O2 and O" are very active and attack hydrocarbon molecules at the regions of high electron densities. Nucleophilic oxygen species, such as lattice oxygen are less reactive and are suitable for partial oxidations. In a reaction with nucleophilie oxygen, activation of the hydrocarbon molecule is the rate-determining step. In a reaction with electrophilic oxygen, adsorption of oxygen to form the electrophilie species is the rate-limiting step. Therefore, a selective oxidation catalyst should be able to adsorb and activate a hydrocarbon molecule for nucleophilic attack by oxygen. It should not generate electrophilic oxygen species. For the NiO-MoO3 system it was found that lattice oxygen is responsible for ODH, while other products of selective and complete oxidation are formed by other species of surface oxygen [ 17]. Moreover, Mazzoeehia et al. [13] found by electrical conductivity measurements the presence of anionic vacancies at the surface lattice of nickel molybdates. They suggest that in propane ODH, propene would be produced via formation of such vacancies resulting from reaction of 0 2- surface anions. The role of gaseous oxygen would be to replenish rapidly the lattice oxygen consumed, reoxidizing the solid. Finally, it must be remarked that gaseous oxygen may have another role: the formation of carbon oxides by attacking the weakened C-C bonds of the adsorbed butane molecules. The kinetic studies found in literature (of. Introduction) suggest that this reaction proceeds via a Mars and van Krevelen mechanism [ 18] and then the following steps are involved: (i) reaction of the hydrocarbon with the oxide to give products and a partially reduced catalyst; (ii) reoxidation of the catalyst by gaseous oxygen to restore the catalyst in its original state. In this mechanism the agents responsible for oxidation are the oxygen ions of the oxide lattice. Therefore, the catalytic behavior of oxide catalysts towards hydrocarbon oxidation should depend on the strength of the metal-oxygen bond of the catalyst [19]. In a previous paper it was found that the catalytic activity for butane ODH over various Cs doped NiMoO4 samples (which decreases when increasing Cs contents) is related to the catalyst reducibility [12]. In fact, Cs promotion increases resistance to reduction, as evidenced by the increase in the temperature of onset of reduction with the Cs loading. This relationship between catalytic activity and reducibility of the catalysts (which gives an idea of the metal-oxygen bond strength) suggests therefore also the existence of a redox mechanism The involvement of lattice oxygen in the reaction is also supported by the fact that no oxygen desorption was observed by O2-TPD and by experiments without oxygen in the reactor feed [20]. In fact, even in the absence of gas phase oxygen, butane can be converted to C4's with a high selectivity, in spite of the low conversion values. The decrease of activity would be due to the decrease of the fraction of active oxidized surface. Also the apparent activation energies for butane consumption for both processes (with and without oxygen in the gas phase) are similar, being the differences due to the different products distribution [21 ]. Taking into account all the above mentioned results, a generalized Mars and van Krevelen model was considered for butane ODH over both catalysts. Consequently, the two steps mentioned above can be generally described by the following redox scheme: butane + OC RC + 02

r, products + RC > OC

where OC and RC are oxidized and reduced active sites, respectively. The reduction and reoxidation rates of catalyst surface earl therefore be expressed as: rr= k ~ P ~ e o

(1)

803

ro = ko Po2"1 p ~ , . 2 Or

(2)

where k~ and ko are kinetic constants, Oo and O~ the fractions of oxidized and reduced sites, respectively, Pi the partial pressure of reactants and ni the partial orders. In steady-state, rr= ro/Ot

(3)

ct is the stoichiometric coefficient of oxygen, i.e., the number of 02 moles reacting per mole of butane. For butenes a = 0.5; for butadiene a = 1; and for the total of C4"s, a is a function of the products distribution. From eqs. (1), (2) and (3) 0o and the global rate can be obtained:

Oo = ko Po2 nl / ( ko P02 nl + o~ kr Pb,aane(n'n2))

(4)

r = ko kr Pbutanen POE"1 / ( ko Po2"1 + ot l~ Pb~.r ("-"2))

(5)

For estimation of the parameters of eq. (5) including the partial orders, the equation was fitted to the experimental rates of each dehydrogenation product by optimization of a multiple correlation coefficient W [20] using non-linear regression analysis. It is noteworthy that for both catalysts, the equation rates that best describe the formation of all dehydrogenation products are all of the same type with n = 1 and nl = n2 = 0. The influence of butane partial pressure in eq. (2) was investigated because a negative effect was previously found for 1butene ODH over bismuth molybdates due to a competition of oxygen and 1-butene over the reduced site [22]. Nevertheless, such order was found to be null in this case. An interesting feature is the fact that Po2 does not affect the oxidation rate of the reduced active sites, i.e., nl = 0. This may evidence that the diffusion of lattice oxygen controls the global process of the catalyst reoxidation. A similar behavior was found for 1-butene ODH over bismuth molybdate at lower temperatures [22]. Therefore, eqs. (4) and (5) can be simplified to 0o = ko / (ko + ct k~ P ~ t ~ )

(6)

r = ko kr Pbutane / ( ko + o~ kr Pbutane )

(7)

The other computed parameters (ko and IQ are presented in Figure 3 as Arrhenius plots. The fitting curves obtained with the computed parameters are included in Figures 1 and 2. The computed values of 1%and k, agree with the Arrhenius law: k i - ki'x exp (- Ea~ / (RT))

(8)

which is additional evidence in favor of eq. (7). From the Arrhenius plots (Figure 3) the preexponential factors k( and the activation energies Eai were computed (Table 1). For the unpromoted NiMoO4 the activation energies for reduction are always higher than Eao. The negative activation energies computed for butenes evidence complexities not taken into account. For Cs doped catalyst activation energies are larger for the reoxidation process, except for 2-butenes. In this case, the difference in the kinetic constants is so high that ko >> ct k~ P ~ and eq. (7) is simplified to a linear variation of the reaction rate with the butane partial

804 1- b u t e n e

e-

1-butene

-7

-9

-11 1,18E-03

-5

~ 1,22E-03 1,26E-03 1 /T (K1)

-11 1,18E-03

I

1,30E-03

trans -2-butene

-7

10

_-

-5

-5

B

-2-butene

tra n s

_-

t ~ 1,22E-03 1,26E-03 1 / T (K "1)

0

J i 1,30E-03

c/s -2-butene

-10 1,18E-03

A

t

I

1,22E-03 1,26E-03 1 / T (K"1)

cis

A

B

-2-butene

t

-7

~

~

I

I

1,22E-03

1,26E-03

1,30E-03

B

~

1 / T (K "~) -3 -5

C4"s

A

I

~

1,22E-03 1,26E-03 1 / T (K "~)

-4

i

1,30E-03

B

C4"s

~ " " ~ 6 _ ~ . . . . . ~ c

-9 1,18E-03

~

-11

1,18E-03

1,30E-03

butadlene

-9 -9

i

1,22E-03 1,26E-03 1 / T (K "1)

-5

.to _r -7

1,18E-03

1,30E-03

0

-10 1,18E-03

I

1,30E-03

butadlene

J 1,22E-03 1,26E-03 1 / T (K 1)

10

_c

-11 1,18E-03

i 1,30E-03

~

,-

-11 1,18E-03

t p 1,22E-03 1,26E-03 1 I T (K "~)

~ I 1,22E-03 1,26E-03 1 I T (K~)

I 1,30E-03

-6

-8

1,18E-03

I

t

1,22E-03 1,26E-03 1 / T (K"I)

~

1,30E-03

Figure 3. Arrhenius plots of the kinetic constants ko (B) and kr (0) for dehydrogenation products over NiMoO4 (A) and 3% Cs - NiMoO4 (B). pressure (r = k~ P ~ ) what is in very good agreement with the experimental data, specially at higher temperatures. In what concerns the other products with the 3% Cs promoted catalyst,

805 Eao > Ear what means that in eq. (7) ko increases stronger with temperature than ~t kr Pb~t~c. In this way, at a temperature of 560 ~ it is necessary a higher butane partial pressure than at 500 ~ in order to run away from the linear variation of r vs. Pbm,o (of. Figure 2). Table 1 Pre-exponential factors and activation energies for formation of dehydrogenation products Catalyst Const? 1-butene t-2-butene c-2-butene butadiene C4"s ko' 2.5• -5 6.1• -7 3.6x10 ~ 6.0 9.6• .3 NiMoO4 kr' 8.8• 5 7.9• 2 1.0• 4 2.4x10 z 3.0• 3 Eao -4.1• 3 -3.0• 4 -1.7• 4 6.7• 4 2.2• 4 Ear 1.3• 5 9.0• 4 1.1• 5 7.8• 4 8.7• 4 ko" 1.8• 1.1• 1.2• 6.0x10 s 5.4x107 3% Cs - NiMoO4 kr' 1.4 8.9 12.8 33.3 22.8 Eao 1.4• 3.4x104 3.4• 1.9x105 1.6xl05 Ear 4.9x104 6.7x104 6.8x104 7.0x104 6.0• a ko" in mol / (h.m2); k~" in mol / (h.m2.bar) and activation energies in J/mol. 5. CONCLUSIONS The negative activation energies obtained with fittings for butenes with the unpromoted ctNiMoO4 catalyst (Table 1) indicate that their formation involves more complex processes than the admitted ones. In fact, in a previous study with the 3% Cs promoted sample [20], when feeding the reactor with 1-butene, selectivities for butadiene around 40% were found while the selectivities of formation of 2-butenes do not exceed 7%. When cis-2-butene is fed, the butadiene selectivity reaches 60% and the selectivity to isomers is around 8-9%. These results point out that isomerization reactions can be neglected and that the subsequent dehydrogenation of each butene to butadiene proceeds in a considerable extension. For undoped catalyst due to its higher acidity, both subsequent dehydrogenation to butadiene and isomerization reactions play a more significant role. However, it must also be taken into account the possible direct formation of butadiene from butane. On the other hand, it is well known that in this ODH reaction n-butane first reacts by cleavage of a secondary C-H bond to form an adsorbed alkyl species. This step is generally accepted to be the rate-determining step. After a second H-abstraction, olefinic intermediates are formed and desorbed as olefins or undergo another dehydrogenation to form butadiene [1,6,7]. Therefore, it is convenient to consider the global process and to use the rate of C4"s formation as a comparative term. It was found that for the dehydrogenation products, and in contrast to expectations, cesium doping decreases the activation energy for the reduction step increasing the corresponding to oxidation (Table 1). Nevertheless, the Cs effect is more pronounced in the pre-exponential factors. In fact the doped catalyst exhibits smaller k~' values and larger ko" than the undoped one. The increase in ko" is so strong that higher kinetic constants for reoxidation are obtained with 3% Cs - NiMoO4 and higher values of kr with NiMoO4 (Figure 3). The increase of the nickel molybdate resistance to reduction after Cs addition explains the smaller k~ values for the doped catalyst. On the other hand, Cs promoted catalysts are much more conducting than unpromoted NiMoO4 due to the contribution of a surface ionic conductivity by mobile Cs + and 02- ions to the overall conductivity [5]. Moreover, solids with high oxygen-

806 ion conductivity have a high capability of transforming any surface oxygen species into lattice oxygen [5]. Consequently, Cs doping favors the reoxidation of the solid because it favors both the oxygen incorporation into the lattice and the diffusion of those species through the solid. Finally it must be pointed out that by using eq. (6), higher 0o values for the promoted catalyst are obtained what means that under any reaction conditions Cs maintains the catalyst in a higher oxidation state allowing the use of higher butane concentrations. REFERENCES

1. H.H. Kung, in D.D. Eley, H. Pines and W.O. Haag (eds.), Advances in Catalysis, Vol. 40, Academic Press, New York, 1994, pp. 1-38. 2. C. Mazzocchia, R. Del Rosso and P. Centola, An. Quim., 79 (1983) 108. 3. R.M. Martin-Aranda, M.F. Portela, L.M. Madeira, F. Freire and M.M. Oliveira, Appl. Catal. A, 127 (1995) 201. 4. F.J Maldonado-H6dar, L.M. Madeira, M.F. Portela, R.M. Martin-Aranda and F. Freire, J. Mol. Catal. A, 111 (1996) 313. 5. L.M. Madeira, J.M. Herrmmm, F.G. Freire, M.F. Portela and F.J. Maldonado, Appl. Catal. A, in press. 6. E.A. Mamedov and V.C. Corberfin, Appl. Catal. A, 127 (1995) 1. 7. M.A. Pepera, J.L. Callahan, M.J. Desmond, E.C. Milberger, P.R. Blum and N.J. Bremer, J. Am. Chem. Sot., 107 (1985) 4883. 8. A. Escardino, C. Sol/Land F. Ruiz, An. Quim., 69 (1973) 385. 9. H.H. Kung, in "Transition Metal Oxides: Surface Chemistry and Catalysis", Studies in Surface Science and Catalysis, Vol. 45, Elsevier, Amsterdam, 1989. 10. G. Centi, G. Fornasari and F. Trifir6, J. Catal., 89 (1984) 44. 11. A. Kaddouri, R. Anouchinsky, C. Mazzocchia, M. Madeira and M.F. Portela, Catal. Today, submitted for publication. 12. L.M. Madeira, M.F. Portela, C. Mazzocchia, A. Kaddouri and R. Anouchinsky, Catal. Today, submitted for publication. 13. C. Mazzocchia, C. Aboumrad, C. Diagne, E. Tempesti, J.M. Herrmann and G. Thomas, Catal. Lett., 10 (1991) 181. 14. F.J.M. H6dar, L.M.P. Madeira and M.F. Portela, J. Catal., 164 (1996) 399. 15. M. Iwamoto, Y. Yoda, N. Yamaoe and T. Seiyama, J. Phys. Chem., 82 (1978) 2564. 16. A. Bielanski and J. Haber, in "Oxygen in Catalysis", Marcel Dekker, Inc., New York, 1991. 17. C. Mazzocchia, R. Del Rosso and P. Centola, Proc. 5th Ibero-American Symp. Catal., 1976, Rev. Port. Quim., 19(1-4) (1977) 61. 18. P. Mars and D.W. van Krevelen, Chem. Eng. Sci. (Spec. Suppl.), 3 (1954) 41. 19. D.J. Hucknall, in "Selective Oxidation of Hydrocarbons", Academic Press, London, 1974. 20. L.M. Madeira, F.J.M. H6dar, M.F. Portela, F. Freire, R.M.M. Aranda and M. Oliveira, Appl. Catal. A, 135 (1996) 137. 21. F.J.M. H6dar, L.M. Madeira, M.F. Portela and R.M.M. Aranda, Proc. 15th IberoAmerican Symp. Catal., (1996) 251. 22. M.F. Portela, M.M. Oliveira, M.J. Pires, F.M.S. Lemos and L. Ferreira, Proc. 8th International Congress on Catalalysis, (1984) 533.

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

807

A COMPARISON OF IRON MOLYBDATE CATALYSTS FOR METHANOL OXIDATION PREPARED BY COPREClPTATION AND NEW SOL-GEL METHOD A. P.

Vieira Soaresa, M. Farinha Portela~9and A.Kiennemannb

aC~ECAT- Gn~ de ~ de Cmtdise~ UniversidadeT~-'nicade Lidx~ ~ Rm4.so3Pais, 1096Lidx~ Cxxt~ Portugal,Fax:351-1-847"/695

SulxriorT6cmco,Avenida

%CPM - D6partement Catalyse, LERCSI, Universite Louis Pasteur, 1 Rue Blaise Pascal, BP296, F 67008 Strasbourg Cedex, France Iron molybdates, with atomic ratio Mo/Fe=3, were prepared by coprecipitation and sol-gel techniques in acid medium. Characterisation results show that sol-gel catalysts have much higher surface areas than coprecipitated catalysts. Study of catalytic activity shows that Fedefective catalysts, prepared by sol-gel technique, perform better than an industrial catalyst in the same conditions. In fact these catalysts achieve higher performances at 50K lower than industrial catalysts. 1. I N T R O D U C T I O N Formaldehyde is nowadays one of the major produced chemicals due to its uses in many fields of chemical industry [ 1]. The commercial production started in 1890 using metallic copper catalysts. In 1910 copper catalysts were replaced by silver catalysts with higher yields [2]. Although the first report of the excellent catalytic behavior of iron molybdates in selective oxidation of methanol to formaldehyde is of 1931, the related industrial process based on them only went into operation in 1940-50 [1]. A recent report [3] shows that iron molybdates and silver catalysts are nowadays equally used as industrial catalysts for formaldehyde production. Many metal molybdates catalyse the reaction under consideration [4] and the active sites are widely believed to be associated with surface Mo atoms in octahedral coordination [5]. Octahedral coordination of Mo is only achieved in Fe-defective iron molybdates what is in accordance with the fact that maximum activity is obtained with catalysts with Mo/Fe atomic ratio greater than 1.5 [6]. The presence of two terminal oxygens double bonded to Mo in such coordination allows the methanol reacting molecules to be bonded simultaneously by two points. The activation of the hydrogen of the hydroxyl group produces methoxy species that are intermediates in the formaldehyde formation. The role of Fe in iron molybdates catalysts would be to act in the transfer of 02 and H20 between surface and gas phase [6] and to hinder the reduction of Mo +6 [7]. Several studies of the catalytic behavior of iron molybdates as a function of their physical and chemical properties show that surface acidity, related with M0 +6 ions, is a necessary condition for the catalyst effectiveness for formaldehyde formation [8,9]. The oxidizing function is also necessary, though the activity is not necessarily determined by this function. In a recently study Sun-Kuo et al [6] demonstrated that catalytic behavior of iron

9To whom correspondence should be addressed

808 molybdates is closely related to atomic ratio Mo/Fe of the catalyst. They found the maximum activity for a catalyst with an atomic ratio Mo/Fe=l.7. This result agrees with the work of Kolovertnov et al 30 years ago [10]. Some researchers claim that the active phase is normal iron molybdate but recognise that an excess of Mo is required to compensate its loss during reaction. Industrial catalysts are always Fe-defective molybdates due to this reason. Up to now several methods have been used to prepare iron molybdates, the most part of them based on coprecipitation techniques. Previous studies [11 ] have evidenced that the catalytic behavior of Mo-Fe oxides depends on many variables of the coprecipitation procedure: starting compounds, concentration of parent solutions, pH and temperature of coprecipitation step, order of addition of parent solutions, ripening etc. In a typical preparation procedure iron molybdate is coprecipitated from solutions of ferric chloride or ferric nitrate and ammonium molybdate [8]. The control of all the above mentioned procedure variables, strongly difficult the preparation of these catalysts and deviations from the preparation recipe can have very adverse effects on the performances of the catalyst from the standpoint of activity, selectivity and stability. Recent sol-gel methods have been recognized as promising procedures to prepare catalysts [12-14]. The sol-gel methods allows a unique way of catalyst design, because they represent an ab initio synthesis of the final solid from well defined molecular compounds [ 13]. By suitable choice of reagents, reaction and drying conditions, such technique allows to predefine pore structure, porosity, composition, surface polarity and crystallinity or amorphicity of metal oxides [12]. In principle, any metal that forms stable oxides can be forced to copolymerise with other metals in sol-gel procedures to provide mixed metal oxides [ 13]. To investigate of influence of preparation methods on catalytic properties of iron molybdates catalysts with controlled excess of MoO3 were prepared by coprecipitation and solgel techniques. Their properties and performances were compared with an industrial catalyst. 2. EXPERIMENTAL

2.1. Catalyst Preparation Iron molybdates with atomic ratio Mo/Fe = 3 were prepared by coprecipitation and solgel techniques. Coprecipitated catalysts were obtained from starting aqueous solution of Fe(NO3)3.9H20 and (NH4)eMo7024.Iron nitrate solution was slowly added to the solution of Mo previously acidified (pH ~2) with HNO3. After the addition of Fe solution the precipitates were ripened in contact with mother liquor at 100~ for 3 hours under vigorous stirring. Afterwards the precipitates were filtrated, dried at 120 ~ overnight and finally calcined. Sol-gel catalysts were prepared in acid medium using appropriate molybdenum precursor and Fe(NO3)3.9H20 or FeCI3. In a typical procedure, first the precursor solutions of Mo and Fe in acid medium were prepared and then ferric solution was added to Mo solution. No precipitate was observed during the addition. The resulting solution was evaporated until dryness. The formed film was removed by adding liquid N2, then was crushed in a glass mortar and dried at 120~ overnight. Calcination, for both techniques, was performed at 648 K during 10 hours. For coprecipitated catalysts calcination was always under air flow, whereas sol-gel catalysts were calcined with or without air flow.

809

2.2. Catalyst Characterisation The BET surface areas were measured using a Perkin-Elmer Shell 212C sorptometer instrument based on the N2 physisorption capacity at 77K. Bulk elemental composition were determined by atomic absorption. X-Ray diffraction patterns were obtained with a D5000 Simens diffi'actometer equipped with a primary beam quartz monochromator (Co Kczl=l.78897A~) at 40kV and 25 mA. The morphology, chemical analysis and homogeneity of the prepared catalysts were examined with a scanning electron microscope (SEM) JOEL JSM840 equipped with a Delta Kevex energy dispersive X-ray analyzer. FT-IR spectra were recorded with a Perkin-Elmer 1600 spectrometer with a range of 4000 - 400 cm-1, and a resolution of 2 cm-1. X-ray photoelectron spectroscopy (XPS) was used for chemical analysis and investigation of reduction state at the catalysts surface. The analyses were performed on a VG ESCA 3 apparatus. The kinetic energy of the emitted photoelectron is given by: E~=1486.6-EBcor, were 1486.6 is the energy of the incident radiation (AI Kcz ray] and Emo~ the corrected binding energy of the electron. The binding energy were calibrated with respect to the signal for adventitious carbon (biding energy: 284.8 eV). Iron phases bulk composition was studied by MOssbauer spectroscopy at room temperature operating in the constant acceleration mode and with a radiation source of 57Co in a Rh matrix. 2.3. Catalytic Tests Methanol oxidation was carried out in a conventional flow apparatus at atmospheric pressure. The feed mixtures were prepared by injecting the liquid methanol into air flow with a Gilson 302 pump. The catalyst was diluted with inert carborundum (1:3 volume ratio) to avoid adverse thermal effects, and placed in a tubular pyrex reactor with a coaxially centred thermowell with thermocouple. The reactor outlet was kept at 403 K, to prevent condensation of liquid products and formaldehyde polymerization, and it was connected with multicolumn Shimadzu GC-SA gas chromatograph with thermal conductivity detector. The column system used (1.5m of Poropak N+l.5m of Poropak T+0.9m of Poropak R) could separate CO2, formaldehyde, dimethylether, water, methylformate, dimethoxymethane and formic acid. The last product was never detected. 3. RESULTS AND DISCUSSION

3.1. Catalyst characterisation BET results in figure 1 show that sol-gel technique yields catalysts with surface areas that are approximately twice of coprecipitated ones. This may be due to formation of a second amorphous phase of MoO3 in sol-gel catalysts [6] instead occupation of lattice interstices, by Mo excess, which occurs in coprecipitated catalysts [ 16]. Industrial catalysts present a lower surface area than coprecipitated catalyst with the same atomic ratio Mo/Fe (=3) which can be attributed to the severe calcination step. The scanning electron micrographs of the catalysts (figure 2) show that catalysts prepared by both methods have the same sponge-like morphology. This result agrees with the fact that Mo excess retards the crystallization of Fe2(MoO4)3 [ 17]. Morphological aspect and surface areas are in good agreement. However our results are in disagreement with recent results of Sun-Kuo et al [6]. These authors have found that highest surfaces areas corresponded to samples formed by associated ordered lamellae. After long activity tests, including catalytic tests with water in reactor feed, catalysts kept their morphological aspect.

810

L

i I

i I

i I

i I

: ::- -:..:. ~_:::-::, ~ :.. :::.--~i :~ :- i~ --:-_-:--.~ i i i

i i i

i i i

i I

i I

i I

_:- :.~: ~.~.- - :_ .-i-~-._-:-.-.~-:~

i [ i

i:ii~::::-!:~: :i-i-~.~:::.~:-"~"-: ~::-.............L3

0

I

I

2

4

I

6

8

10

12

14

16

18

surface area (m~/g)

Figure 1 - Surface area of catalysts prepared by coprecipitation and sol-gel techniques (from no halogenated precursors), and industrial catalyst. a)

b)

Figure 2 - Scanning electron micrographs of flesh catalysts a)Mo/Fe=3 coprecipitated, b)Mo/Fe=3 sol-gel. Bulk elemental analysis of some catalysts was performed by atomic absorption. The results presented in table 1 show that coprecipitation method yields, within the experimental errors, catalysts with the Mo/Fe atomic ratio of preparation. Catalysts prepared by sol-gel method showed lower Mo/Fe atomic ratios than the expected value. Furthermore calcination of these catalysts under air flow enhances the loss of Mo. Table 1 Bulk elemental composition by atomic absorption Catalyst

M o / F e (atomic ratio)

Coprecipitated

2,9

Sol-Gel (calcined without air flow)

2,4

Sol-Gel (calcinated with air flow)

1,9

Industrial

3,1

811 In figure 3 some representative X-ray patterns of the fresh catalysts are presented. All catalysts look crystalline but coprecipitated catalyst seems to be slightly better crystallized than sol-gel one. They present the diffraction lines corresponding to normal iron molybdate and MoOs. Some diffraction lines of these two phases are superimposed. Furthermore catalysts with a large excess of Mo present x-ray lines that cannot be assigned neither to normal iron molybdate nor to MoOs. According to Abaulina et al [16] those lines (6.9, 3.4, 2.3) correspond to another phase, similar to natural ferrimolybdate, formed by MoOs dissolved in iron molybdate. Catalysts prepared by sol-gel method the using halogenated precursors have the same x-ray patterns. Calcination under air flow has no effect on the phase distribution yielded by this technique. Phase composition and degree of the reduction of catalysts were also examined by MOssbauer spectroscopy at room temperature. In table 2 parameters for flesh catalysts are presented. The results show that Fe-defective coprecipitated catalyst has only one type of iron(III) molybdate which correspond to normal iron molybdate. Fresh sol-gel catalyst with Mo excess have a small amount (3.8%) of reduced phase FeMoO4 which is not detected by XRD.

0

I 15

20

! 45

! 60

Figure 3 - X-ray diffraction patterns of fresh catalysts MOssbauer analysis after long catalytic runs evidences that Mo/Fe=3 catalyst prepared by coprecipitation undergo reduction; after 72h of reaction this catalyst had 2% of FeMoO4. Solgel catalyst lose Mo during long activity runs and show MOssbauer characteristic parameters of" Fe203 phase (5%). The presence of FeMoO4 supports the deactivation mechanism proposed by Pemicone et al [19]: Fe2(MoO4)3+CH3OH ~ 2FeMoO4+MoO3+HCHO+H20 All IR spectra in figure 4 show a strong absorption band centered at 830 cml what is associated with stretching of Mo=O in tetrahedral environment [6,16,19]. The shoulder located at 780 cm-1 can be ascribed to stretching of Mo-O bond of heteropolyanions of Mo with octahedral coordination [ 19]. The weak band at 960cm "l is assigned to Fe-O-Mo vibrations in the ferric molybdate phase [19]. The broad band at 610 cm-1 displayed by Mo/Fe=3 coprecipitated, sol-gel and industrial catalyst is assigned to Mo=O in octahedral environment of Fe-defective iron molybdates [ 16]. The intensities of this band and of the band at 995cm l

812 are proportional to the Mo excess of the catalysts and they tend to disappear with time on stream: the industrial catalyst had almost the same composition than Mo/Fe=3 coprecipitated catalysts, but the intensities of those two bands were smaller. This strange result is possibly due to the fact that the industrial catalyst is 7 years old. Table 2 Mrssbauer parameters of fresh catalysts (~i-isomeric displacement, W-band width, Aquadripolar separation). CATALYST Fe species ~i W A I/Io (%) Phase mln.s -1

mln.s -1

m m . s -~

Industrial

Fe 3+

0,41

0,32

0,21

1O0

Fe2(Mo04)3

Mo/Fe=3

Fe 3+ Fe 2§

0,44 0,89

0,26 0,28

O,19 2,34

99 300~ and high level of conversion (>90%) small amounts of CO2 were also observed. Table 3 Atomic ratios and degree of reduction at surface for fresh and used catalysts by XPS (preparation atomic ratio Mo/Fe=3). CATALYST

Coprecipitated

fresh after 72 h of reaction after 24 h of reaction with water

Mo/Fe 2,35 2,60 2,43

Fe2+/Fe3+ 0,19 0,41 0,46

Sol-gel

fresh after 72 h of reaction after 24 h of reaction with water

2,17 2,18 2,43

0,60 0,58 0,72

Catalytic runs for testing the effect of preparation method on catalytic behavior were performed at 573 K with 4% of methanol (in air) in reactor feed. Results (figure 5) show that among tested catalysts, including an industrial catalyst, the more active and selective was the sol-gel catalyst. Comparing the behavior of this catalyst with the industrial catalyst, at several temperatures (figure 6) it is noteworthy that the sol-gel catalyst performs better than industrial catalyst, even at temperatures 50K lower. The calculation of the specific activities, of formation of formaldehyde and CO, from yield versus contact time data provided the results presented in table 4. It is seen that the activity per square meter of surface are for both products, exhibited by sol-gel catalyst is lower than the specific activity of coprecipitated catalyst. Such differences of activity are attributable to the lower Mo contents of sol-gel catalysts (table 1). 100

75

[]Conversion [] Sel CH3OCH3 9 Sel HCOOCH3 [] Sel CO

% 50

25 DII~IUgI~| Igllgll~ll~ll~l 9~ ! ~ !

.

.

.

.

.

.

.

.

i

Mo/Fe=3 coprec.

.

.

.

.

.

.

.

.

.

.

.

.

.

i

Mo/Fe=3 sol-gel no halogenated precursors

i

Industrial

Figure 5- Conversion and selectivities of catalysts prepared by coprecipitation and sol-gel (without halogenated precursors).

814 Results in figure 7 evidence that calcination of sol-gel catalysts under air flow leads to a strong loss of activity, attributable to loss of volatile species of Mo during calcination step. The use of halogenated precursors in preparation of sol-gel catalysts occasions loss of activity, eventually due to the formation of halogenated species on catalysts surface that are inactive or less active. Unfortunately it was not possible to identify such hypothetical species. Table 4 Specific activities (T=598 K) of formation of HCHO and CO reaction mixture

specific activity Otmol mons-tm~ HCHO CO 2,32 0,46 2,48 0,54 1,65 0,36 1,14 0,29 0,74 0,24

0a'a)

Coprecipitated

Sol-Gel

MeOH 0,10 0,18 0,54 0,93 1,30

02 17,74 17,17 18,00 18,25 18,44

H20 3,34 3,73 4,23 4,69 5,02

HCHO 2,33 2,55 2,91 3,27 3,26

0,55 0,60 0,73 0,86

17,70 17,69 17,84 17,94

4,89 4,89 4,57 4,38

2,13 2,14 2,23 2,17

100

0,11 0,49 0,77 0,98

0,25 0,33 0,39 0,43

m

...i -.

_

_

_

~HG'HO

%50

C A - 1 3 O C H 3

~

HO:X)GH3

~CO

m

Industrial T--548K

Industrial T=573K

Imlumial T=623K

Mo/Fc=3

Mo/Fc=3

T---54gK

T=573K

sol-gcl

sol-gel

Figure 6- Conversion and selectivities of industrial and sol-gel catalyst at several temperatures. Runs for testing stability (figure 8) showed that prepared catalysts were stable in used operating conditions. Furthermore, water do not inhibit the reaction and have a benefit effect on the selectivity. Comparing the results of stability and the degree of surface reduction (Fe2+/Fe3+ in table 2) we conclude that there is no correlation between the two set of results and the apparent high degree of surface reduction is probably due to the high vacuum conditions during XPS measurements. 4. CONCLUSIONS Sol-gel techniques provide interesting routes to prepare iron molybdates catalysts for selective oxidation of methanol. Catalysts prepared by this method have higher surface areas than catalysts prepared by coprecipitation techniques what allows to operate at lower temperatures with the advantage of limiting the consecutive oxidation of formaldehyde to CO.

815 In this type of catalysts Mo excess form an amorphous separate phase that provides large contribution to the total surface area, thus increasing the number of methanol adsorption sites. The precursors used in sol-gel methods seem to influence the catalytic properties of the catalyst. In particular halogenated precursors have a beneficial effect on the catalyst selectivity. 1 0 0

_

_

_

_

_

Q

~

9

o/e 50

-

Sel

r

-

-

_

~

-

-

'

9

l

HCHO S

Set

e

l

~

3

HC.OOCH3

CO

IDSel

I

no halog,

no halo~

halog,

precursors

precursors

preempts

preem~ors

halo~

calcination without an"

calcination in air

calcination without air

calcination in air

Figure 7-Effect of precursors on catalytic behavior of catalysts prepared by sol-gel(T=573 K). coprecipitated a) 100

.

.

.

D-

sol-gel a) .

I

100

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r

:

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10 15 Time (h)

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El -conversion O-sel. HCHO A-sel.CO. Figure 8-Stability tests without (a) and with (b) water in reactor feed.

I

I

10 15 Time (It)

I

20

_

816 REFERENCES

1. Reuss, G., W. Disteldorf, O. Grundler, A. Hilt, "Formaldehyde", in Ullmann~Encyclopedia of Industrial Chemistry, VoI.A11, p.619, VCH Publishers, 54 Ed. (1992). 2. Satterfield, C. N., "Catalytic Oxidation Methanol to Formaldehyde", in Heterogeneous Catalysis m Pratice, McGraw-Hill, New York (1980). 3. ECN Process Review, p.30, April 1994. 4. Golodets, G. I., J. R. H. Ross, "Heterogenous Catalytic Reactions Involving Molecular Oxygen", m Studies m Surface Science and Catalysis, Vol. 15, Elsevier Science Publishers B. V., Netherlands (1983). 5. Trifir6, F. S. Notarbartolo and I. Pasquon, J. Catal., 22, 324 (1971). 6. Sun-Kou, M. R., S.Mendioroz, J. L. G. Fierro, J. M. Palacios, A. Guerrero-Ruiz, J. Mater. Sci., 30, 496 (1995). 7. Novakova, J. and P. Jiru, J. Catal., 27, 155 (1972). 8. Pernicone, N., J. Less-Comm. Met., 36, 289 (1974). 9. Ai, M., J. Catal., 54, 426 (1978) 10.Kolovertnov, G. D., G. K. Boreskov, V. A. Dzisko, B. I. Popov, D. V. Tarasova and G. C. Belugina, Kinet. Katal., 6, 6, 950 (1965). 11 .Wilson, J. H., Ph.D. Thesis, University of Wisconsin-Madison (1986). 12.Schneider, M and Baiker A., Catal.Rev.-Sci. Eng., 37(4), 515 (1995) 13.Wilhelm F. Maier, F. M. Bohnen, J. Heilm~n, S. Klein, Hee-Chanco, M. F. Mark, S. Thorimbert, I.-C. Tilgner, M. Wiedorn, in Applications of Organometallic Chemistry in the Preparation and Processing of Advanced Materials, Ed. J. F. Harrod and R. M. Laine, Kluwer Academic Publishers, Netherlands (1995). 14.Mizukami, F., S. Niwa, M. Toba, T. Tsuchiya, S. Shimidzu, S. Imai and J. Imamura, "Preparation and Propreties of the Catalysts by a Chemical Mixing Procedure" in Preparation of Catalysts IV, Ed. B. Delmon, P. Grange, P.A.Jacobs and G. Poncelet, Elsevier Science Publishers B.V., Amesterdam (1987). 15.Soares, A.P.V., Ph.D. Thesis, Technical University of Lisbon, Lisbon (1996) 16.Abaulina,L. I., G. N. Kustova, R. F. Klevtsova, B. I. Popov, V. N. Bibin, V. AMelekhina, V. N. Kolomiichuk and G. K. Boreskov, Kinet.Katal., 17, 5, 1126 (1976). 17.Boreskov, G. K., G. D. Kolovertnov, L. M. Kefeli, L. M. Plyasova, L. G. Karakchiev, V. N. Mastikhin, V. I. Popov, V. A. Dzis'Ko and D. V. Tarasova, Kinet. Katal., 17, 1, 125 (1965). 18.Pernicone N., Catal. Today, 11, 85 (1991) 19.Anagha, A. B., S. Ayyappan and A. V. Ramaswamy, J.Chem.Tech.Biotechnol., 59, 395 (1994). 20.Petrini, G., F.Garbassi, M. Petrera and N. Pernicone, "Study of Iron(H) Molybdate as Precursor of Catalysts for Methanol Oxidation to Formaldehyde", in Chemistry and Uses of Molybdenum, Ed. H. F. Barry and P. C. Mitchell, p.437, Climax Molybdenum Company, Ann Arbor, Michigan, USA (1982).

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

817

Oxidation Catalysts Prepared by Mechanically and Thermally Induced Spreading of SbzO3 and V205 on TiO2

U. A. Schubert a, J. Spenglera, R. K. Grasselli a'b, B. Pillep c, P. Behrens c and H. KnOzingera

alnstitut for Phys. Chemie, Ludwig-Maximilians-Universitat, D-80333 Mimchen, Germany bDepartment of Chem. Engineering, University of Delaware, Newark, DE 19716-3116, USA ~ far Anorg. Chemie, Ludwig-Maximilians-Universitat, D-80333 MOnchen, Germany

Vanadium and antimony oxides are essential parts of some industrial catalysts for the selective oxidation of substituted aromatics to the corresponding anhydrides [1] and the selective oxidation of paraffins to the corresponding unsaturated acids and nitriles [2]. These catalysts are generally prepared by impregnation or coprecipitation methods. It was the objective of this study to investigate an alternate method of catalyst preparation, a method based on a solventless ball-milling technique, aimed at adequately dispersing the active catalyst ingredients on a given support material to yield acceptably effective catalysts. Specifically, the spreading and the dispersion of V-oxide, Sb-oxide, and V-Sboxides on Ti02 supports were investigated by means of the ball-milling technique, and the so prepared materials compared to conventionally prepared materials. The tribo-ehemieal process of the former method was followed by spectroscopic techniques including XPS, XANES, and TPR, which revealed that active phase dispersions comparable to those obtained by conventional preparation techniques can readily be achieved by the milling method. It was further observed that the addition of small amounts of water during the milling greatly enhances the rate of the dispersion process. Catalytic tests of a V-oxide-on-TiO2 composition reveal that the selective oxidation of o-xylene to phthalic anhydride (PA) proceeds at comparable levels for compositions prepared by either the milling or the impregnation method. Therefore, it is concluded that the solventless ball-milling technique is also an effective alternative method for the preparation of selective oxidation catalysts.

1. I N T R O D U C T I O N Partial oxidation of hydrocarbons employing mixed metal oxides as catalysts comprises an economically important class of reactions for the upgrading of base feed stocks [3]. An illustrative example of it is the partial oxidation of o-xylene and/or naphthalene to phthalic anhydride (PA) with a world production of 3.2 million metric tons per year, industrially carried out in shell and tube reactors using air as the oxidizing agent [4].

818 Catalysts employed in the oxidative production of PA are V20~-based compositions which are generally of the monolayer type and are supported on titania (anatase). With such catalysts, selectivities in excess of 80 mole % PA are achieved at essentially complete conversion. The utilization of Sb-V-oxide-based catalysts supported on anatase improves the PA selectivity. However, little is known about the intrinsic chemical or electronic effects of Sb203 in such catalytic systems, as well as the chemical and physical characterization of the supported Sb-oxide or supported Sb-V-mixed metal oxide [ 1]. Industrial catalysts are prepared exclusively by impregnation of the support material with aqueous solutions of the active phase materials, followed by subsequent drying and calcination. In the work described here, an alternative mode of catalyst preparation was chosen, based on a solventless, mechano-chemical method of incorporating the active catalytic components on the support carrier via ball-milling. Mechano-chemical treatment of solid materials leads not only to an alteration of its morphology and texture, e.g. formation of surface solid solutions [5], but also in the case of catalytic compositions to an alteration of their catalytic behavior [6]. For example it was demonstrated in the case of a V-oxide/TiO2 - based PA catalyst, that milling of the catalyst components led to an enhancement of PA yields, concomitant with observable spectroscopic and microscopic changes of the catalyst surface [7]. In our study reported here, the tribochemically induced dispersions of Sb-oxide and Voxide on TiO2 were investigated by means of spectroscopic methods for the binary systems Sb-oxide/TiO2 and V-oxide/TiO2, and the ternary system Sb-V-oxide/TiOv One of our research aims was to prepare a catalyst of the ternary system Sb203/V2Os/TiO2 by ball-milling, followed by calcination, having comparable dispersions of the single oxide components as those obtained by conventional methods employing solution impregnation. Another objective in our current study was to determine what influence the addition of water to the materials to be milled might have on the dispersion of the f'mal product, as compared to the product obtained by dry milling only. The catalysts prepared by ball-milling the oxide components with the carrier were evaluated by a test reaction, the oxidation of o-xylene with air, in a specially designed microreactor, determining their activity and selectivity, and by comparing their so determined catalytic properties with those of industrially prepared catalysts.

2. E X P E R I M E N T A L

For catalyst preparation the mechanical treatment was done in a planetary mill with the dry oxidic compounds or with 10 wt. % 1-120. For that purpose the oxides were mixed carefully in a mortar and then milled (145 rpm) in a agate vessel (250 ml) with six agate balls. The time of milling was chosen to be between l h and 20h. In the case of heat treatment of the materials, 400~ for 5h was chosen for calcination. The impregnated samples were prepared from an aqueous Sb-III-acetate suspension followed by drying at 110~ and calcination at 400~ for 5h.

819 The theoretical monolayer loadings were estimated according the formula of Roozeboom [8]. In the case of 1 ML SbxO3/TiO2, 10 wt. % Sb203 was used and for 1 ML V2Os/TiO2, 6 wt. % V20-yrI'iO2 Was used. The XPS measurements were carried out at a modified VSW ESCA 100 with Mg Ka and AI Ka X-ray radiation. The analyzer was operated in the fixed analyzer transmission mode (FAT) at 22 eV pass energy. The resolution of the spectrometer is given by the linewidth of AU4fT/2. For a sputtered and annealed gold sample, 1.65 eV was obtained. The lines corresponding to Cls, Sb3ds/2, V2p3/2, Ti2p and Ols were analyzed. The pressure in the analyzer chamber during spectra collection was always below 5.0 * 10-8 mbar. To determine charging effects, all signals were referred to the C 1s line corresponding to graphitic carbon at 284.4 eV or to the Ti2p3/2 signal from TiO2 at 258.5 eV [9]. The TPR experiments were carried out with the following parameters: reduction in H2 (0.58 ml/min) in N2 as carrier gas ( 11.5 ml/min) at a heating rate of 10~ The catalytic test reactions were carried out in a microreactor under the following conditions: reaction temperature T = 330~ atmospheric pressure with 0.7 % o-xylene in air, and a space time W/F = 5.0 * 10-5 kgeat * 1-1 * h. Under test conditions the catalysts were diluted in quartz (1 : 7 wt. %).

3. RESULTS and DISCUSSION 3.1.1 SlhOa/TiO2

With the help of X-ray-Photoelectron-Spectroscopy (XPS) it is possible, because of the surface sensitivity of the method, to make in addition to the determination of surface oxidation states also some conclusions about the dispersion of surface species. The ratio of the signal intensities of the support material to those of the supported species is a measure of the dispersion [10]. The ratio of the signal intensities of Ti2pv2 and Sb3d3/2 decrease as a function of milling time. This is shown in Figure 1, for the case of a theoretical monolayer (1ML) of Sb2OflTiO2 as a function of the pretreatment, i.e., the duration of the milling. The decrease of the intensity ratio correlates with the increase of the dispersion of Sb-oxide on TiO:. In all instances of milling it is observed that the signal ratios in the beginning, i.e., at short milling times of l h to 5 h, experience a large change while the differences between 10 and 20 hours of milling are relatively small. It follows from this that for the ball mill employed, the milling times of l h to 20 h are sufficient to achieve dispersions which do not increase any further with milling times exceeding 20 h.

820

8t / n

[

--I---

t"

8

ctL,y ~

--V--" with 10 wt. % H ~ .~11_~_

m

--O--" dry ~ ! , ~

+ ..~ 400~ --&--" ~ " ~ ~ / + ~ 400~

-6

"-,,,... i,

C

0

5

9

10

15

20

L-O

duration of .~1]~r~g in h Figure 1:

XP-signal ratios Ti2p/Sb3d of 1 ML Sb203]TiO2 depending on the milling time.

The signal ratios of the dry milled materials are compared in Figure 1 to those of the calcined dry milled materials, as well as to the 10 wt % H20-milled materials, and the conventionally, by means of wet impregnation, prepared material. Where applicable, the calcination temperatures are indicated in the figure. It is apparent that the dispersion of Sboxide on TiO2 of the 20 h dry milled sample is comparable to that of the conventionally prepared sample. Comparable values are also obtained from the dry milled and subsequently calcined samples (5h 400~ With the impregnated sample, complete dispersion is reached already after 1 h of milling. An increase of the Sb-oxide dispersion as a function of milling time cannot be found with the calcined samples. The signal ratios of the milled samples in the presence of 10 wt. % H20 are also illustrated in Figure 1. The influence of the addition of water is clearly noticeable. The signal ratios of the samples milled in the presence of water are clearly different from those milled dry, when the milling times are kept short. Already after 2 h of milling, the materials milled in the presence of water reveal signal ratios comparable to the impregnated samples, the calcined samples or the 20 h dry milled samples. The signal ratios of the samples milled in the presence of 10 wt. % H20 which were

821 subsequently also calcined are not shown in Figure 1. The signal ratios of the latter samples are comparable to those obtained for the dry milled and calcined samples. Similar conclusions about the increase of the dispersion of reducible species such as Sb203 can be obtained by employing temperature programmed reduction spectroscopy (TPR). For a quantitative interpretation of the TPR profiles it is necessary to integrate the recorded signals. In this manner it is possible to assess the oxidation states of the respective samples through the uptake of the reduction equivalent. The observed oxidation changes of the materials investigated are illustrated in Table 1.

Table 1:

From TPR results derived changes in the oxidation state of Sb-oxide in the systeme 1 ML Sb-oxide/TiO2 depending on the pretreatment.

s~steme 1 ML Sb-oxide/'ri02 after 20h dry milling

chan~es of oxidation state

3.1

Sb(III) -~ Sb(O)

1 ML Sb-oxide/TiO2 after 20h dry milling + 20h 450~ in N2

2.9

Sb(III) -~ Sb(O)

1 ML Sb-oxide/TiO2 after 20h dry milling + 20h 450~ in 02

5.0

Sb(V) -~ Sb(0)

For the system 1ML Sb203/TiO2, milled dry for 20 h, it is possible to assess an oxidation change from Sb 3+ to Sb ~ whereas in the 02 calcined samples an oxidation state of Sb5§ was found, as expected. These changes in oxidation state can be also followed by means of XPS. For the dry milled samples there is no significant change to be observed in the binding energy of Sb3d3r2 signal (1 h milled, 539.4 eV; 20 h milled, 539.6 eV). This agrees well with the literature value for Sb203 of 539.4 eV [ 11]. When milling with 10 wt. % HE0, a change in the binding energy is observed between the 1 h and 2h milled samples. After 1 h of milling the Sb is still found to be in the +III oxidation state (Sb3d3/2 being 539.6 eV), while after 2h of milling a value of 541.8 eV is observed. The latter corresponds to the literature value for Sb205 of 541.6 eV for the Sb3d3/2 signal [12]. The same values for the binding energies are obtained for the Sbsignals of the calcined samples. This implies that no tribo-chemically induced oxidation of the Sb3+-oxide could be observed for the dry milled materials, however, oxidation was definitely observed in the calcined samples irrespective of the milling time. In contrast to these findings is the observation that unsupported Sb203 requires a significantly higher temperature, i.e., . 9 600 ~ to oxidize Sb3+-oxide to Sb4+-oxade. Neat SbS+-oxide can be obtained from Sb3+-oxide under dry conditions only when placed under oxygen at high pressures (sealed tubes, autoclaves). However, this oxidation occurs already at room

822 temperature with TiO2 supported samples milled for 1 h in the presence of water. In this case not only the mechanical activation of the oxide, but also the interaction of the Sb-oxide with the support material (TiO~) is of importance, since this tribo-chemical effect does not occur by milling only the neat Sb-oxide. The change in oxidation state of Sb can also be determined for the dry milled and calcined samples by means of X-ray absorption measurements (XANES), using the L~-edge (beamline E4, DORIS III, HASYLAB, Hamburg). Herewith, it is possible to recognize Sb 3+-, Sb4+- and SbS+-compounds through the shitt of the white line (transformation from 2s to 5p) in the reference compounds Sb203, Sb204, Sb205 and Na[Sb(OH)6], shifting from 4.7023 keV for Sb 3+ in Sb203 to 4.7068 keV for Sb 5+ in Na[Sb(OH)6]. The spectra of the 1ML Sb203/TiO2 dry milled samples exhibit the same white line shift as the Sb203-reference, irrespective if the milling was carried out for l h, 5h or 20h. Conversely, the calcined samples exhibit a distinct double peak structure, which suggests the existence of Sb 3+ and Sb5+ species. Through spectral simulation and comparison with reference compounds it is possible to assess the ratio of Sb 3§ to Sb5+. With about a 45:55 ratio, a stoichiometry of Sb2Oa+xis indicated [13], herewith, confirming by the XANES-investigations the XPS-results discussed above. The apparent differences in the oxidation states of the Sb-oxides, as measured by XPS and XANES, can be explained on the basis that the differentiation of Sb4§ and SbS§ in the XP-spectra is most difficult because of the small difference in the binding energies of only 0.3 eV between the two species [ 11]. However, Sb 5§ species in particular which reside on the surface of these systems lead, because of the surface sensitive XPS method, to an overweighting of the Sb 5+ oxide by this method.

3.1.2. VzOs/TiOz For the system V2Os/TiO2it is also possible to show with the help of XPS that milling leads to a distinct increase of dispersion of the V-oxide on the TiO2 carrier. The ratios of the signal intensities of Ti2p~/2 and V2p3/2 are depicted in Figure 2 as a function of the milling time. It is apparent that under dry milling conditions the dispersion increases significantly with time of milling. When l0 wt. % H20 is present during milling, the dispersion achieved after 1 h of milling is about equal to that achieved after 20 h of milling under dry conditions. The signal ratio of the wet milled sample does not change significantly with an additional 20 h of milling from that achieved after only 1 h of milling. In contrast, calcination of the samples enhances significantly the dispersion, which increases still further when mechanical activation preceded calcination. With long milling times the calcination step outweighs dispersion differences, whether the samples were milled dry or wet. In order to determine if a reduction of the vanadium oxidation state takes place with milling, starting with V205 on TiO2, it was necessary to subtract the interfering x-ray satellites in the region of the V2p signals. This is achieved with the aid of spectral fitting programs, which allow the different binding energies of V-oxides to be considered (reference values are: V2p3~: V2Os, 517.2 eV; V204, 516.0 eV; V203, 515.8 eV).

823

9-

8- \

_._:

" \

7_-

/

i ~ ~%/~

I

-A-: -v-:

m

I-9

1-8

~_-,., ,n.th io ,,t. ~, %0 ,,~m,,,a.th~.o,,e. , ~ o + ~ . , 4 . ~ * c

r

9

7 6 5

2 0

2

4

6

8

10

12

14

16

18

20

d u r a t i o n o f rail]~r~g in h XPS signal ratios Ti2p1/E/V2p3/2 of 1 ML V2Os/TiO2depending on the conditions of preparation.

Figure 2:

In the V2pa/2-signals of the V-oxide/TiO2 systems, a distinct increase in the linewidth as compared to the signals of the reference compositions V205, V204 and V203 with 1.6 eV is observed. This can serve as a basis for the differentiation of V 5§ and reduced V-species, since the difference in the binding energy between V 5+- and V4+-oxide is more than 1 eV, and is therefore easily distinguishable in the simulation of the spectra. Conversely, the differentiation between V4 - and V -oxtdes by simulation is, became of the small differences of the binding energies of only 0.2 eV, nvt+possible. Therefore, in the subsequent discussion, distinction will be made only between - and reduced V-species. Significant amounts of reduced V-oxide species are identifiable already at short milling times for the dry milled Voxide/TiO2 material. The signal ratio VS§ r~ remains independent of the milling time at about 30:70. In the 10 wt. % H20 milled samples, the signal ratio shifts to approximately 50:50, whereby the post calcined samples exhibit a reverse ratio of 70:30, irrespective of the presence or absence of water during the milling operation. Calcination at 450 ~ for 5 h causes a portion of the vr~L-Species in the vicinity of the surface to reoxidize to V 5+. Since the mechanically induced reduction of V205 starts at the surface, the ratios of V 5§ to v~d-species -

9

9

+

3+

9

824

is certainly too low in comparison to the total amount of V-species, i.e., V " ~ is with certainty overweighted.

3.1.3. SbzOz/VzOs/TiOz From the signal intensities it is apparent that the milling conditions employed significantly influence the V/Sb surface properties. If the activation is carried out in the dry state, then the Sb-content increases significantly with milling time. With th el0 wt. % H10 containing samples, a constant V/Sb-ratio sets in already atter short milling time, albeit at a lower level than with samples which were ball-milled dry. That implies that a still greater enrichment of Sb occurred on the surface. Comparable results are obtained when the samples are post-treated through calcination after milling, whereby the dry milled and calcined samples exhibit a constant, but smaller V/Sb ratio than the wet milled samples. These results are illustrated in Figure 3.

14 'U m

B

12

z ~

~o..~2%/']:~io 2

-II-

:d~

,,i 11 _,~

12

--V--" dry .~11~,.1 + 5h 400~ --O--: with 10 wt. % ~ O ~ 1 ~ = ~

10"

--A--: with 10 wt. % ~ O .~11=~ + 5h 400~

E 0

.r'l

-10

s~

4~ .

m.

20

Figure 3:

14

V ~ V

6 i :-----_* '

I

5

V

6

-* '

I

10

'

I

15

'

duration of milling in h

I

20

XPS signal ratios V2p3/z/Sb3d3/2of 1 ML Sb203N2Os/TiO2 depending on the conditions of preparation.

In the case of the dry milled and calcined samples, no V/Sb values could be determined for the lh to 3h milled samples, because the signals of V2p and Sb3d are not sufficiently large to be evaluated. Therefore, only the differences of the signal ratios of the individual samples are summarized in Table 2.

825 Ti2pl/z/Sb3d3/2 signal ratios of the Sb203N2Os/TiO2systeme depending on the preparation.

Table 2:

duration of milling

dry milled

dry milled + 5h 400~

milled with 10 wL % HzO

milled with 10wt. % H20 + 5h 400~

lh

9.1

--

1.5

1.7

20h

3.6

7.3

2.0

2.1

From the observed differences of the signal ratios it is recognizable that the treatment of the specific samples has a significant influence on the dispersion of active components on the surface of the carrier material. Particularly for the dry milled and calcined samples, the active components appear after calcination not to be accessible any more for XPS, because of the small signal intensities. This effect can be avoided through the addition of 10 wt. % H20. For the determination of the oxidation states of Sb and V, the spectra were analyzed as discussed above through curve fitting. The binding energies so detemined for the Sb3d3/2 signals are given in Table 3.

XPS Sb3d3/2 binding energy of the Sb203/V2Os/TiO2systeme depending on the preparation.

Table 3:

duration of milling

dry milled

dry milled + 5h 400~

milled with 10 wt. % H20

milled with 10wt. % H20 + 5h 400~

lh

539.6

--

540.0

539.9

20h

539.9

539.8

540.0

539.9

With the exception of the dry milled samples for short periods, the binding enel'gy of 3+ the Sb-signal rises, and values are found which lie between the Sb -oxide (539.6 eV) ~ the 4+ 9 Sb -oxide (540.3 eV) [12], independent of the preconditioning of the samples. For the V2p3a signal it is possible to show through fitting, that in all samples in addition to vS+-species also reduced V-species exist.

826 3.1.4 Catalyst Evaluation

As known from the literature, V2Os/TiO2 is, besides the Sb203/V2Os/TiO2 system, a well established catalyst in the PA synthesis [14, 15].Therefore two supported V-oxide-onTiO2 catalysts, one preparedby the milling method and the other by the conventional impregnation method, were tested for the oxidative conversion of o-xylene. The reaction was carried out in a microreactor under the conditions described above and the results are summarized in Table 4. Table 4:

Sample

results from o-xylene oxidation of a milled V2Os/TiO2- and a conventionally prepared V2Os/TiO2 catalyst.

Conversion in %

Y(o-tol.) in %

Y(PA) in %

Y(phthlide) in %

S(PA) in %

S(total OX.

prod. in % 1ML VEOs/TiO2 lhdry milled 1ML V2Os/TiO2 impregnated

Y(o-tol.): Y(PA): Y(phthalide): S(PA): S(total oxid. prod.):

15

2.8

8.0

1.2

53

80

20

2.9

8.9

1.3

45

65

yield of o-tolualdehyde. yield of phthalic anhydride yield of phthalide selectivity referring to phthalic anhydride selectivity referring to the total amount of oxidation products

It is apparent that the two catalysts give comparable yields of the desired PA product, under comparable reaction conditions. Therefore, it is concluded that the solventless ballmilling method described herein is a viable method of preparing selective oxidation catalysts.

4. S U M M A R Y In this study it was demonstrated that with the aid of solventless ball-milling of catalyst components, dispersions of active components on carrier materials could be achieved, as measured by TPR, XPS and XANES, that are comparable to the dispersions achieved through conventional impregnation techniques of catalyst preparation. Comparable catalytic results are obtained by both preparation methods. Specifically, catalysts supported

827 on TiO2 prepared by ball milling techniques yield phthalic anhydride in the oxidation of oxylene at comparable levels to conventionally prepared catalysts. Therefore, based on catalytic and spectroscopic results presented in this study, it is concluded that selective oxidation catalysts can be prepared by solventless ball-milling techniques which give comparable catalytic properties to those prepared by conventional solution impregnation methods. Addition of small amounts of water during the milling process enhances the rate of active phase dispersion on the catalyst carrier and thereby shortens the milling time required to attain the desired catalytic properties.

ACKNOWLEDGMENTS This work was financially suplx~ed by the Bayerische Forschungsverbund Katalyse FORKAT, by the Deutsche Forsehungsgemeinschaft (SFB 338) and by the Fond der Chemischen Industrie. R.K. Grasselli gratefully acknowledges the Alexander von Humboldt Stiftung for receipt of a Humboldt Research Prize.

REFERENCES

[1] [2]

[3] [4]

[5] [6] [7]

[8] [9] [lO] [11] [12] [13] [14] [15]

S.E. Golunski and D. Jackson, Appl. Cata., 48 (1989) 123. a. A.T. Guttmann, R:K. Grasselli, J.F: Brazdil and D.D Suresh, US Patent No. 4 746 641 (1998). b. A. Andersson, S.L.T. Andersson, G. Centi, R.K. Grasselli, M. Sanati and F. Trifiro, Proc. 10th Int. Congr. Catal., Budapest (Eds. L.Guczi et al.) 1992 A, 691. a. R.K. Grasselli and J.D. Burrington, Adv. Catal., 30 (1981) 133. b. R.K. Grasselli, J. Chem. Ed. 63 (1986) 216. K. Weissermel and H.-J. Arpe, Industrielle Organische Chemie, VCH Verlagsgesellschaft, Weinheim, 1994, 415. P.A. Zielinski, R. Schulz, S. Kaliaguine and A. Van Neste, J. Mater. Res., Vol. 8, No. 11 (1993) 2985. H.S. Horowitz, C.M. Blackstone, A.W. Sleight and G. Teufel, Appl. Catal., 38 (1988), 193. V.A. Zazhiggalov, J. Haber, J. Stoch, L.V. Bogutskaya, I.V. Bacherikova, Proc. 1lth Int. Cong. Catal., Baltimore (Eds. J.W. Hightower, W.N. Delgass, E. Iglesia and A.T. Bell), Elsevier, 1996 B, 1039. F. Roozeboom, T. Fransen, P. Mars and P.J. Gellings, Z. Anorg. Allg. Chem. 449 (1979) 25. D. Briggs and M.P. Seah, Practical Surface Analysis, John Wiley & Sons, Chichester, 1994, Part I: Auger and X-ray Photoelectron Spectroscopy. J.W. Niemantsverdriet, Spectroscopy in Catalysis, VCH Verlagsgesellschaft, Weinheim, 1993, 51. B. Visswanathan, S. Chokkalingam, T.K. Varadarajan and S. Badringarayanan, Surf. Coat. Technol. 28 (1986) 201. R. Izquierdo, E. Sacher and A. Yelon, Appl. Surf. Sci., 40 (1989) 175. U.A. Schubert et al., to be published in J. Phys. Chem.. M.S. Wainwright and N.R. Foster, Catal. Rev. -Sci. Eng., 19 (1979) 211. V. Nikolov, D. Klissurski and A. Anastasov, Catal. Rev.-Sci. Eng. 33 (1991) 319.

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3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

829

The Effect of Preparation Parameters on the BET Surface Area of Z r O 2 Powder YuanYang Wang YanZhen Fan

YuHan Sun*

SongYing Chen

State Key Laboratory of Coal Conversion, Institute of Coal Chemistry, Chinese Academy of Sciences, Taiyuan 030001, P. R. China

1. INTRODUCTION Zirconium oxides are extensively used in many fields (such as ceramic, refractory, sensor and catalysis) due to their propel~es[ 1]. Many methods have been developed for the production of such materials[2]. Among them, Rapid Thermolysis Approach (RTA) is a safe, simple and instantaneous route. Kingsley firstly reported the preparation of alumina and related oxides by RTA[3]. However, the samples derived showed a relative low BET surface area. ZrO 2 soproduced has been shown to be highly active towards CO oxidation and methane combustion. In order to have a full understanding of so-derived zirconia, the preparation of zirconia is in the first instance investigated in the present work. 2. E X P E R I M E N T A L

Zirconium oxychloride (or zirconium oxynitrate) was mixed with urea, nitric acid and water in an evaporating dish, and then thermolyzed in a muffle furnace at a fixed temperature about 10 minutes. This led to samples with a foam structure. The samples were then characterized by physical adsorption of N 2 at 77K (ASAP-2000), TEM (H-600 II) and XRD (D/max-TA). 3. RESULTS AND DISCUSSION In the RTA process, large amount of gases are produced by the decomposition of precursor and fuel (i.e. urea and nitric acid), which leads to a foam structure. Figure 1 illustrates the photograph of so-produced samples. The ZrO 2 foam is observed to brim over the container. 3.1. Effect of Precursor and Fuel Composition on BET Surface Area of ZrO 2 Powder

Table 1 shows that BET surface areas of samples change with the precursor and the fuel composition (i.e. urea, nitric acid or their mixture). Samples from zirconium oxychloride show * To whom all correspondence should be addressed

830 higher surface area than those from zirconium oxynitrate. Furthermore, BET surface area for ZrO(NO3) 2+ urea system are lower than those for ZrOCI~ + HNO 3+ urea system. This is in a good agreement with the reported by kingsley[3], in which BET surface area of pure A1203 sample was 8.30 mVg, and that of 50 wt% ZrO 2 in A1203 sample only 3.12 m2/g when zirconium oxynitrate was employed. These indicate that nitric acid plays important roles in the formation of the foam. Obviously, the foam structure and high BET surface area should be attributed to the synergetic effect of the decomposition of precursor and the combustion of fuel[4]. Figure 1. The photograph of ZrO 2 with foam structure Table 1 The effect of precursor and fuel composition on BET surface area of ZrO 2 powder* Sample No.

Precursor

H20/Zr 4§ HNO3/Z1"4§

CO(NH2)2/Zr 4§

SBzr (m~/g)

P1

ZrO(NO3) 2

10.0

0

4.0

3.82

P2

ZrO(NO3) 2

10.0

0

1.0

6.25

P3

ZI~)C12

27.5

4.0

4.0

8.33

P4

ZrOC12

13.8

4.0

1.0

35.31

* thermolyzed at 773K.

3.2. Effect of Solution Composition on the BET Surface Area of ZrOz Powder The effect of solution composition on the BET surface area of ZrO 2 powder is shown in Table 2. Obviously, with the decrease of the amount of urea or water in the presence of HNO 3, BET surface area ofZrO 2PoWder increases; and the concentration of nitric acid shows a optimum value. This indicates that less urea and water, with moderate nitric acid, are important for higher BET surface area. Figure 2 indicates the difference of two types of ZrO 2 samples in their morphologies. Although the particle size for both are in nanometer-size, their morphologies and BET surface

831 Table 2 The effect of solution composition on the BET surface area of ZrO 2 powder* Sample No.

H20/ZrOC12

HNO3/ZrOC12 CO(NH2)2/ZrOC 89 SBrr (m2/g) i

S1

0

4.0

0

$2

0

4.0

0.5

150.46 51.22

$3

0

4.0

1.0

40.15

$4

0

2.0

1.0

22.20

$5

0

8.0

1.0

24.10

$6

13.8

4.0

1.0

35.31

$7

27.5

4.0

1.0

27.47

* thermolyzed at 773K. area differ from each other seriously. Samples produced in the presence of urea consist of 50 nm of particles with crystal structure whlie that made without urea is highly-dispersed. It is clear that these defferences are closely related to the praparation mechanism. At high temperature, the following reactions take place in the system of precursor and fuel[5]: CO(NH2) 2 + H20 ~ HNO 3

> CO2(g) + 2NH3(g)

> NO2(g) + NO(g) + O2(g) + H20(g)

Figure 2. The TEM graphs of samples S 1 and $7

(1) (2)

832 CO(NH2) 2 + HNO 3

> CO(NH2)2 HNO 3

(3)

CO(NH2) 2 HNO 3 + 3HNO3 > CO2(g) + 4NO2(g) + 2NH3(g) + H20(g) + O2(g) ZrOC12 + 2HNO 3 > ZrO(NO3) 2 + 2HCI(g)

(4)

ZrO(NO3) 2

(6)

> ZrO 2 + 2NO2(g) + 1/2 O2(g)

(5)

In the present of urea, large amount of gases are produced to form a foam structure, and the process is then controlled by combustion mechanism, the reactions (5) and (6) hardly influence the process. Without urea in the sample (see S 1 in Table 2), only reactions (2), (5) and (6) occur, ZrO~ powder is mainly made by the decomposition of ZrO(NO3) 2 and no foam structure appears, the process is controlled by decomposition mechanism. When the samples are prepared through combustion route, the reactions between nitric acid and urea gives rise to a very high temperature which even reaches 1873K or so[3]. Obviously, this high temperature results sintering and aggregation of ZrO2 powder, leading to large particles and then low surface area. With the decomposition route, ZrO2 powder can be produced by the decomposition of zirconium oxynitrate at about 773K. At such a mild temperature, small particles in a highly-dispersed state, of course, show high BET surface area, which is even higher than that of ZrO2 aerogels ( about 100 m2/g) prepared by sol-gel method involving supercritical drying at the same temperature[6]. The amount of water also influences the BET surface area. The BET surface area of ZrO2 is found to decrease with increasing the amount of water (see TaNe 2). This may be due to that water might retard the release of gases and the rapid expansion of the mixture. 3.3. Effect of Thermolysis Temperature on the Texture of ZrO 2 Powder

The thermolysis temperature is found to have a strong influence on the BET surface area ofZrO 2powder. Samples (Tl~and T2) thermolyzcd at 573K and 773K almost show the same BET surface area, and their surface areas are much higher than that of sample T3 produced at 973K (see Table 3). However, it is interesting that sample T2 shows a concentrated pore distribution around 3.5nm while the bimodal pore distribution is observed for samples T1 and T3 at 3.5nm and 6.5nm (see Figure 3). Furthermore, the pore volume of T3 is much lower than those of T1 and T3 (see Table 3). These imply that their BET surface area originates differently from each other. Considering that the preparation of the samples follows the decomposition mechanism, the difference should be related to their dispersion degree and crystal structure. As mentioned above, the sample produced via thermolysis at 773K consists of homogeneous highly-dispersed particles (see Figure 2). Consequently, sample T2 displays a concentrated pore distribution around 3.5nm, which produces high surface area. With the sample thermolyzed at high temperature(i.e. 973K), however, the sintering and aggregation of particles

833 Table 3 The effect of thermolysis temperature on the texture of ZrO 2 powder* Therm.

SB~

Pore

Pore vol.

Crystal size**

t/m**

temp.(K)

(m2/g)

size(nm)

(cm3/g)

(nm)

ratio

T1

573

146.16

7.31

0.27

. . . . . . . .

T2

773

150.46

5.19

0.20

7.60

3.13

T3

973

44.31

9.71

0.11

12.40

2.52

Sample no.

* H=O:HNO3:CO(NH2)2:ZrOC12= 0:4.0:0:1.0 ** calculated based on the XRD results. takes place. This leads to inhomogeneous growth of highly-dispersed particles and then a bimodal pore distl'ibution with most of pores around 6.5nm, which causes the decrease of the BET surface area of the sample. In the meantime, the phase transformation of tetragonal to monoclinic produces more monoclinic phase in samples T3 (see figure 4 and Table 3), and leads to the formation of large partiicles due to volumetric expansion[7]. Both result in a sharp decrease of BET surface area.

.....

O -- T e t r a g o n a l X -- M o n o c l i n i c

TI

. . . . . . . . . T2 ~

0.6

E 0,4

."" .,

.,'\

-!

(}.2

/"~

~

:-,'; : i,' ', : ," '~

" t

;; I'

'. '.

T3

o

,.

~. I

o

A

o

[

I

I T3

|~

1o

20

30

40

Pore diameter(nm)

Figure 3. The pore distribution of samples with different thermolysis temperature

50

6O

70

20

Figure 4. The XRD patterns for samples with different thermolysis temperature

In the case of sample T1, no such an aggregation occurs because of its lower thermolysis temperature of 573K even if it shows a similar pore distribution to sample T3. This may be caused by the inhomogeneity of amorphous ZrO2 particles due to the thermolysis temperature less than the crystallization temperature of 743K (see Figure 4). Therefore, the high BET surface area should be attributed to the high dispersion of amorphous ZrO2 particles.

834 4. CONCLUSION Rapid Thermolysis Approach is one of best routes to produce ultrafine ZrOr The samples so-derived are of nano-sized particles with high BET surface area. BET surface area of powder changes with the system of precursor and fuel. Synergetic effect of the decomposition of precursor and the combustion of fuel leads to foam-structured powder with high BET surface area. The presence of urea and nitric acid is very important. Thermolysis temperature has a strong influence on the texture of ZrO~ powder. 773K is the proper temperature for the preparation of ZrO~ powder with high BET surface area and concentrated pore distribution. ACKNOWLEDGEMENT

The authors acknowledge the National Natural Science Foundation of China for its financial supports. REFERENCES

1. S. S. Prakashi, C. J. Brinker, A. J. Hurd, Nature, 374(1995)439. 2. D. A. Ward, E. I. Ko, Chem. Mater., 5(1993)956. 3. J. J. Kingsley, K. C. Patil, Mater. Lett., 6(1988)427. 4. J. J. Kingsley, K. Suresh, K. C. Patil, J. Mater. Sci., 25(1990)1305. 5. A. M. Wynne, J. Chem. Educ., 64(1987) 180. 6. H. W. Xiang, B. Zhong, S. Y. Peng, Mole. Catal.(China), 8(1994)263. 7. P. D. L. Mercera, J. G. V. Ommen, E. B. M. Doesburg, Appl. Catal., 78(1991)79.

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

Preparation of VOHPO4.0.5H20 and (VO)2P207 performance for maleic anhydride synthesis

835

and

their

catalytic

T. Miyake and T. Doi Tosoh Corporation, Yokkaichi Research Laboratory, 1- 8 Kasumi, Yokkaichi- shi, Mie 510, JAPAN

Firstly, influence of alcohol solvents on characters of VOHPO4.0.5H20 prepared from V409 and ortho-HaPO4 was studied by valence of vanadium, XRD, SEM, TG/DTA, ~ P-NMR and Raman spectra. It was revealed that morphology and bulk characters of VOHPO4.0.5H20 differed among alcohols used. Then, oxidation of n-butane to maleic anhydride on (VO)2P207 obtained by transformation of VOHPO4-0.5H20 was investigated. Selectivity of maleic anhydride was independent of the ratio of surface areas; namely, so-called selective to non-sective surface of (VO)2P207.

1.INTRODUCTION Vanadium-phosphorus oxides are known as the effective catalysts for the production of maleic anhydride by oxidation of n-butane [1-3]. Especially a crystalline (VO)2P207 is claimed to be the active and selective phase [4,5], though some research groups insist that VOPO4 plays an important role [6]. Generally three methods are known to synthesize (VO)2P207; aqueous solvent method [4], organic solvent method [5] and VOPO4.2H20 reduction method [7]. It is reported that VOHPO4.0.5H20, precursor of (VO)2P207, and (VO)2P207 prepared by these methods have different surface and bulk characters. For example, Horowitz et al. [7] reported that VOHPO4-0.5H20 of different morphologies were obtained by varying P/V ratio and organic solvents in the VOPO4-2H20 reduction method. In the organic solvent method, Kesteman et al. [8] concluded that the primary factor which influenced crystallinity and morphology of precursor was the choice of alcohol used during dehydration-condensation of vanadium compound and ortho-H3P04. From an industrial viewpoint, as the selectivity of maleic anhydride is around 60%, the improvement is strongly desired. Recently, Igarashi et al. [4] and Bordes [9] proposed that the selectivity depended on the ratio of selective crystalline face to non-selective one of (VO)2P207. More precisely, on (200) mainly maleic anhydride was produced and on (042) thus-produced maleic anhydride was over-oxidized to carbon oxides. When this hypothesis is accepted, the morphology of (VO)2P207, and therefore that of the precursor VOHPO4"0.5H20, might play a very important role for the catalytic performance. In this study, at first in order to obtain a pure and crystalline precursor having various (001)/(220) ratio, precursors were synthesized using various alcohol solvents

836 during dehydration-condensation of V409 and ortho-H3P04. Then, the relation between the (200)/(042) ratio of (VO)2P207 and the selectivity of maleic anhydride was investigated with the pure (VO)2P207 transformed from the precursors.

2. EXPERIMENTAL 2.1. Preparation of VOHPO4.0.5I-hO Typical preparation procedure was as follows [10]: Firstly, V205 (91 g) was reduced to V409 in isobutyl alcohol (280 g) under reflux for 48 hours in N2 stream. After filtration and washing with acetone, the black solid V409 was dried at 313 K for 15 hours. Then the solid V409 (20 g) was suspended in 100 g of various flesh alcohols (isopropyl alcohol, isobutyl alcohol, 2-butyl alcohol, 2 - m e t h y l - l - b u t y l alcohol, cyclohexyl alcohol or 2 - e t h y l - 1 - h e x y l alcohol). The suspension was heated until 383K. (In the cases of isopropyl alcohol and 2-butyl alcohol, the suspension was heated until reflux.) Then, here added an ortho-H3P04 (18.2 g) in the same alcohol (7.2 g) dropwisely for 30 minutes. The ratio of P to V was 1.13 for preparation of the precursor. After H3P04 addition, this condition was kept for 5 hours (in the case of isopropyl alcohol, for 24 hours). Then the obtained slurl~ was filtered. Thus obtained solid was washed with acetone and was dried at 313 K for 15 hours. Precursors were also prepared by the aqueous solvent method and VOPO4.2H20 reduction method according to the literatures [4, 7]. To obtain pure precursors, special care was given to mixing, washing and so on. 2.2. Characterization Elemental analysis for V and P of the precursor and (VO)2P207 was carried out using inductively coupled plasma spectroscopy (ICP) on Kyotokouken UOP-2. According to the method of Hodnett [11], the average valence of vanadium was determined by the double titration method using KMn04. Powder X - r a y diffraction patterns were recorded with Mac 'Science M18XHF diffractometer using C u - K c~ radiation (40 kV, 100mA). Crystallite size of VOHPO40.5H20 and (VO)2P207 was calculated by Scherrer equation. Scanning electron microscopy (SEM) analysis was carried out with Shimadzu TPM 810. Thermogravimetfic (TG) analysis was carried out using SEIKO TG/DTA-220. 31 p_ NMR measurements were carried out with Nippondenshi JNM-GSX-270WB. Raman spectrum was recorded on I. S. A. Jobin Yvon Ramnor-U1000 spectrophotometer. The emission line at 514.35 nm from Ar laser was used for excitation. 2.3. Catalytic performance Precursors were transformed to (VO)2P207 by heat-treatment in N2 stream or n - b u t a n e - a i r stream. Into the tubular reactor, a 10 g portion of (VO)2P207 was placed. When gas component was arranged to 1.5% n-butane at space-velocity of 1500 h - 1 , the temperature was raised to the desired one. Products were analyzed by GC and LC.

837 3. RESULTS AND DISCUSSION 3.1. Influence of alcohol solvents on characters of VOHPO4-0.5H20 The P/V ratio and the average valence of vanadium of precursors prepared by the organic solvent method in various alcohols are shown in Table 1. In every case, the P/V ratio was 1.09 + 0.08. The average valence of vanadium of precursors were substantially 4+ irrespective of the alcohol used. Although there remained a small portion of alcohol between the layers (vide infra), these alcohols had substantially no influence on the measurement of valence (" Or-5 "~ Or-4

L.

o

-I

_ ~,t ~

10

_ 20

~.~ ..........

~

30

40

50

2 0/degree

Figure 1. XRD patterns of precursors prepared by the organic solvent method. (Notation; see Table 1)

838

Figure 2 shows the result of TG/J3TA measured in N2 stream. It is well known that transformation of VOHPO4-0.5H20 to (VO)2P207 occurs between 520 to 770 K [12]. Although this weight decrease is clearly seen in Figure 2, the rate differed among precursors. In the cases of precursors prepared in isobutyl alcohol, isopropyl alcohol and 2 - m e t h y l - 1 - b u t y l alcohol, the weight decrease curve was not sharp and this suggested that crystallinity of the precursor was not so good. To the contrary, in the cases of precursors prepared in cyclohexyl alcohol, 2 - e t h y l - 1 - h e x y l alcohol and 2-butyl alcohol, the weight decrease curve was very sharp and this suggested that the crystallinity was rather high. Differences in sharpness of the TG curves were well related to those of XRD peak intensities; the stronger the XRD peak intensity was, the sharper the weight decrease curve was. In Table 1, the TG result is summarized. Calibration was carried out by neglecting the weight decrease below 520 K and the theoretical weight decrease which corresponds to the transfomation is 10.5 wt%. T h e calibrated weight decrease for the latter three alcohols (namely, cyclohexyl alcohol, 2 - e t h y l - 1 - h e x y l alcohol and 2-butyl alcohol) was slightly higher than the theoretical value and this suggests that a small portion of the alcohol solvent remained betWeen the layers.

r~ r162

95

-o 90

Or-2

",

Or-3

~ 85 [-.,

Or-4

80

273

473

673

873

273

473

673

873

1073

Temperature / K

Figure 2. TG/DTA analysis of precursors prepared in various alcohols by the organic solvent method. (Notation; see Table 1)

Figure 3 shows the 31 p_ NMR spectra of precursors prepared in various alcohols. The precursors prepared in alcohols other than isobutyl alcohol showed the characteristic peak at - 1 3 8 ppm which could be assigned to VOHPO4-0.5H20 and no peak assigned to VO(H2P04)2 was observed. Therefore, we could say that purity of these precursors was high and short range order of VOHPO4-0.5H20 crystal was good, especially for 2-butyl alcohol and 2 - e t h y l - 1 - h e x y l alcohol. When intensities of the peak are compared, they are proportional to the crystallinity estimated from XRD. Here we would like to consider 3, P - M R result together with those from XRD and TG/DTA. Although we ignored the TG weight decrease below 520 K, especially in cases of isobutyl alcohol, isopropyl alcohol and 2 - m e t h y l - 1 - b u t y l alcohol, substantial weight decrease seems to exist. If we estimate that these weight decreases are from the alcohols remained between layers, precursors prepared in these alcohols could have

839

2 - butyl alcohol

LI~[~

2 - m e t h y l - 1 - butyl alcohol

.

j

isopropyl alcohol -. J . ~.L ,i , k...ua..,,.lt,dll.lk.lu~ld.klhahlh,di~k~ljlhk~l~iAl~lJd~htt.,.,I,,jLti,,,

~,~w'T"""*q"'"'"'~'"r ......r r l ~ , , W g ~ ~

,

isobutyl alcohol

300

200

100

0 -100 Chemical shift

-200

-300

Ppm

Figure 3. a, p_ NMR spectra of precursors prepared by the organic solvent method.

isobutyl alcohol

]

cyclohexyl alcohol

2 - e t h y l - I - hexyl

rather poor XRD peak intensities and ~1 P - N M R peak intensities. Figure 4 shows the Raman spectra of precursors prepared in various alcohols. In the case of 2 - e t h y l - 1-hexyl alcohol, it was difficult to measure Raman spectra of high resolution because of the fluorescence from the precursor. From the spectra observed, it should be emphasized that all the peaks ascribed to the precursor [13] are observed and no other peaks stem from impurities such as VOP04 phases were seen. Generally speaking, the intensity of the main peak assigned to V=O was stronger for alcohols which gave poor XRD peak intensities and ~ P - N M R resolution and vice versa. For example, Raman peak intensity of precursors prepared in isobutyl alcohol or cyclohexyl alcohol was strong while 31 P - N M R peak intensity of the precursor prepared in either of these alcohols was weak. In case of 2-butyl alcohol or 2 - e t h y l - 1 - hexyl alcohol, the opposite relation was observed. Considering that both XRD and ~1 p _ NMR are characterizing the bulk and that Laser Raman is sensitive to the surface, this phenomenon might be reasonably explained; the precursors having intensive XRD peak are thick and gives weak Raman intensity because of relatively low abundance of V=O alcohol on the surface.

q,)

2- methyl- I-

[

. . . . 2 - butyl alcohol

T

~a~k_-~,, ,-~,.,-,~ L ~ , . ,I ~',,~-~,wr~,~'~r't IW~"q~il

r,.,~r r ~

Figure 4. Raman spectra of precursors prepared by the organic solvent method. 100

700

1300 100

Raman shift / cm -'

700

1300

840

.............

....

;:~::~::~;~:;%~::::::~ii!:~i'.:,ii i ;:.4ii!~!ii i:~:~ii i;~i~F~i~.::::Niiiii? ili~i ::~U i~ ~!i~i ' ~iiii!iiiii',~.... iii~ii ...... ~,~

i-~-

i~iiiiN,

"

.......

isobutyl alcohol

cyclohexyl alcohol

isopropyl alcohol

2- ethyl- 1- hexyl alcohol

,.

.

.

.

m

..........~..-.

.

...... i:ii~!i ::i:~i:}~i:!i!:;ii~.~{::!~iii '?ii:i!i~:.:i~:~ ........

~:~:~,. :

~:::

2 - methyl- 1- butyl alcohol

i

m

2- Butyl alcohol

Figure 5. SEM photographs of precursurs prepared by the organic solvent method.

841 Figure 5 shows SEM photographs of precursors prepared in various alcohols. It is apparent that in every case the particles are of the same form and no other forms from impurities were observed. In the cases of precursors prepared in isobutyl alcohol, isopropyl alcohol and 2 - m e t h y l - 1 - b u t y l alcohol, the morphology of the precursor was petal- or rose-like. On the other hand, the morphology of the precursors prepared in cyclohexyl alcohol, 2 - e t h y l - 1 - h e x y l alcohol and 2-butyl alcohol was plate-like and rather flat. This was in accordance with the XRD results. In addition, it should be noted that the morphology of precursors were very similar to that of precursors prepared in the corresponding alcohol by the VOPO4-2I-hO reduction method [7]. From these results, it can be said that precursors are pure and uniform.

3.2. Relation between the ratio of surface area of selective face to non-selective one and maleic anhydride selectivity It is sometimes reported that to obtain steady catalytic performance hundreds of hours are necessary for activation of a catalyst under n-butane oxidation reaction [14]. Therefore, firstly the selectivity of the catalyst whose activity became stable at high n-butane conversion (>90 %, ca.100 hours) was compared with that of the catalyst activated under conventional procedure for about 500 hours (Figure 6). As the same maleic anhydride selectivity was obtained between these two catalysts at the same 60% n-butane conversion, we adopted this short-time activation procedure.

80

Sel.

6o

.s

--0--,-0-

;>

40

0

~0~.0---0

2o

m

--0--'0-

Conv.

m,,

320

q)

0 0

340 ~

~

360 ~

I

I

I

I

100

200

300

400

Reaction time / hr Figure 6. Activity and selectivity change during reaction Pressure; 101.3 kPa, n-Butane; 1.5%, GHSV; 1500 hr 1

--~

500

842

In Table 2, the catalytic performance at 60% n-butane conversion and the properties of catalysts after the reaction are summarized.

Table 2 Properties of catalysts and their catalytic performance Preparation Solvent TransV a l e n c e Surface Surface-area Temp. Conv. Sel formation of V Area(m2/g) ratio(-) ~ % % Aqueous H20 N2, 650~ 4.0 15.2 1.08 410 65.1 59.9 Organic isobutyl alcohol Butane-air 4.3 29.1 0.64 400 64.6 63.5 solvent 2-Butyl alcohol Butane-air 4.5 17.3 0.52 410 62.3 65.7 c-Hexyl alcohol Butane-air 4.1 20.2 0.83 400 59.5 63.0 MOP04 Benzyl alcohol N2, 800~ 4.0 7.2 0.82 440 61.0 58.0 Reduction Benzyl alcohol N2, 650~ 4.0 29.7 1.45 360 60.3 59.3 Benzyl alcohol N2, 650~ 4.0 28.2 1.44 380 60.3 61.4 Benzyl alcohol N2, 650~ 4.0 24.6 1.35 375 58.1 60.1 2-Butyl alcohol N2, 650~ 4.0 14.8 1.20 430 58.8 53.0 Surface-area r a t i o Crystal was tentatively considered to be a disc whose diameter and thickness were equal to the size of (042) and (200), respectively. Then, surface-area ratio is calculated as follows. Surface-area ratio = 2 * n ((042)/2)- / n 9(042) * (200) = (042) / 2 9(200)

As is shown in Figure 7, conversion was dependent on preparation procedure; more (Figure 8). This indicates that

it was indicated that the temperature for 60% n-butane the surface area of (VO)2P207 irrespective of the catalyst precisely, dependent on the surface area of (200) face (200) face is active and other faces are rather inactive.

440 420 ~

400

~ 380 => 360

~ 8:340

A O; aqueous solvent method r-l; organic solvent method

320 - A; VOPO4"2H20 reduction method 300

I

5

I

I

I

10 15 20 Surface area / m2g1

I

25

Figure 7. Relation between the temperature for 60 % n-butane conversion and surface area of the catalyst. Pressure; 101.3 kPa, n-Butane; 1.5%, GHSV; 1500 hr -1

30

843

440 420 ,.Q

a~

~176 400

a

0

0 ~ -~ 380

r

~>360 g

~

340

~

320 30O

D

O; aqueous solvent method !-1; organic solvent method A. VOPO4. 2H20 reduction method I

!

~"~'-

!

5 10 15 Surface Area of (200) face / m2g-1

20

Figure 8. Relation between the temperature for 60 % n-butane conversion and surface area of (200) face of the catalyst. Pressure; 101.3 kPa, n-Butane; 1.5%, GHSV" 1500 hr -1

No correlation between the maleic anhydride selectivity and the ratio of surface areas of selective to non-selective faces was observed (Figure 9). This suggests again that maleic anhydride is formed only on the so-called selective (200) face and it is not possible to improve the selectivity simply by increasing the surface area of (200) face.

844

4. CONCLUSIONS Depending on the alcohol solvent used for preparation, VOHPO4.0.5H20 of different characters were obtained from V409 and ortho-H3P04. Especially the precursor having stronger XRD peak intensities gave sharper TG weight decrease curve and higher 3~ P - N M R peak intensity, which meant that the precursor of this kind was pure, uniform and highly crystalline. Raman spectra indicated purity of precursors was high. Maleic anhydride selectivity did not depend on the ratio of the surface area of selective to non-selective faces and it was suggested that the maleic anhydride selectivity is not improved simply by increasing this ratio.

9 70

>

65

-~ 60

O

~.

ZX

55 O; aqueous solvent method .~ 50 -[2]; organic solvent method ZX; VOP04" 2H20 reduction method I 45 0.5

a

1.0 Ratio of selective surface area to non-selective one / -

Figure 9. Relation between maleic anhydrie selectivity and the ratio of selective surface area to non-selective one.

1.5

845 REFERENCES

1. 2. 3. 4. 5. 6. 7.

E. Bordes and P. Courtine, J. Catal., 57(1979)236. G. Centi and F. Trifiro, Chem. Rev., 88(1988)55. G. J. Hutchings, Appl. Catal.,72(1991)1. H. Igarashi, K. Tsuji, T. Okuhara and M. Misono, J. Phys. Chem., 97(1993)7065. G. Busca, F. Cavani, G. Centi and F. Trifiro, J. Catal., 99(1986)400. G. Centi, Catal. Today, 16(1993)5 and references therein. H. S. Horowitz, C. M. Blackstone, A. W. Sleight and G. Teufer, Appl. Catal., 38(1988)193. 8. E. Kesteman, M. Merzouki, B. Taouk, E. Bordes and R. Contractor, Scientific Bases for the Preparation of Heterogeneous Catalysts 6th. Intl. Symp., Poster Session No. 1, 301 9. E. Bordes, Catal. Today, 16(1993)27. 10. T.Miyake and T.Doi, Appl. Catal., 131(1995)43. 11. B. K. Hodnett, Ph. Permanne and B. Delmon, Appl. Catal., 6(1983)231. 12. F. Cavani, G. Centi, F. Trifiro and G. Poll, J. Them. Anal., 30(1985)1241. 13. F. B. Abdelouahab, R. Olier, N. Guilhaume, F. Lefebvre and J. C. Volta, J. Catal., 134(1992)151. 14. M. Abon, K. E. Bere, A. Tuel and P. Delichere, J. Catal., 156(1995)28.

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3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 1997 Elsevier Science B.V.

847

H y d r o x y l a t i o n o f B e n z e n e on Z S M 5 T y p e Catalysts M.H~ifele, A.Reitzmann, E.Klemm and G.Emig Lehrstuhl for Technische Chemie I, Universit~it Erlangen-Ntimberg, Egerlandstr. 3, 91058 Erlangen, Germany

ZSM5 type zeolites were used as catalysts for the one-step synthesis of phenol by benzene partial oxidation with nitrous oxide. Isomorphous substitution of A13+ ions by other trivalent metal ions revealed a high catalytic performance of the H-Ga-ZSM5 in a wide temperature range. Systematic variation of the partial pressures of the reactants led to satisfactory preliminary kinetic models. Deactivation could be reduced by postsynthetic catalyst silylation which is believed to block the strongest acid sites responsible for coke formation. 1. INTRODUCTION With worldwide phenol consumption exceeding 5 million tons in 1995, optimizing production routes of this essential chemical becomes very important. As an alternative to the traditional cumene process, a one-step-synthesis of phenol from benzene is highly desirable. With a ZSM5 type zeolite in its acid form as catalyst and nitrous oxide as oxidant, benzene may be directly oxidized to phenol [ 1-4]: C6H6 + N20

ZSM5)C6HsOH + N2

(Eq. 1)

The vast economical potential has led Monsanto to consider commercializing the process [ 17]. Discovering the high potential of a one step phenol synthesis (Eq. 1), Ono and coworkers [2] improved the original process that was based on vanadium pentoxide catalysts (comparing nitrous oxide and oxygen as oxidants [8]). It is generally agreed upon that ZSM5 type catalysts are outstanding among a wide range of metal oxide and zeolite catalysts in this gasphase-reaction [9-11 ]. Besides H-A1-ZSM5 [ 1] especially H-Fe-ZSM5 and H-A1/Fe-ZSM5 [4] zeolites were used as catalysts. Only little attention has been paid to H-Ga-ZSM5 [5,12] even though element electronegativity and ion size of Ga 3+ suggest a similar behaviour of H-GaZSM5 zeolites compared to H-A1-ZSM5. Depending on the catalysts used, several quite different mechanisms of the hydroxylation reaction have been suggested: The group of Ono [2], Burch and Howitt [1] and Tirel et al. [3] using H-A1-ZSM5 proposed an acid catalyzed mechanism. Panov et al. [4] using iron containing ZSM5 zeolites advocate a redox mechanism mainly catalyzed by iron species. While the suggested mechanisms are quite different, similar experimental results were obtained concerning the dependence on temperature and partial pressures of the reactants [ 1,2,13]. Furthermore, appreciable coke formation led to a loss in catalyst activity in all cases [2,14,15].

848 In addition to the effects of different catalysts on the hydroxylation reaction, which are still discussed controversially in the literature [ 16], the knowledge of the influence of the reaction conditions, especially temperature and reactant partial pressure, is also important. In this paper detailed reaction engineering investigations with a H-Ga-ZSM5 are hence presented [5-7]. 2. EXPERIMENTAL All zeolite samples were synthesized by the hydrothermal method described in detail in [6]. The experiments were performed in a completely automated laboratory setup including an integrally operated plug flow tubular reactor. Reaction components were analyzed by on-line gas chromatography with FID and TCD [5-7]. Table 1 summarizes the reaction conditions for the benzene hydroxylation on the H-Ga-ZSM5 catalyst. Nitrogen was used as balance.

Table 1." Reaction conditions (W=weight o vst, F=total molar flow) catalyst modified residence time ~ total reactor benzene nitrous oxide composition W/F pressure temperature concentration concentration [g'min/mol] [Pa] [~ [%] [%] SiO2/Ga203=80

0-126

103

350--450

2-12.5

4-26

3. RESULTS AND DISCUSSION Similar observations were made with all catalysts used. Hydroxylation occurred with high selectivities and yields to phenol. Except phenol only benzoquinone was produced in higher yields. Formation of CO2 was not observed below 400~ Initial phenol yields were about 25% but deactivation by coke formation quickly led to a decreasing activity. 3.1 Comparing different ZSM5 zeolites Hydroxylation of benzene with nitrous oxide on ZSM5 type zeolites was strongly influenced by catalyst composition and modification, e.g. type of framework metal ion, type and strength of acid sites, molar ratio of Si/A1 and pretreatment conditions [1-5]. Substitution of the framework aluminum by other trivalent metal ions (especially gallium and iron) had a significant influence on reaction performance. Activity increased in the sequence H-AIZSMS.. 0

0 325

350

375

400

425

450

475

T e m p e r a t u r e [~

Figure 4: Influence of reactor temperature on product distribution Phenol was the main product over the whole temperature range, butthe formation of byproducts became important above 400~ (Fig. 4). At this temperature the expected exponential increase of phenol yield versus temperature did not occur because consecutive reactions of phenol lead to the formation of di-hydroxy-benzenes (catechol, resorcinol) and benzoquinone as well as by total oxidation. Surprisingly, no hydroquinone was found, suggesting a fast consecutive oxidation of the para-isomer to benzoquinone [5-7]. Additionally, above 400~ the conversion of nitrous oxide increased because secondary products generated by these further reactions need higher stoichiometric amounts of nitrous oxide (e.g. total oxidation requires 15 molecules of nitrous oxide) [6]. The increase in undesired products, especially the steep rise of total oxidation, confined the maximum operating temperature to 400~ 3.3 Effect of feed concentrations An increase of nitrous oxide feed concentration led to an increase inreaction rate and benzene conversion. On the other hand benzene conversion decreased with increasing benzene feed concentration (Table 2). Phenol yield basically changed in the same way as benzene conversion in all cases. Upon variation of the feed concentrations, the highest phenol production was obtained at 26% nitrous oxide and 12.5% benzene at T=400~ W/F=92 gmin/mol and a catalyst time on stream of 40 minutes [5-7].

851

Table2." Benzene conversion as a function of nitrous oxide and benzene partial pressures (temperature=400 ~ W/F = 92 ~.min/mol, catalyst time-on-stream=40 minutes) feed concentration of nitrous oxide [%] 8.3 16.7 26.0

feed concentration of benzene [%] 4.2 36.1 % 19.3 % 6.6 %

2.1 4.2 12.5

39.8 % 20.7 % 8.4 %

42.6 % 23.8 % 10.7 %

46.8 % 29.6 % 11.7 %

Rather than the relative magnitudes of conversion and yield, the absolute values prove more helpful in understanding the experimental observations. Both the absolute amount of converted benzene and formed phenol increased with increasing nitrous oxide feed. In contrast to decreasing relative benzene conversion in Table 2 the amount of absolute converted benzene increased with increasing benzene feed concentration. Phenol formation displayed an even stronger dependence on the benzene feed concentration [6]. [] 4.2% benzene

912.5% benzene

.2.1% benzene

100 90 ,,...-

H-Ga-ZSM5 W/F: 92 g min/mol T= 400~

80

r o

70

..r

60

o

50

._~ >

4o

"6 "~

3o 20 lO I

I

I

I

I

10

15

20

25

30

Feedconcentration of nitrous oxide [%]

Figure 5 Selectivity to phenol as a function of feed concentration of nitrous oxide at different benzene feed concentration These effects were reflected by the change of selectivity to phenol upon a variation of the feed concentrations in Figure 5: Benzene selectivity to phenol shows the strong influence of benzene feed concentration on the product distribution. Upon increasing the benzene feed concentration from 2.1% to 12.5%, the selectivity to phenol increased from about 45% to nearly 95%. The influence of nitrous oxide feed concentration on selectivity to phenol is less pronounced. An increase of the nitrous oxide partial pressure led to a decreased selectivity to phenol. In the same way selectivity to benzoquinone increased. But due to a stoichiometric consumption of three molecules nitrous oxide per molecule of benzoquinone it is obvious that its selectivity is even more strongly dependent on the nitrous oxide partial pressure [6]. In summary, increasing feed concentration of nitrous oxide led to an increase of reaction rate with only little loss of selectivity to the desired product phenol. Absolute benzene conversion increased with increasing benzene feed concentration especially at high nitrous oxide

852 concentrations. Selectivity to phenol increased with increasing benzene feed concentrations. In the same way this led to a rise in phenol production exceeding the increase of converted benzene. Upon proper variation of these parameters space-time-yields of more than 1 kg phenol per kg catalyst and per hour may be reached. These observations can be explained by considering adsorption effects. Sorption simulation calculations (Cerius 2) resulted in higher benzene adsorption amounts compared to nitrous oxide. Excess benzene concentration on the catalyst surface especially at low nitrous oxide partial pressure leads to a lower influence of benzene partial pressure variation on the reaction rate compared to nitrous oxide. The increase of phenol selectivity with rising benzene feed concentration on the other hand can be explained by competitive adsorption between benzene and phenol. Increasing benzene partial pressure changes the sorption equilibrium, and forces desorption of phenol to play an increasing role. Further hydroxylation and polymerisation reactions are suppressed and selectivity increases. This explanation is supported by adsorption simulation calculations. The effect of benzene concentration on selectivity to benzoquinone is controlled by two competing effects. High benzene feed concentration suppresses benzoquinone formation by forced desorption of phenol, while low benzene feed concentration leads to consecutive reactions of benzoquinone. Both have a comparable influence on phenol selectivity. 4. KINETIC MODELING Based on these reaction engineering results several of kinetic models were developed. The first step was to develop a power law model describing the experimental data. Especially, the activation energy of benzene hydroxylation to phenol was of particular interest. The following reaction scheme was suggested C6H6 + N20 --> C6HsOH + N2

(Eq. 2)

C6HsOH + VN20 N20 --> consecutive products

(Eq. 3)

resulting in a set of differential equations (Eq. 4-6). Because of the measured higher consumption of nitrous oxide for consecutive reactions of phenol, the parameter VN20 was introduced in the modell (Eq. 2, Eq. 5). It is a global factor for NEO-Consumption for all consecutive products and not correlated with the order q refering to xmoin Eq 5. Table 3 summarizes the results of the parameter estimation. All values of the kinetic parameters - reaction orders, rate constants and activation energies - were estimated by nonlinear regression analysis based on numerous experiments. -EA,I

O~C6H 6

&

(~2 O

-

--

O~C6H5OH

&

-

kol

kol

--

-

--

kol

9e

RT

. X m C6H6

(Eq. 4)

n

"

XN20

-EA,1 e RT 9

-EA,2 rn XC6H6 9

n XN20 9 -- VN20

ko2 9

-EA,1 9e

RT

e9

RT

. Xp C6H5OH

Xl~12 9 O

(Eq. 5)

-EA,2 . Xm

C6H6

N20 -- ko2

. X n

e9

RT

. X p

C6HsOH

Xl~12 9 O

(Eq. 6)

Using experimental data at the beginning of the reaction (time-on-stream = 15 min) it can be assumed that deactivation is negligible.

853

This power law model is able to describe the experimental results with good significance (standard deviation in Table 3).

Table 3

Valuesof parameters for the power law model

parameter

value

standard deviation

n

0.26

0.03

m

0.3

0.02

kol

0.955 mol/(g min)

0.136 mol/(g min)

EA,1

40.6 kJ/mol

3.2 kJ/mol

P

0.18

0.08

q

0.43

0.07

k02

55 mol/(g min)

32 mol/(g min)

EA,2

65.4 kJ/mol

7.3 kJ/mol

VN20

3.03

0.36

9benzene

+10%

m "0

o

9

0.25

o

E

0.20

v

t-

s9 ~6 o 0

E

9phenol

a nitrous oxide

0.03

0.30

0.02

t-

0.15

O ,m

~6

0.10

0.01

o

o.o5

E

0.00 0.00

0.10

0.20

mole fraction (experiment)

0.30

0.00 0.00

0.01

0.02

mole fraction (experiment)

0.03

Figure 6 Power law modell" Parity plots of calculated and observed mole fractions Except for a few points, the parity plot of benzene and nitrous oxide mole fractions displays a satisfactory agreement with a maximum deviation of + 10% (Fig. 6 left). The higher deviation between calculated and measured phenol values (Fig. 6 right) stems from the simplicity of the model for the description of more complex consecutive reactions of phenol. Values of reaction orders (Table 3) and competitive sorption effects mentioned in Section 3.3 could be quantified more satisfactorily using a adsorption kinetic model. The model was derived under the assumptions of constant volume reaction, nitrous oxide reacting from the gas phase (because of the small influence of nitrous oxide feed concentration on phenol selectivity), three times higher phenol sorption constant than benzene sorption constant (Kc6HSOH~3KC6H6, as derived from sorption simulation calculations), but without considering the dependency on temperature (reaction temperature 400~ [5,7]. In addition, the consecutive oxidation of phenol to benzoquinone was regarded as a two step reaction with hydroquinone as intermediate, thus avoiding the the need to postulate formulation of a trimolecular collision. It was further assumed that hydroquinone was

854

immediately oxidized to benzoquinone because no hydroquinone was found in the product stream. When modeling the reaction rate at 400~ total oxidation can be neglected. Similar to the above mentioned power law rate model, only the first data points were used for parameter estimation, assuming no deactivation at this time-on-stream. ~PCsH6

&

= -k 1

aPN20 o~

= -kl

9

K C6H6 Pc6H6 9 1 + K C6H6 p C6H6 9

"1-

PN20 9

(Eq. 7)

3 K C6H6 p C6H5OH 9

KC.H. PC6H. PN20 "

1 + Kc.H.

PC.H6 + 3 KC6He PC6H5OH

(Eq. 8)

3. K C6H6 9PC6H5OH-PN20 - 2k2 aPc6H5OH c~

Pc6H6 -I" 3

K C6H6 " P C 6 H 5 O H

9

KC6H6 PC6H6 PN20 9 1 + K 06H6 9Pc6H6 q- 3. K C6H6 9PCsHsOH

= +k 1 -k

1 + K C6H6

2

3. K C6H6 .PC6HsOH

9

9

(Eq. 9)

PN20

1 + K C6H6 "Pc6H6 -'l- 3. K C6H6 "PC6H5OH

With the values from parameter estimation given in Table 4 a satisfactory description of the experimental data with the kinetic model at a temperature of 400~ can be achieved (Fig.7), though the number of parameters is lower compared to the power law model. This demonstrates the particular effect of the competetive adsorption of the aromatic components on the reaction resulting in increasing selectivity at high benzene feed concentrations.

Table 4: Results of parameter estimation at 400~ for the adsorption model parameter

value

standard deviation

kl [mol/(g min)]

1.7E-03

9.1E'05

k2 [mol/(g min)]

1.3E-03

1.5E-04

KC6H6 [1/bar]

176

66

KC6H5OH [1/bar]

528 9phenol

9 benzene & nitrous oxide 30 1 m ll) "O O

+

"o o vE

25

E

3.0 2.5

20

v

~

15

(/) ,.~ ~.__,

I~.

10

a. (~

1.0

"t::

5

1:::

0.5

.~_

,0 ,._,

0 0

10

20

30

partial pressure (e~enment) [kPa]

2.0

0.0 0.0

1.0

2.0

3.0

partial pressure (expedment) [kPa] Figure 7 Adsorption model: Parity plots of calculated and observed mole fractions

855 Especially, the satisfactory prediction of phenol values confirms that the adsorption model provides a better description of the kinetics than the power law model. (Fig. 7, right). 5. REDUCING DEACTIVATION For an industrial application it is necessary to achieve stable space time yields of about 1 kg phenol per hour per kg of active catalyst material [6]. Suppressing the strong deactivation, by adding of oxygen to the feed is not successful because the phenol yield decreases with increasing oxygen concentration [5]. To reduce coke formation a modification of the catalyst properties is necessary. It is assumed that stronger acid sites are responsible for deactivation [18]. A selective poisoning of the strongest acid sites by silylation should hence improve long term stability with little loss of activity towards benzene hydroxylation [6]. Figure 8 shows the improvement in long term stability for H-A1-ZSM5 zeolite with a SiO2 to A1203 ratio of 100 at 425~ As expected the initial activity of the silylated zeolite was reduced compared with the untreated catalyst but phenol yield was higher after 50 minutes time-on-stream. Optimizing this pre-treatment procedure could probably lead to a catalyst with strongly reduced deactivation. [] non-silylated

9silylated

20 18~

14

0 t'-

12

"o9 ._~ >..

8 6

W/F= 92 gmin/mol

06H6; 4.2 % N20:26 %

T= 425 ~

(i.) 10 t"

4

2+ ot 0

t

t

t

I

I

t

I

50

1O0

150

200

250

300

350

400

time-on-stream [rain] Figure 8 Effect of silylation on long term stability of H-A1-ZSM5 (SIO2/A1203 ratio=100) 7. CONCLUSIONS Our investigations showed that H-Ga-ZSM5 is an active catalyst for benzene hydroxylation with nitrous oxide. The high activity observed does not support the hypothesis that the presence of iron, possibly included as impurities during synthesis of zeolite, is necessary for benzene hydroxylation. On the other hand the acid strength of Br~nsted acid sites decreases in the sequence H-A1-ZSM5>H-Ga-ZSM5>H-Fe-ZSM5 and this is in contrast to the activity gain by exchanging aluminum for gallium or iron [5,6]. So in mechanistic questions no conclusive answer is possible. Further investigations to explain the nature of the active sites in the complex zeolitic system are necessary.

856 In contrast to the indistinct nature of active sites, our investigations gave detailed information about the reaction pattern. The influence of reaction temperature on product distributions and phenol yield could be quantified. The appearance of total oxidation products confined the maximum operating temperature to 400~ in the case of H-Ga-ZSM5 catalysts. It seems to be advantageous to produce phenol working at high feed concentrations of nitrous oxide and benzene. High nitrous oxide partial pressure leads to high reaction rate with only little loss in selectivity towards the main product phenol. An increase in benzene partial pressure leads to an increase in both reaction rate and selectivity to phenol. High selectivity at high reaction rate improves the space time yield important for industrial application. In [6] promising space time yields of 1 kg phenol per hour per kg of active catalyst material are reported. Kinetic modeling resulted an initial satisfactory approache to describe the experimental data. While with a power law model, activation energy could be estimated, with the adsorption model the competitive sorption effects at various feed concentrations could be described. The main problems of this new route to phenol are the high costs of nitrous oxide and the strong deactivation of zeolite catalysts due to coke formation. Nevertheless less expensive nitrous oxide may be recovered from the waste gas stream of adipic acid plants, attempts to reduce deactivation by a modification of the zeolite catalysts appear promising. REFERENCES

[1] R.Burch, C.Howitt, Appl.Catal.A, 103 (1993) 135 [2] E.Suzuki, K.Nakashiro, Y.Ono, Chem.Lett.,6(1988)953 [3] M.Gubelmann, P.Tirel, J.Popa, 9th International Zeolite Conference, Montreal, July 1992. [4] G.Panov, A.Kharitonov, V.Sobolev, Appl.Catal.A, 98 (1993) 1 [5] M.H/ifele, A.Reitzmann, D.Roppelt and G.Emig, Erdrl Erdgas Kohle, 12 (1996) 512. [6] M.H/ffele, A.Reitzmann, D.Roppelt and G.Emig, Appl.Catal.A, accepted for publication (1996). [7] A.Reitzmann, M.H~ifele and G.Emig, Trends in Chemical Engineering, Research Trends, Council of Sci. Res. Trivandrum, Vol.3 (1996) 63. [8] M.Iwamoto, J.Hirata, K.Matzukami and S.Kagawa, J.Phys.Chem., 87 (1983) 903. [9] M.Gubelmann and P.-J.Tirel, EP 341165 (1989). [10]R.Burch and C.Howitt, Appl.Catal.A, 86 (1992) 139. [ 11] G.I.Panov, G.A. Sheveleva, A.S.Kharitonov, V.N.Romannikov and L.A.Vostrikova, Appl.Catal. A,82(1992)31. [ 12] M.Gubelmann, J.-M.Popa and P.-J.Tirel, EP 406050 (1990). [ 13] G.I.Panov, A.S.Kharitonov and G.A.Sheveleva, WO 95/27691 (1994). [14]R.Burch, and C.Howitt, Appl.Catal.A, 106 (1993) 167. [15] A.S.Kharitonov, G.A.Sheveleva, G.I.Panov, V.I.Sobolev, Y.A.Paukshtis and V.N.Romannikov, Appl.Catal.A, 98 (1993) 33. [16] V.I.Sobolev, K.A.Dubkov, E.A.Paukshtis, L.V.Pirutko, M.A.Rodkin, A.S.Kharitonov and G.I.Panov, Appl.Catal.A, 141 (1996) 185. [17] Petrochemical News, 1996, 35 (53), 2. [ 18] R.Barrer, R.Jenkins, G.Peeters, Amer.Chem.Soc., Molecular Sieves II, 258-270,1977

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

857

Direct Hydroxylation of Benzene to Phenol by Nitrous Oxide Anthony K. Uriarte a, Mikhail A. Rodkin a. Michael J. Gross a, Aleksandr S. Kharitonov b and Gennady I. Panov b a Fibers Strategic Business Unit, Monsanto, P.O. Box 97, Gonzalez, FI 32560, USA b Boreskov Institute of Catalysis, Novosibirsk, 630090, Russia

1. ADIPIC ACID MANUFACTURE AND NITROUS OXIDE BYPRODUCT About 5 billion p o u n d s per year of adipic acid are m a n u f a c t u r e d worldwide by the nitric acid oxidation of cyclohexanone a n d / o r cyclohexanol (KA). The KA to adipic acid yields are near 94% of theory. Glutaric and succinic acids are the major byproducts and account for most of the yield loss. Monsanto and some other adipic acid producers recover or upgrade these to salable byproducts resulting in an overall KA utilization efficiency that approaches 99%. However, the nitric acid efficiency is lower because approximately 1 mole N20 is produced per mole of adipic acid, in addition to the easily recyclable NOx that is generated as a result of nitric acid reduction. In the early 1990's it was reported that N 2 0 emissions from adipic acid producers could contribute to atmospheric ozone depletion and global warming [1]. It was estimated that adipic acid production may account for up to 10% of the a n n u a l increase in the atmospheric N20. This report sparked an a b a t e m e n t initiative among the major adipic acid producers. Successes have been a n n o u n c e d and implementations are scheduled for 1996-98 [2]. However, Monsanto was already practicing complete a b a t e m e n t by a thermal reduction process and elected to p u r s u e a more desirable path of value-added utilization. Two general areas of utilization were considered: I) oxidation of N20 to NO and s u b s e q u e n t conversion to nitric acid; and 2) use of N20 as a selective oxidant. The latter had the potential of satisfying the criterion of value addition. There were several reports on the selective hydroxylation of aromatics. Because of the economy of scale, the full use of the available N20

858 would be an important factor in the consideration of the potential options. Monsanto's adipic acid production at Pensacola, Florida is over 600 million p o u n d s per year which equates to almost 200 million p o u n d s of N20. For the hydroxylation of aromatics at stoichiometry, this a m o u n t is equivalent to 400 million p o u n d s of phenol. Based upon a projected a n n u a l growth in the phenol m a r k e t of 3-5%, paper studies were initiated to evaluate the use of N20 to produce phenol. 2. HYDROXYLATION OF BENZENE AND NEW ADIPIC ACID P R O C E S S CONCEPT Reports on use of N20 for hydroxylation of benzene to phenol appeared as early as 1983 [3] - M. Iwamoto used V2Os/SiO2 as a catalyst and at 550oC achieved 10% conversion and 70% selectivity of benzene hydroxylation. The other reports that followed concentrated on usage of ZSM-5 type catalyst for this transformation [4-7]; see also a comprehensive review on the subject [8]. A brief s u m m a r y of catalysts used and process p a r a m e t e r s is given in Table 1. Fast catalyst deactivation was reported for most cases and the values of productivity usually refer to initial stages of the reaction. Table 1 Catalyst

V205/SIO2 H-ZSM-5 H-ZSM-5 H-ZSM,5 Fe-ZSM-5

SiO2/A1203 Selectivityof benzene conversion to phenol, % 70 85 NA >90 95 33 98 100 99

Productivity,mmole phenol per gram of catalyst per hour 1.0 0.3 3.2 1.8 3.0

Ref.

[3] [4] [5] [6] [7]

In summary, benzene can be reacted with nitrous oxide in the vapor phase at elevated temperatures over ZSM-5 or similar catalysts to give phenol and nitrogen (eq. (I)).

300-500oc + N20

7_,SM-5

~~,/OH + N2

(1)

859 The reaction has very high selectivity of benzene conversion to phenol (>99%). A further step was taken to incorporate the phenol scheme into an overall adipic acid process. Eq. (2) summarizes one such possibility.

%,.iC00Ii + N20

A--

2}Iz

i (2}

The process would use N20 to hydroxylate benzene to phenol. The phenol would be hydrogenated to cyclohexanone using available technology. The final step is the currently practiced nitric acid oxidation of cyclohexanol and cyclohexanone, which r e t u m s N20 for use in the front end of the process. The stoichiometric balance is close; however, either some additional on-purpose N20 or KA would likely be required for a stand alone plant. The successful commercialization of the overall process concept depended on the viability of the first step which is a breakthrough technology. The data reported in the literature showed high selectivity of benzene conversion to phenol and good productivity. However, the catalyst coked quickly - in most reported cases the catalyst lost its activity in a matter of a few hours. Another problem of the reported chemistry is the low N20-to-phenol selectivity. In fact, the stoichiometry of benzene oxidation to CO2 by N20 implies that 1% of benzene selectivity loss to deep oxidation is accompanied by 15% selectivity loss in N20 conversion. Considering that the supply of nitrous oxide is limited, the efficiency of its utilization is very important for the commercial operation. To develop a viable commercial process, Monsanto and The Boreskov Institute of Catalysis (BIC) formed a joint R&D team in 1994. Understanding of the fundamental aspects of the reaction played an important role in bringing this concept from a lab curiosity to what, we hope, will become an industry standard in short time.

860 3. REACTION MECHANISM The heart of the benzene hydroxylation is the catalyst active site known as alpha-site (a term coined by BIC). The proposed mechanism is outlined is equations (3-5) below:

N20 + ( )a

~

~~,,~OH +

~ O H

(3)

N2 + (O)t:t

a

(o)a

a

._

~~,~O]H

+ ( )a

(4)

{5}

At the initial stage nitrous oxide is believed to decompose on an alpha-site, loading it with a unique form of oxygen, called alpha-oxygen, and releasing dinitrogen. This active alpha-oxygen then reacts with a benzene molecule inserting oxygen into the C-H bond and yielding adsorbed phenol. Desorption of the product is the final stage, which frees the active site for further reaction. How real is the proposed mechanism? It turns out that its stages can be modeled separately. In the absence of benzene, decomposition of nitrous oxide in a closed system at temperatures below 300oC leads only to evolution of nitrogen - all of the released oxygen under these conditions is left on the catalyst in the form of alpha-oxygen. This process was used as a basis for one of the procedures developed for measurement of alpha-site concentration [9]. When a catalyst loaded with alpha-oxygen is isolated from nitrous oxide, cooled and reacted with benzene vapors at or below room temperature, phenol is extracted from the catalyst as the only product, thus supporting a model for the second and third steps of the proposed mechanism [10].

861 4. P R O C E S S CHEMISTRY AND E N G I N E E R I N G

4.1 Catalyst Deactivation Deactivation of the catalyst and benzene selectivity loss to coking were major concerns at the initial stages of the program. Dramatic improvements have been achieved in this a r e a - see Fig. 1

t

Productivity, mmoles of phenol

Time

Fig. 1. Change in catalyst deactivation profile The activity of the first generation catalysts dropped to one-half of their initial value within 3-5 hours. With catalyst and process improvements the catalyst half life has been increased to 3-4 days. Fig. 1 also shows that we were able to change the original exponential decay into linear deactivation with the new catalyst system and process design.

4.2 Catalyst Productivity In the early experiments, the system showed a good productivity of ca. 1 mmole of phenol per gram of catalyst per hour. Further process development led to more t h a n 10-fold increase in productivity. In the current design base case we a s s u m e the average productivity of the catalyst over 48 hours to be above 4 mmole of phenol per gram of catalyst per hour, which is at the top end of the best industrial catalysts.

862 4.3 C a t a l y s t R e g e n e r a t i o n

Once the catalyst activity goes below the acceptable level, the catalyst can be regenerated by passing oxygen-containing gas through the catalyst bed at elevated temperatures, which completely restores the catalyst activity. No performance deterioration has been observed after multiple reactionregeneration cycles (over 100) and longer term effects are under investigation. 4.4 R e a c t o r D e s i g n

Based on the early performance data (rapid catalyst coking and high reaction exothermicity), the initial choice for the process design was a fluidized bed reactor. However, further studies revealed that the reaction selectivity is remarkably insensitive to the temperature rise present in an adiabatic reactor. This observation, along with a substantial improvement in catalyst stability and the need for quick scale-up, led to the selection of a simple adiabatic plug flow reactor design. The pilot plant, Fig 2, has been in operation since May 1996. The unit design includes continuous recycle of the vapor and liquid streams with a full complement of on-line analyzers. (column bypass), _ .......

-~--~I!

!

Air

9 ~.-

Reactor [

Vent

I

:~..~ ....

Nitrogen'-" Feed Phenol Storage

Benzene Storage

Compressor , (Vent Gas Recycle! |

IIR~alyz~l :, |

,

N20 Storage Preheate

____~4 ..........

|

!

-~ ...... ~ ...........

[02 Analyzer I : .,~ . . . . . . . . . . .~

Flg. 2. Simplified scheme of the benzene-to-phenol Pilot Plant in Pensacola, Florida

863

4.5 Overall P r o c e s s P e r f o r m a n c e Table 2 shows typical performance p a r a m e t e r s achieved in the pilot plant d e m o n s t r a t i o n runs. Table 2 Typical performance d a t a Performance p a r a m e t e r s Reaction Temperature, ~ C

400-450

Contact Time, seconds

1-2

Benzene to Phenol, mol %

97-98

Benzene to COx, mol %

0.2-0.3

Benzene to Diols, mol %

1

Nitrous oxide to Phenol, mol %

85

Productivity, mmole P h e n o l / g c a t a l y s t / h r

4

4.6 P r o c e s s Safety The described reaction Offers o u t s t a n d i n g safety features for process design. It is a g a s - p h a s e reaction with very short residence time, therefore there is a minimal inventory of flammable material. Operating conditions have been defined t h a t e n s u r e the whole process is non-flammable throughout: the reaction step, the separation systems a n d the recycling loops. Another safety feature of the proposed technology is the absence of any highly-reactive intermediates. All active catalyst surface species are immediately c o n s u m e d a n d their concentration is minuscule.

4.7 E n v i r o n m e n t a l F e a t u r e s The process h a s a very high selectivity toward the target phenol. Unreacted benzene is completely recycled. The separation is a simple distillation - there is minimal a q u e o u s waste, no inorganic salts a n d some of the by-products can be isolated a n d sold. The total organic waste is expected to be less t h a n 2% of the phenol m a n u f a c t u r e . And last, b u t not least - the process u s e s waste nitrous oxide, a b a t e m e n t of which currently c o n s u m e s n a t u r a l gas a n d emits m u c h more CO2 t h a n is expected to be emitted by the benzene to phenol process.

864

REFERENCES 1. M.H.Thiemens, W.C. Trogler, Science, 251 (1991) 932. 2. R.A. Reimer, C.S. Slaten, M.Seapan, M.W.Lower, P.E.Tomlinson, EnvironmentaI Progress, 13(2)(1994) 134. 3. M. Iwamoto, J.I. Hirata, K. Matsukami, S. Kagawa, J. Phys. Chem., 87 (I 983) 903. E. Suzuki, K. Nakashiro, Y. Ono, Chem. Lett., (1988) 953. 5. G u b e l m a n n , P.J. Tlrel, Eur. Pat. Appl. EP 341165 A1 8 Nov 1989, 4 pp (US Patent No. 5,001,280); M. Gubelmann, J.M. Popa, P.J. Tirel, Eur. Pat. Appl. EP 4 0 6 0 5 0 A2 2 J a n 1991, 9 pp (US Patent No. 5,055,623). 6. R. Burch, C. Howitt, Appl. Catal., A, 86 (1992) 139. 7. A. S. Kharitonov, T.N. Aleksandrova, L.A. Vostrikova, K.G. lone, G.I. Panov, USSR Pat. No. 4 4 4 5 6 4 6 (1989); A. S. Kharitonov, G. I. Panov, K. G. Ione, V. N. Romannikov, G. A. Sheveleva, L. A. Vostrikova, V. I. Sobolev, U.S. US Patent No. 5110995 A 5 May 1992, 8 pp. 8. G. I. Panov, A. S. Kharitonov, V. I. Sobolev, Appl. Catal., A, 98 (1993) 1. 9. G. I. Panov, V. I. Sobolev, A. S. Kharitonov, J. Mol. Cata/., 61 (1990) 85. 10. V.I. Sobolev, A.S. Kharitonov, Ye. A. Paukshtis, G.I. Panov, J. Mo/. Cata/., 84 (I 993) 117. .

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

865

R a p i d c a t a l y t i c o x y g e n a t i o n of h y d r o c a r b o n s w i t h perhalogenated ruthenium porphyrin complexes

John T. Groves*, Kirill V. Shalyaev, Marcella Bonchio and Tommaso Carofiglio Department of Chemistry, Princeton University, Princeton, NJ 08544, USA

Abstract. Perhalogenated ruthenium porphyrins were found to be efficient catalysts for the oxygenation of hydrocarbons including secondary alkanes and benzene in the presence of 2,6-dichloropyridine N-oxide under mild conditions in aprotic media. Up to 15,000 turnovers and rates of 800 TO/min were obtained. A mechanism where Ru(III) - Ru(V) intermediates play an important role is proposed and discussed. The search for new methods for the catalytic oxygenation of hydrocarbons is one of the most important directions of modem chemistry [1]. Among the metalloporphyrinmediated oxidations, ruthenium catalysts display remarkable activity for aerobic oxidations [2] (Fig. 1) and promising reactivity with N20 [3]. O

O O

O

o

o

[ ~ )

Figure 1. Catalytic cycle for olefin epoxidation with dioxygen using ruthenium porphyrins.

~-~r 2

866 Highly efficient oxygenation reactions with ruthenium porphyrin complexes and aromatic N-oxides in the presence of strong mineral acids have been described by Hirobe et al. [4]. We have recently reported that electron deficient perhalogenated ruthenium porphyrins catalyze the oxygenation of a variety of even unreactive substrates under mild conditions (40- 65~ in the presence of 2,6-dichloropyridine N-oxide in aprotic media [5]. Unusually high rates and turnover numbers (TO) were obtained. Carbonyl (5,10,15,20-tetrapentafluorophenylporphyrinato) ruthenium(H) Run(TPFPP)(CO) has shown unusually high activity with 2,6-dichloropyridine N-oxide as the oxygen donor (Table 1). Here we describe further studies of the mechanism of this remarkable process. CO

~

oo

PFPP

,,

~

F Adamantane and cis-decalin were hydroxylated with high selectivity, complete stereoretention, extraordinarily high rates (up to 800 turnovers/min), and high efficiency (up to 15,000 turnovers, Table 1, entries 1-3). Similar conversions were obtained when RuVI(TPFPP)(O)2 and RuVI(TPFPBrsP)(O)2 were used as catalysts. Oxygenation of less reactive substrates such as benzene and cyclohexane proceeded with lower but still significant turnover numbers (100-3,000, Table 1, entries 4-6). Tertiary versus secondary selectivity in adamantane oxidation adjusted for the number of carbons-hydrogen bonds of each kind was above 210. No rearrangement products were detected in cis-decalin hydroxylation. The kinetics of product evolution in a typical reaction of adamantane hydroxylation showed an initial induction period followed by a fast, apparently zero-order phase with the maximum rate and highest efficiencies (Fig. 2). Deviation from linear behavior took place only after 90% oxygen donor and 80% of the substrate had been consumed. When II RuVI(TPFPP)(O)2, prepared by reaction of Ru (TPFPP)(CO) with 3-chloroperbenzoic acid was used as the catalyst, no induction time was detected and zero-order kinetics were observed as well. The well defined and characteristic UV-vis spectra of metalloporphyrins provide an invaluable tool for the mechanistic studies. Thus, monitoring the state of the metalloporphyrin catalysts during the course of both model reactions by UV-vis spectroscopy revealed that the initial form of the catalyst remained the predominant one throughout the oxidation, i.e. in the Run(TPFPP)(CO) catalyzed reaction c.a. 80% of the porphyrin catalyst existed as Run(TPFPP)(CO) and in RVI(TPFPP)(O)2 catalyzed reaction more than 90% of

867 the catalyst was still in the form of RuVI(TPFPP)(O)2 despite the high turnover numbers reached (-- 400 TO). The fact that Run(TPFPP)(CO) demonstrates similar and even higher maximum turnover rate of 4.9 TO/min in adamantane hydroxylation versus 4.0 TO/min for RuVI(TPFPP)(O)2 under the same conditions indicates that an active catalyst species other than RuVI(TPFPP)(O)2 is involved in the fast catalytic hydroxylation. Table 1. Hydrocarbon oxidations a catalyzed by [Run(TPFPP)(CO)]. Time (min.)

Product (% conv.) b

adamantane

20

1-adamantanol (76.2) adamantane- 1,3-diol (7.3)

91

72

2e

adamantane

120

1-adamantanol (61.0) adamantane-l,3-diol (6.0)

97

800

3

cis-decalin

25

(Z)-9-decalol (79.6) (Z)-decal-9,10-diol (4.2)

90

64

4

trans-decalin

60

(E)-9-decalol (25.8) secondary alcohols (4.3) ketones (13.9)

70

4.4

5e

cyclohexane

180

cyclohexanol (1.6) cyclohexanone (6.7)

95

22

6f

benzene

12h

1,4-benzoquinone (13.3)

40

7

1-octene

60

1,2-epoxyoctane (96)

96

11 (36)g

8

1-octene/adamantane

60

1,2-epoxyoctane (54) 1-adamantanol (28)

90

9.5 4.8

9h

cyclohexene

320

cyclohexene oxide (18.2)

-

0.38

#

Substrate

1

Yield c Max. rate d (%) (TO/min)

ii a [substrate] - [pyC12NO] - 0.02 M, [Ru (TPFPP)(CO)] = 50 I~Jl. All reactions in CH2C12 at 65 ~ in sealed containers, b% conversion based on substrate consumed. Products were identified by GC-MS and compared to authentic samples, c % yield based on pyC12NO consumed. d maximum oxidation rate measured as the slope of the zero order phase of the kinetic plot. e [substrate] = [pyC1ENO] = 0.2 M, [Run(TPFPP)(CO)] = 10 ~ M . f [benzene] = 2 M, [pyC1ENO] = O.i02 M, [RulI(TPFPP)(CO)] = 50 ~/I. g [1-octene] = 0.O1 M, [pyC1ENO] = = 0.04 M, [Ru (TPFPP)(CO)] = 50 ~lM, maximum rate = 36 TO/min. h [cyclohexene] = 0.04 M, [pyCl2NO] - 0.02 M, [Run(TPFPP)(CO)] = 50 ~M, reaction was not complete at 320 min.

868

Figure 2. Adamantane Hydroxylation Catalyzed by RuI~(TPFPP)(CO) and RuVI(TPFPP)(O)2, [adamantane] = [pyC1ENO] = 0.02 M, [catalyst] = 50 ~tM, CHEC12, 40~

TO = moles of product/moles of catalyst

400

~

u

v

~

0

300TO CO

200-

100-

0

50

100 Time, min

150

200

Therefore, the classical t r a n s - d i o x o R u ( V I ) - oxoRu(IV) catalytic cycle [2] (Fig. 1) can be ruled out as the primary reaction pathway in case of rapid catalytic oxygenation. The apparent zero-order kinetics observed are consistent with a steady-state catalytic regime accessible from different initial states of ruthenium metalloporphyrin. Indeed, common oxidants, other than aromatic N-oxides, such as iodosylbenzene, magnesium monoperoxyphthalate, Oxone | and tetrabutylammonium periodate produced the t r a n s dioxoRu(VI) species from RuII(TPFPP)(CO) under reaction conditions but were ineffective for the rapid catalysis. A two-electron oxidation of Run(TPFPP)(CO) would produce oxoRu(IV) porphyrin and eventually dioxoRu(VI). What is the alternative pathway for Run(TPFPP)(CO) activation? It is known that ruthenium(II) ~-cation radicals are formed from the corresponding carbonyl compounds by chemical or electrochemical one-electron oxidation [6]. Such species have been shown to undergo intramolecular electron transfer upon axial ligation and removal of CO to give ruthenium(III) porphyrins [6c, 7]. An emerald green solution of Run(TPFPP)(CO) § radical cation with a strong EPR signal (g = 2.00) was quantitatively obtained when Run(TPFPP)(CO) was oxidized with ferric perchlorate in methylene chloride. An EPR signal typical of a ruthenium(HI) species (gll = 2.55, g_l_= 2.05)

869 [5, 6d,e] was detected after the addition of 2,6-1utidine N-oxide to the solution of the radical cation. Photo-stimulation of RuII(TPFPP)(CO) catalyzed reactions with red-orange light (> 560 nm) which shortened the induction period was observed. Interestingly, RuII(TPFPP)(CO) has only residual absorbance in this region of the spectrum but RuH(TPFPP)(CO) +" radical cation shows a strong band at 635 nm. We conclude that the effect of the red-orange light is consistent with photoejection of the carbonyl ligand from the radical-cation to produce a Ru(III) species. We propose that Ru(III) and oxoRu(V) species are the key intermediates in the catalytic cycle of "rapid oxygenation" which can be viewed as a part of the general scheme of diverse oxidative chemistry of ruthenium porphyrins (Fig. 3). The aerobic oxygenation pathway (involving the even oxidation states Ru(II), Ru(IV) and Ru(VI)) we have previously described [2] is shown in the left half of Figure 3. The new fast catalytic process on the right half of the figure reveals chemical interconnectivity between the fast and slow catalytic regimes.

Figure 3. Mechanisms of Ruthenium Porphyrin Oxidation Catalysis

-e-

~

7+"

~

CO

CO

OD 02 D X

SO

OD

so

OD

slow

fast

0

s

~ 1

RDS 0

X

SO O

7+"

S - substrate OD - oxygen donor O Thus, one-electron oxidation of RuVI(TPFPP)(O)2 would give the known dioxoRu(VI) cation radical. Oxygen transfer from this species could enter the fast cycle by producing the proposed transient oxoRu(V) complex [8]. Likewise, a one-electron reduction of dioxoRu(VI) porphyrin would also result in formation of the same reactive species.

870 The competitive oxidation of a 1:1 mixture of adamantane and adamantane-d16 catalyzed by RuVI(TPFPP)(O)2 showed a kinetic isotope effect, kn/kD = 4.8 at 40~ (Fig. 4A). Significantly, the deuterated and undeuterated substrates displayed similar turnover rates (kH/kD= 1.2) for hydroxylation in separate reactions (Fig. 4B). 120

120

A

B

100

100-

80

8060 -

60 TO 40

TO 40-

20

~

d16

II

20-i ~"

I

0

i

10

i

i

I

0

i

i

i

i

i

20 30 40 50 0 10 20 30 40 50 Time, min Time, min Figure 4. (A) Competitive hydroxylation of adamantane/adamantane-d16. [ad.] = [ad.-d16] = 0.01 M, [PyC12NO] = 0.02 M, [RuVI(TPFPP)(O)2] = 50 IxM. (B) Separate hydroxylation [substrate] = 0.02 M, [PyC12NO] = 0.02 M, [RuVX(TPFPP)(O)2] = 50 lxM, CH2C12,40~

m

AH~ = 19 kcal/mol

~

R = 0.9967

9 ~

f

4-

I

n

2

I

2.9

3.0

I

I

3.1 3.2 1/T x 1000, 1/K

I

3.3

3.4

Figure 5. Eyring plot of adamantane hydroxylation catalyzed by RuVI(TPFPP)(O)2.

871 Similar results were obtained with Run(TPFPP)(CO) as the catalyst [5]. A kinetic scheme consistent with this observation is as follows: as long as the oxidant OD is present in excess, the active Ru(III) catalyst will exist as the adduct Ru(III)-OD. The reactive oxoRu(V) would then be formed in the rate determining step, which is independent of the concentration of the oxidant (OD) (Fig. 3). The temperature dependence of adamantane hydroxylation yielded a linear Eyring plot for the observed rate constants determined over the range of 25 - 65~ with an apparent AH-~ = 19 kcal/mol (Fig. 5). Olefins such as 1-octene and cyclohexene demonstrated unusually low turnover rates compared to adamantane and cis-decalin (Table 1, entry 7), although in competitive oxidations 1-octene was twice as reactive as adamantane (Table 1, entry 8). Apparently, the catalyst is inhibited in the presence of olefins. Interestingly, upon a stoichiometric reaction of RuVI(TPFPP)(O)2 with cyclohexene in methylene chloride, rapid transformation of the ruthenium porphyrin into Run(TPFPP)(CO) was observed (Fig. 6). 1.0

o

0.9

co

I< p,pp

0.8 0.7

o

0.6 O

0.5

< 0.4 0.3 0.2 0.1 0.0 450

470

490

510

530

550

570

590

610

630

650

Wavelength, nm

Figure 6. Transformation of RuVI(TPFPP)(O)2 into Run(TPFPP)(CO) in the presence of cyclohexene, [cyclohexene] - 0.04 M, [RuVI(TPFPP)(O)2 ] - 50 I.tM, CH2C12, 40~ 30 min. Trans-dioxoRu(VI) complexes are known to react with olefins according to the classical oxo-transfer mechanism [2] (Fig. 1). The oxoRu(IV) intermediate produced in this process disproportionates readily to give dioxoRu(VI) complex and Ru(II) porphyrin which has strong affinity even towards trace amounts of carbon monoxide. A similar process realized as a side reaction in the "rapid oxygenation" system would constantly and effectively tie up the catalyst in the catalytically inactive form of Run(TPFPP)(CO). Indeed, no noticeable changes had been detected in the UV-vis spectrum of the ruthenium porphyrin during the course of Run(TPFPP)(CO) catalyzed oxidation of cyclohexene.

872

Conclusions High selectivities, rates and yields demonstrated in hydrocarbon oxygenation make the catalytic system of perhalogenated ruthenium porphyrins and 2,6-dichloropyridine N-oxide promising for a number of practical applications. Unusually high potency of the active oxidant observed in the substrate studies and results of the kinetic experiments clearly indicate that an active species other than RuVI(TPFPP)(O)2 is implicated in catalysis. We propose that Ru(III) and oxoRu(V) species are the key intermediates in the catalytic cycle. Relatively low rates of olefin epoxidations are explained by the constant buildup of inactive RuII(TPFPP)(CO) during the reaction. Further studies on the mechanism of this remarkable oxygenation reaction and a search for access to the rapid catalytic cycle from other oxidants are in progress.

Acknowledgements Partial support of this research by the Monsanto corporation and the National Science Foundation for the purchase of an NMR spectrometer are gratefully acknowledged.

References 1. (a) Selective Hydrocarbon Activation: Principle and Progress; Davies, J. A., et al., Eds; VHC: New York, 1994. (b) Activation andfunctionalization of alkanes; Hill, C. L., Ed.; John Wiley & Sons: New York, 1989. (c) Metalloporphyrins in Catalytic Oxidations; Sheldon, R. A., Ed.; Marcell Dekker, Inc.: New York, 1994. 2. (a) Groves, J. T.; Quinn, R. J. Am. Chem. Soc. 1985, 107, 5790-5792. (b) Groves, J. T. and Han, Y. Z. in Cytochrome P450: Structure, Mechanism, and Biochemistry; Ortiz de Montellano, P. R., Ed.; Plenum Press: New York, 1995. 3. Groves, J. T.; Roman, J. S. J. Am. Chem. Soc. 1995, 117, 5594-5595. 4. (a) Ohtake, H.; Higuchi, T.; Hirobe, M. Heterocycles 1995, 40, 867-903. (b) Ohtake, H.; Higuchi, T.; Hirobe, M.J. Am. Chem. Soc. 1992, 114, 10660-10662. 5. Groves, J. T.; Bonchio, M.; Carofiglio, T.; Shalyaev, K. V. J. Am. Chem. Soc., 1996, 118, 8961-8962. 6. (a) Dolphin, D.; James, B. R.; Leung, T. Inorg. Chim. Acta 1983, 79, 25-27. (b) Leung, T.; James, B. R.; Dolphin, D. Inorg. Chim. Acta 1983, 79, 180-181. (c) Barley, M.; Becker, J. Y.; Domazetis, G.; Dolphin, D.; James, B. R. Can. J. Chem. 1983, 61, 2389-2396. (d) James, B. R.; Dolphin, D.; Leung, T. W.; Einstein, F. W.; Willis, A. C. Can. J. Chem. 1984, 62, 1238-1245. (e) James, B. R.; Mikkelsen, S. R.; Leung, T. W.; Williams, G. M., Wong, R. lnorg. Chim. Acta (B) 1984, 85, 209-213. (f) For a review see James, B. R. in Fundamentals of Research in Homogeneous Catalysis, Shilov, A. E. Ed. Gordon Breach, NY, 1986, 309-324. 7. Barley, M. H.; Dolphin, D.; James, B. R. J. Chem. Soc., Chem. Commun. 1984, 1499-1500. 8. (a) Che, C. M.; Ho, C.; Lau, T. C. J. Chem. Soc. Dalton. Trans. 1991, 1259-1263 and references therein. For leading references on non-porphyrin ruthenium oxidation catalysts see also: (b) Dobson, J. C.; Helms, J. H.; Doppelt, P.; Sullivan, P.; Hatfield, W.; Meyer, T. J. Inorg. Chem. 1989, 28, 2200-2204.

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

873

Ethanol Oxidation Using Ozone over Supported Manganese Oxide Catalysts: A n In Situ Laser R a m a n Study Wei Li*, and S. Ted Oyamat, Laboratory for Environmental Catalysis and Materials, Departments of Chemical Engineering and Chemistry, Vir~nia Polytechnic Institute and State University, Blacksburg, VA 24061-0211 Ethanol oxidation using ozone was investigated over ahmina and silica supported manganese oxide catalysts at temperatures t~om 300 to 550 K by in situ laser Raman spectroscopy coupled with reactivity tests. Ethanol was found to react with ozone at lower temperatures than with oxygen, and also with a lower activation energy. This is in agreement with the stronger oxidizing ability of ozone compared to oxygen. The oxidation reactivity was found to be closely related to that of ozone decomposition, suggesting an important role of ozone decomposition in the reaction mechanism In situ laser Raman spectroscopic studies showed the existence of adsorbed ethoxide species on the catalyst surface under reaction conditions, however, at a much lower concentration than when oxygen alone was used as the oxidant. Transient experiments provided direct evidence that surface peroxide (an adsorbed species due to ozone) and surface ethoxide (an adsorbed species due to ethanol) reacted with each other on the catalyst surface.

1. INTRODUCTION Ozone has been attracting increasing attention recently as an alternative oxidant in the oxidation of volatile organic compounds (VOCs) due to its strong oxidizing ability and hence lower reaction temperatures [1-7]. Ozone was generally found to be effective at enhancing the conversion of VOCs, especially at low temperatures. The kinetics of complete oxidation of benzene by ozone on MnO2 was investigated by Naydenov and Mehandjiev [7]. They found that the activation energy for benzene oxidation with ozone (30 kJ mol"1) was much lower than that with oxygen (88 kJ moll), but was similar to that of ozone decomposition (32 kJ moll). It was concluded that the rate determining step for benzene oxidation by ozone was ozone decomposition. Klimova et al. [4] studied the oxidation of lower aliphatic alcohols by ozone over silica and ahnnina. It was found that the main products of the oxidation reactions were: acetaldehyde and carbon dioxide from ethanol, propionaldehyde and carbon dioxide from n-propanol, and acetone l~om isopropanol. It was determined that for the oxidation of 1 mole of the alcohols about 1 mole of ozone or less was consumed. By varying the reactor vohnne while keeping a constant surface area of the catalyst it was demonstrated that the reaction proceeded primarily on the catalyst surface. Spectroscopic

Department of Chemical Engineering $Departments of Chemical Engineeringand Chemistry

874 study of ozone oxidation reactions has been rare (8), and no work has been done under reaction conditions. This paper reports a comparative study of ethanol oxidation reaction with an ozone/oxygen mixture or oxygen alone over supported manganese oxide catalysts using/n situ laser Raman spectroscopy coupled with reactivity measurements. Alumina and silica supported manganese oxide catalysts were chosen because manganese oxide is an excellent catalyst for complete oxidation, while the alumina supported catalyst has a si~ificantly higher activity for ozone decomposition than the silica supported sample. 2. EXPERIMENTAL 2.1. Catalyst Preparation and Characterization

Ahtmina and silica supported MnO2 catalysts were prepared by incipient wetness impregnation of supports (Degussa, Aluminoxide C and Cabosil L-90) using manganese acetate (Aldrich) as the precursor and were calcined in air at 773 K for 3 h prior to use The crystal phases of MnO2 in the catalysts were identified using X-ray diffraction (XRD) and their surface areas were determined by N2 physisorption using the BET method. Ozone Generator 02~

~ Purifier ~

~

Check Valve

Vent MFC

He ~ Purifier MFC

ValveCheck

Carrier GaSGas Chromatography Sample ]Check Valve

r'-~-~l]o'ns]~ [aere .......... ' , ! Porap.ak .Q.S E Vent~9. . . . . . . . . . . . . . . . . . . . . . . . "

Ozone Analyzer Motor Ethanol (Syringe Pump) Heating Wire

Collecting Notch Focusing Lens Filter Lens

Temperature Ar§ Ion Laser Controller (514.5 nm)

Figure 1. Schematic of the in situ laser Raman system

CCD Detector

875

2.2. In Situ Laser Raman Spectroscopy The in situ laser Raman spectroscopic studies were carded out with a high throughput spectrometer using a Raman sample cell (Fig. 1), which allowed the spectrum to be acquired under reaction conditions. The catalysts (200 mg) were pressed into thin wafers 15 mm in diameter and about 1 mm in thickness, and were held on a ceramic rod by a stainless steel cap The rod was spun at 1800 rpm to avoid local overheating by the laser. The temperature was controlled by a programmable temperature controller (Omega, CN2010), and was measured by a thermocouple placed 3 mm away from the sample The laser Raman spectrometer was equipped with an Ar ion laser (Lexel, Model 95, wavelength = 514.5 nm) as the exciting source, a single-stage monochromator (Spex, 500 M) and a charge-coupled device (CCD) detector (Spex, Spectrum One). A key feature of the Raman system was a holographic notch filter (Kaiser, Super-Notch Plus), which effectively rejected Rayleigh scattering, while allowing > 80% of the Raman signal to pass through. Three sets of reactivity tests were performed on the catalysts: ethanol oxidation using an ozone/oxygen mixture, using oxygen alone, and ozone decomposition. The samples were pretreated in situ at 773 K in oxygen for 2 h before each set of measurements. The reaction feed for oxidation using an ozone/oxygen mixture contained 7.8 mol% oxygen, 0.16 mol% ozone, 0.8 tool% ethanol with helium as the balance gas, and the total flow rate was 110 cm3 rain"~ (82 ~ o l s'~). For the oxidation using oxygen alone, the feed composition was essentially the same except that no ozone was used and the total flow rate was kept at 110 cm3 rain"~ by increasing the flow rate of oxygen. For ozone decomposition, no ethanol was injected into the stream and the total flow rate was kept at 110 cm3 rain ~ by increasing the flow rate of hefiun~ Some bypassing of the gas around the sample likely occurred, but the rate data should be accurate for conversions of 10% or less (differential conditions). Ozone was produced by passing oxygen (Air Products, Extra Dry, > 99.6%) through a corona-discharge ozone generator (OREC, Model V5-0), and the inlet, and outlet ozone concentrations were measured using a UV absorption ozone monitor (Safety Caution" Ozone is highly toxic, hence leak checking and purification of the exhaust stream with an ozone decomposition filter should be carried out.) The reaction products were analyzed by an on-line gas chromatograph equipped with thermal conductivity and flame ionization detectors. Separation of 02, CO, and CO2 was achieved by a Carbosphere column (Alltech) while separation of organic compounds was carried out with a Porapak QS column (Alltech).

3. RESULTS

3.1. Catalyst Characterization Although both alumina and silica supported catalysts had 10 wt% loading of manganese oxide, they exhibited considerably different MnO2 crystallinity (Fig. 2). The alumina supported catalyst essentially only had a well-dispersed manganese oxide phase, the silica supported sample showed strong di~action peaks due to crystalline manganese oxide (13MnO2, JCPDS 24-735). The surface area of the ahxmina and silica supported catalysts was found to be 88 and 75 m2 g-1 respectively.

876

3.2. Ethanol Oxidation Reactivity As expected, ethanol was found to be more reactive with ozone than with oxygen (Fig. 3), especially at low temperatures. The reactivity difference became less si~ificant at higher temperatures, and eventually disappeared above 500 K. MnO2/SiO2 showed fimilar behavior (Fig. 4), however, the reactivity difference between ozone and oxygen was more pronounced in this case, and only disappeared around 530 K. For both catalysts CO2 and H20 were the main products, with a small amount of CO produced at higher temperatures (> 450 K), and no organic products were detected.

02 9

I

,

I

,

I

9

'

,

I

SiO2 I

I

I

I

I

10% MnO2/AI203

o

I

I

I

I

I

20

40

60

80

100

20/Degree Figure 2. XRD patterns of the supported MnO2 catalysts and the supports. The measurements of inlet and outlet ozone concentrations allowed the calculation of the ratio of converted ozone to converted ethanol (Fig. 5). Converted Ozone

%Ozone Conversion x Ozone Feed Pate

Converted Ethanol

%Ethanol Conversion • Ethanol Feed Rate

(1)

This ratio were found to decrease from 14 - 10 at 300 K to over unity at > 500 K (not shown). The ratio on the silica-supported sample was slightly higher than that on the alumina-supported catalyst. Because of the absence of crystalline manganese oxide, the reaction turnover rates on MnO2/Al2Eh were calculated as~ming 100% dispersion of manganese oxide. Activation energies of ethanol oxidation using oxygen and an ozone/oxygen mixture were calculated from the Arrhenius plots of the turnover rates (Fig. 6). The activation energy was much higher (89 kJ molq) when oxygen alone was used as the oxidant than when an ozone/oxygen

877 mixture was used. In the latter case the activation energy was only 3.7 kJ mol"1 at lower temperatures (< 400 K), and 48 kJ tool~ at higher temperatures (> 400 K). 40 40 o

~o 30

"~ 30 O

20

-~ 20

Ozone/O~ge~....~

,

gen

n_m__n_____._._i~ ~m~~'O/

300

350

400 450 500 Temperature / K

550

300

Figure 3. Ethanol oxidation reactivity using 03 or 02 on 10% MnO2/AI203.

O

gen

350

400 450 56o Temperature / K

Figure 4. Ethanol oxidation reactivity using 03 or 02 on 10% ]VlllO2/SiO2.

20

n/U~n,.,

~

MnO2/AI203

MnO2/SiO 2

O

~e

O

rj

0

,, , , ,

360

3;0 460 Temperature / K

4~o

Figure 5. Converted ozone to ethanol ratio on the supported MnO2 catalysts

-7a = 48 kJ

-S

b~ ~-9

~~ta

mol-1

= 89 kJ mo1-1

---10

-11 -12 1.8

.... 2.0 2.2

;~,.,. 2.4

EaT3:7.kJ,m~

2.6 2.8 3.0 103 T-1 / K-1

3.2

' 3.4

5;0

3.6

Figure 6. Arrhenius plots of the turnover rate of ethanol oxidation on MnO2/A1203

878 3.3. In Situ Laser Raman Spectroscopic Studies The Raman spectra of the catalyst samples were compared under various conditions (Fig. 7). The spectrum under oxygen exhibited a broad signal peaking at 658 cm-1 (Fig. 7a). With the introduction of ozone the 658 cm"l peak ~hiRed to 634 cmq and a new sharp signal at 878 cmq appeared (Fig. 7b). When ethanol was introduced in the absence of ozone, a new species with Raman signals at 884, 2878, 2930, 2970 cmq was observed which was assigned [9] to an adsorbed ethoxide species (Fig. 7c). Under the coexistence of ethanol and ozone (Fig. 7d), the intensity of both the new signal at 878 cmq and the ethoxide species decreased dramatically, and the signal at 658 cmq shifted to 640 cmq.

2936

640

2930

d

2970 A

c ~

b b a I

I

I

,

I

3000 2900 2800 R a m a n S h i f t / c m -1

I

,

I

~

I

~

I

,

I

1000 800 600 400 200 Raman S h i f t / c m -1

Figure 7. Raman spectra of 10% MnO2/AI203 sample under various conditions (a. oxygen; b. ozone; c. ethanol; d. ethanol and ozone). In addition to steady state experiments, transient measurements were also performed on the ahtmina-supported catalyst to investigate the interaction between the adsorbed species. The transient experiments were carded out by starting with a surface with preadsorbed ethoxide species. Ozone was then suddenly introduced and a set of Raman spectra was acquired at regular time intervals. The intensity of the 878 cm"l peak increased with time, while the adsorbed ethoxide peaks in the higher wavenumber region (2800 - 3200 cm"l) decreased with time (Fig. 8).

879

150 rain 120 rain

150rain ~

90rain

120 rain

r~

60 rain

90 rain

30rain

60rain 30rain

0min

0min

31'00 ' 3 0 ~ " 29~ ' 2s'o0

'12100'1000'

Raman Shift/G~I1-1

800 ' 600 ' 400

Raman Shift/cm-1

Figure 8. Transient experiment Raman spectra on 10% MnO2/A1203. 3.4. Ozone Decomposition Ozone decomposition reactivity was measured to investigate its role in the ethanol oxidation reaction (Fig. 9). On MnO2/AI203, with increasing temperatures ozone conversion increased slowly at temperatures < 400 K, then increased sharply above 400 K, and finally approached 100% around 500 K. While on MnO2/SiO2, ozone conversion was lower at low temperatures (< 320 K), but steadily increased with increasing temperatures, and reached 100% at a similar temperature (500 K) as on MnO2/A1203. The remover rates on MnO2/AI203 were also calculated assuming 100% dispersion of manganese oxide, and the Arrhenius plot showed two kinetic regions (Fig. 10). The lower temperature region which was dominated by the catalytic decomposition on MnO2 gave an activation energy of 3.2 kJ tool1, while the higher temperature region which was dominated by the gas phase decomposition gave an activation energy of 41 kJ m o r I. 100

-~r

o 80

-7.6

/;;:F

-7.8 a

41 kJmolq

b~ -8.0 ~

_~ -s.2

- . -

20

0

o~

- - - ~o~/~a~o~

./ 3;0

' 3;0

4;0

' 4;0 ' 560 Temperature / K

' 5;0

Figure 9. Ozone decomposition reactMty on supported Mn02 catalysts.

-8.4

E a = 3.2 kJmo1-1

-8.6 212'214'2:6'218'310'312'314' 103 T -1 / K d

Figure 10. Arrh~ttius plots of turnover rate of ozone decomposition o n MnOjA1203.

3.6

880 4. DISCUSSION

4.1. Reaction Stoichiometry For most reactions, reaction stoichiometry is easy to determine, however, this is not the case for oxidation reactions using ozone. One reason is that it is not clear whether oxygen is also involved in the reaction as ozone is usually used as an ozone/oxygen mixture. Another reason is that it is uncertain how many oxygen atoms of each ozone molecule contribute to the oxidation reaction. If only ozone is involved in the reaction, the reaction equation can be written as (assuming only CO2 and H20 are produced):

GH, OH + 203 = 2CQ + 3H20

(2)

However, when each ozone molecule only contn'butes one oxygen atom while producing an oxygen molecule, the reaction equation becomes: q

r,o/7 + 603 = 2 c Q + 6Q +

(3)

In addition to these reactions the ozone decomposition reaction may occur independently. 03 = 3/202

(4)

By measuring the inlet and outlet ozone concentrations, the converted ozone to ethanol ratio was determined, which may give insights on the reaction stoichiometry. The low ratio value at high temperatures (> 400 K) was consistent with a si~ifieant involvement of molecular oxygen in the reaction. In other words, ethanol could initially be activated by ozone to form intermediate oxidation products (aldehydes, earboxylic acids, etc.), and these could subsequently be oxidized by oxygen. At even higher temperatures (> 500 K), the oxidation was likely dominated by oxygen because of the rapid decomposition of ozone and because of the higher activation energy of ethanol oxidation by oxygen compared to that by ozone. At lower temperatures when the oxygen involvement was probably low, the measured ratios ranged fIom 10 to 20, which were closer to the stoichiometry of equation 3. The measured ratio on the silica supported catalyst was found to be slightly higher than that on the alu~mina catalyst, which is consistent with the observation that the silica catalyst was not as active as the a~mina one for ozone decomposition. In a study on the oxidation of lower aliphatic alcohols by ozone on alumina and silica from 293 and 363 K [4], it was found that the converted ethanol to ozone ratio was about unity. However, in that study the main oxidation products were acetaldehyde and ketones, while the only products observed on the supported manganese catalysts of this study were carbon dioxide and water.

4.2. In Situ Laser Raman Spectroscopic Studies Spectroscopic studies of the interaction between ozone and organic molecules on catalyst surfaces have been very rare. Mariey et al. [8] reported a Fourier transform in~ared (FFIR) study of ozone interaction with phenol adsorbed on silica and celia. Ozone was found to be reactive toward phenol and carboxylic acids and aldehydes were detected as possible intermediates. However, the study was carded out from 77 to 220 K, which is far from

881 reaction conditions (usually > 300 K) for those surfaces. Recent work from this laboratory has demonstrated that in situ laser Raman spectroscopy is an excellent tool to study surface intermediates at reaction conditions [10,11]. The signal at 650 cm-1 under oxygen can be ascribed to Mn304, however, this observation does not mean it is the only phase present. Other phases like, and Mn20 3, MnO show only very weak Raman signals while MnO 2 is completely Raman inactive. The introduction of ozone formed an adsorbed peroxide species with a Raman signal at 878 cm"l, while the introduction of ethanol generated an adsorbed ethoxide species with Raman signals at 884, 2878, 2930, 2970 cml. Under reaction conditions with the coexistence of ozone and ethanol, the intensities of both these adsorbed species dramatically decreased, indicating that these two species reacted with each other on the catalyst surface. This was also supported by the transient experiment results. When ozone was introduced on a surface preadsorbed with ethoxide species, the intensity of the ethoxide species decreased gradually due to the reaction with ozone (gas phase or adsorbed), and that of the peroxide species increased with time due to the removal of ethoxide species from the surface. However, if the reaction of ethoxide species was mainly due to gas phase ozone, under steady state conditions, the surface should be covered by adsorbed peroxide species. The in situ Raman spectra indicated that the reaction of ethoxide species was primarily due to reaction with adsorbed peroxide species because the concentrations of both adsorbed species decreased dramatically in the presence of both ethanol and ozone. Thus, a Langmuir-Hinshelwood type mechanism appears to be operating:

o3 + 2 " EtOH + 2 " E t o * + o~ *

2 0 * (or 0 2 . ) O*(orQ*)

+ 2H*

) o~*+o*

(5)

> EtO * + H *

(6)

:~" > c 0 2

) 02 + * (or Z*)

:,,n ." 1_120

(7) (8) (9)

Overall the results are consistent with both ethanol and ozone adsorption being the slow steps in the overall reaction. The surface reaction and desorption steps are fast, and result in a surface that is close to bare.

4.3. Role of Ozone Decomposition As indicated in equation 3, ozone decomposition can be closely associated with the oxidation reactions. It is well accepted that ozone decomposes to produce active oxygen species, which can activate the organic molecules at lower temperatures. It has also been reported by several groups [7, 12, 13] that at low temperatures the activation energy for oxidation reactions by ozone are similar to that of ozone decomposition, which suggests that the rate determining step for the oxidation reaction is probably ozone decomposition. The results from this study are consistent with that conclusion. However, ozone decomposition likely involves steps (5) and (8). Since at low coverage step (5) will be the rate-determining process, the measured activation energy probably corresponds to that step. In addition, the in situ laser Raman spectroscopic study provides direct evidence that surface ethoxide reacted with peroxide species formed by ozone decomposition

882 5. CONCLUSIONS Ethanol oxidation using ozone was investigated over supported manganese oxide catalysts at temperatures from 300 to 550 K by in sire laser Raman spectroscopy coupled with reactivity measurements. Ethanol was found to react with ozone at lower temperatures than with oxygen, and also with a lower activation energy. This is in agreement with the stronger oxidizing ability of ozone compared to oxygen. The oxidation reactivity was found to be closely related to that of ozone decomposition, suggesting an important role of ozone decomposition in the reaction mechanism In sire laser Raman spectroscopic studies provided direct evidence for the reaction between the peroxide (due to ozone) and the ethoxide species (due to ethanol) on the catalyst surfaces. ACKNOWLEDGMENT We gratefully acknowledge the financial support for this work by the Director, Division of Chemical and Thermal System of the National Science Foundation, under Grant CTS9311876. REFERENCES 1. A. Gervasini, G.C. Vezzoli, and V. gagaini, Catal. Today, 29 (1996) 449. 2. A. Gervasini, C.L. Bianchi, and V. gagaini, in Environmental Catalysis, J.N. Armor (ed.), ACS Symp. Set. 552; ACS: Washington, DC, 1994, 352. 3. W. Li and S.T. Oyama, in Heterogeneous Hydrocarbon Oxidation, B.I~ Warren and S.T. Oyama (eds.), ACS S3anp. Set. 638; ACS: Washington, DC, 1996, 364. 4. M. N. Klimova, B.I. Tarunin, and Yu.A. Aleksandrov, Kinet. Katal., 26 (1988) 1143. 5. K, Hauffe and Y. Ishikawa, Chem. Ing. Techn. 5 (1974) 1035. 6. V. Ragaini, C.L. Bianchi, G. Forcella, and A. Gervasini, in Trends in Ecological Physical Chemistry, Bonati, L. (eds.) Elsevier: Amesterdam, 1993, 275. 7. A. Naydenov and D. Mehandjiev, Appl. Catal., A 97 (1993) 17. 8. L. Mariey, J. Lamotte, J.C. Lavalley, N.M. Tsyganenko, and A.A. Tsyganenko, Catal. Lett., 41 (1996) 209. 9. W. Zhang and S.T. Oyama, J. Phys. Chem., 99 (1995) 19468. 10. W. Zhang and S.T. Oyama, J. Phys. Chem, 100 (1996) 10759. 11. W. Zhang and S.T. Oyama, J. Am. Chem See., 118 (1996) 7173. 12. N.A. Kleimenov and A.B. Nalbandian, Proceedings of Academy of Science of USSR, 122 (1958) 635. 13. 1L Del gosse, C. Mazzocchia, and P. Centoh, React. Kinet. Catal. Lett., 5 (1976) 245.

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

883

Generation o f singlet oxygen from the catalytic system H202/Ca(On)2 and applications to the selective oxidation o f tmsaturated compounds J. M. Aubry* and V. Nardello Equipe de Recherches sur les Radicaux Libres et l'Oxyg~ne Singulet, URA CNRS 351 Facult~ de Pharmacie de Lille, 3 rue du Professeur Laguesse, BP 83 F-59006 Lille Cedex, France

A new chemical source of singlet molecular oxygen (Io2, lAg), H202/Ca(OH)2, has been investigated in detail. First, the formation of ~O2 has been proved unambiguously by resorting both to the specific detection of the IR luminescence at 1270 nm of Ioz and to the chemical trapping of the excited species with a new cationic water-soluble trap. The process has been shown to be catalytic and the influence of several parameters (pH, concentrations and purities of reactants) on the initial rate of ~O2 formation has been examined. Finally, the ability of the system H20~/Ca(OH)2 to oxidize various water-soluble electron-rich substrates has been assessed.

INTRODUCTION

Excited molecular singlet oxygen, (IO2, ~Ag) is a powerful and highly selective oxidant very useful in organic synthesis. ~'2 The usual way to generate IO2 is photochemical but many chemical reactions are also able to produce this excited species by decomposition of a peroxo compound. 3 In 1985, we showed that about thirty inorganic oxides, hydroxides or oxo-anions induce the decomposition of hydrogen peroxide into ~O2.4 Among all these new chemical sources of ~O2, the homogeneous system H202/MoO42 appeared particularly efficient since it generates quantitatively a high flux of 102 at room temperature (Eq. 1) :

2H2Oz

MoO~" water, pH 10

>

2H=O + 'O2

(1)

Since then, much work has been devoted to this system and both the mechanism of the reactions and the ab'~ty to oxidize various organic substrates, either in water6 or in microemulsions,7 have been investigated in detail.

884 The screening experiments of the periodic classification revealed another particularly interesting chemical source of IO2 9the system H202/Ca(OH)2 (Eq. 2) which appears both attractive for its environmental friendly feature and amazing with regard to the simplicity of the catalystinvolved.

2 H20~,

Ca(OH)= water >

2 H20 + ~ ~0 2 +(1 - (x) z02

(2)

For these reasom, we have undertaken a detailed study of this system by confirming, in a first step, the generation of 102 by chemical trapping and by detection of its specific IRluminescence. Then, the kinetics of the reaction were examined by studying the influence of several parameters such as pH and concentrations and purities of reactants on the initial rate of IO2 formation. Finally, the ability of this system to oxidize various water-soluble electron-rich substrates such as polycyclic aromatic, cyclohexadienic and acrylic derivatives was assessed.

1. EXPERIMENTAL PART 1.1. Reagents Calcium hydroxide (98 %), calcium oxide (99.9 and 99.995 %), and tiglic acid 7 were purchased from Aldrich Chemie. Calcium peroxide was l~om Air Liquide. Hydrogen peroxide (50 % Rectapur) was from Prolabo, Paris. Deuterium oxide (98 % D) was l~om CEA (Commissariat ~ l'Energie Atomique, Saclay). bis-(4'-trimethylphenylammorfium)-9,10anthracene dichloride 1 (BPAA), s potassium 9,10-anthracene dipropionate 3 (ADP) 9 and sodium 1,3-cyclohexadiene-l,4-diethanoate (CHDDE) ~~5 were prepared according to known procedures. 1.2. Instrumentation Disappearance of BPAA was monitored by UV/visl'ble spectroscopy at 373 nm with a Milton Roy Spectronic 3000 spectrophotometer. High performance liquid chromatography (HPLC) analyses were carried out with a CAIson pmnp model 303 by using a 25-cm cohmm packed with Spherisorb RP18-50DS. A mixture of H20 and CH3OH was used as eluent and UV detection was performed with a variable-wavelength monitor (Gilson Holochrom H/MD). The nuclear magnetic resonance spectra (~H and 13C NMR) were obtained on a Bruker AC 300P FT-spectrometer. 1.3. Singlet Molecular Oxygen Monomol Emission (1270 nm) Infrared emission of 102 was measured with a liquid nitrogen cooled germanium photodiode detector (Model EO-817 L. North Coast Scientific co., Santa Rosa. CA) sensitive in the spectral region from 800 nm to 1800 nm with a detector of 0.25 cm2 and a saphire window. Measurements were carried out in a cuvette with mirrored walls (35 m m x 6 m m x 55 nun).

885

1.4. Chemical Detection of Singlet Molecular Oxygen 90 ~tl H202 30 % (0.2 M) were added, under stirring, to an aqueous solution (I)20) of BPAA 1 (10.2 M) and CaO 99.995 % (0.1 M). The disappearance of BPAA was monitored by UV spectroscopy at 373 nm from acidified (H3PO4) samples and the final reaction medium was analyzed by ~H and ~3C NMR spectroscopy alter centrifugation. 1.5. Deuterium Solvent Effect 25 ~tl H202 50 % (0.1 M) were added under stirring to an aqueous solution (I)20 or H20) of BPAA (10"3 M) and CaO 99.995 % (0.05 M). The disappearance of BPAA was monitored by UV spectroscopy at 373 nm from acidified (H3PO4) samples. A similar procedure was used for the study of the influence of pH, concentrations and purities of reactants. 1.6. Peroxidation of Tigfic Acid 7 5.75 ml H202 50 % (10.0 M) were added under stirring to an aqueous solution (H20) containing 1 g tiglic acid 7 (1 M) and 1.136 g CaO 99.99 % (2 M) at 35 ~ The pH of the heterogenous reaction medium was equal to 7.75. The course of the reaction was followed by HPLC at 210 nm from filtered aliquots (1 rrd) and the final reaction medium was analyzed after 4 h by 1HNMR spectroscopy. Comparisonwith a genuine sample obtained photochemicaUy showed that 50 % of the corresponding hydroperoxide 8 was formed. A similar procedure was used for the peroxidation of ADP 3 and CHDDE 5. The experimental conditions are given in Table 1.

2. RESULTS AND DISCUSSION. 2.1. Evidences for the generation of singlet oxygen (~Oz, tag) from the system HzO2/Ca(OH)2 The first aim of our study was to prove unambiguously the generation of 102 from the system H202/Ca(OH)2. One method used was the chemical trapping of this excited species. Since the current water-soluble chemical traps of IO2 bear carboxylate or sulfonate functions, II which are likely to interact with the calcium ions, a new cationic water-soluble trap was designed:the bis-(4'-trimethylphenylammonium)-9,10-anthracene dichloride 1 (BPAA). This trap reacts efficiently with IO2 giving the endoperoxide BPAAO2, 2 (Eq. 3). This trapping process com~tes with the solvent-induced quenching of 102 (Eq. 4).

886

N(CH3) cr

N(CH3)"cr

(3)

,~ (so oc)

! N(CH3) cr

(CH3) cr

BPAA, 1

BPAAO2, 2

102

water

>

302

(4)

The formation of BPAAO2 was confirmed both by ~H and ~3C NMR spectroscopy by comparison with a genuine sample photochemically obtained and by thennolysis at 80 ~ of the final reaction mixture which led to the regeneration of the initial trap according to a specific property of some polycyclic aromatic endoperoxides. 12However, although the formation of the endoperoxide BPPAO2 constitutes a convincing proof of 102 involvement, it is useful to strengthen this result by another independent test based on the longer lifetime of the excited species in deuterated water (67 I~s in D20 instead of only 4.4 ~ in H20). ~3 In order to understand an increase in rate in D~O, it is a s s u n ~ that the main process is the deactivation of ~O2 by solvent i.e. BPAA must be introduced at a concentration well below [3 = l~/k, where Iq is the pseudo-first order rate constant for the deactivation of IO2 by water (Eq. 4) and 1~ is the second-order rate constant for the chemical reaction between '02 and BPAA (Eq. 3). Therefore, two experiments were carded out by using 10.3 M BPAA to the system H202/Ca(OH)~ in D20 or in H20. The disappearance of BPAA, monitored by UV spectroscopy at 373 nm, is reported in Figure 1. Comparison of the initial rates of the disappearance of BPAA and of the final amounts of BPAAO2 leads to an increase in rate by a factor of 13 when the reaction is conducted in deuterated water, giving support for a process involving 102.

887

1 0,8 rb

----"--'-*

A

E e.

,.'~ 0,6 t~ v 0


200 12.1 35.5 16.0 0.6

Mo-Ce(sg) 20-40 19.0 34.0 37.0 0.9

Mo-Ce oxide particles were not changed. This suggested that the ultrafine complex oxide particles were quite stable during the catalytic oxidation. After reaction for 7h, the catalytic oxidation reached a steady state. The catalytic properties of CeO2, MoO3 and two Mo-Ce oxides catalysts (Ce/Mo atomic ratio =1) are given in Table 1. It can be seen from these results that under the same reaction conditions, the toluene conversion on the above oxides catalysts showed the order as CeO2 >Mo-Ce(cp)-~Mo-Ce(sg) >MOO3. The oxidation products on mono-component CeO2 catalyst were CO,CO2 and H20 but without benzaldehyde or other selective oxidation products, indicating that CeO2 is an active component for complete oxidation of toluene. By adding ceria to MOO3, however, the selectivity to benzaldehyde was remarkably improved, so that the complex oxides showed higher catalytic selectivity to benzaldehyde than the mono-component MoO3 catalyst. Interesting, the conversions of toluene on both complex Mo-Ce oxides were very similar, but the benzaldehyde

906

Cl

b .I

300

I

500

1

700 Temperature, ~

I

900 ~.

t

1100

Figure 3. TPR profiles of the mono-component MoO3 and the complex Mo-Ce oxide. (a) MoO3, (b) complex Mo-Ce oxide.

selectivity of the ultrafine particle catalyst was much higher than that of the larger particles. These results have revealed that ultrafine oxide particle catalysts have unique catalytic properties for the selective oxidation of toluene to benzaldehyde. In particular, the specific activity (benzaldehyde yield) of the ultrafine complex oxide particles was higher than those of the larger complex oxide particles catalyst and the mono-component MoO3 catalyst, which

907 can not be explained by the effect of mere particle size. Apparently, the differences in the nature of active species in the oxide catalysts should be taken into consideration. As pointed out by Haber[1-2], the lattice oxygen ions in molybdenum oxides are the main active species for selective oxidation of lower olefins to aldehydes. The state of lattice oxygen species in the mono-component MoO3 oxide and Mo-Ce complex oxides were studied by using TPR and LRS. As can be seen in Figure 3, which shows the TPR profiles of MoO3 and Mo-Ce complex oxide, the hydrogen consumption peaks for MoO3 appeared at 670, 755 and 990~ By adding ceria to MOO3, the hydrogen consumption peaks shift to lower temperatures: 510, 715 and 860 ~ These results suggest that due to interaction of Ce with Mo in the complex oxides, the molybdenum oxides are easier to reduce to lower valance. By using IR spectroscopy, Jonson et al.[13] showed that the vibrational frequency of Mo=O in MoO3 was 995cmq. Figure 4 shows the Raman bands of Mo=O species in the complex MoCe oxides. For the complex Mo-Ce(cp) oxide, the vibrational frequency of Mo=O was 953cmq, which is lower than that of Mo=O in MOO3. The red shift of vibrational frequency indicates a weaker bonding between Mo=O in the complex Mo-Ce oxide. This is consistent

953

1100

1000

900

850

Rarnan shift, cmq

Figure 4. Laser raman spectra of Mo=O species in the complex Mo-Ce oxide catalysts prepared by different methods. (a) Mo-Ce(cp), (b) Mo-Ce (sg).

908 with the TPR result. As the lattice oxygen species in the complex Mo-Ce oxides have higher mobility, their reactivity for selective oxidation of toluene are increased. This can account for the higher specific activity of complex Mo-Ce oxide catalysts for oxidation of toluene to benzaldehyde. Moreover, it can be seen from Figure 4 that the vibrational frequency of Mo=O species in Mo-Ce(sg) was much lower than in Mo-Ce(cp), indicating that Mo=O chemical bonding in the ultrafine oxide particles was even weaker and the lattice oxygen ions in ultrafine complex oxide particles have a higher mobility. The higher reactivity of lattice oxygen in the matrix of ultra:fine Mo-Ce complex oxides can account for the reason that ultrafine complex oxide particles exhibit unique catalytic properties for selective oxidation of toluene to benzaldehyde. Our results have clearly confirmed the reactivity of lattice oxygen ions can be improved by decreasing the oxide particle size to nano-scale. These results suggest that the ultrafine complex oxide particles are potentially new catalytic materials for selective oxidation reactions. Reference

1. J. Haber, in Proceedings of the 8th Intemational Congress on Catalysis, July, 1984, Berlin (Wes0, Vol. 1, Verlag Chemie, 1984, P. 85. 2. J. Haber, in Heterogeneous Hydrocarbon Oxidation, ACS Symposium Series 638, American Chemistry Society, 1996, P. 20. 3. Y. Moro-oka and W. Ueda, Adv. Catal., 40 (1994) 233. 4. K. Van der Wide and P. J. Van den Berg, J. Catal., 39 (1975) 473. 5. J. Buiten, J. Catal., 21 (1968) 188. 6. S. Tan, Y. Moro-oka and A. Ozaki, J. Catal., 17 (1970) 125. 7. M. Ai and T. Ikama, J. Catal., 40 (1975) 203. 8. K.A. Reddy and L. K. Doraiswamy, Chem. Eng. Sci., 24 (1969) 1415. 9. Z. Yan and S. Lars T. Andersson, J. Catal., 131 (1991) 350. 10. C. Simon, R. Bredesen, H. Gronadal, A. G. hustoit and E. Tangstad, J. Mater. Sci., 30(1995) 5554. 11. K. Maede, F. Mizukami, M. Watanabe, N. Arai, S. Niwa, M. Toba and K. Shimizu, J. Mater. Sci. Lett., 9 (1990) 522. 12. J. Y. Guo, F. Gitzhofer and M. I. Boulos, J. Mater. Sci., 30 (1995) 5589. 13. B. Jonson, R. Larsson and B. Rebenstoff, J. Catal., 102 (1986) 29.

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

909

Liquid Phase Oxidation of Aikylaromatic Hydrocarbons over Titanium Silicalites Georgi N. Vayssilov a, Zdravka Popova a, Stefanka Bratinova b and Alain Tuel c aFaculty of Chemistry, University of Sofia, 1, J. Bourchier avenue., Sofia 1126, Bulgaria e-mail: [email protected] bNational Centre for Environment and Sustainable Development, Sofia 1453, Bulgaria CCNRS, Institut de Recherches sur la Catalyse, 2, avenue A.Einstein, 69626 Villeurbartr, e Cedex, France 1.INTRODUCTION Partial catalytic oxidation of alkylaromatic hydrocarbons is interesting both from the industrial and the scientific point of view. The industrial interest is due to the availability of these substances from the petrochemical industry and to a number of applications for the possible oxidation products. Conventional gas phase oxidation concerns the side chain and leads mainly to benzoic acid or even to destruction of the aromatic ring. Oxidation under mild conditions could cease the reaction at earlier stages and reduce the number of the products formed. However, the appropriate catalyst for such partial oxidation has not been found yet. A promising step in this direction is the oxidation with hydrogen peroxide over titanium silicalite molecular sieves TS-1. This catalytic system was successfully applied in the hydroxylation of phenol, epoxidation of alkenes, and partial oxidation of alkanes and other organic substances [1-4]. The advantages of the TS-1 catalyst reside in the selectivity of the reaction and the long duration of the activity of the catalytic samples [1-5]. The catalytic oxidation of toluene and ethylbenzene using H20 2 over TS-1 was mentioned in one of the initial works on titanium silicalites [ 1]. In both cases hydroxylation of the aromatic ring was observed while only the alkyl chain of the ethylbenzene was oxidised to 1-phenylethanol and acetophenone. So far, only Khouw et al. [6] and Mal et al. [7] reported the quantitative data for ethylbenzene oxidation. Under the reaction conditions used they achieved quite different turnover numbers (TON) - 52 and 8.4, respectively. The product distribution also depends considerably on the reaction conditions. Clerici [8] reported formation of acetophenone and a minor amount of ethylphenols, while the other authors found between 23 and 40 % ethylphenols [1,6,7]. Another difference was the conversion to 1-phenylethanol - from 4% to 33% [6,7]. From the other monoalkylbenzenes, only 2-propyl - and 1-butylbenzene were tested in liquid phase oxidation with H20 2 over TS-1 [8]. The former hydrocarbon was not reactive under the reaction conditions, while the side chain of the latter one was oxidised to corresponding alcohols and ketones at ot and ), positions to the phenyl group. It is clear that there are only occasional studies on the partial oxidation of alkylaromatic hydrocarbons with hydrogen peroxide in the presence of titanium silicalites. This is the first

910 work directed exclusively to the investigation of toluene, ethylbenzene, 1- and 2propylbenzene oxidation over TS-1 and TS-2 catalysts under various reaction conditions. Attention is mainly focused on the evolution of the product distribution (ring hydroxylation vs. side-chain oxidation, alcohol vs. ketone in the side chain) during the initial hours of the catalytic process. 2.EXPERIMENTAL

2.1.Catalyst Titanium silicalite-1 (TS-1) was synthesised following Example 1 of the original patent [9] and titanium silicalite-2 (TS-2), according to ref. [10]. The removal of occluded organics was performed by calcination of the sample at 550~ in air for about 6 h. Chemical analysis of the solid materials (Si/Ti= 50) was obtained by atomic absorption after solubilization of the sample in HF-HCI solution. UV-vis spectrum of the calcinated material showed an adsorption band at about 205 nm, characteristic of titanium silicalites. The absence of signal beyond 275 nm indicated that the material was free from extraframework oxide species. S.E.M. pictures revealed that the sample was in the form of very small uniform crystals of about 0.3 ~tm in size. Before each catalytic run the catalyst sample was heated at 820 K for 5 h.

2.2.Catalytic experiments Hydrocarbons used in catalytic experiments were obtained by Fluka - ethylbenzene and 1propylbenzene, and M e r c k - toluene and 2-propylbenzene. Catalytic experiments were performed in a 50-ml round-bottomed flask equipped with magnetic stirrer and condenser. The initial mixture containing 0.1 g catalyst, 5.0 ml solvent and 2.0 ml substrate was stirred and heated up to the reaction temperature. Then 1.0 ml 30% aqueous solution of hydrogen peroxide (Merck) was added. Small samples (0.1 ml) of the reaction mixture were taken after 1, 2 and 4 h, diluted in 4.0 ml methanol, dried and analysed by GC-MS (Hewlett-Packard). Veratrole was used as internal standard. Hydrogen peroxide conversion was followed by standard iodometric titration. 3.RESULTS AND DISCUSSION

3.1.General observations There are two routes to the partial oxidation of alkylaromatic hydrocarbons - oxidation of the side chain or of the aromatic ring. In both cases the oxidation could proceed at different positions (except for the methyl group of toluene) and to different extents. Titanium silicalites were found to activate the oxidation of the secondary or tertiary saturated carbon atoms, the terminal methyl groups remaining unaffected [2,6-8]. This behaviour limits the number of the

911

CH3

CH3

CH3

CH2OH

HO

CHO

HO-CHCH3 O--CCH3 + ~~/OH

C2H5.......~ .~

~o

~m m o H~~m c H o

HO'cHCH2CH3 O=CCH2CH3 OH CHEC2H5 CH2CHCH3 I

O CHECCH3 II

HO

OH H3c-C-CH3 ~ I

CH(CH3)2 CH(CH3)2 ~ ( C ~ +

" ~ HOHEC-CHCH30HC-CCH3

Scheme 1. Expectedreactionroutesto the partial oxidationof alkylaromatichydrocarbons.

912 expected products from the side-chain oxidation (Scheme 1). The aromatic ring itself could be oxidised at o-, m-, and p- positions with respect to the alkyl group, but due to the electrophilic mechanism of the reaction [4,11,12] the m-isomer can hardly be expected. Although the obtained alkylphenols should be more reactive than the initial alkylbenzenes, their further hydroxylation and oxidation to alkylbenzoquinones is hindered for steric reasons. Since hydroxyl and methyl groups have similar dimensions, the lower conversions of p-xylene and p-ethyltoluene compared to toluene and ethylbenzene [6] suggest the effect of the steric restrictions in catalytic oxidation over TS-1. Products of oxidation both in the ring and in the aliphatic chain, could also be formed in principle. The results of the catalytic experiments follow in general the above assumptions. However, some of the expected products are not observed. Conversions and selectivity in the oxidation of the alkylbenzenes studied are presented in Table 1. In all cases hydroxylation in the aromatic ring is observed, mainly to p-isomers. This p-selectivity is due not only to the substrate molecules used but also to the properties of the solvent molecules. It is known that phenol and anisole hydroxylation in methanol and ethanol exhibits substantial p-selectivity while in non-alcoholic solvents, such as acetonitrile and acetone, o-selectivity is found [5,1315]. The p/o ratio increases with the bulkiness of the alkyl group but, surprisingly, even 2propylbenzene is hydroxylated to some extent at the o-position. Table 1 Conversion and selectivity in oxidation of alkylaromatic hydrocarbon with H202 over TS-1. Alkyl Group

Conversion, %

Oxidation in the ring, %

Side-chain oxidation, %

methyl

1.4

100

0

ethyl

3.5

11

82

1-propyl

2.8

53

47

2-propyl

2.1

100

0

Reaction conditions as described in the Experimental section, solvent- ethanol. Although the intermediates formed during the oxidation of the methyl group in toluene could easily stabilize by resonance with the aromatic ring, the corresponding benzyl alcohol and benzaldehyde are not found. This implies that the simple energetics of the reaction does not determine the absence of terminal methyl group activation over titanium-containing zeolites [6]. Some other metallo-silicate molecular sieves, such as VS-1 and Sn-Sil-1, can activate oxidation of primary carbon atoms under the same reaction conditions [7,16]. Mal et al. [7] suggest that this is due to the different reaction mechanisms over these metal centers via peroxo-radical intermediates over V and Sn and via (electrophilic attack and) carbenium ion formation over Ti. Following this scheme, toluene should be oxidised in the methyl group but, since this is not the case, one should look for another explanation. Simple geometrical factors can also be excluded because Sn4+ and Ti 4+ have similar ionic radii - 71 and 68 pm,

913 respectively [17], and all three catalysts have MFI structure. A difference between these ions resides in the occupation and possibly spatial orientation of the frontal orbitals, especially dorbitals. Another reason could be the relatively higher stability of the doubly reduced Sn2+ and V 3+ metal ions compared to Ti 2+. Side-chain oxidation is observed with ethylbenzene and 1-propylbenzene, while the aliphatic part in 2-propylbenzene is not affected. The lack of side-chain oxidation at the tertiary C atom is probably due to steric hindrance by the two nonreactive methyl groups in addition to the bulky benzene ring. Only the first (benzylic) carbon atom in the ethyl and 1propyl groups undergoes oxidation. The preferential activation of this C atom in the 1-propyl group, compared to the other secondary carbon atom, can be explained by the weak C-H bond at benzylic position and/or by mesomeric stabilization of the intermediate formed (probably carbenium ion [7]). Unfortunately, our investigation cannot answer the hot question why titanium silicalites do not activate terminal methyl groups. However, it clearly shows that although the energetic factors could not explain this peculiarity, they play an important role in the discrimination of different secondary C atoms available for activation. As seen in Table 1, the lower conversion of 1-propylbenzene, compared to ethylbenzene, is due to a lower oxidation of the alkyl chain. While the ratio between the alcohol and ketone formed from 1-propylbenzene is about 1, in the case of ethylbenzene the alcohol represents more than 90 % of the side-chain oxidation product.

3.2.Oxidation of ethylbenzene Ethylbenzene was chosen for the further detailed study of the influence of various reaction conditions (amounts of oxidant, catalyst or solvent (methanol or ethanol), and the framework structure of the catalyst - TS-1 or TS-2) on the conversion and selectivity of the reaction. Ethylbenzene is convenient because the three types of selectivity could be followed (ring vs. side-chain oxidation, p/o ratio in the ring hydroxylation, alcohol/ketone ratio in the sidechain oxidation) and the conversion is the highest. Table 2 shows the reagent concentrations and the product distribution in some of the experiments after 4 h reaction. Turnover numbers found are about one-half of the result of Khouw et al. [6] - TON=52 for 12 h reaction, and higher than the TON=8.4 reported by Mal et al. [7]. The highest conversion is achieved in methanol or without solvent - 5.1 and 4.4 molar %, respectively. However, without solvent the reaction proceeds much faster (Vin= 21 h -1) and after the first hour only 0.4 % of the substratum is additionally oxidised to acetophenone or p-ethylphenol. The higher reaction rate in the beginning and the subsequent interruption of the processes are probably due to the higher temperature - 357 K, while in methanol the reaction was carried out at 342 K. The smaller TON (calculated for ethylbenzene oxidation) and the lower conversion in ethanol is due to the oxidation of the solvent itself, as already reported [18]. The main product of ethanol oxidation is diethylacetal formed by initial oxidation of the alcohol to acetaldehyde and further nucleophilic addition and substitution with other alcohol molecules. Some amount of acetic acid and acetaldehyde are also observed in the reaction mixture. However, the

914 Table 2 Reagent concentrations and product distribution for ethylbenzene oxidation with hydrogen peroxide over TS-1 and TS-2. TS-1 a

TS-1

TS-1

TS-1

TS-2

Mperoxide

mol/1

1.1

0.7

1.1

2.9

1.1

Msubstrate

mol/1

2.1

1.3

2.1

5.6

2.1

ml- 1

3.3

2.1

3.3

8.8

3.3

K

350

349

342

357

350

ethanol

ethanol b

methanol

-

ethanol

19.4

16.7

26.7

24.0

13.9

h- 1

12.2

4.6

11.3

21.1

5.3

1-Phenylethanol

%

77.6

80.3

91.5

83.6

51.6

Acetophenone

%

4.3

11.7

3.8

8.7

7.9

o-Ethylphenol

%

0.8

0.8

0.6

0.4

0.2

p-Ethylphenol

%

10.4

6.7

2.4

6.4

2.8

Others

%

6.8

0.5

1.8

0.8

37.6

Mperoxide/Ti ions Temperature Solvent TON c Initial rate Vind Product distribution

a The standard reaction condition described in the Experimental section. The molar ratio substrate/peroxide=1.9 and substrate/Ti ions=510 are the same in these experiments. b An experiment in diluted solution c TON achieved after 4 h experiment, mol ethylbenzene converted per mol Ti ions d Initial reaction rate during the first hour, mol ethylbenzene converted per mol Ti ions for 1 h. ethylbenzene oxidation in ethanol is a good example that the TS-1 catalyst activates oxidation of hydrocarbons even in excess of alcohol molecules, which are usually much more reactive. Moreover, two-fold dilution of the initial reaction mixture leads to similar TON and conversion as under the standard conditions, described in the experimental section. Of course, in ethanol, the yield with respect to the converted hydrogen peroxide is lower than in methanol (about 5 % in ethanol and 50 % in methanol). Figure 1 shows the evolution of the selectivity to oxidation in the ring and in the alkyl chain during the first hours of the reaction. A very high selectivity to side-chain oxidation is observed in methanol - up to 97 % at the 4th hour. The main product is 1-phenylethanol which amounts to 93 % of the entire conversion. The maximal conversion of ethylbenzene to 1phenylethanol achieved after 4 h reaction is 4.7 % (Fig.2). The ring hydroxylation selectivity

915

100 8O o

o

m Others

60

D Side-chain oxidation

r

40

m Ring oxidation

.,,-4

20

a

b

c

d

Solvents Figure 1. Evolution of the selectivity (1, 2 and 4 h reaction) of ring oxidation, side-chain oxidation and products of deeper oxidation over TS-1. Experiments in ethanol (a), two-fold diluted in ethanol (b), in methanol (c), and without solvent (d).

5-

40

o

-40

30

30 A

z 9 20

20 ._~

o2

.~_

10--

10

O

"~ 1 ~D

>

o0

'

1

I

2

'

I

'

3

i

4

Reaction time, h Figure 2. Time dependence of the ethylbenzene oxidation to 1-phenylethanol over TS-1 in ethanol (rhombus), two-fold diluted in ethanol (squares) and in methanol (triangles).

0

' I ' u ' u ' 1

2

3

4

0 5

Ti ions concentration, mmol/1 Figure 3. TON (open symbols) and initial reaction rate (filled symbols) for ethylbenzene oxidation in ethanol (squares) with variation of the catalyst amount. Rhombus correspond to two-fold diluted solution and triangles to methanol.

916 decreases with time and the portion of acetophenone in the side-chain oxidation product slightly increases. Ring hydroxylation leads preferably to the p-isomer with 88 % selectivity for the first hour and 81% for 4 h reaction. This decrease of the p/o ratio with time is opposite to the trend observed in phenol and anisole hydroxylation over TS-1 [13,15] both in alcoholic and non-alcoholic solvents. Qualitatively, the selectivity of ethylbenzene oxidation in ethanol is the same as in methanol. However, in ethanol the ring hydroxylation selectivity increases up to 12 % (with more than 90 % p-selectivity), and selectivity to 1-phenylethanol decreases, especially for the reaction in dilute solution. The catalytic reaction is performed by varying the peroxide and catalyst concentrations. The TON (Fig. 3) increases using a smaller amount of the catalyst and for 1.0 mmol/l Ti ions (1/4 of the standard amount) TON 35.6 is greater than in methanol (with 4.1 mmol/1 Ti ions). Ring hydroxylation selectivity in the experiments with lower catalyst concentration is lower, compared to the reaction under the standard conditions, but higher than that in methanol. In addition to the lower conversion of the substrate when the amounts of the catalyst and oxidant are reduced, the main difference in these cases is the change in the selectivity to ring oxidation vs. side-chain oxidation (Fig.4). Higher ring-oxidation selectivity is achieved at lower peroxide concentrations - up to 15 % of the product. The factor determining this increase could not be the peroxide concentration in the reaction mixture itself since in the experiment with two-fold diluted reaction mixture the ring-oxidation selectivity is much lower. Probably the important factors for ring hydroxylation are the ratio between peroxide and titanium concentrations or the changes of the solvent properties due to the smaller amount of water in the mixture. 15

,.

~2 0 0

..

.

I I I I I I 2 4 6 8 10 12 Peroxide concentration/Ti ions, 1/ml

'~ 14

Figure 4. Selectivity to ring oxidation in the experiments with variation of the peroxide concentration per Ti ion (1/ml) for 1 h reaction (rhombus) and 4 h reaction (triangles). The TS-2 catalyst shows a lower total conversion than TS-1 with the same titanium content but the selectivity to side-chain oxidation and especially to acetophenone is much higher - up to 35 % of the product after 1 h reaction. While 1-phenylethanol concentration increases with

917 time, the amount of acetophenone remain the same. This is related to the substantial increase of the products of further oxidation of acetophenone (see "others" in Table 2) - Fig.5. A small amount of such products is observed also over the TS-1 catalyst. 1.0 0.8 ~J

U 1-Pheny lethan~ ~

0.6

o

l Acetophenone

r/) ~D ""

o

0 .4

U Others

0.2 0.0 lh

2h

4h

Figure 5. Conversion of ethylbenzene to 1-phenylethanol, acetophenone and products of deeper oxidation (others) over TS-2 catalyst after 1, 2 and 4 h reaction. The deeper oxidation of ethylbenzene over TS-2 can be explained with the slower diffusion of 1-phenylethanol and acetophenone formed in the zeolite pores where they could undergo additional oxidation to acetophenone or other products, respectively. Another possible reason could be some differences in the local geometry of the titanium sites due to the different framework structure of the two titanium silicalites. In addition to the above mentioned products, traces of benzaldehyde are detected. The amount of benzaldehyde increases in the course of the reaction and is higher over TS-2 catalyst. It is probably formed after a C-C bond break in the side chain of 1-phenylethanol or acetophenone. Another by-product - ethylbenzoquinone, is observed in the experiments without solvent and in ethanol under standard conditions. The ethers derived from 1phenylethanol and the solvent- methanol or ethanol, are also found.

4.SUMMARY Titanium silicalites TS-1 and TS-2 catalyze hydroxylation in the aromatic ring of the monoalkylbenzenes studied to corresponding alkylphenols, using hydrogen peroxide as oxidant. Para-isomers are mainly formed in methanol or ethanol as solvents. In the case of ethyl- and 1-propylbenzenes, the first carbon atom of the aliphatic chain is also oxidised both to alcohols and ketones. As expected, the terminal methyl groups in all hydrocarbons are not oxidised. The probable reasons for this behaviour of titanium silicalites are discussed. The influence of the reaction conditions on the conversion and selectivity is studied by ethylbenzene oxidation. Methanol leads to higher conversion than ethanol under the same

918 reaction conditions and after 4 h the reached value is 5.1% with TON 26.7. In both solvents the reaction is selective to side chain oxidation. The highest selectivity to ring hydroxylation (15 %) is found in ethanol, at low peroxide concentration per unit titanium ion. The molar ratio of the alcohol to ketone formed is higher than 9 over TS-1 while over TS-2 this ratio is 1.5. The decrease of the amounts of oxidant and catalyst leads to an almost proportional decrease of the conversion to all products. Suggestions concerning the explanations of the observed correlations are made. Liquid phase catalytic oxidation of ethylbenzene with hydrogen peroxide over TS-1 molecular sieves is most appropriate for the production of 1-phenylethanol with high selectivity (up to 93 % of all the oxidation products in methanol) under the reaction conditions studied here. An additional increase of the 1-phenylethanol selectivity could be achieved with smaller amounts of the catalyst. The highest conversion to acetophenone is found over TS-2 zeolites but further oxidation easily takes place in this case.

Acknowledgement This work was supported in part by the Bulgarian National Science Foundation.

REFERENCES 1.U.Romano, A.Esposito, F.Maspero, C.Neri, M.G.Clerici, Chem.&lnd 72 (1990) 610 2.D.R.C.Huybrechts, L.de Bruycker, P.A.Jacobs, J.Mol. Catal. 71 (1992) 184 3.S.Gortier and A.Tuel, Appl.Catal.A 118 (1994) 173 4.B.Notari, Adv.Catal. 41 (1996) 253 5.G.N.Vayssilov, M.Yankov, Z.Popova, L.Dimitrov, Proc. 15th Conference on Cataysis in Organic Reactions, Phoenix, USA, 1994, p.443 6.C.B.Khouw, C.B.Dartt, J.A.Labinger and M.E.Davis, J. Catal. 149 (1994) 195. 7.N.K.Mal and A.V.Ramaswamy, Appl. Catal. 143 (1996) 75. 8.M.G.Clerici, Appl. Catal. 68 (1991 ) 249 9.M. Taramasso, G. Perego and B. Notari, US Pat 4 410 501 (1983). 10.J.S. Reddy and R. Kumar, Zeolites 12 (1992) 95. 11 .J.S.Reddy, S.Sivasanker, P.Ratnasamy, J.Mol. Catal. 71 (1992) 373. 12.G.Bellussi and M.S.Rigutto, Stud.SurfSci.Catal. 85 (1994) 177. 13.A.Tuel, S.M.-Khouzam, Y.Ben Taarit, C.Naccache, J.Mol. Catal. 68 (1991) 45. 14.J.A.Martents, Ph.Buskens, P.A.Jacobs, A.van der Pol, C.Ferrini, H.W.Kouwenhoven, P.J.Kooyman, H.van Bekkum, Appl. Catal. A 99 (1993) 71. 15.G.N.Vayssilov, Z.Popova and A.Tuel, Chem.Eng. Technol., in press 16.P.R.Hari Prasad Rao, A.V.Ramaswamy, P.Ratnasamy, J. Catal. 143 (1993) 275. 17.F.A.Cotton and G.Wilkinson, Advanced Inorganic Chemistry, Wiley-Interscience, New York, 1972, vol.II. 18.F.Maspero and U.Romano, J.Catal. 146 (1994)476.

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

919

Coupled Vanadyl Centres in Vanadium Phosphorus Oxide Catalysts: Essential Structural Units for Effective Catalytic Performance in the Ammoxidation of Methylaromatics A. BrOckner *, A. Martin, B. LOcke and F. K. Hannour

+

Institut lhr Angewandte Chemie Berlin-Adlershof e. V., Rudower Chaussee 5, D-12484 Berlin, Germany

ABSTRACT The ammoxidation of substituted toluenes over differently prepared (NH4)2(VO)3(P207)2and VO(PO3)2-phases as well as over (VO)2P207 has been studied by catalytic and in s i t u - E S R measurements. For effective catalytic performance at least two structural properties were found to be essential: i) Closely neighbouring VO 2+ centres must be exposed at the surface which enable the simultaneous adsorption and conversion of the substrate and ii) the catalyst structure must contain building blocks of exchange-coupled VO 2+ ions e. g. in the form of chains or layers which support the electron transport during the redox process.

I. INTRODUCTION In addition to a variety of transition metal oxide catalysts also vanadium phosphorus oxides (VPO) proved to be suitable catalysts for the ammoxidation of substituted toluenes to the corresponding benzonitriles [1 ]. Although numerous papers on the mechanism of the ammoxidation reaction have been published in the past the various reaction steps, the nature of the active sites and, in particular, structural properties required for effective catalytic performance are still ambiguous. I n s i t u - i n v e s t i g a t i o n s under working conditions should lead to a better understanding of the relation between structural peculiarities and catalytic properties. However, only few methods such as electron microscopy [2], Raman [3] and FTIR spectroscopy [4] have been developed to perform such measurements so far. In this study we demonstrate how m situ - electron spin resonance (ESR) can be used to study the behaviour of unsupported VPO catalysts during the ammoxidation of substituted toluenes in a special self-constructed in situ - ESR flow reactor. From the results, conclusions on the arrangement of structural units required for effective catalytic performance were derived.

+

On leave from the Department of Chemical Engineering, University of Technology of Compi6gne, France

920 2. EXPERIMENTAL

2.1. Catalysts (VO)2P207 was prepared by calcining the precursor VOHPO4 0.5 H20 in nitrogen atmosphere for 2 hours at 753 K. The precursor was obtained by evaporating an aqueous solution ofV2Os, H3PO4 and oxalic acid as described elsewhere [5]. Pure crystalline a-(NH4)2(VO)3(P2002 (NVP~) was prepared by calcining a mixture of V205 and (Nt~)2HPO4 (V:P:N - 1:3:6) at 573 K in air [6]. An equilibrated catalyst (N ,o) consisting of ~- 75 mol-% a-(NI-hh(VO)3(P207h and = 25 mol-% of an amorphous vanadium oxide phase was obtained after treating VOHPO4 0.5 H20 under ammoxidation conditions for 40 h [e. g. 6]. Two further catalysts were prepared by mixing NVP~ mechanically with solid NH4VO3 (NVPm) and by impregnating NVP~ with an aqueous solution of NH4VO3 (N i) both in a molar ratio of NVP~/NH4VO3 = 1 followed by 3h formation under ammoxidation conditions. Two VO(PO3)2-samples were obtained by calcining the precursor VO(H2PO4)2 [7] in air at 673 K for 6 h (amorphous phase) and at 1073 K for 1 h (crystalline t-phase). Crystalline aVO(PO3)2 was prepared by heating V205 and H3PO4 at 643 K for 5 h [8]. As a reference sample, V409 was prepared by heating a finely divided stoichiometfic mixture of VzO5 and sulfur at 453 K for 65 h in an autoclave [9]. The vanadium valence state was determined by potentiometric titration [ 10]. 2.2. Methods Catalytic activities and selectivities were determined separately in a fixed-bed U-tube quartz-glass reactor (8 mm i. d.) under similar conditions. Toluene conversion and benzonitrile formation were followed by on-line GC using a FID detector. The carbon oxides were measured by non dispersive infrared photometry. ESR spectra were recorded with the cw spectrometer ERS 300 in X-band (Zentrum fttr WissenschafUichen Geratebau Berlin). For in situ - measurements, i. e., investigation of the catalyst under reaction conditions in the cavity of the spectrometer, a self-constructed ESR flow reactor equipped with a bifilar heating winding of Pt wire and connected to a gas/liquid supplying system was used [11]. In each run 400 mg catalyst particles (0.5 - 1 mm) were treated with a gas flow of 980 ml h"~ (molar ratio : aromatic hydrocarbon : air : NH3 : H20 = 1 931 94.5 924.3) at atmospheric pressure. 3. RESULTS AND DISCUSSION Suitable VPO-catalysts for the ammoxidation of toluene can be obtained on dehydrating the precursor VOHPO4 0.5 H20. When this dehydration is done in inert atmosphere crystalline (VOhPzO7 is obtained as the only product. However, when the precursor is calcined under the ammoxidation feedstock (air, NH3, hydrocarbon, HzO vapour) a catalyst is formed which contains crystalline (NH4)2(VO)3(P2OT)2 comprising about 75 % of the overall vanadium content along with an amorphous vanadium oxide phase. Both catalysts were found to catalyze the ammoxidation of toluene effectively [e. g. 1]. However, in the latter case it was not clear whether crystalline (NI-h)2(VO)3(P2OT)2 or the additional amorphous phase is the catalytically active component. Therefore, in addition to pure crystalline (NI-h)2(VO)3(P2OT)2 two further samples were tested catalytically in which NVP~ had been modified by mixing with solid NH4VO3 and by impregnating with an aqueous solution of NH4VO3, respectively, to generate an additional amorphous vanadium oxide phase as formed in NVP,o, too. In the following

921 sections the results of catalytic and ESR-spectroscopic investigations of these catalysts and three different VO(PO3)2 samples are described.

3.1. Catalytic Results The benzonitrile selectivities for the various catalysts proved to be rather similar. They varied slightly between 80 % and 100 % depending on the degree of conversion. However, marked differences were found in the activities. From the plot of the area-specific toluene conversion as a function of reaction temperature (Fig. 1) it can be seen that (NH4h(VO)3(P2OTh catalysts are most active when they contain an additional amorphous vanadium oxide phase generated either by in situ-formation during catalysis or by treatment with NH4VOa while pure crystalline (NH4)2(VO)a(P207)2 is rather inactive. This suggests that not the crystalline but the amorphous phase in (NH4)2(VO)a(P207)2-based catalysts is the catalytically active 400 componem. (VO)2P207 is a good catalyst, too, even though E 300not as active as the three e-

"~

200-

~

1000 6O0

(Nn4)2(VO)3(e2OT)2-basea catalysts. In the case of vanadyl

I

620

I

640

I

660

I

680

I

700

I

720

I

740

T/K

Fig. 1 Area specific toluene conversion, Cto~,for NVPi (I),

NVPm (#), NVPao (0), NVPsyn (A), (VO)2P207 (O), a(A), fl- ( ) and amorphous VO(POa)2 (r'l)

polyphosphates it is interesting to note that amorphous VO(POa)2, in particular at higher temperatures, is markedly more active as the respective crystalline a- and t3VO(PO3)2. To get more information on the reasons of the different catalytic behaviour the catalysts were investigated under reaction conditions using in situ-ESR.

3.2. ESR-Investigations The ESR spectra of all samples consist of one single isotropic line. It arises from the VO 2§ centres of the structure and is narrowed due to effective spin-spin exchange interactions between them. For the same reason, the anisotropic g and his tensors are not resolved. The efficiency of this spin-spin exchange depends strongly on the arrangement of the neighbouring VO 2§centres within the catalyst structure as well as on changes of their electron density caused by the catalytic reaction. As we demonstrated recently [12], evaluation of the temperature dependence of the ESR signal intensity and the line shape leads to two parameters which can be used to characterize quantitatively the exchange efficiency: i) the exchange energy, AE, and ii) the quotient of the 4th and the square of the 2nd moment of the ESR signal, /2 Both parameters have been used in this work to analyze structural differences between the various catalysts and their behaviour under working conditions. In the crystal structure of the precursor, VOHPO4 0.5 H20, the vanadyl ions are incorporated as exchange-coupled dimers of VO6 octahedra. During the dehydration in inert or ammoxidation atmosphere, respectively, these dimers are cleaved and the VO6 octahedra rearrange to form infinite ladderlike double chains in (VO)2P207 (formed in inert atmosphere) and single chains in (NH4)2(VO)a(P207)2(formed in ammoxidation atmosphere). When this trans-

922 formation is followed by in situ-ESR [12, 13] a strong line broadening is observed at the transition temperature which is caused by the temporal collapse of the spin-spin exchange. The moment quotient, /2, passes a minimum before it increases again when new exchange pathways are established during the formation of the catalyst structures. Thus, CoNaY, Co(N03)2 > CoPc.

100

Figure 3. Kinetic curves of trans-stilbene epoxide accumulation for aerobic oxidation of trans-stilbene (0.3 mmol) in the presence of IBA (1.14 mmol) and Co(II) catalysts (6.10 .3 mmol): COW12 (A), PWllCO (O), Co(NO3)2 (11), CoPc ((3) and CoNaY (~) (11 mg)

80 60 O

40

20 0 0 1 2 3 4 5 6 7 8 9 10 11 Reaction time, h

It was assumed that high activity of Co(II) compounds in alkene-aldehyde co-oxidation is due to the ability of cobalt to form complexes with dioxygen, thus providing its activation [ 14, 26]. Indeed, numerous Co(II) chelates are known to form stable superoxo and pperoxocomplexes by reaction with dioxygen [ 1-4]. Note that among the catalysts studied, only CoPc is known to coordinate 02 [1, 3 and references therein]. Nevertheless, it showed less catalytic activity as compared to the other Co(II) catalysts employed (Fig. 3). By analogy with manganese and iron complexes, the formation of active CoIV--o or CoV=O species by the twoelectron oxidation of Co(II) or Co(III) ions with peroxy acid was suggested by few authors [15, 26]. It is well known that appropriate ligands are needed to stabilize higher oxidation states of transition metals [1-4]. Alkene epoxidation by PhIO in the presence of PWllCO and PWllMn was proposed to proceed most likely via the formation of the high valent metal-oxo species [2, 18, 32]. The formation of such species may in principle be expected for CoPc, but

952 it seems to be unlikely for Co(N03)2, CoNaY and impossible at all for CoWl2, in which Co(II) occupies the center of the HPA and hence is not capable to form oxo-species. Remember that CoWl2, contrary to PWllCO, was found to be inactive in the alkene epoxidation by PhIO [18]. UV-Vis data confirm the stability of the Keggin structure of the HPA during the reaction with O2/IBA [33]. The fact that Co(NO3)2, CoNaY and CoWl2 appeared to be good catalysts for the epoxidation with OJIBA indicate that the ability of Co(II) to form complexes with dioxygen or to form metal-oxo species is not the determinant reason of the high catalytic activity of cobalt in Mukaiyama's system. Moreover, the pronounced activity of CoWl2, the well-known reagent of the outer-sphere electron transfer processes [34], shows that preliminary coordination of dioxygen and/or aldehyde is not necessary for the reaction to proceed successfully. Table 3 PIBA decomposition in the presence of Co(II) catalysts (a) Catalyst

PWllCO

COW12

Co(NO3)2 CoNaY (c)

CoPc

PIBA conversion(b) (%)

68

73

39

23

40

(a) Reaction conditions: PIBA 0.28 mmol, catalyst 6.10 .3 mmol, MeCN 6 ml, 24~ Cb)determined by iodometric titration after 2 h; (c) CoNaY 11 mg (Co 3.29%). To clarify the mechanism of the catalytic action of Co(II) compounds we have studied their activity in decomposition of perisobutyric acid (PIBA), formation of which was detected by IH NMR during alkene-IBA co-oxidation [16, 17, 19]. The ability of Co(II) compounds to mediate homolytic decomposition of peroxy acids was mentioned earlier [19, 25, 26, 35, 36]. The data on the activity of Co(II) compounds in the PIBA decomposition are summarized in Table 3. PWllCO and CoWl2 display the highest activity, while CoPc does the lowest one. The catalytic activity of Co(II) compounds in the PIBA decomposition seems to correlate with the rates of the epoxidation in O2/IBA/Co(II) systems. We have recently found an analogous correlation for O2/IBA/PWllM systems, where M = Co(II), Mn(II), Cu(II), Pd(II), Ti(IV), Ru(IV) and V(V) [19]. Taking into account the radical chain mechanism of the alkenealdehyde co-oxidation, we concluded that superior catalytic activity of cobalt compounds, at least in part, arises from their ability to catalyze the chain branching via the homolytic decomposition of peroxy acid formed during aldehyde autoxidation. Moreover, cobalt most probably takes part in the chain initiation via the reaction with aldehyde. The initiation step was proposed to proceed via the formation of the oxygenated adduct, like CoIII-O-O~ which reacts with aldehyde in the rate-determining step [26]. However, the fact that the induction time is more for CoPc as compared to CoWl2 (the latter is unable to form any adducts), shows that species different from oxygenated cobalt adducts may promote the chain initiation. We believe that for most cobalt catalysts studied these species are Co(III) forms of the catalyst produced by one-electron oxidation of Co(II) initial forms with peroxy acid, which in turn is produced in the course of aldehyde autoxidation. Usually, the end of the induction period coincides with the change of the reaction mixture colour expected for Co(III) appearance. The redox potentials (E) for the majority of the studied Co(II) catalysts are unknown for MeCN medium. In aqueous solution E is known to be higher for PWllCO as compared to CoWI2 (1, 39 and 1.0 V vs SHE, respectively) [37, 38]. Among the catalysts studied, CoPc

953 most likely has the lowest E (0.61 V vs SHE in DMF [39]). Note, that the induction time falls in the order: PWllCO > CoW~2 > CoPc. Thus, we may assume that the greater E is, the higher is the rate of the chain initiation. On the other hand, the chain termination via the reaction of RCO; radicals with Co(II) (vide supra) seems to be more favorable for compounds with low E. Coordinated acylperoxy radicals were proven to act as the epoxidizing species when Mn(III) Salen complexes were used as the catalysts [11 ]. We suppose that similar species may be involved in the epoxidation process in the case of CoPc or some other chelate complexes with low E. In this case one might expect some effects of a catalyst on the stereoselectivity of the epoxidation as it was observed in [11, 14] and in this work (Table 2). At the same time, acylperoxy radicals are expected to be the main epoxidizing species for cobalt compounds having relatively high values of E. Electrochemical and UV-Vis studies, which are in progress now, should provide further understanding of the mechanism of the alkene-aldehyde catalytic co-oxidation. The results obtained allow us to propose the reaction mechanism comprising the following elementary steps of the chain radical process leading to epoxide and isobutyric acid formation: RCHO+Co

3+

RCO + 02 ~

~RCO+Co

2+ H+ +

RCO 3

(1)

B

RCO~ + -- ~

RCOf/

(6)

(2)

RCO3+ RCHO ---~ RCO3H+ RCO

(3)

RCO~/ - - " ---~ , / ~

RCO3H+ Co22---~ RCO)+ Co3+ + OH"

(4)

RCO2 + RCHO ~

RCO~H + Co3+ ----~ RCO ~+ C2++ H+

(5)

2RCO~ ~

+ RCO~

(7)

RCO 2H + RCO

(8)

termination

(9)

3. CONCLUSIONS The results obtained in this investigation prove that alkene epoxidation by 02 in the presence of IBA and cobalt catalysts proceeds via radical chain mechanism. Acylperoxy radicals most likely act as the main epoxidizing species although some other species, e.g., coordinated to the metal center acylperoxy radicals, may contribute into the epoxidation process when catalysts with low redox potentials are used. Superior catalytic activity of cobalt compounds in alkene epoxidation by O2/IBA system is due to the high ability of cobalt to catalyze the chain branching and promote the chain initiation rather than the ability of cobalt to activate dioxygen via its coordination. The presence of chelating organic ligands is not necessary to provide the efficient alkene epoxidation in O2/IBA/Co(II) systems but such ligands are needed when the effect on the epoxidation stereochemistry is desired.

4. EXPERIMENTAL Catalysts. Co(NO3)2.6H20 was of pure grade. CoNaY zeolite and tetra-n-butylammonium salts of PWllCO and CoWl2 heteropolyanions were obtained as described in [19] and [25], respectively. The content of cobalt in the CoNaY zeolite was 3.29 % wt. The formation of the Keggin structure and the purity of CoW~2 and PWllCO were confirmed by 170 and 31p NMR, respectively. Materials. Trans-stilbene (Fluka AG) and (+)-o~-pinene (Aldrich Chemical Company) were used as received. Cis-stilbene and perisobutyric acid (PIBA) were prepared as described

954 in [19]. (-)-Caryophyllene (>99%) was isolated from the oil of Eugenia caryopyllata by vacuum rectification. (+)-3-Carene (95%) was prepared by rectification of the Pinus sylvestris turpentine. Oxidation procedure. Alkene oxidation was carried out in a thermostated 20 ml Pyrexglass reactor equipped with a stirring bar and a reflux condenser. Isobutyraldehyde was added to a solution of alkene (0.1 M) in a solvent (3 ml) containing a catalyst, and the reaction mixture was vigorously stirred. Product analysis. The oxidation process was monitored by GLC ("Tsvet-500", 2mx3mm Carbowax 20M on Chromaton N-AW-HMDS for stilbenes and 15mx0.3mm SE-30 for other alkenes, Ar, FID). The reaction mixture was percolated through alumina (/=2 cm, Q=I cm), concentrated at reduced pressure and the crude product was analyzed by ~H and ~3C NMR on a Bruker AM 400 instrument. The reaction products were identified from their NMR and GCMS spectra. The yields of epoxides were determined as described in [25]. ACKNOWLEDGMENTS

This work was supported by Russian Basic Research Foundation (Grant N 96-03-34215). We thank Prof. M. A. Fedotov and Dr. A. V. Golovin for NMR measurements. A generous gift of cobalt phtalocyanine from Prof. E. N. Savinov is highly appreciated. REFERENCES

1. R.A. Sheldon and J.K. Kochi, Metal-Catalyzed Oxidations of Organic Compounds, Academic Press, New York, 1981. 2. C.L. Hill. (ed.), Activation and Functionalization of Alkenes, Wiley, New York, 1989. 3. K.A. Jorgensen, Chem. Rev. 89 (1989) 431. 4. B. Meunier, Chem. Rev. 92 (1992) 1411. 5. N. M. Emanuel, E. T. Denisov and Z. K. Maizus, Chain Reactions of Hydrocarbon Oxidation in Liquid Phase, Moscow, Nauka, 1965. 6. T. Yamada, T. Takai, O. Rhode and T. Mukaiyama, Bull. Chem. Soc. Jpn., 64 (1991) 2109. 7. T. Takai, E. Hata, T. Yamada and T. Mukaiyama, Bull. Chem. Soc. Jpn., 64 (1991) 2513. 8. T. Mukaiyama and T. Yamada, Bull. Chem. Soc. Jpn., 68 (1995) 17. 9. R.V. Kucher and I.A. Opeida, Co-oxidation of Organic Compounds in Liquid Phase, Kiev, Naukova Dumka, 1989. 10. T. V. Filippova and E. A. Blyumberg, Uspekhi Khimii, 51 (1982) 1017 (in Russian). 11. Y. Katsuki, Coord. Chem. Reviews 140 (1995) 189. 12. S. Bhatia, Y. Punniyamurthy, B. Bhatia and J. Iqbal, Tetrahedron, 49 (1993) 6101. 13. P. Mastrorilli and C.F. Nobile, J. Mol. Catal., 94 (1994) 19. 14. P. Mastrorilli, C.F. Nobile, G.P. Suranna and L. Lopez, Tetrahedron, 51 (1995) 7943. 15. N. Mizuno, T. Hirose, M. Iwamoto, in V.C. Corberan and S.V. Bellon (Eds.), New Developments in Selective Oxidation II, Elsevier Science B.V., 1994, p. 593. 16. N. Mizuno, T. Hirose, M. Tateishi and M. Iwamoto, Chem. Lett. (1993) 1839.

955 17. M. Hamamoto, K. Nakayama, Y. Nishiyama and Y. Ishii, J. Org. Chem., 58(1993) 6421. 18. O.A. Kholdeeva, V.A. Grigoriev, G.M.Maksimov and K.I. Zamaraev, Topic in Catalysis, 3 (1996) 313. 19. O.A. Kholdeeva, V.A. Grigoriev, G.M. Maksimov, M.A. Fedotov, A.V. Golovin and K.I. Zamaraev, J. Mol. Catal. A, 114(1-3) (1996), 123. 20. S.-I. Mirahashi, Y. Oda and T. Naota, J. Am. Chem. Soc., 114 (1992) 7913. 21. S.-I. Mirahashi, Y. Oda, T. Naota and N. Komiya, J. Chem. Soc. Chem. Commun., (1993) 139. 22. E. Bouhlel, P. Laszlo, M. Levart, M.-T. Montaufier and G.P. Singh, Tetr. Lett., 34 (1993) 1123. 23. P. Laszlo and M. Levart, Yetr. Lett., 34 (1993) 1127. 24. A. Atlamsani, E. Pedraza, C. Potvin, E. Duprey, O. Mohammedi and J.-M. Bregeault, C.R. Acad. Sci. Paris, ser. II, 317 (1993) 757. 25. O.A. Kholdeeva, A.V. Ykachev, V.N. Romannikov, I.V. Khavrutskii and K.I. Zamaraev, Stud. Surf. Sci. Catal., 1997, in press. 26. J. Haber, Y. Mlodnicka and J. Poltowicz, J. Mol. Catal. 54 (1989) 451. 27. P. I. Valov, E. A. Blyumberg and N. M. Emanuel, (a) Bull. Acad. Sc. USSR (1966) 1283; (b) ibid. (1969) 718. 28. A.D. Vreugdenhil and H. Reit, Recl. Trav. Chim. Pays-Bas, 91 (1972) 237. 29. F. Marta, E. Boga and M. Matok, Disc. Far. Soc., (1968) 173. 30. Y. Nishida, T. Fujimoto and N. Tamake, Chem. Lett., (1992) 1291. 31. K. Kaneda, S. Haruna, T. Imanaka, M. Hamamoto, Y. Nishiyama and Y. Ishii, Tetr. Lett., 33 (1992) 6827. 32. C.L. Hill and R.B. Brown, J. Am. Chem. Soc., 108 (1986) 536. 33. Y. Shimura and R. Tsuchida, Bull. Chem. Soc. Jpn., 30 (1957) 502. 34. L. Eberson and L.-G. Wistrand, Acta Chem. Scand. B34 (1980) 349. 35. T. I. Ikawa, T. Fukushima, M. Muto and T. Yanagihara, Can. J. Chem., 44 (1966) 1817. 36. M.A. Brook, J.R. Landsay Smith, R. Higgins and D. Lester, J. Chem. Soc. Perkin Trans., II (1985) 1049. 37. B.C. Rong and M.T. Pope, J. Am. Chem. Soc., 114 (1992) 2932. 38. Z. Amjad, J.-C. Brodovitch and A. McAuley, Can J. Chem., 55 (1977) 3581. 39. M.R. Tarasevich, K.A. Radushkina, Catalysis and Electrocatalysis by Metalloporphyrines, Nauka, Moskow (1982) 35.

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3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

957

Epoxidation o f olefins over thermally stable polyimide-supported M o ( V I ) complexes J.H. Ahn a, J.C. Kim a,

S.K. Ihm b and D.C. Sherrington~

"Department of Chemical Engineering and RECAPT, Gyeongsang National University, 900, Kajwa-dong, Chinju 660-701, Korea* bDepartment of Chemical Engineering, Korea Advanced Institute of Science and Technology, 373-1, Kusung-dong, Taejeon 305-701, Korea* ~Department of Pure and Applied Chemistry, University of Strathclyde, 295 Cathedral Street, Glasgow G1 1XL, United Kingdom

Polyimide particulates carrying a functional group have been prepared by non-aqueous suspension polycondensation. Molybdenum(VI) complex has been supported on a functional polyimide bead and used as a catalyst in the liquid-phase epoxidation of cyclohexene with tert-butylhydroperoxide (TBHP), as oxygen source. The polyimidesupported Mo catalyst was highly active and selective, and has been recycled 10 times with no detectable loss of Mo from the support.

I. INTRODUCTION Polystyrene-based resins have been used widely as supports for metal complex catalysts and other reactive species. These polymers, however, have a drawback in their limited thermo-oxidative stability [1,2]. The scope for application is therefore restricted, particularly in polymer-supported transition metal complex oxidation catalysts [3]. Consequently there is a need for the development of polymer supports with a much higher intrinsic thermo-oxidative stability. Polybenzimidazoles and polyimides are likely candidates in this respect. Polybenzimidazole resin in a porous bead has been applied as a polymer support for homogeneous metal complexes [4-6]. The polybenzimidazole-supported Mo(VI) species showed retention of high activity in the epoxidation of propylene, but progressive loss of activity on recycling in the epoxidation of cyclohexene.

*This work was carried out by G-7 Environmental Technology Development Program and the financial support is gratefully acknowledged.

958 Unlike polybenzimidazole-based thermo-oxidatively stable supports, polyimides can be prepared under relatively mild conditions from starting materials of only low or modest cost [7,8]. Polyimide particulates were prepared in a bead form without functional [9,10]. We now report the synthesis of functional polyimide beads and their use as epoxidation catalyst supports. The presence of the functional group in the polyimides allows further chemical exploitation, particularly as a catalyst support capable of operating under rather severe oxidative conditions. In this work, polyimide-supported Mo(VI) complexes were prepared and employed as heterogeneous catalysts in the epoxidation of cyclohexene using t-butylhydroperoxide (TBHP), as the oxidant.

2. EXPERIMENTAL

2-1. Materials N,N'-Dimethylacetamide (DMAc) (Aldrich, HPLC grade) was used without further purification. Acetic anhydride (BDH) was pre-dried over anhydrous sodium acetate. Pyridine (Aldrich, anhydrous) was distilled from KOH prior to use. Poly(maleic anhydride-co-octadec-l-ene)(l:l) (Polysciences), as a polymeric stabilizer, was used as supplied. Paraffin oil (A.J. Beveridge Ltd., liquid paraffin 5LT) was used as a suspending medium. Pyromellitic dianhydride (1,2,4,5-benzenetetracarboxylic dianhydride) (Aldrich) was recrystallized from butan-2-one before use. p-Phenylenediamine (Aldrich) and 2,6-diaminopyridine (Aldrich) were recrystallized from ethanol. 3,5-Diaminobenzoic acid (Aldrich) and 2,5-diaminobenzene sulfonic acid (Aldrich) were recrystallized from water and heated at l l0~ under vacuum to remove water. 3,5-Diamino-l,2,4-triazole (Aldrich) and tris(2-aminoethyl)amine (Tokyo Kasei) were without further purification. 2.2. Suspension Polycondensation A procedure similar to that which we have already reported was employed [9,10]. This involves the preparation of a pre-polymer poly(amic acid) (PAA) solution in DMAc, followed by imidization in suspension in paraffin oil. A typical procedure for the preparation of linear functionalized spherical polyimide particulates was as follows. A round-bottomed 3-necked flask was flushed with N2 and charged with a diamine in DMAc. The diarnine was completely dissolved in DMAc. While solution was mechanically stirred, finely ground pyromellitic dianhydride (PMDA) was added to the mixture on an ice bath in small portions, and then stirring continued overnight at room temperature. Paraffin oil with poly(maleic anhydride-co-octadec-l-ene)(l:l) (0.5wt% in oil) as a suspension stabilizer was added to the flask. The PAA solution was suspended for 2hr at 60~ at the speed of 400rpm. After that, imidization was initiated by dropwise addition of a mixture of acetic anhydride (4.0 molar excess of PMDA used) and pyfidine (3.5 molar excess of PMDA used). After 24hr, the polyimide particulates were filtered, washed with dichloromethane and then dried at 80 ~ in a vacuum oven. To obtain crosslinked polyimide particulates tris(2-aminoethyl)amine as a crosslinking agent was added to the PAA solution suspended in paraffin oil. After 24hr, a mixture of acetic anhydride and pyridine was added to give crosslinked polyimide.

959

2.3. Preparation of polyimide-supported Mo complex The polyimide bead bearing a triazole residue was used to immobilize a stable and active Mo(VI) epoxidation catalyst. The polyimide was refluxed with molybdenyl acetylacetonate in ethanol for 3days. Upon completion, a polyimide-Mo complex catalyst was filtered and extracted with ethanol in a Soxhlet apparatus for 3 days. The supported complex was dried thoroughly under vacuum. The molybdenum content was measured by inductively coupled plasma (ICP) to be 1.08mmolg~ for PI-DAT.Mo and 1.10mmolg~ for CPI-DAT.Mo. 2.4. Catalytic Epoxidation Catalyst (0.08g), cyclohexene (7.5ml), and bromobenzene (0.5ml) were placed in a three-necked thermostated reaction vessel equipped with condenser, septum cap, and stirrer, and left to thermally equilibrate at 60~ for 20min. Anhydrous TBHP solution (2ml, 5mmol TBHP) was added. Samples were withdrawn by syringe, and the concentration of cyclohexene oxide was monitored by gas chromatography (HP5890 Series II plus) with a capillary column (Ultra 2). 2.5. Analytic Methods Particle size distribution of functional polyimide particulates was determined by sieving: mesh 38/~m, 75/~m, 106/~m, 212/~m and 425/~m. Each size fraction was represented by wt%. FTIR spectra were recorded on a Nicolet SX20B instrument with KBr discs.

3. RESULTS AND DISCUSSION The suspension polycondensation methodology adopted has already been reported [9,10]. In this instance the pre-polymer poly(amic acid) solution in N,N'-dimethylacetamide was prepared from pyromellitic dianhydride and the functional diamines, 3,5-diamino- 1,2,4-triazole; 2,5-diaminobenzoic acid; 2,5-diarninobenzene sulfonic acid or 2,6-diaminopyridine. Each pre-polymer solution was then dispersed as droplets in paraffin oil containing poly(maleic anhydride-co-octadec-l-ene)(l:l) as a suspension stabilizer, and imidization induced at 60~ by addition of a mixture of acetic anhydride and pyridine (Figure 1). For comparison a nonfunctional polyimide (PI) was prepared using p-phenylene diamine and a crosslinked analogue of this (CPI) also produced by inclusion of tris(2-amino ethyl)amine. Typically 90---100% of mainly spherical polyimide particulates (---20g) were obtained after washing and drying (Figure 2). The results of elemental analyses and particle size distributions were shown in Table 1 and 2. For all species the H content found is higher than expected, probably reflecting trapping of solvent, fragments from the dehydrating agents, and moisture. Figure 3 shows the FTIR spectra of the triazole-containing polyimide bead (PI-DAT). The characteristic absorption bands were obtained at 1780, 1720(heterocyclic carbonyl vibration), 1348(C-N stretch vibration) and 720cm~(imide ring deformation). The thermogravimetric analysis curves for the polyimide beads are shown in Figure 4. The progressive loss below ~-300~ almost certainly corresponds to the physically trapped components mentioned above.

960 Serious degradation of all the functional polyimide particulates in air does not happen until --400~ All the polyimides therefore show good prospects for high temperature application as supports, certainly in reaction up to 300~ The homogeneous Mo complex, MoO2(acac)2, was supported on the PI-DAT bearing a triazole group. The FTIR spectra of PI-DAT.Mo (Figure 3) showed a band attributable to Mo=O stretching mode at 960cm1, indicating the presence of oxomolybdenum centers. There was no evidence in the IR spectrum of PI-DAT.Mo for an Mo-O-Mo bridge and so the structure of the Mo center in this polymer catalyst remains unclear. However, it is considered that all Mo derivatives on the supported catalysts probably contain oxomolybdenum centers from the initial rate data with no induction period (Figure 5). One of the authors has already reported on the possible structures of polymer-supported Mo complexes [11].

0~~

-t-

O

.o o

H2N-Ar-NH 2

A&.

-'~ N H - - ~ ~ O OH O PAA solution

Paraffin Liquid with a suspension stabilizer

PAA droplets

in oil

Acetic anhydride / Pyridine - H20 PI particulates

H PI--DAT

PI

COOH PI-COOH

SO3H PI-SO3H

PI-Py

Figure 1. Schematic synthesis of polyimide particulates. Table 1 Elemental analysis of polyimide particles CODE

Calculated(%)

Found(%)

C

H

PI

66.21

2.07

N 9.66

C 64.72

H 3.10

N 8.97

PI-DAT

58.85

1.59

10.16

45.79

2.93

12.63

CPI-DAT

49.56

3.05

17.15

52.86

2.99

16.80

PI-COOH

61.08

1.80

8.38

58.21

3.26

8.85

PI-SO3H

51.89

1.62

7.51

52.50

3.20

8.06

PI-Py

61.86

1.73

14.43

49.71

3.59

9.75

961

Table 2 Particle size distribution of polyimide particles Particle size fraction(wt%) a CODE A

B

C

D

E

F

PI

7.7

2.3

7.1

38.2

44.7

0

PI-DAT

0

3.0

5.6

34.8

41.4

15.2

CPI-DAT

0

0

0.3

1.5

80.4

17.8

PI-COOH

0

7.4

20.0

33.4

39.2

0

PI-SO3H

0

1.5

2.3

3.8

16.0

76.4

PI-Py

2.4

10.8

7.0

6.9

22.3

50.6

aA; < 38/~ m, B; 38-75/~ m, C; 75-105/1 m, D; 106-212/~ m, E; 212-425/~ m, F; >425/~ m

Figure 2. Optical Photograph of PI-DAT beads.

962

PI-DAT.Mo

PI-DAT

I

I

2000

I

1600

I

1200

I

800

400

Wave number(cm")

Figure 3. FTIR spectra of PI-DAT and PI-DAT.Mo.

100~

80

o~' v

60

rr

9

40

9

PI 9 PI-DAT PI-COOH

9 PI-SO3H 9 PI-Py

20

0

0

i

1

200

400

600

800

Temperarure(C) o

Figure 4. TGA analysis of functional polyimide particulates.

963 It is well known that the soluble species MoO2(acac)2 is a potent catalyst in the epoxidation of olefinic compounds. Figure 5 shows conversion curves for the epoxidation of cyclohexene by TBHP catalyzed by homogeneous MoOz(acac)2 and heterogeneous CPI-DAT.Mo. The initial activity of homogeneous MoO2(acac)2 was higher than that of CPI-DAT.Mo. However, the activity of CPI-DAT.Mo should be comparable with that of the homogeneous analogue after 30min.

CH3 I + H3C--C--OOH I CH3

Mo(VI)

,~-

[ ~

CH3 I O + H3C--C--OH I CH3

100

o-e,

~

60

0

o

40

~-

2o

o

0

20

40

60

80

1 O0

120

Time(min)

Figure 5. Epoxidation of cyclohexene with TBHP using CPI-DAT.Mo and homogeneous Mo. The prolonged activity of polymer-supported heterogenized catalysts is probably the most important factor in their performance. The deactivation of the catalyst by either degradation of the polymer-supported itself or by leaching of active species from the catalyst is unfavorable. Figure 6 shows the yield of cyclohexene oxide after 120min. The triazole-containing polyimide-supported Mo catalysts were recovered at the end of each run and used repeatedly under identical conditions. The catalyst shows substantial retention of activity over 10 recycles unlike an earlier polybenzimidazole-supported Mo complex, where the latter displayed rapid deactivation on recycling. The presently reported retention of activity is most encouraging, and suggests that catalysts based on functional polyimide particulates might form the basis of a range of stable polymer-

964

\ 60

\

/

o o

"6

9 9 9

40

"o

~--

C P I - D A T . M o 700C C P I - D A T . M o 60~ PI-DAT.Mo 60"C

20

I

i

I

I

I

I

I

I

i

1

2

3

4

5

6

7

8

9

10

R e c y c l e number

Figure 6. Recycling of polyimide-supported Mo catalysts in the epoxidation of cyclohexene by TBHP at 120min supported metal complex catalysts, where the support is readily synthesized and is highly cost-effective. Application on the both a laboratory and a technical scale also seems feasible. REFERENCES

1. P. Hodge and D.C. Sherrington (eds.), Polymer-supported Reactions in Organic Synthesis, Wiley-Interscience, Chichester, 1980. 2. Y.I. Yermakov, B.N. Kuznetsov and V.A. Zakharov, Catalysis by Supported Complex, Elsevier, Amsterdam, 1981. 3. D.C. Sherrington, Pure Appl. Chem., 60 (1988) 401. 4. H.G. Tang and D.C. Sherrington, J. Catal., 142 (1933) 540. 5. M.M. Miller and D.C. Sherrington, J. Catal. 152 (1995) 377. 6. M.M. Miller, D.C. Sherrington and S. Simpson, J. Chem. Soc. Perkin Trans. 2 (1994) 2091. 7. D. Wilson, H.D. Stenzenberger and P.M. Hergenrother (eds.), Polyimides, Chapman and Hall, New York, 1990. 8. L.H. Lee (ed.), Adhesives, Sealants and Coatings for Space and Harsh Environments, Polymer Science and Techology, Plenum, New York, 1988. 9. T. Brock and D.C. Sherrington, J. Mater. Chem., 1 (1991) 151. 10. T. Brock, D.C. Sherrington, and J. Swindell, J. Mater. Chem., 4 (1994) 229. 11. M.M. Miller and D.C. Sherrington, J. Catal., 152 (1995) 368.

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

965

Selective Partial O x i d a t i o n of P r o p y l e n e to P r o p y l e n e Oxide on A u / T i - M C M Catalysts in the Presence of H y d r o g e n and O x y g e n Yuri A. Kalvachev* ", Toshio Hayashi b, Susumu Tsubota a and Masatake Haruta" aOsaka National Research Institute, AIST, Midorigaoka 1-8-31, Ikeda 563, Japan bNippon Shokubai Co., Ltd, Research Division, Nishi-Otabi 5-8, Suita 564, Japan 1. INTRODUCTION The epoxidation of alkenes is attracting increasing attention of researchers, from both academia and industry. The epoxides are one of the most useful intermediates in organic synthesis as they are versatile products that easily undergo ring-opening reactions to form bifunctional compounds. In the chemical industry, propylene oxide (PO) is mainly used for producing resins. Major conventional manufacturing methods for the synthesis of PO require a two-stage processes, using chlorhydrin or hydroperoxides. The direct synthesis of PO, by the use of oxygen, has long been considered desirable and is one of the most important reactions still not solved by catalysis [ 1]. The direct vapour-phase oxidation of propylene to PO in the presence of oxygen and hydrogen is an environmentally friendly process, or as J. Thomas once stated : "The name of the game is to get mild conditions and cheap oxidants that are environmentally friendly" [2]. Since direct oxidation is difficult to achieve, several approaches have been made to attempt to produce PO by oxidation with hydrogen peroxide over titalaosilicate catalyst [3,4]. In liquid phase Pd-Fe zeolite [5] and Pd-titanosilicate [6] are known to oxidize alkanes and alkenes to oxygenates by hydrogen peroxide generated in situ. New materials consisting of amorphous silica with regular pore structure, therefore called mesoporous molecular sieves, have recently been described [7]. Isomorphous substitution of Si by Ti has been attempted by performing the synthesis in the presence of titanium compounds. Ti-MCM have been tested for oxidation of hydrocarbons in liquid phase, using H202 or hydroperoxides as oxidants [8-10]. Because gold has long been regarded as being catalytically less active than platinum group metals it has attracted little attention in the development of heterogeneous catalysis. The basic reason is that gold catalysts are highly sensitive towards the preparation methods and it is normally impossible to prepare active gold catalysts with classical impregnation methods. However, when gold is dispersed as fine particles over suitable support by coprecipitation or deposition-precipitation methods, it has been found that the supported gold exhibits exceptionally high catalytic activity for such reactions as : CO oxidation [ 11-14], CO~ and CO hydrogenation [ 15], hydrocarbon combustion [ 16], the reduction of NO to N 2 [ 17]~ and the water-gas shift reaction [ 18,19]. Quite often reactions occur at low temperature. In previous work [20] the first evidence of the direct vapour-phase oxidation of propylene to propylene oxide was presented, using a catalyst, comprised of gold deposited on TiO 2.

*permanent address : Institute of Catalysis, Bulgarian Academy of Sciences, 1113 Sofia, Bulgaria

966 In continuation with our work, we now investigate the catalytic activity of gold deposited on titanium-containing MCM-41 for the partial oxidation of propylene to PO, in the presence of oxygen and hydrogen. In the present report, we have studied the influence of the amount of titanium in Ti-MCM-41 and gold loading over this reaction.

2. EXPERIMENTAL 2.1. Sample preparation Ti-MCM-41 samples were synthesized by using dodecyl trimethyl ammonium chloride as a template following the procedure described in [8]. The MCM supported gold catalysts were prepared by deposition-precipitation [21] on:pure-silica MCM-41 ; Ti-MCM-41 with ratio Ti/Si=l/100 ; Ti/Si=3/100 ; Ti/Si=6/100 by using HAuC14, followed by washing, drying and calcination in air at 673 K for 4 h. TEM micrographs showed that gold particles are homogeneously dispersed on the support, with an average diameters around 2 nm. The particle size distributions were obtained by TEM (Hitachi H-9000NA). Au-TiOJTi-MCM-41 were prepared by the same method by replacing Ti-MCM with TiOJTi-MCM, which was prepared by impregnating Ti-MCM with titanyl acetyl acetonate, followed by drying and calcination ha air at 773 K. The same procedures were also applied to Au-TiO2/SiO 2 catalysts. 2.2. Catalytic activity The catalytic activity of the samples was measured in a flow reactor at atmospheric pressure over the temperature range of 323-393 K. A gaseous mixture of propylene, oxygen, hydrogen and argon as diluent was passed through a fixed bed containing catalyst sample (Ar:O::H2:Pr = 7:1:1:1). Samples were sieved between 70 and 120 meshes and weighed to 0.5 g. Catalysts were pretreated in mixture of Ar:O 2 (7:1) at 573 K for 1 hour. The space velocity was 4000 h1. Reaction products were analyzed by gas-chromatography. 3. RESULTS AND DISCUSSION Table 1 lists the measured values of selectivity to PO, propylene conversion and hydrogen conversion for the reaction of propylene with hydrogen and oxygen over AuFFi-MCM catalysts at 323 and 373 K. The conversion of propylene is about 1-2 %, but the selectivity to PO is very high, in most cases above 90 %. Carbon dioxide is the main by-product. The most active catalyst in this reaction was Au/Ti-MCM with a Ti/Si ratio of 3/100. A sample with the same gold loading on silica MCM-41, which did not contain titanium, is catalytically inactive. Ti-MCM (31100) sample without gold loading is also inactive under these reaction conditions. This means that both components - gold and titanium are necessary for the oxidation of propylene, and that there is most probably a specific interaction between gold and titanium. A synergistic behavior of gold deposited on metal oxides is observed for lowtemperature oxidation of CO [22] and low-temperature water-gas shift reaction [ 19]. Fig.1 (a,b) shows the results for propylene oxidation over 8 wt%AtffTi-MCM with Ti/Si ratio 1/100 and 3/100, respectively. In Fig. 2 (a,b) the results from oxidation of propylene at 373 K on Ti-MCM (Ti/Si=3/100) with 4 and 12 wt % gold loading are presented. The PO yield is the highest on 8 wt % Au/Ti-MCM with aPO selectivity of 95 % (Fig.lb). On 12 wt % Au/Ti-MCM conversion of propylene is higher, but PO selectivity is 80 % . Thus the yield of propylene oxide is lower than that observed on 8 wt % Au sample. The catalytic activity increased with time over 2-3 hours (at 323 K) and 70-90 min (at 373 K) followed by a period of relative stability. Naito et al. [23] have suggested that oxygen molecules on small gold particles behave as a peroxo-like adsorbed species, which enhances the dissociation of hydrogen molecules. Moreover, the products are similar to those obtained

967 TABLE 1 Oxidation of propylene over Au/Ti-MCM catalysts Ti/Si ratio

T, K

gold

conversion

PO

loading

selectivity % propylene

323

1/100 3/100

9

18

4%

96

0.65

9

96

0.74

9

1.21

28

2%

91

0.36

11

4%

84

1.01

27

8%

91

0.36

13

8%

97

1.34

30

4%

94

0.96

27

8%

95

1.75

36

12%

80

1.89

50

6/100

a

0.45

90

373

1.20

97

8%

3/100

1.40

8%

12% 6/100

1/100

2%

80

0.88

47

4%

69

1.20

55

8%

80

1.13

26

9 9

9

uJ

1.80 -1.60 1.40 :

9 9 9

1.00 o _1 0.80

9

9 ua

* 9

>" 0.60 O

hydrogen

~

~.oo

b 000 9149149149149 O0

9

0.80

o 0.60

o. 0.40

0.40

0.20

0.20 +O

0.00

0.00 0

50

I O0 TIME

(MIN)

150

200

0

50

I O0 TIME

150

(MIN)

Figure 1. PO Yield in the oxidation of propylene at 373 K over 8 wt % Au/Ti-MCM as a function of time a) Ti/Si=l/100 ; b) Ti/Si=3/100.

200

968 in the oxidation of hydrocarbons by H202 on Ti-MCM [8]. We can thus speculate that the active species are a similar hydroperoxo-species, located at the boundary between gold and titanium. 1.00 T

1.40

a

0.90 -

9

0.80

9

0.70

1.20

9

9

9

uJ >" 0.50 o 0.40 " 0.30

9

1.00

9

o 0.60

41,

9

9 1 4 9 1 4 9c~ 0.80 u.i ~ 0.60 o

9

0

O~'OO0 0 9

0

~ 0.40

0.20

0.10

9

9

0.20

o.oo

0.oo o

50

~o 9 TIME

~5o

z0o

o

~

~

~

~

5o

~o 9

~5o

zoo

TIME

(MIN)

(MIN)

z50

Figure 2. P 9 yield in the reaction of propylene at 373 K over Au/Ti-MCM (Ti/Si=3/100) as a function of time a) 4 wt%Au ; b) 12 wt %Au. Fig.3 shows the P 9 yield as a function of time at different ratio of hydrogen : propylene in the gaseous mixture. When the concentration of H 2 is lower (H2:Pr=-5/10), the catalytic activity decreases and the shape of the curve with time is different- the activity increases over 6 hours. This may suggest that hydrogen is involved in forming the active species.

~.ZOT a

9

1.00

-o~ ~

OOOOOO0

0.80 "~ 0.60 o "

o

0.20 0.00

0.50 0.40

~176

0.30

***

0.40

b

"

o o.zo -

9

,,~

0.10

9

o

0.00 50

_O~

100

150

TIME (MIN)

200

250

f

0

100

200 TIME (MIN)

Figure 3. Oxidation of propyleneat 323 K over 12 wt % Au/Ti-MCM (Ti/Si=3/100) with a) H2:Pr=10/10 ; b)H~:Pr=5/10 vol.%.

300

400

969 The results presented in Fig.4 show that at the beginning of the reaction, the consumption of hydrogen is higher, whereas PO yield is low. Moreover, at low temperature (323 K) an induction period has been observed, during which there is not a formation of PO. We speculate that during this period a formation of hydroperoxide species takes place. 20.0

-

,r.,,

18.0 16.0 N -1-

14.0

"

12.0

0

9 10.0

8.0 bO

z O u

6.0 4.0

2.0 0.0 0

I

I

I

I

I

100

200

300

400

500

TIME

(MIN)

Figure 4. The consumption of hydrogen in the reaction of propylene with oxygen and hydrogen at 323 K over 2%Au/Ti-MCM (Ti/Si=6/100) In order to confirm this possibility, the following experiment was carried out - introducing only hydrogen and oxygen over the catalyst for the first 80 minutes. It turns out that before adding propylene, the hydrogen is oxidized 100 %. As it is seen in Fig.5 after the addition of propylene, the consumption of hydrogen decreases and PO yield increases with time. There is a competition between hydrogen and propylene in the oxidation reaction. The gradual increase in PO yield over a period of 70 min indicates that in the absence of propylene, hydroperoxidelike species are not formed or decomposed very rapidly. 0.60 T

I00.0 90.0

0,0

0.50

80.0 e,i 70.0 "-r" u. 60.0 0 50.0 40.0 30.0 Z

0.40 r~ _I

ua 0.30 O

~. 0.20

0

0.10

-4,0,,,,,, __ b

9

00000,0

20.0 10.0 o.o

0.00 5o

~o o TIME

(MIN)

~s o

zoo

o

so

~o o TIME

~s o

zoo

(MIN)

Figure 5. PO yield (a) and hydrogen consumption (b) at 373 K on 4%Au/Ti-MCM (Ti/Si=6/100).

970 Fig.6 shows PO yield over 8 wt % Au/TiOJSiO 2 as a function of time. The catalytic activity of Au/TiOJSiO 2 catalyst is not stable. Water is continuously formed during the oxidation of propylene and the oxygenated intermediates may block the active sites and depress the adsorption of propylene on the surface of the catalyst. MCM materials have hydrophobic character and Ti-MCM preferentially adsorbs less polar olefin molecules. This decreases the competition from water and probably avoid the accumulation of the oxygenated intermediates to lead to more stable catalytic activity. 2.00

9 9

1.80

Oo

1.60 1.4o

00

000 O0

o, 1.20 "' 1.00

O0 0

0.80 0.60 0.40 0.20 0.00 0

I

I

I

I

so

lOO

1 so

zoo

TIME

(MIN)

Figure 6. PO yield over 8 wt%Au/TiOJSiO 2 at 373 K. Table 2. Oxidation of propylene on Au/1 wt%Ti02/Ti-MCM-41 catalysts Ti/Si ratio*

T, K

gold loading

PO selectivity %

conversion % propylene

silica MCM

323

8%

hydrogen

inactive

1/100

8%

90

1.10

22

3/100

8%

90

0.65

16

8%

87

0.93

17

1/100

8%

74

2.45

53

3/100

8%

84

1.33

36

silica MCM

373

*this ratio is for Ti-MCM support

971 In Table 2 the results for the reaction of propylene with hydrogen and oxygen over AtffTiO2/Ti-MCM are presented. The most active sample was Ti/Si with a ratio of 1/100. From obtained data it can be concluded that for oxidation of propylene there exists an optimum amount of titanium in the catalyst. This may suggest that the density of the active sites influences the catalytic activity and PO selectivity. The results show that the optimum temperature for selective pal~ial oxidation of propylene to propylene oxide is about 100~ With increasing temperature, the conversion of propylene increased but PO selectivity was then lower. The conversion of propylene on these catalysts is about 1-2 %, while conversion of hydrogen, in all cases, is high (Tables 1 and 2). Probably the oxidation of propylene occurs at the perimeter interface between gold and the support, whereas hydrogen is activated on the surface of the gold particles. 4. C O N C L U S I O N S There exist typical characteristics features in the reaction of propylene with hydrogen and oxygen over Au/Ti-MCM catalysts: - an increase of PO yield over 2-3 hours (at 323 K) and 70-90 min (at 373 K), followed by a period of relative stability ; - increasing PO selectivity and decreasing hydrogen consumption as a function of time ; - PO selectivity decreases with increasing the amount of titanium in the catalysts. - in all cases the conversion of hydrogen is high (10-50%). The results point to the formation of hydroperoxide species that are the active oxidant agents. The reaction occurs at the interface perimeter of the support around the gold particles. The activity of Au/Ti-MCM-41 in the reaction of propylene with hydrogen and oxygen can be attlibuted to synergetic effects between gold and titanium in these catalysts. The hydrophobic character of MCM molecular sieves is the probable reason for the stability of the activity with time.

A C K N O W L E D G M E N T

Yu.A.K. gratefully acknowledges financial support by the Science and Technology Agency of Japan.

R E F E R E N C E S

1. 2. 3. 4. 5. 6. 7. 8. 9. 10. 11.

J.F.Roth, Chemtech, (1991) 357. M. Freemantle, C&EN, 74(31) (1996) 47. M.Clerici, G.Bellussi, and U.Romano, J.Catal., 129 (1991) 159. B.Notari, Adv. in Catal., 41 (1996) 253. N.Herron and C.Tolman, J.Am.Chem.Soc., 109 (1987) 2837. T.Tatsumi, K.Yuasa, and H.Tominaga, J.Chem.Soc., Chem.Commun., (1992) 1446. C.Kresge, M.Leonovicz, W.Roth, J.Vartuli, and J.Beck, Nature, 359 (1992) 710. A.Corma, M.Navarro, and J.Perez-Pariente, J.Chem.Soc.,Chem.Commun., (1994) 147. P.Tanev, M.Chibwe, and T.Pinnavaia, Nature, 368 (1994) 321. T.Blasco, A.Corma, M.Navarro, and J.Perez-Pariente, J.Catal., 156 (1995) 65. M.Haruta, S.Tsubota, T.Kobayashi, H.Kageyama, M.Genet, and B.Delmon, J.Catal., 144 (1993) 175. 12. M.Bollinger and M.Vannice, Appl.Catal., B:Environmental, 8 (1996) 417. 13. S.Gardner, G.Hoflund, B.Upchurch, B.Schryer, E.Kielin, and J.Schryer, J.Catal., 129 (1991) 114.

972 14. 15. 16. 17. 18. 19. 20. 21. 22. 23.

S.Tanielyan and R.Augustine, Appl.Catal. A:General, 85 (1992) 73. H.Sakurai and M.Haruta, Appl.Catal. A:General, 127 (1995) 93. M.Haruta, A.Ueda, S.Tsubota, and M.R.Torres Sanchez, Catal.Today, 29 (1996) 443. A.Ueda, T.Oshima, and M.Haruta, Appl.Catal. B:Environmental, in press. D.Andreeva, V.Idakiev, T.Tabakova, A.Andreev, and R.Giovanoli, Appl.Catal. A:General, 134 (1996) 275. D.Andreeva, V.Idakiev, T.Tabakova, and A.Andreev, J.Catal., 158 (1996) 354. T.Hayashi, K.Tanaka, and M.Haruta, Preprints of Symposia on Heterogeneous Oxidation, 21 lth National meeting of Amer.Chem.Soc., New Orleans, 1996, pp. 71-74. S.Tsubota, D.Cunningham, Y.Bando, and M.Haruta, "Preparation of Catalysts VI", G.Poncelet, J.Martens, B.Delmon, P.A.Jacobs and P.Grange (eds.), Elsevier, (1995) 227. M.Haruta, S.Tsubota, A.Ueda, H.Sakurai, "New Aspects of Spillover Effect in Catalysis", T.Inui, K.Fujimoto, T.Uchijima, M.Masai (eds), Elsevier (1993) 45. S. Naito and M. Tanimoto, J.Chem.Soc., Chem.Commun., (1988) 832.

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 1997 Elsevier Science B.V.

973

Immobilization of t r i a z a c y c l o n o n a n e - t y p e metal complexes on inorganic supports via covalent linking: spectroscopy and catalytic a c t i v i t y in olefin o x i d a t i o n Y.V. Subba Rao, D.E. De Vos,* B. Wouters, P.J. Grobet and P.A. Jacobs Center for Surface Chemistry and Catalysis, Katholieke Universiteit Leuven, Kardinaal Mercierlaan 92, 3001 Heverlee (Belgium) Different approaches are tested in the covalent linking of the triazacyclononane (tacn) macrocycle to amorphous or mesoporous siliceous supports. The best catalytic results for the epoxidation of olefins are obtained with a tacn, attached via a 3-oxypropyl-2-hydroxypropyl spacer to the support. The organic structures on the surface are studied with TGA, TPD-MS, 13C-NMR and sorption measurements. ESR is used to probe the details of the metal binding on these surfaces. 1. I N T R O D U C T I O N Catalytic mono-oxygen transfer from first row transition metals to nucleophilic substrates has been the subject of intensive studies since the late seventies [1-2]. The classic procedures of porphyrin-catalyzed oxidations have however obvious disadvantages [3-6]. Chlorinated solvents are often used, either in a two phase system or as co-solvents to dissolve the porphyrin. The reaction mixtures are heavily colored. Catalyst recuperation is not obvious, and often the porphyrin doesn't even survive a single catalytic run. Several groups have attempted with varying success to inlprove the usability of porphyrins by diverse heterogenization techniques [7-10]. As an alternative to porphyrin and phthalocyanine catalysts, complexes of Mn and the cyclic triamine 1,4,7-trimethyl-l,4,7-triazacyclononane (tmtacn) clearly deserve more attention [11]. In acetone and at subambient temperature, the activity of Mn-tmtacn matches that of the more active porphyrins, with 1,000 turnovers within a few hours in the styrene epoxidation [12]. Moreover, Mntmtacn is colorless after reaction, and because of its relatively moderate price, it has even been commercialized for a short while in laundry powders [13]. A heterogeneous version of Mn-tmtacn would obviously offer even more advantages. We have proposed an immobilization of Mn-tmtacn based on zeolite We acknowledge support from K.[LL~dven (YVSR) and F.W.O. (DEDV and PJG). This work was performed in the frame of an interuniversitary attraction pole (I.U.A.P.) program Supramolecular Catalysis. E-mail: [email protected]

974 Y [14]. A major problem is however that this hydrophilic matrix attracts H202. This is a drawback from the peroxide efficiency viewpoint. Therefore a purely siliceous matrix seems more attractive. As pure SiO2 lacks ion exchange capacity, one has to revert to other immobilization strategies, such as the covalent route. The present paper investigates various routes for covalent a t t a c h m e n t of the tacn macrocycle to a pre-formed support matrix. Two different spacers are used to link the surface and the tacn: propyl (P), and glycidoxypropyl (GP). The affinity of the modified surface for metals is probed with the test ions Cu 2§ and Mn 2§ and styrene is the test substrate for the selective hydrocarbon oxidation. A preliminary note on this work has appeared [15]. 2. E X P E R I M E N T A L

MCM-41 was prepared following an existing procedure [16]. The quality of the synthesis was evaluated based on the diffractogram and the N2 sorption isotherm. The material was calcined at 823 K and stored in a desiccator to avoid rehydration. Silica was purchased from Fluka (70-230 mesh) and pretreated under vacuum. 3-Chloropropylsilica was from Aldrich. For the anchoring of the organosilane on the Si matrix, 4.5 mmol of (3glycidyloxypropyl)trimethoxysilane was reacted during 10 h with 3 g of dry support material in 25 ml pre-dried and refluxing toluene. Excess silane was removed by toluene soxhlet extraction. For the reaction of tacn with GP- or Pbearing materials, 80 mg tacn was reacted overnight with 1 g of the vacuumdried support at 323 K in 50 ml toluene, followed by another toluene extraction. Eventually the remaining secondary amine groups on the bound tacn were alkylated with an estimated 5-fold excess of propylene oxide (ethanol, 293 K, 24 h). An overview is given in Scheme 1. H I

S c h e m e 1.

~o

,",,,~o

H i

H

H,,~'-~~H

/'~,/NCl

H i

~OH

H I

..

.

HOL

975 For TGA, a S e t a r a m TGA-DTA 92 a p p a r a t u s was used. Alternatively, a homebuilt a p p a r a t u s was employed for TPD-MS. GC analysis was on a Chrompack CP-Sil-5 column, eventually coupled to a Fisons mass spectrometer. ESR spectra were recorded with a Bruker ESP-300 and a TEl04 cavity at t e m p e r a t u r e s between 130 and 300 K. N~ sorption experiments were performed with an Omnisorp-100 i n s t r u m e n t . The t-plot method was applied for the analysis of the pore volume. Solid state 13C NMR spectra were recorded using a B r u k e r MSL 400 spectrometer at a resonance frequency of 100.61 MHz. Cross polarization was optimized with glycine as a reference. For the m e a s u r e m e n t of liquid samples, a B r u k e r AMX 300 system was used, operating at 300.13 and 75.47 MHz for 1H and ~3C, respectively. 3. R E S U L T S

Characterization of the functionalized surfaces T h e r m o g r a v i m e t r i c analysis is a basic technique for quantifying surface loading with organic groups. Samples were heated at 5 K per minute up to 1073 K in a He/O~ atmosphere. The weight loss (in %) above 453 K is given for all samples in Table 1. The organic weight increase after the reaction with tacn shows t h a t the linking is successful both for 3-chloropropyl (P) and for glycidyloxypropyl (GP) residues. Tacn surface concentrations are highest with MCM-GP and lowest for the commercial Sil-P. The TGA profiles are highly similar for tacn-containing samples (Sil-P-tacn, Sil-GP-tacn, MCM-GP-tacn) on one h a n d and the tacn-free precursors (Sil-P, Sil-GP, MCM-GP) on the other h a n d (Figure 1). With the latter materials, one main and sharp exotherm is observed around 473 K. With all tacn-containing samples, the combustion occurs over a much broader t e m p e r a t u r e interval (473-823 K).

~

,

Exo

_

-5

-5

Exo I

a

a

-15

2o0

40o

600 (c~

-35 ~,,

200 I

400 I

600 I

(el

F i g u r e 1. Weight loss (%, a) and heat flow (b) for MCM-GP (left) and MCM-GPtacn (right).

976 T a b l e 1. Surface loading (wt. % or tacn concentration) as d e t e r m i n e d by TGA.

Sample

% wt. loss

Sil-P Sil-P-tacn Sil-GP Sil-GP-tacn MCM-GP MCM-GP-tacn

[tacn] (retool / g) 0.20 0.28 0.40

5.9 8.5 12.3 15.9 12.2 17.3

W h e n this t h e r m a l decomposition is assessed by mass spectroscopy, typical f r a g m e n t s of decomposition of e.g. GP are detected. For instance, m/z = 57 is probably due to a 2,3-epoxypropyl group. 13C-NMR can be applied to check the intact nature of the surface groups after the anchoring and the extractions. As an example, we discuss the d a t a for the glycidylated MCM-41 (MCM-GP). *

a

~t

i

,!!

'i

I,,~~i,

!

!

,

i

t I :O 2-Cyclohexene-l-one

7-Oxabicyclo[4.1.0]heptan-2-one

Scheme 1

1004 cyclohexene during this period to a 3:1 mixture of 2-cyclohexen-1-one and 2-cyclohexen-1ol. No epoxide was formed under these conditions. Table 2 shows the results for the oxidation of cyclohexene catalyzed by a series of titanium modified hexagonal NaY zeolites. The epoxide is produced in all cases but under the present conditions is transformed to the diol as the major product. Additionally, there are smaller amounts of autoxidation products (K + A). This product distribution was to be expected given the acidity of these zeolites and has certainly been noted before with other titanium modified zeolites [24,25]. After aluminum removal from the zeolites, the side reactions and selectivity improves [26]. Most of the catalyst activity occurs in the first few hours since after one day there is only a small increase in the conversion of substrate or peroxide. There is an apparent improvement in selectivity for the epoxide after a day but the rate of reaction is quite low. This would suggest clogging of the zeolite pores which retards the reaction chemistry. The spent catalyst is typically pale yellow in color while the starting zeolite is white. The MCM-41 modified with titanium in a similar fashion to our study also turns yellow in a peroxide based cyclohexene oxidation but deactivates after only 90 minutes [ 12]. The used catalyst may be calcined at 500~ in flowing oxygen to remove the color and restore the original catalytic activity. Controlling tile catalyst acidity as well as evaluating the influence of solvent will be important factors to consider in improving catalyst lifetime. Table 2. Results for cyclohexene oxidation catalyzed by titanium modified hexagonal NaY.

% Conversion Sample

1 2 3 4 5

% Selectivitya'b

Hours

CY

H202

CYO

diol

K

A

% Efficiencyc

5 24 5 24 5 24 5 24 5 24

5.7 6.4 8.2 10.5 10.8 12.8 7.2 9.5 4.5 6.4

50 54 69 75 56 65 36 50 30 47

5.0 4.8 6.1 5.5 5.7 4.3 4.3 4.0 6.3 5.7

58.1 60.0 57.9 50.2 69.3 53.1 59.1 68.1 52.5 51.8

24.2 22.9 31.1 31.0 15.4 25.8 27.5 19.9 32.0 31.9

12.6 12.2 16.8 13.2 9.3 1.6.7 8.9 7.8 9.2 10.6

36 37 38 45 59 61 64 62 54 53

CY = cyclohexene; CYO = cyclohexene oxide; diol = 1,2-cyclohexanediol; K - 2-cyclohexen- 1-one; A = 2-cyclohexen- 1-ol. b Selectivity = (mmol product / mmol total products) x 100. c Efficiency = (retool of converted cyclohexene /mmol of converted H202) x 100 a

1005 Table 2 indicates that the conversion of cyclohexene as well as the peroxide efficiency increases as the titanium content increases for the first three samples. Recall samples 4 and 5 showed evidence of bulk TiO2 occlusion which must have adverse effects on the reaction including pore blockage. Sample 5 is the dealuminated zeolite which should have fewer but stronger acid sites. In spite of having the highest titanium loading this sample exhibits the lowest conversion of both substrate and peroxide. The peroxide efficiency is comparable to the better catalysts which may reflect the lower acid site density. A slightly lower level of epoxide hydrolysis may also represent a more hydrophobic environment at the active site. These result certainly warrant further investigation of the high silica hexagonal Y type zeolites as support materials for isolated titanium species. 4. CONCLUSIONS We have grafted titanocene derived species onto the surface of hexagonal NaY before and after dealumination resulting in effective epoxidation catalysts. The organometallic was intended to reduce the possibility of forming oligomers and bulk titania but it is clear that at high enough loadings this is unavoidable. However, this method of zeolite surface modification has resulted in epoxidation activity one would associate with isolated titanium centers which then represents a viable alternative to framework titanosilicate catalysts. Clearly more work is needed to better define the optimum zeolite composition and reaction conditions that will stabilize the system and maximize selectivity. ACKNOWLEDGMENTS The support of the Robert A. Welch Foundation and the National Science Foundation are gratefully acknowledged. REFERENCES

.

4. 5. .

7. 8.

M. Taramasso, G. Perego and B. Notari, Preparation of porous crystalline synthetic material comprised of silicon and titanium oxides, U. S. Patent No. 4,410,501 (1983). M. Taramasso, G. Manara, V. Fattore and B. Notari, Silica based synthetic material containing titanium in the crystal lattice and process for its preparation, U.S. Patent No. 4,666,692 (1987). B. Notari, Catal. Today, 18 (1993) 163. B. Notari, Stud. Surf. Sci. Catal., 37 (1987)413. D.E. De Vos, P.L. Buskens. D.L. Vanoppen, P.P. Knops-Gerrits and P.A. Jacobs in Comprehensive Supramolecular Chemistry 7 (1996) 647 and references there in. 1L Millini and G. Perego, Gaz. Chim Ital., 126 (1996) 133. B. Notari, Adv. Catal., 41 (1996) 253. K~J. Balkus, Jr., A.G. Gabrielov and S.I. Zones, Stud. Surs Sci. Catal., 97 (1995) 519. I~J. Balkus, Jr., A.A. Khanmamedova, A. Gabrielov and S.I. Zones, Stud. Surf Sci. Catal., 101 (1996) 1341.

1006 10. 11. 12. 13. 14. 15. 16. 17. 18. 19. 20. 21. 22. 23. 24. 25. 26.

A. Corma, M.T. Navarro and J. Perez-Pariente, J. Chem. Soc., Chem. Commun., (1994) 147. P.T.Tanev, M. Chibwe and T.J. Pinnavaia, Nature, 368 (1994) 321. T. Maschmeyer, F. Rey, G. Sancar and J.M. Thomas, Nature, 378 (1995) 159. P.G. Pries de Oliveira, J.G. Eon and J.C. Volta, J. Catal., 137 (1992) 257. A. Corma, A. Fuerte, M. Iglesias and F. Sanchez, J. Mol. Catal., 107 (1996) 225. N. Ichikuni, M. Shirai and Y. Iwasawa, Catalysis Today, 28 (1996) 49. IL Butch, N. Cruise, D. Gleeson and S.C. Tsang, Chem Commtm., (1996) 951. S. Schwartz, D.1L Corbin and A.J. Vega, Mater. Res. Soc. Symp. Proc., 431 (1996) 137. M.W. Anderson, K.S. Pachis, F. Prebin, S.W. Cart, O. Terasaki, T. Ohsuna and V. Alfredsson, J. Chem Soc., Chem_ Commun., (1991) 1660. J. Klaas, IC Kulawik, G. Schulz-Ekloff and N.I. Jaeger, Stud. Surf. Sci. Catal., 84 (1994) 2261. M.1L Boccuti, I~M. Rao, A. Zecchina, G. Leofami and G. Petrini, Stud. Surf. Sci. Catal., 48 (1989) 133. J.S. Reddy, 1L Kumar and P. Ratnasamy, Appl. Catal., 58 (1990) L1. A. Zecchina, G. Spoto, S. Bordiga, A. Ferrero, G. Petrink G. Leofanti and M. Padovan, Stud. Surf. Sci. Catal., 69 (1991) 251. G.W. Skeels and E.M. Flanigen, ACS Symp. Set., 398 (1989) 420. C.B. Dam and M.E. Davis, Appl. Catal., 143 (1996) 53. A Corma, P. Esteve, A. Martinez and S. Valencia, J. Catal., 152 (1995) 18. M.A. Camblor, M. Constantini, A Corma,L. Gilbert, P. Esteve, A. Martinez and S. Valencia, Chem Commun. (1996) 1339.

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

1007

OXIDATIONS CATALYZED BY ZEOLITE Ti-UTD-1 Kenneth J. Balkus, Jr.* and Alia K. Khanmamedova

University of Texas at Dallas, Department of Chemistry, Richardson, TX 75083-0688, United States

SUMMARY Titanosilicates have been synthesized which have the large pore structure of UTD-1. These molecular sieves were prepared using the metal complex Cp*2Co§ (bis(pentamethylcyclopentadienyl)cobalt(lll) ion) as the template. Ti-UTD-1 has been found to be an effective catalyst for the oxidation of alkanes, alkenes and phenols using hydrogen peroxide as well as the larger oxidant t-butylhydroperoxide. The channel structure defined by 14 membered rings in Ti-UTD-1 also allows the conversion of larger substrates such as 2, 6- di-t-butylphenol. 1. INTRODUCTION UTD-1 (University of Texas at Dallas) is a novel large pore high silica zeolite [1-4] with many promising applications in the area of catalysis. The structure ofUTD- 1 is based on one dimensional channels nmning in parallel [5]. The elliptical shaped channels are defined by 14 tetrahedrally coordinated silicon atoms with pore opening of 10 x 7.5 A as shown in Figure 1. Silicon in the UTD-1 framework may be replaced (using direct synthesis) by many

Figure 1. View ofUTD-1 along the [010] direction transition metals, including titanium [3,6,7]. The Ti-UTD-1 preparation involves the novel use of a metal complex template such as bis(pentamethylcyclopentadienyl) cobak(m) hydroxide as well as titanium ethoxide as the source of titanium As much as 3.5% titanium by weight has been incorporated into Ti-UTD-1 during synthesis. It was anticipated that Ti-

1008 UTD-1 may have catalytic properties such as those observed with the commercially sucsessful titanium silicalite TS-1 catalyst which is effective for alkane and alkene oxidation as well as phenol hydroxylation in the presence of hydrogen peroxide [8]. The large pore nature of Ti-UTD-1 should allow the reaction of large substrates such as 2,6-di-tertbutylphenol as well as the use of oxidants such as tert-butylhydroperoxide (t-BHP) which are too large for the medium pore TS-1 zeolite. Ti-UTD-1 offers an opportunity to examine reactivity in pore space greater than Ti-beta but less than the mesoporous Ti-MCM-41 type molecular sieves. In the present study results for the peroxide based oxidation of cyclohexane, cyclohexene and 2, 6- di-tert-butylphenol will be presented. 2. EXPERIMENTAL Ti-UTD-1 (3.5 % Ti by weight) was prepared and characterized as previously described [3]. The catalytic oxidations were carried out as batch reactions in sealed glass vials (15 ml) at 25, 50 and 60~ In a typical reaction the reactor was charged with 0.1 grams of Ti-UTD-1 catalyst, 6 mmol of substrate, 3 mmol (or 12 mmol) of 90% aqueous tert-butyl hydroperoxide (t-BHP) and 20 mmol (or 27 retool) of acetone or tert-butanol as a solvent. In the case of cyclohexene oxidation using hydrogen peroxide (H202), 8 mmol of olefin and 2 mmol of H202 in t-butanol as well as acetonitrile solvent (80 mmol) were used. All reactions were sampled by syringe through a rubber septum and analyzed by gas chromatography using an HP 5840A capillary GC equipped with a 15 m AT-1 capillary column and a flame ionization detector. Products were verified by known standards and/or GC-MS. 3. RESULTS AND DISCUSSION

3.1 Cyclohexane oxidation The oxidation reactions involving hydrocarbons are frequently important components in many industrial processes yet can be among the most difficult of reactions to accomplish selectively. Cyclohexane serves as a good model substrate but itself is commercially important as a precursor to adipic acid. We have previously reported that Ti-UTD-1 is effective catalyst for the oxidation of cyclohexane using t-butylhydroperoxide [7]. H202 which can donate 47% of its oxygen in comparison with only 17.8% for t-BI-IP [9], is the preferred oxidant. However, the results for cyclohexane oxidation using H202/Ti-UTD-1 are disappointing. Although, Ti-UTD-1 should provide a hydrophobic environment at the active sites, the large pore nature of this material may allow more access to water. Even TS-1 exhibits very low activity for cyclohexane conversion using H202 as the oxidant [10]. The tBI-IP oxidations are more efficient but produce a range of products. Scheme 1 illustrates the oxidation products generally observed in cyclohexane oxidation, where the preferred product is adipic acid. The major product in the case of Ti-UTD-1 catalyzed reactions is generally cyclohexanone with smaller amounts of the alcohol and adipic acid. It should be noted that decomposition of the cobalt complex used as the template generates nanoclusters of a cobalt oxide on the surface of the zeolite [4]. Cobalt oxide and cobalt complexes are effective as peroxide decomposition catalysts [11,12] and should be removed from the :~,eolite to better evaluate the role of the titanium sites.

1009 -~

I V

copt Adipic acid

\

o _ , . Q"-

CO2H Glutaric acid CO2H

~~~C02H 0

Succinic acid

Scheme 1 Table 1 shows some representative results for cyclohexane oxidation with t-BHP over calcined Ti-UTD-I(Co) and Ti-UTD-1 (acid treated to remove cobalt [4]). In acetone the cyclohexane is converted to cyclohexanone as the major product with trace amounts of adipic acid as well as the other carboxylic acids and diketones in Scheme 1. Changing the solvent from acetone to t-butanol improves the conversion and the yield of adipic acid, however, the peroxide efficiency decreases. Increasing the amount of peroxide in the system also improves conversion but the peroxide efficiency is decreased. The product distribution observed with the Ti-UTD-l(Co) samples suggets a homolytic process catalyzed by cobalt. Removal of the cobalt lowers the activity in t-butanol but improves the peroxide efficiency. Table 1. Results for cyclohexane oxidation using t-BHP at 60~ Catalyst

Hours

% Conversion

% Yield

% Efficiencyr

K e A e AA ~ 24 10.8 7.2 3.5 0 36 24 22.2 11.3 4.0 1.7 20 24 29 12.5 3.5 0.4 17 24 13.8 1.5 4.3 trace 34 50 20.5 7.2 3.7 1.2 32 a. acetone solvent, b. t-butanol solvent; c. cyclohexane:t-BHP = 2:1, d. cyclohexane:t-BHP : 1:2; e. K = cyclohexanone, A = cyclohexanol and AA = adipic acid; f. efficiency = (mmol cyclohexane converted/retool t-BHP converted) x 100

Ti-UTD- 1 (Co)a'c Ti-UTD- l(Co) b'c Ti-UTD- I(Co) a'd Ti-UTD- 1b,c

1010 3.2 Cyclohexene oxidation

Oletins are generally much easier to oxidize than alkanes but again the selectivity can be a challenge. Cyclohexene was employed as a model substrate with the preferred product being the corresponding epoxide. Table 2 shows the results for cyclohexene oxidation catalyzed by Ti-UTD-1 (cobalt free) using H202 in t-butanol solvent. At 50~ a reasonable conversion of cyclohexene was observed with reasonable selectivity for the epoxide. Even though the substrate/peroxide ratio was 4:1 there was a fairly high efficiency for hydrogen peroxide utilization. The diol product is the result of epoxide hydrolysis while the allylic products (K + A) reflect a competing homolytic process. In acetonitrile solvent the Table 2 Results for cyclohexene oxidation using H202 in t-butanol solvent at 50~ Hours

% Selectivitya'b

% Conversion

% Efficiency

CYO A K diol 1 1.5 33.3 78 11.1 22.2 33.3 3 3.4 15.0 78 10.0 25.0 50.0 27 11.0 1.5 79 9.4 23.4 65.6 a. CYO = cyclohexene oxide, A = 2-cyclohexen-1-ol, K = 2-cyclohexen-1-one, diol = 1,2cyclohexendiol.; b. Selectivity was calculated as ratio (mmol product/mmol total product) xl00. conversion improved over the same time period in t-butanol. Additionally, the selectivity for cyclohexene oxide improved with very little hydrolysis early in the reaction. The peroxide efficiency improved in acetonitrile as well. These results suggest that there is a strong solvent dependency for this process. Therefore, the oxidation of cyclohexene was run in several solvents and mixtures including acetone and methanol but the best results so far for conversion and selectivity to the epoxide were obtained in acetonitrile. Polar solvents such as water, acetone and methanol have been previously been noted to have an inhibitory effect on the epoxidation reaction in mesoporous titanosilicates [ 11] and Ti-beta [ 13]. In these cases it is proposed that the solvent coordination at the oxide surface blocks the active sites. Table 3 Results for cyclohexene oxidation using H202 ill acetonitrile solvent at 50~ Hours

1.5 3 23

% Conversion

9.5 12 26

% Efficiency

% Selectivity CYO 79.3 70.3 45.1

A 6.9 8.1 12.1

K 13.8 16.2 23.1

diol tr. 5.4 19.7

,

,,

79 84 99

1011 This trend in large pore materials appears to be the opposite from that observed in the medium pore TS-1 catalyst [ 14]. The effect of acetontrile maybe to poison acid sites and/or to allow a greater concentration of substrate near the active site. The later has been proposed to be the dominant factor in Ti-beta [ 13]. The aprotic acetonitrile does not form the extensive hydrogen bonded networks expected for water or the alcohols. However, the acetonitrile is polar enough to allow more substrate into the relatively hydrophobic zeolite. A hydrogen bonded titanium hydrogen peroxo complex has been suggested as an important intermediate where a silanol, titanol or protic solvent stabilizes the structure. In the case of Ti-beta it was proposed that acetonitrile solvent allowed water to function in this capacity where as in other solvents the water would be displaced [13]. As a donor the acetontrile may also bind weakly enough to the titanium to affect the peroxide decomposition. Scheme 2 shows titanol stabilized titanium hydrogen peroxo species in equilibrium interacting with acetonitrile. This would acount for in part why the efficiency is lower and level of autoxidation products is higher in protic solvents. It is interesting that the large pore titanosilicates seem to exhibit similar solvent behavior.

I

H a C - CN:.- -~Ti

/

"o

\ H

o

/n...

\

I H a C - CN:- 9.,..-Ti

O/

+

H+

~O

~0 Scheme 2

A disadvantage of using n202 as the oxidant is low stability with respect to radical decomposition especially in presence of the catalytically active Ti(IV) species. Solvent can certainly influence the homolytic process but we may also add radical traps to supress the formation of allylic products. If one adds hydroquinone to the H202/Ti-UTD-1 system in aeetonitrile then the yield of cyclohexene oxide improves to over 85% while the allylic ketone and alcohol are reduced to trace impurities. The amount of diol formed also decreases but the presence of water in the system is unavoidable and some hydrolysis of the epoxide is expected. Changing the oxidant to t-butylhydroperoxide is expected to improve selectivity. The large pores of Ti-UTD-1 allow the use of t-BHP in contrast to TS-1 where this oxidant is too big. Figure 2 shows the results of cyclohexene oxidation using t-BHP (80 wt.% in ditert-butylperoxide) and t-BHP (90 wt.% in water) after two days of reaction at 60~ where the solvent was acetonitrile. Clearly the presence of water lowers the cyclohexene conversion and the selectivity for epoxide formation.

1012

70

%

60-

50-

~d

40-

30"

20

1

2

3

4

5

6

7

CONVERSION, % Figure 2. Selectivity for cyclohexene oxide versus conversion of cyclohexene for 1 - t-BHP in di-t-butyl peroxide (80%) and 2 - t-BHP in water (90%).

3.3 Oxidation of 2,6-Di-tert-butyl phenol A characteristic feature of Ti-UTD-1 which we would like to exploit is the 14 membered ring system where substrates too large for other zeolites can effectively be oxidized. The large 2,6-di-tert-butyl phenol (2,6-DTBP) has been used as a susbtrate for titanium containing mesoporous molecular sieve catalyzed oxidations [15] to yield the corresponding quinone as shown in Scheme 3 below. Table 4 shows the results for 2,6-DTBP oxidation at 65~ using H202 as the oxidant in acetone solvent. The selectivity to the quinone

-'~OH

+ H20~

0 = = ~

\

N Scheme 3

0

1013 is quite high over both Ti-UTD-1 and the mesoporous titanosilicates with no evidence of the dimer product after 5 hours. The conversion wth the Ti-UTD-1 catalyst is comparable to TiMCM-41 over a similar time frame while the Ti-HMS is somewhat more active. Clearly these results indicate that the large pore Ti-UTD-1 is capable of addressing large substrates. Table 4 Results for the oxidation of 2,6-di-tert-butyl phenol Samples

Pores size (A)

Hours.

% Conversion

Ti-UTD-1

10 x 7.5

Ti-MCM-4 la Ti-HMS a a. Reference 15

- 30 - 30

4.5 40 2 2

20 90 20 83

% Selectivity > 99 95 >98 >95

4. CONCLUSIONS The novel zeolite UTD-1 with titanium in the framework (up to 3.5% by weight) is an effective catalyst for the oxidation of cyclohexane, cyclohexene and 2,6-di-tert-butyl phenol. The catalytic behavior is similar to that of other large pore zeolites and mesoporous molecular sieves modified with titanium which includes solvents effects. Additionally, TiUTD- 1 allows the use of oxidants and substrates too large for the commercial TS- 1 catalyst. We are currently evaluating further the role of solvent and oxidant in an effort to improve selectivity as well as expand the utility of this material in oxidation catalysis. ACKNOWLEDGMENT The financial support of this work by the National Science Foundation, the Robert A.Welch Foundation and Chevron Research and Technology Company is gratefully acknowledged. REFERENCES

.

K. J. Balkus, Jr. and A. G. Gabrielov, The Synthesis of Novel Molecular Sieves using Metal Complexes as Templates, U.S.Patent No. 5,489,424 (1996). K.J. Balkus, Jr., A.G. Gabrielov and N. Sandier, Mater. Res. Soc. Symp. Proc., 368 (1995) 369. K. J. Balkus, Jr., A. G. Gabrielov and S. I. Zones, Stud. Surf. Sci. Catat, 97 (1995) 519. K. J. Balkus, Jr., M. Biscotto and A. G. Gabrielov, Stud. SurE. Sci. Catal., 105 (1997) 415. C.C. Freyhardt, M. Tsapatsis, R.F. Lobo, K.J. Balkus, Jr. and M.E. Davis, Nature, 381 (1996) 295.

1014

,

9. 10. 11. 12. 13. 14. 15.

K. J. Balkus, Jr. and A. G. Gabrielov, The Synthesis of Novel Molecular Sieves using a Metal Complex as Template, U.S.Patent No. 5,603,914 (1997). K, J. Balkus, Jr., A. Khanmamedova, A. G. Gabrielov and S. I. Zones, Stud. Surf. Sci. Catat, 101 (1996) 1341. B. Notari, Adv. Catal. 41 (1996) 253 and references therein. G. Strukul in Catalytic Oxidations with Hydrogen Peroxide as Oxidant, Kluwer, Boston (1992), 6. U. Schuchardt, H.O. Pastore and E.V. Spinace, Stud. Surf. Sci. Catal., 84 (1994) 1877. Z. Liu, G. M. Crumbaugh and 1L J. Davis, J. Catal. 159 (1996) 83. C.B. Roy, J. Catal.,12, (1968), 129. A. Corma, P. Esteve and A. Martinez, J. Catal. 161 (1996) 11. M.G. Clerici, Appl. Catal., 68 (1991) 249. P.T. Tanev, M. Chibwe and T. J. Pinnavaia, Nature, 368 (1994) 321.

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 1997 Elsevier Science B.V.

1015

Zeolite T i t a n i u m Beta" A selective catalyst in the M e e r w e i n - P o n n d o r f V e r l e y - O p p e n a u e r reactions. J.C. van der Waal a, P.J. Kunkeler a, K. Tan b and H. van Bekkum a a Laboratory of Organic Chemistry and Catalysis, Delft University of Technology, Julianalaan 136, 2628 BL Delft, The Netherlands b Tianjin Research Institute of Chemical Industry, Dingzigu, Hongqiao, 300131 Tianjin, PR China Zeolite titanium beta has been tested in the liquid- and gas-phase Meerwein-PonndorfVerley reduction of cyclohexanones and the Oppenauer oxidation of cyclohexanols. A high selectivity towards the thermodynamically unfavourable cis-alcohol was observed, which has been ascribed to transition-state selectivity in the pores of the zeolite. Under gas-phase conditions the dehydration of alcohols to cycloalkenes is observed as a side reaction. The catalyst was found to be active even in the presence of water and ammonia.

1. INTRODUCTION The Meerwein-Ponndorf-Verley reduction of carbonyl compounds and the Oppenauer oxidation of alcohols, together denoted as MPVO reactions, are considered to be highly selective reactions. For instance, C =C double bonds are not attacked. In MPV reductions a secondary alcohol is the reductant whereas in Oppenauer oxidations a ketone is the oxidant. It is generally accepted that MPVO reactions proceed via a complex in which both the carbonyl and the alcohol are coordinated to a Lewis acid metal ion after which a hydride transfer from the alcohol to the carbonyl group occurs (Fig. 1) [1]. Usually, metal sec-alkoxides are used as homogeneous catalysts in reductions and metal t-butoxides in oxidations [1]. RI-_

"')==o+ ;XR, Fig. 1

u

R3

-

.,

The Meerwein-Ponndorf-Verley-Oppenauer reaction [1].

Zeolites are crystalline metal oxides which have potential as regenerable heterogeneous catalysts in various organic reactions [2]. Because of their unique microporous structure, zeolites are especially promising in the field of shape-selective

1016 catalysis. As far as we know, only very few examples of the use of zeolites in MPVO reactions have been reported [3-7,9]. The reactions were carried out in the gas-phase over zeolites A, X and Y, exchanged or impregnated with alkali or alkaline-earth cations [3-5]. Shape-selectivity was only observed by Shabtai et aL[4] in the conversion of citronellal over zeolite X. It was shown that selectivity could be tuned by the size of the exchanged metal ion. In NaX there was enough space for the substrate to perform an intramolecular ring closure to isopulegol, whereas over CsX reduction to the linear citronellol was observed. Similar steric effects were also found for various other substrates. Recently Creyghton et aL [6,7] reported the use of zeolite beta in the MPVO reduction of 4-t-butylcyclohexanone. The high selectivity towards the thermodynamically less favoured cis-alcohol is explained by a restricted transition-state around a Lewis-acidic aluminium in the zeolite pores. When using an aluminium-free zeolite, titanium beta, in the epoxidation of olefins, we have shown that Ti-beta has acidic properties when alcoholic solvents were employed [8]. This was ascribed to the Lewis-acidic character of titanium in the zeolite framework. As we reported very recently [9], Ti-beta is found to be an excellent catalyst in MPVO reactions with a tolerance for water. Here, results are presented on the high selectivity, stability and low by-product formation of the catalyst, Ti-beta, in both the liquid-phase and gas-phase MPVO reactions.

2. EXPERIMENTAL

Zeolite titanium beta (Ti-beta) was synthesized according to van der Waal et al. [8,9] using di(cyclohexylmethyl)dimethylammonium hydroxide (DCDMA.OH) as the template. For a typical synthesis, 0.25 g titanium(IV) ethoxide (TEOT) was added to 30.0 g of a 19.5 % w/w DCDMA.OH solution and the mixture was stirred until all TEOT was dissolved. To facilitate the dissolution of TEOT, 1 ml of H20 2 (30 % w/w aqueous) was added. To the resulting clear solution 3.0 g Aerosil 200 (Degussa), 0.15 g seeds (all-silica beta [10] and 11.8 g water were added and the gel was aged for at least 24 h. After crystallisation (14 days at 140 ~ the zeolite (1.4 g) was filtrated, washed, dried and calcined at 540 ~ in air. Elemental analysis, performed using a LINK EDX system, showed a Si:Ti molar ratio of 76 and confirmed the absence of oligomeric titanium dioxide phases. Zeolite aluminium beta (Al-beta, Si:AI = 10) was prepared according to Wadlinger and Kerr [11] and all-silica beta (Si-beta, Si:A1 > 5000) was prepared according to van der Waal et aL [10]. Liquid-phase MPVO reactions were performed in 25 ml isopropanol (reductions) or 25 ml 2-butanone (oxidations) at 85 ~ using 2.5 mmol of the appropriate substrate: 4-t-butylcyclohexanone (4-Bu-ONE), 4-methylcyclohexanone (4-Me-ONE) or 4-t-butylcyclohexanol (4-Bu-OL, cis/trans mixture); 0.5 g zeolite or 0.25 mmol aluminium isopropoxide as the catalyst and 1,3,5-tri-t-butylbenzene as the internal standard. Samples were taken at regular intervals and analyzed by GC on a Carbowax CP-52 column and GC/MS. Gas-phase MPVO reactions were performed at 85 to 400 ~ in a fixed bed continuous down-flow reactor operated at atmospheric pressure under plug flow conditions. The catalyst, Ti-beta or N-beta (0.30 g), was diluted with 1.20 g a-quartz powder and processed to pellets then crushed to particles with a diameter of 0.7 - 1.0 mm. Reactant mixtures were pumped into a stream of preheated carrier gas (usually

1017 nitrogen) by means of a motor-driven syringe pump. The gas flow contained 10 vol.% isopropanol (reductions) or acetone (oxidations) and 1 vol.% of the appropriate substrate: 4-methylcyclohexanone (4-Me-ONE) or 4-methylcyclohexanol (4-Me-OL). The total gas flow was 50 ml/min and the molar gas flow of 4-Me-ONE or 4-Me-OL was 2.04 10-5 mol/min (WHSV = 2.9 gtotal goat-1 h'l) 9Samples of the reactor effluent were taken regularly and analysed on a CP-Sil-19 column. 3. RESULTS

3.1 Liquid-phase MPVO In the liquid-phase, the MPVO reactions catalysed by Ti-beta are highly selective and have low by-product formation. In the reduction of 4-t-butylcyclohexanone and 4-methylcyclohexanone a high selectivity (99 % and 98%, respectively) towards the thermodynamically unfavourable cis-alcohol is observed, similar to that observed with A1beta (Table 1). The high selectivity towards the cis-alcohol over beta-type zeolites was explained by Creyghton et aL [6] on the basis of transition-state selectivity. It can be seen in Fig. 2, that the transition-state leading to the cis- and trans-alcohol differ substantially in spatial requirements. The transition-states leading to the cis-isomer is aligned with the zeolite channel while that leading to the trans-alcohol occupies a much more axial position. Table 1 MPVO reduction of cyclohexanones over beta-type catalysts in refluxing isopropanol. Substrate

Catalyst

TOF a

Conversion b (%)

Selectivityb cis : trans

4-Me-ONE

Ti-beta

1.04

33.7

99:1

4-Me-ONE

M-beta

> 12

100.0

98:2

4-Me-ONE

AI(OPr)3

16.7

99.8

27:73

4-Bu-ONE

Ti-beta

2.26

64.9

98:2

4-Bu-ONE

Al-beta

> 12

100.0

95:5

4-Bu-ONE

Si-beta

4-Bu-ONE

AI(OPr)3

0.0 8.72

100.0

36:64

a Initial turn-over-frequency in mol ketone per mol of titanium or aluminium per hour. b Conversion and selectivity after 6 h reaction based on the initial amount of ketone added.

1018

H-~Me Me

....

/ Fig. 2

O~ /0

/

Proposed transition-states for the formation of cis-4-t-butylcyclohexanol (left) and trans-4-t-butylcyclohexanol (right).

In the Oppenauer oxidation of an equimolar mixture of cis- and trans-4-tbutylcyclohexanol, the cis-alcohol is converted selectively over both Ti-beta and Al-beta (Table 2), which is probably due to the same spatial restriction on the transition-states as for the reduction depicted in Fig. 2. In this way, the MPVO reduction and oxidation are in harmony as to the high transition-state selectivity. Table 2 Liquid-phase Oppenauer oxidation of 4-t-butylcyclohexanol (1:1 cis/trans mixture)with butanone as the oxidant. Catalyst

Conversion a of cis-alcohol (%)

Conversion a of trans-alcohol (%)

By-products a'b (mg)

Ti-beta

52.1

4.6

22.4

Al-beta

98.6 49.4 d

47.3 5.7c

317.3 63.8c

Ti(OiPr)4

6.1

12.9

22.7

Al(OiPr)3

5.3

7.4

56.5

a After 6 h reaction time. b Not MPVO related, predominantly from 2-butanone via aldol condensation, c After 15 min reaction time. d After 15 min. The most important side reaction in heterogeneously catalysed MPVO reactions is the acid-catalysed aldol condensation. Aldol products are usually observed during the Oppenauer oxidation of alcohols, when a surplus of ketone or aldehyde is used as the oxidizing agent and the solvent. The low amount of by-products formed when Tibeta was used as the catalyst, demonstrates the advantage of the titanium system over A1beta. This is probably caused by the much weaker BrCnsted acidity of the solvated titanium site [8] compared with the strong H+-acidity of the aluminium site in Al-beta. As we have shown earlier Ti-beta has a high tolerance towards water, which further shows the catalytic potential of Ti-beta in MPVO reactions [9].

1019

3.2 Gas-phase MPVO From an industrial point of view, gas-phase reactions are often preferred due to their ease of operations. The reduction of 4-methylcyclohexanone and the oxidation of cis- and trans-4-methylcyclohexanol over Ti-beta and Al-beta in the gas phase were studied at 100~ As can be seen from Fig. 3, both Ti-beta and Al-beta are active, but Tibeta has a considerably lower rate of deactivation. The deactivation of Al-beta is probably caused by the higher acidic strength of the protonic aluminium site compared to the non-protic titanium site. Another difference between the two catalysts is the pronounced dehydration of the alcohols formed over Al-beta. Two important differences between the gas-phase and liquid-phase reaction were observed. The most striking is the selectivity towards the cis-alcohol. Under liquid phase conditions (Table 1.) a selectivity of 99 % towards the cis-alcohol was observed, while in the gas-phase over Ti-beta only 53 - 62 % cis-alcohol was obtained. Furthermore, dehydration of the alcohols formed to 4-methylcyclohexene (4Me-ENE) was observed in the gas phase, while no trace of dehydrated products could be detected in liquid phase MPV reductions. The somewhat higher temperature of 100 ~ compared with 85 ~ for the liquid-phase experiments, could not explain this behaviour, since lowering the temperature to 85 ~ still resulted in a 4-Me-ENE selectivity of approximately 10 % over Ti-beta. The remarkable differences in selectivity between the gas- and liquid-phase MPVO reaction induced us to study the evolution of products as a function of the temperature, in the reduction of 4-methylcyclohexanone. From Fig. 4 it can be seen that the selectivity to the cis-alcohol decreases at higher temperatures; initially the selectivity to the trans-alcohol increases but at higher temperatures dehydration to 4-methylcyclohexene becomes the major reaction. At still higher temperatures (> 200 ~ the 4-methylcyclohexene is subsequently isomerised to the more stable 1-methylcyclohexene. I00

80

'N

60

20

0

2

4 Time on stream [h]

6

8

a) Ti-beta Fig. 3a

The gas-phase MPV reduction of 4-methylcyclohexanone with isopropanol over Ti-beta at 100~ 9 = conversion; 9 = cis-4-Me-OL; x = trans-4-Me-OL; n = 4-Me-ENE.

1020

100

80

~.43

------'~

!o A

o

-

i

A

4 ' Time on stream [h]

i

b) Al-beta Fig. 3b

The gas-phase MPV reduction of 4-methylcyclohexanone with isopropanol over Al-beta at 100~ 9 = conversion; 9 = cis-4-Me-OL; x = trans-4-Me-OL; [] = 4-Me-ENE. 100

80 ,___,

c~

.~ 40

0

Fig. 4

50

100

150 200 250 Temperature [ * C]

300

350

Influence of the temperature on the selectivity in the reduction of 4-methylcyclohexanone with isopropanol. Temperature increment was 0.2 ~ 9 = conversion; 9 = cis-4-Me-OL; x = trans-4-Me-OL; [] = 4-Me-ENE.

Since the 4-methylcyclohexene can be formed via dehydration of either of the two alcohols formed, both alcohols were tested in the Oppenauer oxidation, using acetone as the oxidant at 100 ~ It was observed that the cis-alcohol (Fig. 5a) is easily oxidized to the corresponding ketone, while the trans-alcohol showed a much lower activity. This is in accordance with the transition-states assumed for liquid phase MPVO reactions (Fig. 2). It can be observed further that the dehydration of the trans-alcohol only occurred to very limited scale (Fig. 5b) whereas dehydration of the cis-alcohol was an important side-

1021 reaction. This can be understood by assuming an E2-mechanism, in which the axial OH group in the cis-isomer is in the ideal position for water elimination. Another important side reaction for both alcohols was isomerization. It is therefore proposed that the alkene is formed mainly from the cis-alcohol and that at low temperatures the trans-alcohol can only be dehydrated if it is first isomerised to the cisalcohol via an MPVO transition-state. Comparing the deactivation of the Ti-beta catalyst in the oxidation (Fig. 5) and reduction reactions (Fig. 3) shows that the deactivation is more pronounced during oxidative conditions. This is probably caused by the high amount of ketones present, which easily form aldol condensates which may plug the zeolite channels, thus inhibiting access to the micropore system. 100

80

r~

..~ 40

i

r,.)

0

2

4 Time on stream [h]

6

a) Oxidation of cis-4-Me-OL 100

80

"~

60

.~ 40

I 2o

r.,)

0

2

4 Time on stream [h]

6

8

b) Oxidation of trans-4-Me-OL Fig. 5

The gas-phase Oppenauer oxidation of the 4-methylcyclohexanols with acetone over Ti-beta at 100~ a) cis-4-Me-OL and b) trans-4-Me-OL. 9 = conversion; 9 = cis-4-Me-OL; x = trans4-Me-OL; 9 = 4-Me-ONE. [] = 4-Me-ENE.

1022 The commonly used MPVO catalysts consist of metal alkoxides, which are easily hydrolysed to inactive oxides in the presence of water. Since the proposed catalytic species for the MPVO reaction also consists of an alkoxide intermediate [6,9] the influence of water and strong Lewis bases on the catalytic activity and selectivity was investigated. As already reported for the liquid-phase reaction [9], Ti-beta has a high tolerance for water due to its hydrophobic interior. As can be seen from Fig. 6a the presence of water is not detrimental to the activity of Ti-beta in the MPV reduction of 4-methylcyclohexanone. The temperature of 110 ~ at which a ketone conversion of 50% is measured, is identical to the temperature required for 50% conversion in the absence of water (Fig. 3), i.e. water has no effect whatsoever on the overall MPVO activity of the titanium site. 100

50

100

150 200 250 Temperature [ * Q

300

350

a) Reduction in the presence of water 100

9

--

w--r-~

80

60

.~

40

i 0 100

150

200 250 Temperature * C]

300

350

400

b) Reduction in the presence of ammonia Fig. 6

Temperature-programmed gas-phase MPV reduction of 4-methylcyclohexanone over Ti-beta in the presence of a) 2.66 vol.% water or b) 5 vol% ammonia. 9 = conversion; 9 = cis-4-Me-OL; x = trans-4-Me-OL; [] = 4-MeENE.

1023

The selectivity towards alcohols increased from 8 1 % to 97% at 85~ This enhanced selectivity can be ascribed to either a kinetic suppression of an irreversible dehydration or to a change in the alcohol/olefin equilibrium due to the higher amount of water present. In the case of a shift in equilibrium, it should be possible to oxidise olefins with ketones in the presence of water via in situ formed alcohols. In an attempt to oxidise cyclohexene with acetone in the presence of a water, no conversion to cyclohexanone was observed between 85 and 400 ~ It is therefore concluded that under the experimental conditions used, 4-methylcyclohexene is formed irreversibly from the cisalcohol (Fig. 7) and the increased selectivity towards alcohols should therefore be ascribed to a deactivation of the dehydration sites in the presence of water. OH

cis-4-Me-OL

4-Me-ENE

1-Me-ENE

Ti 4-Me-ONE

~

OH

trans-4-Me-OL

Fig. 7

Proposed reaction scheme for the MPV reduction of 4-methylcyclohexanone.

For liquid-phase reactions at 85 ~ we reported that small amounts of a strong base, e.g. pyridine, completely poisoned the catalyst [2]. It can be seen from Fig. 6b that, in the presence of ammonia, higher temperatures are required to reduce 4-methylcyclohexanone. The higher temperatures are required for the desorption of ammonia from the catalytically active site. The relatively low temperature of 305 ~ at which 50 % conversion is observed, suggests that the ammonia is not bonded to a strong acidic site. Since Br0nsted-acidic aluminium sites desorb ammonia at about 480 ~ this confirms that the MPVO reactions proceed via the titanium sites and not via any residual aluminium sites (Si:AI > 5000). The latter can also be concluded from Table 1; the all-silica analogue of zeolite beta [10] was found to be completely inactive even though it has a Si:A1 ratio similar to Ti-beta. It was also observed that in the presence of ammonia, no isomerisation of 4-Me-ENE to 1-Me-ENE occurred even at 400 ~ suggesting that the isomerisation requires strong Brcnsted acid-sites, most probably the residual aluminium sites. As was already shown in Fig. 3 and 5a, Ti-beta exhibits a low deactivation rate in the reduction of 4-methylcyclohexanone and a significantly higher deactivation rate in the oxidation of cis-4-methylcyclohexanol. In both cases, frequent regeneration of the catalyst will be necessary. Catalyst stability was tested by regeneration at 480~ in air after each run. No significant loss in activity, selectivity or product distribution was observed, even after 35 consecutive runs. Similar results were also obtained for the Tibeta used in liquid-phase reductions; after regenerating the Ti-beta catalyst 5 times, the same initial catalytic activity per gram of catalyst was observed.

1024 4. CONCLUSION Ti-beta is found to be an excellent catalyst in MPVO reactions under both liquid- and gas-phase conditions. Under liquid-phase conditions, a very high selectivity in the reduction of 4-substituted cyclohexanones towards the thermodynamically unfavourable cis-alcohols was observed. By-products were observed only during the oxidation of alcohols using ketone solvents and consisted primarily of aldol condensation products. Remarkable differences exist between the liquid-phase and gas-phase reactions under otherwise similar conditions. The selectivity towards the cis-alcohol is still above the thermodynamically expected value but significantly lower than under liquid-phase conditions. In contrast to the liquid-phase reactions, dehydration of the alcohols to the corresponding alkene is an important side-reaction. The oxidation of both the cis- and the trans-alcohol clearly showed that the olefin is exclusively formed from the cis-alcohol. Dehydration of the trans-alcohol is assumed to proceed by isomerisation via a MPVO mechanism to the corresponding cis-alcohol. The catalytic potential of the titanium-based catalyst is shown by the low amount of by-products formed over Ti-beta compared with Al-beta, the high resistance to water and the excellent stability of the catalyst with respect to regeneration.

ACKNOWLEDGEMENT Dr. Eddy Creyghton is thanked for valuable discussion and the Dutch Institute for Scientific Research (NWO/SON) for financial support. REFERENCES

1

For a review, see : C.F. de Graauw, J.A. Peters, H. van Bekkum and J. Huskens, Synthesis, 10 (1994) 1007. 2 P.B. Venuto, Microporous Mater., 2, (1994) 297; W. H61derich and H. van Bekkum, Stud. Surf Sci. CataL, 68, (1991) 631. 3 J. Shabtai, R. Lazar and E. Biron, J. Mol. Catal., 27, (1984) 35. 4 M. Huang, P.A. Zielinski, J. Moulod and S. Kaliaguine, Appl. CataI. A, 118, (1994) 33. 5 M. Berkani, J.L. Lemberton, M. Marczewski and G. Perot, Catal. Lett., 31 (1995) 405. 6 E.J. Creyghton, S.D. Ganeshie, R.S. Downing and H. van Bekkum, J. Chem. Soc. Chem. Commun., (1995) 1859. 7 E.J. Creyghton, S.D. Ganeshie, R.S. Downing and H. van Bekkum, J. Mol. Catal., in press (1997). 8 J.C. van der Waal, P. Lin, M.S. Rigutto and H. van Bekkum, Stud. Surf Sci. Catal., 105, 1093 (1997). 9 J.C. van der Waal, K. Tan and H. van Bekkum, Catal. Lett., 41 (1996) 63. 10 J.C. van der Waal, M.S. Rigutto and H. van Bekkum, J. Chem. Soc., Chem. Commun., (1994) 1241. 11 R.L. Wadlinger and G.T. Kerr, US patent, Appl. 3.308.069 (1967).

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 1997 Elsevier Science B.V.

1025

Selective Oxidation of Cyclohexane over Rare Earth Exchanged Zeolite Y

Emerson L. Pires, Martin Wallau, Ulf Schuchardt; Instituto de Quimica, UNICAMP, Caixa Postal 6154, 13083-970 Campinas-SP, Brasil; e-mail: [email protected]

Rare earth oxides are known to exhibit strong redox properties and to catalyze, in the presence of oxygen, the oxidation of hydrocarbons like n-butane, propylene and benzene to carbon dioxide and water. Recently the selective oxidation of cyclohexane to cylohexanol and cyclohexanone in the liquid phase catalyzed by SmC13 was reported. 1 Also Ce(IV) impregnated on a cation exchange resin was used as catalyst for the oxidation of alcohols to the corresponding carbonyl compounds. 2 Although these results demonstrate the performance of rare earth cations as catalysts for the oxidation of hydrocarbons, no attention has been given to the redox properties of rare earth cation exchanged zeolites, which are very well studied as acid catalysts in petrochemical processes. In this work we report first results on the redox properties of rare earth exchanged zeolites with FAU structure as catalyst for the oxidation of cyclohexane with tert- butyl hydroperoxide (TBHP) in the liquid phase. The catalysts were prepared from_ a sodium zeolite Y (NAY) with a Si/Al ratio of 3, by solid state ion exchange with CeC137H20, NdCI36H20, SmCI3"6H20 or YbCI36H20, respectively. The NaY was mixed with the rare earth chloride (REel3) in a molar ratio AI/REC13 of 3. The homogenized mixture was placed in a reaction tube, evacuated and heated to 450~ for 6 h. After cooling to room temperature, the solid was separated, carefully washed with water and dried at 120~ for 12 h. The catalysts were characterized by X-ray diffraction (XRD), FTIR spectroscopy and elemental analysis. The oxidation of cyclohexane was carried out in a suspension of the REY (200 mg) in 15.6 g of cyclohexane. 10 mmol of TBHP dissolved in 1.1 g of cyclohexane was added and the reaction mixture was stirred for 24 h under reflux. After filtering off the solid catalyst, the reaction products were quantified by gas chromatography, using internal standard and calibration curves. All peaks observed in the XRD patterns of the REY can be assigned to the FAU

1026 structure and the unit cell parameter ao, calculated by linear regression, does not vary compared to the parent NaY. The absence of a broad reflection between 20 and 25 o (20), which would indicate the presence of amorphous silica, demonstrates that the FAU structure remains intact after the solid state ion exchange. The decreased intensity of the reflections observed in the patterns of the REY may be attributed to an enhanced adsorption of the X-rays in the presence of the rare earth cations, rather than to a decreased crystallinity. It can be observed in the FTIR spectra that the ratio of intensities between the bands around 1000, 710 and 460 cm "1, attributed to internal tetrahedra vibrations (structure insensitive) 3 and the bands around 1130, 780 and 570 cm"1, attributed to the vibrations of the external linkages (structure sensitive), 3 do not differ between the parent NaY and the REY. This is a further confirmation that the FAU structure remains intact after the solid state ion exchange. The results of the catalytic oxidation of cyclohexane with TBHP are given in Table 1. The amounts of oxidized products obtained in a blank experiment are subtracted. Besides the main products cyclohexyl hydroperoxide (chhp), cyclohexanol (ol) and cyclohexanone (one), small amounts of other still unidentified products were also obtained. While NaY is inactive for the oxidation of cyclohexane, the activity of the REY increases in the following order: SmY NdY < YbY < CeY. An ol/one ratio close to unity, observed for NdY, SmY and YbY, indicates that both products are formed simultaneously from a common intermediate as already claimed for the cyclohexane oxidation catalyzed by SmC13.1 A probable reaction mechanism is the hydrogen abstraction catalyzed by the REY. The resulting cyclohexyl radicals react with molecular oxygen to cyclohexylperoxo radicals, which decompose in a bimolecular reaction to cyclohexanone, cyclohexanol and oxygen, or are reduced and protonated to cyclohexyl hydroperoxide. The enhanced cyclohexanol selectivity of CeY may be due to a selective decomposition of cyclohexyl hydroperoxide to cyclohexanol. The relative low activity of the REY is due to its high hydrophilicity, which leads to a rapid deactivation by adsorption of the polar reaction products. Studies to increase the activity of REY by using NaY with a higher Si/AI ratio are in progress. To the best of our knowledge, this is the first report on the redox activity of rare earth cations supported on zeolites in liquid phase oxidation.

1027

Table 1- Results of the oxidation of cyclohexane with TBHP catalyzed by REY catalyst NaY CeY NdY SmY YbY

ol (mmql) traces 1.14 0.22 0.39 0.39

.

one (mmol) traces 0.44 0.21 0.30 0.43

chhp (mmol) traces 0.74 0.56 0.22 0.46

ol/one 2.6 1.0 1.3 0.9

I. Yamanaka, K. Otsuka, J. Mol. Catal. A 1995, 95, 115 S. Kanemoto, H. Saimoto, K. Oshima, H. Nozaki, Tetrahedron Lett. 1984, 25, 3317 D.W. Breck, Zeolite Molecular Sieves, Wiley & Sons, New York, 1974

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3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 1997 Elsevier Science B.V.

1029

RATIONALLY DESIGNED OXIDATION CATALYSTS : FUNCTIONALIZED METALLOPORPHYRINS ENCAPSULATED IN TRANSITION METAL-DOPED MESOPOROUS SILICA (abstract) Lei Zhang, Tao Sun and Jackie Y. Ying Dept. of Chemical Engineering Massachusetts Institute of Technology, Cambridge, MA 02139, USA Catalytic oxidation by metalloporphyrins plays an important role in the conversion of both saturated and unsaturated hydrocarbons to valuable fine chemicals. The advantages in the development of metalloporphyrin systems that mimic cytochrome P-450 mono-oxygenases include substrate specificity, chemoselectivity and high catalytic activity under mild reaction conditions. Substantial efforts have been devoted to the development of an effective supported metalloporphyrin heterogeneous catalyst system that prevents self-oxidation of the active centers and allows for facile recovery of the catalyst. Mesoporous materials with a well-defined pore structure have recently attracted a great deal of research attention in catalytic applications. In this study, a rational design strategy has been established to prepare heterogeneous metalloporphyrin oxidation catalysts encapuslated in a hexagonally-packed mesoporous molecular sieve, providing (i) superior catalytic performance through a rational design of the matrix structure, and (ii) improved metalloporphyrin fixation through a tailored interaction between the catalyst and support. A series of mesoporous silicas known as MCM-41 has been synthesized with selective dopants for the metalloporphyrins. The microstructure of the materials was found to depend strongly on processing parameters, such as dopant concentration, aging time and temperature. In the case of Nb-doped silica (Nb/Si-TMS8,Nb/Si-TMS9), the formation of covalent bonding between the surface-exposed niobium sites and the functionalized groups of the iron porphyrin is crucial for immobilizing the iron complex within the mesoporous structure~ The fixation mechanism is different from the Coulombic forces and hydrogen bonding interaction involved in the conventional supported system. It effectively prevents leaching of the porphyrin and guarantees continued usage of the heterogeneous catalyst system. The well defined spacious mesoporous channels further allow for free diffusion of reactants and products. The functionalized metalloporphyrins, such as iron(III) meso-(tetra-aminophenyl)porphyrin bromide (FeTrva2PPr), encapsulated in transition metal-doped mesoporous silica showed high catalytic activity and improved selectivity for the epoxidation of olefins and hydroxylation of alkanes at ambient conditions. In the hydroxylation of cyclohexane, exclusive formation of cyclohexanol was observed with a yield of 53% after 6 hr. In epoxidation of cyclooctene, a single product of cyclooctene oxide was obtained with a conversion of 57% after 5 hr. No catalyst leaching was detected during the reaction over the supported catalystys, FeTNlI2PPBr/Nb/Si-TMS8 and FeTrcmPPBr/Nb/Si-TMS9. This novel rational design strategy provides a significant improvement on the fixation mechanism to achieve a better heterogneous catalyst. By manipulating the structural characteristics, such as the surface area, pore size, nature of the dopant and dopant concentration of the support material, as well as the functional groups of the iron porphyrin, the catalytic behavior can be controlled systematically via synthesis parameters.

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3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 1997 Elsevier Science B.V.

1031

Catalytic Oxidations with Biomimetic Vanadium Systems I.W.C.E. Arends, M. Pellizon Birelli and R.A. Sheldon Laboratory for Organic Chemistry and Catalysis, Delft University of Technology, Julianalaan 136, 2628 BL Delft, The Netherlands.

ABSTRACT Schiffs base complexes of vanadium were encapsulated in the supercages of zeolite NaY and their catalytic activities were tested in the epoxidation of several alkenes and allylic alcohols with tert-butylhydroperoxide The complexes investigated were VO(HPS) (vanadylN-(2-hydroxyphenyl)salicylideneimine) and VO(salen) (vanadyl-N,N'(bis)salicylideneimine). Particular attention was devoted to the question of leaching of vanadium during reaction. 1. INTRODUCTION The widespread presence of vanadium-containing enzymes, such as bromoperoxidase, has only recently been appreciated 1. For example, vanadium-bromo-peroxidase (V-BrPO) catalyzes oxidative bromination of compounds in seawater. This action can be mimicked effectively to halogenate compounds for synthetic purposes 2. The mechanism involves oxidation of Br- to the brominating species in media with vanadium and hydrogen peroxide. The identity of the intermediate is still poorly understood 3. On the other hand peroxocomplexes of vanadium have long been known for their oxidative capabilities in organic media 4. Mo, W, Ti, and V are effective for the epoxidation of alkenes with ROzH. This process forms the basis for the commercial production of propylene oxide using either tert-butyl hydroperoxide or ethylbenzene hydroperoxide and Mo or Ti as catalysts 5. There is general agreement that the mechanism involves an alkylperoxometal complex which transfers an oxygen atom to the olefinic double bond. For a particular alkylperoxovanadium(V) complex the crystal structure has been determined, which showed coordination by both oxygen atoms 6. As to the intricate details of the oxygen transfer step, analogous to molybdenum, both direct nucleophilic attack of the olefin on the electrophilic oxygen of the alkylperoxometal species, and coordination of the olefin to the vanadium, followed by rate limiting insertion into the metal oxygen bond seem possible (Figure 1, pathways (a) and (b) resp.) 7. As epoxidation of unfunctionalized alkenes catalyzed by vanadium is rather slow compared to e.g. molybdenum, the ligand environment is critical in allowing the above described heterolytic mechanism to be dominant 5,7,8 Epoxidation of allylic alcohols by peroxometal complexes has been demonstrated to be even faster and more selective due to coordination of the allylic alcohol through its alcohol group.

1032 Vanadium shows exceptional reactivity in this respect 9, which is consistent with the strong affinity of vanadium(V) for alcohol ligands.

/,,,,,O

But

vV_o 0

"'"'~

Figure 1

But

\But

Our aim was to design a highly selective heterogeneous stable vanadium catalyst, optimal for epoxidation, which can be fine-tuned via its ligands. A biomimetic approach consists of encapsulating a complex in the cages of a zeolite. The inorganic backbone in this sense mimics the environment in the enzyme 1~ Recent examples of these so called "zeozymes" comprise the encapsulation of iron-phthalocyanines 11 and manganese. b~pyndlnes . The principle is that owing to space restriction the complex stays in the zeolite and cannot diffuse out, provided that the complex is stable. This is essentially different from isomorphously substituted vanadium molecular sieves, like vanadium-silicalites and V-APO's. Especially, the latter ones have demonstrated problems in leaching 13'14. Few examples of vanadium encapsulated complexes are known. Recently VO(bipyridyl)215 and VO(salen) ( 1 ) 16 - a Schiffs base complex - were incorporated in zeolite NaY and studied at room temperature. VO(salen) however is known to be not very selective for catalytic epoxidation of alkenes. Mimoun showed that complex (2) vvO(HPS)(OOtBu) is a stable and much more efficient epoxidation catalyst 3 and HPS (N-(2-hydroxyphenyl)salicylideneimine) is, therefore, an interesting choice as a ligand. In this study VO(HPS) was synthesized within the cages of zeolite NaY and tested as a catalyst for oxidation with TBHP. In order to reach significant reaction rates, our studies were performed at 70~ Different substrates were studied, including allylic alcohols, and an attempt was made to distinguish between catalysis in homogeneous solution, versus catalysis in such a catalytically constrained environment within the cage of a zeolite. VO(salen) was studied as a comparison. 9

9

12

O

.o

o. II/o \

!

(1)

(2)

1033 2. E X P E R I M E N T A L

2.1 Materials NaY was acquired from AKZO, with a Si/A1 ratio of 2.44, washed with aq. NaOAc and water before use. For a typical synthesis vanadium was introduced into the zeolites via ion exchange in 2 mM VOSO4 solutions. The zeolite was dried under nitrogen at 200~ and the dry material was mixed under nitrogen with a 2-5 fold excess of ligand and incubated for 6 h at 125~ (salen) or 190~ (HPS). Incubation with HPS gave a brown material, whereas with salen a yellow-orange material was obtained. The zeolite was soxhletted with acetonitrile or acetone and consequently with dichloromethane. Differences in treatment are given in Table 1. HPS (N-(2-hydroxyphenyl)salicylideneimine) was synthesized via condensation of 2aminophenol and salicylaldehyde in boiling ethanol and recrystallized from methanol. Homogeneous vvO(HPS)(OEt)[EtOH] and vVo(salen)(OEt) complexes were synthesized by complexation of VO(i-OPr)3 and the ligand in ethanol 17 Complexes were checked with 51V and 1H-NMR. Materials; anhydrous TBHP was made by extracting a 70% TBHP solution (in water) in chlorobenzene and drying over 3~, molecular sieves, to result in a 5 M solution. Cyclohexene and cyclooctene were distilled and were passed through a column of basic alumina before use. Vanadyl(IV) sulphate hydrate, Salen (N,N'-bis(salicylidene)ethylenediimine), 1,2,-dichloro-ethane (DCE), dichloromethane (DCM), acetone, and acetonitrile were reagent grade and used as received. (-) Carveol was purchased as a 1.48:1 mixture of trans- and cis-carveol and used as such. Cumylhydroperoxide was purchased as a 80% solution in cumene and used as received.

2.2 Catalyst characterization The crystal structure was checked by X-Ray diffraction using a Philips PW 1877 automated powder diffractometer with CuK~ radiation. AAS and ICP-AES measurements were carried out with Perkin Elmer instruments type 1100, 5000 Z and Plasma-40. UV/VIS measurements were carried out on a Varian Cary spectrometer. Diffuse reflectance spectra were recorded against a barium sulphate reference on the Varian Cary-3 and reflectance spectra were converted according to the Kubelka-Munk equation.

2.3 Catalytic oxidation experiments The catalytic oxidation experiments were carried out in a round bottom flask equipped with condenser and stirrer. Typically, 6 or 12 mmol of substrate, 2 mmol bromobenzene (internal standard), 9 ml solvent (1,2-dichloroethane) and 100 mg zeolite (which contains typically around 2.9 lamol V (0.15wt%), TBHP/V ratio = 2070) were heated to 70~ after which 6 mmol of TBHP in a 5.3 mmol chlorobenzene solution (which at the same time can function as an internal standard) were added to start the reaction. A sample was taken immediately afterwards. Before and during the reaction the mixture was purged with nitrogen for oxidations with cyclohexene and cyclooctene. Samples were filtrated over cotton wool and/or alumina, and triphenylphosphine was added to remove TBHP. In case of acetone as the solvent at 70~ reactions were performed in a 50 ml autoclave and the reaction mixture was only purged with nitrogen before heating. After reaction TBHP was determined by iodometric titration.

1034 Quantitative analyses were carried out by GC with a semi-capillary column, CP-Sil-5 CB (50 m x 0.53 mm) for cyclooctene or on a CB wax 52 CB (50 m x 0.53 mm) for cyclohexene and carveol. Qualitative identification of peaks was made by reference samples and GC/MS analyses. Also cyclohexenyl-tert-butylperoxide could be analyzed in this way. Quantitative analysis was performed by using molar responses with respect to the internal standard. 3. RESULTS 3.1 Characterization Schiffs base complexes exhibit very intense UV absorptions. In Figure 2, UV spectra in Nujol and/or DRS spectra are shown for the synthesized materials, VO(HPS)-Y and VO(salen)-Y and their homogeneous counterparts. The spectra for the homogeneous complexes were recorded in DMF, to which two drops of a 0.01 M NaOH solution was added. Apparently in this matrix ethoxy ligands are replaced by OH as a ligand and the characteristic absorption maxima at 424 nm for VO(HPS) and 360 nm for VO(salen) can be seen. This environment closely resembles the ligand environment in the zeolite, where vanadium is surrounded by oxygens from the framework. In dichloromethane for VO(HPS)(OEt) only shoulders at 309 and 370 nm, and for VO(salen) at 260 and 290 nm were observed. Spectra of in-situ generated complexes of vIVoso4 and HPS or salen look identical to those of the preformed V v complexes. Figure 2 clearly indicates that the complexes are present within the zeolites, and for VO(salen) a 20 nm shift occurs upon complexation. Separate Shiffs bases display maxima at 350 and 317 nm for HPS and salen resp.

~(iii)

VO(HPS)

~,i)

VO(salen)

(i) i

200

400 600 wavelength (nm)

800

200

400 600 wavelength (nm)

800

Figure 2a + b: UV/VIS spectra for VO(HPS) and VO(salen). (i) homogeneous complex, 0.1 mM dissolved in DMF + OH added (ii) DRA spectra for VO(HPS)-Y [B] and VO(salen)-Y [HI (c) Solution spectra in nujol for VO(HPS)-Y [D].

ICP analysis indicates that in a typical procedure the V content decreases from approximately 0.23-0.30 wt% before ligand incubation, to 0.15-0.29 wt% V after incubation with HPS and soxhlet extraction, apparently depending on the efficiency of the incubation. Washing with NaOAc, followed by soxhlet extraction with acetone removes another 0.03 0.13 wt% of vanadium. Data are given in Table 1. VO(salen) [J] was prepared with a 50 fold excess of salen and 24 h incubation, approaching as much as possible the conditions in ref. 16.

1035

One unit cell in the faujasite structure contains 8 supercages and in our case 58 A1 atoms. A ratio of 58/1 for A1/V therefore corresponds to 1 V atom per 8 supercages. Most materials used contain 1 vanadium per 8 to 16 supercages. Table 1. Vanadium content as determined by ICP, after several treatments.

VO-Y [A] VO(HPS)-Y [B] VO(HPS)-Y [C] VO(HPS)-Y [D] VO(HPS)-Y [E] VO-Y [F] VO(HPS)-Y [G] VO(HPS)-Y VO(salen)-Y [H] VO(salen)-Y [I] VO(salen)-Y [J]

details treatment

V(wt%)

AI/V (mol/mol) a

2mM exchange at RT [A] + acetonitrile soxhlet [B] + NaOAc wash + soxhlet [A] + acetone soxhlet [D] + NaOAc wash + soxhlet 4mM exchange at 80~ [F] + acetone soxhlet, + NaAc wash + soxhlet [G], washed TBHP/70~ low wt% VO-Y b + acetonitrile soxhlet [A] + acetonitrile soxhlet [A] + 24h incubation + acetone soxhlet

0.23 0.15 0.12 0.29 0.14 1.2 0.61 0.36 0.017 0.076 0.30

68 n.d. 142 62 132 n.d. 32 47 957 251 62

(a) In some cases no reliable data were obtained for the AI/V ratio, and could therefore not be determined (n.d.). (b) VO-Y synthesized by exchange of NaY with 2mM solution of VO(i-OPr)3 in demineralized water.

3.2 Oxidation experiments with cyclohexene In Table 2, the first experiments with cyclohexene as a substrate are given. The necessary blanks and homogeneous experiments are given as a comparison. As can be seen, it is very difficult with vanadium to get 100% selectivity towards the epoxide. Especially at the reaction temperatures of 70~ allylic oxidation interferes. In an attempt to reduce its influence, reactions were performed under nitrogen. As a consequence the main product becomes the cyclohexenyl-tert-butylperoxide, according to reactions 1-7. The In (itiator) could either be tBuOo, tBuOO~ or VOO~ /---X In.

+

+ InH

~

(1)

tBuOOH

+ In.

~

~~

+ 02

~

~~'-OO 9

(3)

~ .

+tBuOO"

,~

~-OOtBu

(4)

~

~--O

~---~-OO-

~-OO.

2 tBuOO.

+tBuOO"

tBuOO 9 + InH

+ ~~-OH

~ ~ = O

~

tBuO. + 02

(2)

+ O2

(5)

+tBuOH + O2

(6)

(7)

1036

Radicals are formed upon one-electron oxidation of tBuOOH 5, but also a v~Vooo diradical has been suggested as an intermediate 6. As can be seen from the formation of ketones and alcohols, it is impossible to completely eliminate the influence of oxygen, also because oxygen is formed in situ according to reaction 7. Table 2. Oxidation of cyclohexene catalyzed by vanadium a catalyst...... time yield epoxyb selectivity product selectivities %c (%) epoxide (on TBHP epoxide, CyOOtBu, Cy=O+CyOH consumed) no VO(acac)2 VO(HPS) VO(HPS) VO-Y [A] VO(HPS)-Y (B) VO(HPS)-Y (D) VO(HPS)-Y (D) VO(salen) VO(salen) VO(salen)-Y [I]

lit e g h i j RT lit e

23h 5h 3h 5h 24h 48h 5h 24h 3h 5h 21h

0 25 32 55 18 (148 TON) 23 5 0.3 (6 TON) 0.1 4 0.6 (12 TON)

23 43 n.d. n.d. n.d. n.d. n.d. 1 2 n.d.

46

44

100d 0

43 f

1f

1f

57 34 33 72 100 1f 5 19

17 58 28 19 77 50

25 7 39 9 10f 18 31

-::iai-i~eaction condiiions; as in experimental, 70~ i00 mg cat (zeolite) 0r 0.06 mmol complex, solvent~gml DCE, 0.55 M TBHP, TBHP:cyclohexene = 1:1. n.d. = not determined (b) Yield on intake cyclohexene. (c) Defined as product/total products. CyOOtBu, refers to cyclohexenyl-tert-butylperoxide, Cy=O + CyOH comprises cyclohexenone (major), cyclohexenole, epoxycyclohexenole (major) and epoxycyclohexenone as products. Mass balance was usually between 60 and 80% due to vaporization of cyclohexene. (d) 3% cyclohexenole (on intake cyclohexene) formed as sole product. (e) Ref.7, conditions; 60~ solvent DCE, 1.40 M cyclohexene, 1.43 M tBuOOH, cat. 0.07 M. (f) Products in % on TBHP consumed. (g) After 0.75h, 14% yield of epoxy, and 67% product selectivity to epoxide, 17% CyOOtBu and 16% Cy=O+CyOH.(h) After 5h, 14 TON and 26% selectivity to epoxide.(i) only 50 mg cat. (j) ratio cyclohexene:oxidant = 2:1. The oxidations catalyzed by VO(HPS)-Y, gave slow reactions, with rather low selectivities. The latter phenomenon is likely a direct result of the lower rate, giving allylic oxidation more chance to interfere. At room temperature almost no reaction was observed. The VO(salen)-Y material tested here gave a limited activity, although the selectivity towards the epoxide was higher than that obtained with the homogeneous VO(salen) complex.

3.3 Oxidation experiments with cyclooctene Conversion of cyclooctene to epoxycyclooctene was studied with different catalysts. As can be seen products of homolytic oxidation are much less prominent compared with cyclohexene. Heterolytic epoxidation of cyclooctene compared to cyclohexene is a faster reaction, and in Table 3 it is shown that whereas VO(salen) did not give significant epoxide yields with cyclohexene, with cyclooctene 78% epoxide selectivity was found. The cyclooctene oxidation is therefore less sensitive to ligand effects, and the reaction was mainly used to study catalyst stability. Experiments to test for leaching were performed in two ways. In the first method the reaction mixture was filtered after 1 h, and the filtrate allowed to react further (Table 3, note g). It is important to filter the solution at the reaction temperature because readsorption can take place on cooling. In the second method the zeolite is incubated

1037

with T B H P for 1 h at 70~ filtered and the reaction started by adding c y c l o o c t e n e to the filtrate (note h) 18. As s h o w n in Table 3 the results were very similar, although in general a slightly l o w e r selectivity to the epoxide was observed after filtering the catalyst as well as a lower selectivity on T B H P consumed. Recycling or extensive incubation with t B u O O H , gave d i m i n i s h e d activity. A p p a r e n t l y T B H P is capable o f extracting the v a n a d i u m from the m o l e c u l a r sieve. Table 3. Oxidation o f cyclooctene, catalyzed by v a n a d i u m a catalyst time yield epoxy sel. epoxlde (on (%) TBHP consumed) no d 24h 7e 23 VO(HPS) 5h 62 70 VO(HPS) d 5h 53 63 VO(HPS)-Y [B] d 24h 28 n.d. VO(HPS)-Y [C] 24h 26 49 VO(HPS)-Y [D] 24h 50 75 VO(HPS)-Y [D] recycle f 24h 26 59 VO(HPS)-Y [D] leach g 24h 43 62 VO(HPS)-Y [D] leach h 24h 33 n.d. VO(HPS)-Y [D] RT ~ 24h VO(HPS)-Y [a] 24h 48 65 VO(HPS)-Y [G] leach h 24h 50 52 VO(HPS)-Y [G] leachj 24h 29 38 VO(HPS)-Y [E] CHP k 24h 5 n.d. VO(HPS)-Y [E] CHP k 24h 10 n.d. VO(salen) d 5h 26 25 VO(salen)-Y [I] l d 24h 1.5 18 VO(salen)-Y [J] 24h 0 "

-

.

.

.

.

.

.

.

.

.

.

~ . . . . . . . . . . .

" .......

13

. . . . . . . . . . .

' .

.

.

.

.

.

.

.

9

'~

....

'

........

product selectwmes (%) epoxide Cy=O CyOOtBu 52 5.0 33 100 97 1.4 0.7 95 2.3 0.7 96 0.7 1.3 97 C6HllOH + H20 Cyclohexanol

(1)

A continuous supply of electrons and protons to the catalyst makes the biological oxidation highly selective and efficient. The electrons are provided by nucleosides such as NADH and NADPH. The nucleosides donate a pair of electrons to co-factors such as flavins. The flavins are 2 electrordl electron switches that transfer electrons one at a time to the Fe llI porphyrin catalyst which is enclosed in a protein matrix. The reason why the P-450 system is so selective, efficient and elegant is due to its ability to cleave the O-O bond of oxygen. On receiving an electron, the Fe III porphyrin is reduced to its Fe H state. The catalyst is now activated and binds an oxygen and a proton to form FeW-O-OH. On receiving a second electron and a second proton, it is converted to FeV=O (the "oxo" species) and water. The O-O bond has been broken. The heat of formation of water (54.63 kcal/mol) provides the necessary energy to break the O-O bond. The FeV=O species is in essence an atomic oxygen (with 6 electrons) partially stabilized by iron in various oxidation states 2 as shown in Equation 2 below. FeV=O

~

FeIV-o.

~

Fe III [.O. ]

(2)

The "oxo" species abstracts a hydrogen from cyclohexane to form FeIV-oH and cyclohexyl radical which rebounds at rates 3 as high as 101~ to form cyclohexanol and Fem porphyrin is regenerated 4. Even though free radicals are involved, the oxidation is not a chain reaction and it does not involve alkyl peroxide radicals as in the commercial processes described earlier. Though elegant, biological oxidations occur within the cell in micromolar quantities in aqueous medium. They are too complex and totally unsuited for large-scale industrial productions. 3. DESIGNING NOVEL INDUSTRIAL REDOX CATALYSTS

3.1 Approach Our approach is to adopt the beautiful chemistry of P-450 in an innovative way to large-scale autoxidation of hydrocarbons using conventional reactors. This would require that we design a totally new composite catalytic system and identify inexpensive electron sources that will activate the catalysts continuously during autoxidations. In other words, we are searching for a new paradigm in catalyst design.

1092 3.2 Objective The work presented here does not provide a specific approach to large-scale selective oxidation of cyclohexane. The objective is to probe several new approaches to hydrocarbon oxidations building on the collective experiences of chemists and chemical engineers with large-scale industrial processing and adopt the elegant chemistries of biological oxidations. Technology is currently available for immobilizing appropriate microbes on solid supports and use them in conventional reactors to convert hydrocarbons to their oxidized products. While such processes are highly selective, they would need an aqueous medium in which the hydrocarbon is present at high dilution. Further, the microbes may have to be replaced every few weeks. Such a process will not be efficient. Scale is an important issue in the manufacture of bulk chemicals. It is the Achilles heel in biochemical processing. Our approach is to adopt conventional large-scale reactors (preferably operating in a continuous mode) and use novel sturdy composite catalytic systems which are different from biological catalysts but will still carry out the elegant chemistries of the P-450 enzyme. This approach will represent the launching of an evolutionary process to achieve the next level of sophistication in catalyst design. When we study and generalize the P-450 oxidations, three important issues stand out. (1) We need an inexpensive electron source for continuously activating the redox catalyst; (2) we need a redox transition metal catalyst capable of forming the "oxo" species and (3) we need an electron mediator, usually selected from noble metals, capable of transferring electrons from its source to the redox catalyst. These three issues are critical in adopting biological oxidation chemistries to large-scale hydrocarbon autoxidations. 3.3 Electron sources and reactor configurations

The electron source will dictate the reactor configuration. In equation (1), the two electrons and protons (needed for activating the catalyst) are equivalent to one mole of hydrogen. (Equation 3.) 2H + + 2e

)

(3)

H2

The two protons, the two electrons and oxygen taken together are equivalent to a mole of hydrogen peroxide. (Equation 4.) 2H + + 2e- + 0 2

)

H202

(4)

Hydrogen is the least expensive and highly concentrated source of electrons and protons. However, it forms an explosive mixture with oxygen and therefore is not safe to be mixed with it. Hydrogen peroxide and its related reagents such as HOC1 and t-butylhydroperoxide are convenient to use. More importantly, they help to overcome oxygen transport problems inherent in autoxidations. Further, peroxides are known to convert redox catalysts such as Fe +3, Mn +3, Ti +4 etc directly to their "oxo" species. The electron mediator is no longer needed. Using peroxides, simplifies the design of the reactor.

1093 Isopropanol (and other secondary alcohols) are also excellent electron sources when used with Ru, Pd and Rh as electron mediators. After donating two electrons and two protons (ie. a hydrogen equivalent) the secondary alcohols are converted into ketones. The ketones are readily reduced back to the secondary alcohols by catalytic hydrogenation and recycled. Using secondary alcohols will therefore add an extra step to the process but they are safe in the presence of oxygen. Thus secondary alcohols can be considered as safe sources of hydrogen. 3.4 Catalytic systems

A composite catalytic system will consist of a redox transition metal catalyst and an electron mediator. The redox catalyst is selected from metal ions such as V +5, Ti +4, Mn+3 Fe +3 etc based on their ability to form the "oxo" species. The electron mediator is selected from Ru, Rh and Pd based on their well established ability to remove and transport electrons from hydrogen and secondary alcohols. 4. RESULTS 4.1 Extracting electrons from hydrogen with catalytic membranes

We developed a membrane system to demonstrate the feasibility of using hydrogen as an electron source and at the same time preventing it from coming in contact with oxygen. The membrane was prepared by mixing a solution of silicone rubbers with 5% Pd on carbon and spreading this mixture over a polysulfone membrane which is not permeable to hydrogen but will allow electrons and protons to pass through. The coated membrane was cured at 100 ~for an hour. Hydrogen at 10 psig pressure was converted to electron and protons and transferred to a buffer solution (at pH 8.5) containing methylene blue on the other side of the membrane. The methylene blue was reduced to its colorless leuco form. On exposing to oxygen, the blue dye was regenerated and oxygen was converted to hydrogen peroxide. However, Mn +3 tetraphenylporphyrin was reduced preferentially over oxygen to its Mn +2 form as evidenced by UV spectrum. These experiments show that Pd is capable of extracting electrons from hydrogen and transporting them through the membrane without hydrogen contacting the redox catalyst or oxygen on the other side of the membrane. However, the membranes developed leaks over a period of a week and they were not compatible with organic solvents. 4.2 Peroxides as source of electrons and oxygen

Peroxides are easy to handle. A large number of papers report using hydrogen peroxide, HOC1 and t-butylhydroperoxide in hydrocarbon oxidations. As discussed earlier, the peroxides are sources of electrons, protons and oxygen circumventing the need for an electron mediators. The peroxides directly convert redox catalysts to their the "oxo" species. We used 30% hydrogen peroxide to oxidize cyclohexane (50 mM) dissolved in dichloromethane containing 5 mM Mn +3 tetraphenylporphyrin as the redox catalyst. The biphasic system was stirred for 24 hours at room temperature. Cyclohexane was converted

1094 to 48% of cyclohexanol and 10% of cyclohexanone. The catalyst was completely destroyed during the course of the reaction. t-Butylhydroperoxide was more effective since it is soluble in organic solvents. A mixture (5 mL) of cyclohexane and benzene (80:20 volume ratio) containing 1 mM Mn +3 porphyrin was stirred at ambient temperatures. Different amounts of t-butylhydroperoxide were added slowly over several hours using a syringe pump. The products were analyzed after 24 hours. About 60 - 80% peroxide was consumed. Cyclohexanol and cyclohexanone were formed in 9:1 ratio along with t-butanol. Trace amounts of side products were also identified by gc/MS to be cyclohexyl t-butylperoxide, cyclohexyl t-butylether and di tbutylperoxide. Even though the reaction is selective at ambient temperatures, it is slow and separation of the catalyst from the reaction mixture was difficult. Further, about 20 - 25% of the catalyst was destroyed in 24 hours. The recovery of t-butanol and its conversion to tbutylhydroperoxide is a costly operation.

4.3 Use of secondary alcohols as electron donors Secondary alcohols are excellent hydrogen donors. In the0Presence of RuH2(PPh3) 3 secondary alcohols such as isopropanol produce hydrogen 5 at 150 with high turn-over numbers. In this process, the alcohols are quantitatively converted to the ketones. Ru, Rh and Pd catalysts are excellent in "transfer hydrogenation" reactions using secondary alcohols as hydrogen sources to reduce other substrates such as the imines. We were convinced that the above catalysts are also potential electron mediators in transferring electrons and protons from secondary alcohols to the redox catalyst. We were able to demonstrate that Pd and Ru are excellent electron mediators that converted cyclohexanol (a secondary alcohol) in 98% selectivity to cyclohexanone. Four electron mediator catalysts (Pd on carbon, PdC12, RuEC1E(p-cymene)2 and a trinuclear Ru carboxylate) were tested. The conversions were poor with Pd catalysts even though the selectivities were high. The RuEC12(p-Cymene)2 was more efficient 6. However, it required MnO 2 as an electron sink. THF was used as the solvent. Cyclohexanone was formed in 36% yield in 60 hours at 70 ~ The most promising results were obtained with the trinuclear Ru carboxylate prepared by Wilkinson and used by Drago 7. Cyclohexanone was formed in 60% yield in 6 hours at 400 in 98% selectivity. In this reaction, cyclohexanol was used neat and oxygen was used as the electron sink and was converted to hydrogen peroxide. This was a surprising result since Tang s has shown that the oxidation of cyclohexanol by oxygen in the presence of RuC13 is extremely slow.

4.4 Composite redox catalysts for selective hydrocarbon autoxidations Titanium silicalite has been recognized as an efficient redox catalyst in a number of industrial processes. Enichem has a process in which Ti silicalite catalyzed the conversion of cyclohexanone by ammonia and hydrogen peroxide to cyclohexnone oxime 9. The mechanism appears to involve the "Ti oxo" species. A number of procedures are available from literature to make the Ti silicalite catalyst and it will be easy to incorporate electron mediators such as Ru or Pd into the silicalite matrix. The performance of such composite

1095 catalysts in selective autoxidations can be evaluated using cyclohexane containing a small amount of cyclohexanol. The autoxidation will be initiated by cyclohexanol by transfering its electrons to activate the composite catalyst. In this step, cyclohexanol will be converted to cyclohexanone. The activated catalyst will react with oxygen and cyclohexane to form water and generate more cyclohexanol thereby starting a self-perpetuating reaction sequence in which the net reaction is the conversion of cyclohexane to cyclohexanone and water occurring in the same reactor. The reaction sequence is summarized in Scheme 1 below. Scheme 1 Cyclohexane autoxidation Step 1. C6HllOH Cyclohexanol Step 2.

Net reaction

C6H100 + H2/Catalyst Cyclohexanone

H2/Catalyst + 02 + C6H12 ~ Cyclohexane C6H12 + 02 Cyclohexane

C6H11OH + H20

) C6H100 + H20 Cyclohexanone

AGo= +6.49 AS~ 29.89 (kcal/mol) ( eu ) AGo= -90.41 AS~ -28.06 (kcal/mol) (eu) AG~

AS~ (kcal/mol)

1.83 (eu)

In Scheme 1, the free energy and entropy for step 1, for step 2 and for the net reaction were calculated from thermodynamic data l~ It is clear that step 1 is endothermic with a positive entropy. Higher temperatures will favor this step. Step 2 is overwhelmingly exothermic; but the entropy is unfavorable. Higher temperatures will affect this step only slightly. The entropies of these two steps counteract each other and as a result the free energy of the net reaction will not change much with temperature. At 25 o and 150 ~ the free energies for the net reaction are respectively -83.92 kcal/mol and -81.83 kcal/mol. Based on this analysis, we decided to run the reaction at 150~ Clerici and Bellussi 11 have shown that hexane in methanol can be selectively oxidized to 2-hexanol, 3-hexanol, 2-hexanone and 3-hexanone using a mixture of oxygen and hydrogen at 25 o - 30 ~ The reactions were run in the presence of HC1 for 20 - 24 hours. Several titanium silicalite catalysts containing Pd (0.01 mol ratio to TiO2) were prepared and used in these reactions. Presumably, hydrogen plays the role of an electron-donor activating the Ti catalyst. However, no explanations were offered. We prepared composite catalysts containing Ru in Ti silicalite (molar composition SiO2:0.025TIO2 : 0.001 Ru). The catalyst was suspended in 3 mL of a mixture (1:4 weight ratio) of cyclohexanol and cyclohexane in a 100 mL steel autoclave and stirred under oxygen at 150 o for 24 hours. Cyclohexane was converted to cyclohexanone to an extent of 15 - 20%. The reaction was highly selective for cyclohexanone. Products other than unreacted cyclohexane and cyclohexanol were not observed. More than the calculated amount of oxygen was consumed. The experimental set up was similar to the one described by Clerici and Bellusi ll.

1096 The use of secondary alcohol for activating redox composite catalysts in hydrocarbon autoxidations is attractive. Isopropanol is excellent as an electron source. Acetone will be produced as the byproduct in these autoxidations. It is easily separated and reduced back to isopropanol by catalytic hydrogenation and recycled. Several composite redox catalysts must be prepared with various combinations and ratios of electron mediators (such as Ru, Pd, Rh etc) with redox catalysts (such as V +5. Ti § Mn +3 , Fe+3 etc) and evaluated for their ability to selectively oxidize hydrocarbons. 5. C O N C L U S I O N A new paradigm in the design of catalysts for selective autoxidation of hydrocarbons is presented. The catalyst will consist of two components: a redox component and an electron mediator. The redox component is selected from transition metal ions. These ions are capable of breaking the O-O bond of oxygen to form the "oxo" species provided they are continuously activated by a supply of electrons. The second component of the catalyst is an electron mediator (selected from noble metals) which help to transfer electrons from a suitable source to the redox component. Hydrogen and secondary alcohols are the best sources of electrons. For safety reasons, the mixing of hydrogen and oxygen is not desirable. This complicates the reactor design. However, secondary alcohols are convenient and safe to use. Buttheir regeneration will add an extra step. Peroxides can serve as sources of electrons, protons and oxygen. Further, they help to overcome oxygen transport problems inherent in autoxidations. However, peroxides also destroy the catalysts particularly at higher temperatures. We have shown that cyclohexane can be selectively converted to cyclohexanone. In this reaction, the intermediate (cyclohexanol) serves as the electron source. What is important is that this selective oxidation was achieved in a conventional batch reactor. We have a long way to go. But we have made the start. We have managed to "lift" the elegant chemistry of the P-450 enzymes out of the confines of biological cells into a larger arena of conventional man-made reactors. REFERENCES

1. R.A.Sheldon, Chemistry & Industry (1992) 903. 2. D.H.R.Barton, F.Halley, N.Ozbalik and E.Young, New J. Chem. No.3 (1989) 177. 3. V.Bowry, J.Lusztyk and K.U.Ingold, Pure & Appl. Chem. Vol. 62, No.2 (1990) 213. 4. M.Hirobe, Pure & Appl. Chem. Vol. 66, No.4(1994) 729. 5. J-E.Backvall, R.L.Chowdhury and U.Karlsson, J.C.S., Chem.Comm. (1991) 473. 6. U. Karlsson, G-Z.Wang and J-E. Backvall, J.Org.Chem.Vol.59 (1994) 1196. 7. C.Bilgrien, S.Davis and R.S.Drago, J.Am.Chem.Soc. Vol. 109 (1987) 3786. 8. R.Tang, S.E.Diamond, N.Neary and F.Mares, J.C.S. Chem.Comm. (1978) 562. 9. S.Tonti, P.Roffia and V.Gerasutti, Ammoimation, US Patent No.5 227 525 (1993). 10. D.R.Stull, E.F.Westrum and G.C.Sinke, The Thermodynamics of Organic compounds, Wiley, New York, 1969. 11. M.G.Clerici and G.Bellussi, Oxidating Paraffins, US Patent No. 5 235 111 (1993)

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

1097

Selectivity of Active Sites on Oxide Catalysts C. Batiot, F.E. Cassidy, A.M. Doyle and B.I~ Hodnett Dept. of Chemical and Environmental Sciences, University of Limerick, IRELAND.

ABSTRACT The selectivity of active sites on oxide catalysts have been assessed by comparing their ability to selectively activate a C-H bond in a reactant rather than a similar C-H or C-C bond in a selective oxidation product. Active sites on oxide catalysts are capable of activating target bonds in reactants that are up to 30-40 kJ mole 1 weaker than similar bonds in the selective oxidation product. Good selectivities are always recorded provided that selective oxidation reactions attempted do not exceed the discriminating capacity of the active site. Evidence is also presented that C-C bonds, which are generally weaker than C-H bonds, are protected from rupture by steric factors. 1. INTRODUCTION The essential value of a selective oxidation catalyst can be represented by a simple selectivity-conversion plot, the best catalysts being those that give the highest selectivity at a given conversion [1]. Other factors such as activity and deactivation are less important, because the former can always be boosted by increasing the W/F ratio or the temperature and generally deactivation phenomena in oxidation catalysis are not severe. Htmdreds of examples of selective catalytic oxidation have appeared in the open and patent literature, but to date we do not know the factors which determine the upper limits of selectivity that can be attained for a given reaction. One factor which is well known is that a given selective oxidation catalyst has to be tailored for a specific reaction and generally a great deal of effort is required, in the first instance by way of screening a wide range of materials and, secondly, in fine tuning the best of these with additives, promoters, dopants and supports to arrive at the best catalyst for a particular reaction and reactor configuration. The concept of active sites on heterogeneous catalysts was first introduced by H.S. Taylor [2] and has been widely used since, although not always fully defined. A meaningful assessment of the selectivity of active sites on oxide catalysts should be carried out on a comparative basis. The ultimate goal in the design of heterogeneous catalysts is to attain a specificity or selectivity typically achieved by enzymes. Hence figure 1 presents a comparison of enzymes and heterogeneous catalysts on the basis of four criteria common to both: These are (i) operating temperature, (ii) molecular mass of the active site, (iii) turnover frequency and (iv) selectivity. The normal definitions of terms (i), (iii) and (iv) have been used for figure 1 and

1098 the molecular mass of the active site for the enzyme was estimated from molecular mass of the enzyme, assuming just one active site per enzyme molecule [3]. The lower limit of active site size for heterogeneous catalysts has been estimated at the size of a single metal atom on a support whereas the upper limit is calculated from the concentration of acid sites in high silica to alumina ratio zeolites. This measure is intended to determine the number of active sites per unit volume that could be fired into a reactor so that the contribution by the support or the underlying bulk is taken into account in making the calculation. Realistically, on this basis, the active site size range in oxide catalysts lies in the range 103-104 amu, the lower value referring to unsupported catalysts and the latter to supported or diluted systems. Temperature (~

Selectivity (%)

Molecular Mass of Active site (Daltons)

Turnover Frequency

(s-b

106

100

1000

200,000 95

95 60,000

100

102

20,000

37 30 m

10.4

-20

10-4

100 H.C.

ENZ

H.C.

ENZ

H.C.

ENZ

H.C.

ENZ

Figure 1: Comparison of active sites in hetergeneous and enzymatic catalysis. (H.C. = Heterogeneous catalysis; ENZ = Enzymatic catalysis). The general overview presented in figure 1 is sufficient to indicate that in the majority of cases active sites on heterogeneous catalysts can be bigger than conventionally represented and overlap to some extent the size of corresponding sites in enzymes. The normal operating temperature range for enzymes is restricted by comparison with heterogeneous catalysts but the turnover frequencies that can be achieved on enzymes far outperforms that achieved by heterogeneous catalysts. In addition the selectivity achieved is normally very dose to 100% for enzymes and in addition the transformations that can be achieved are chemically much more complex than over heterogeneous catalysts [3]. Literature descriptions of active sites on oxide catalysts are oiten speculative and very often just generate a picture of the surface active site by extrapolation of the bulk structure. In general they envisage approach of the starting material to the active site in a preferred orientation without any indication of how the preferred orientation is established. In addition, the description of the active site is usually restricted to a small number of molecules. For example, the vanadium phosphorus oxide catalysts used for n-butane oxidation to maleic anhydride is based on the vanadyl pyrophosphate structure and an active site architecture is

1099 often presented which involves just two vanadium ions surrounded by a first coordination sphere of oxygen anions, without si,~nificant reference to the role of the pyrophosphate species [4]. Other active sites on oxide catalysts have been described where the site is strongly diluted on a molecular scale by a non-reactive species [5]. The central question is how selective can we expect heterogeneous catalysts to perform in the light of their rather simple chemical nature and the simple chemical transformations that they are designed to mediate. In the past attempts have been make to determine genetic factors which appear to be important in all selective oxidation reactions. Some of these factors are dearly related to catalyst structure, such as the requirement to dilute the active site at the surface and the primal role of lattice oxygen. Other factors such as the need to kinetically isolate the reaction products relate more to reactor engineering [5]. This work concentrates on another genetic factor which is related to the structure of the species appearing in the gas phase, emphasising the role of substrate structure, particularly the bond dissociation energies of the weakest C-H bonds in the reactants and products in determining selectivity in oxidation and ammoxidation reactions [ 1]. 2. RESULTS AND DISCUSSION Most selective oxidation reactions can be treated on the basis of a sirnple kinetic scheme as follows: Scheme I

HC

kl

~-

S.O.P

k2

=

COx

t A basic operating principle of selective oxidation catalysis is the need to minimise the contact time between the selective oxidation product and the catalyst to prevent conversion of the product, typically, into oxides of carbon. Whereas this aspect of selective oxidation catalysis is well recognised, it has never been put on a quantitative basis, so that the ability of a particular active site to activate a target bond in a reactant in preference to a similar bond in the product. Two aspects of scheme 1 will be addressed here, namely the factors which determine the limiting selectivity in terms of the ratio of k l to k2 and secondly the factors which determine the limiting selectivity in terms of the ratio of kl to k3. The selectivity at 30% conversion will be taken as measure of k l to k2 and the selectivity at zero to 10% conversion will be taken as a measure o f k l to k3. Towards these ends 14 selective oxidation reactions and two ammoxidation reactions have been evaluated through the use of selectivity-conversion plots, constructed from literature data [ 1]. Two examples of these plots are presented in figure 2 for ethylbenzene oxidation to styrene and methane oxidation to ethane. These selectivity-conversion plots were generated for a variety of catalysts for each reaction over a range of temperatures and space velocities. It should be stressed that the objective of this exercise was not to determine a reaction pathway or network, but simply to evaluate the best performance which has been achieved for any given reaction, hence the use of data from different catalysts and operating conditions.

1100

ethylbenzene - - > styrene 100 4: ' " ' ~

methane ---> ethane 100

"- ......

80 60 40

sel

,."

20

.,..:~.:~.%.'~

2o

0

0

t

I

I

I

20

40

60

80

100

0

20

40

60

80

100

conversion (o~

conversion (%)

Figure 2: Selectivity-conversion plots for the oxidation of ethylbenzene to styrene [6,7] and methane to ethane [8-25]. In all cases studied an upper limit could be identified beyond which experimental studies have not yet progressed and data points which fall below the upper limit are assumed to arise from operation with poor catalysts or in non-optimised conditions [ 1]. The first part of our analysis concentrates on the properties of the species appearing in the vapour phase, namely the reactants and the selective oxidation products and how well an active site distinguishes between these two species. A striking example is the oxidation of methane to ethane [8-25]. The structures of reactant and product are very similar as are their reactivities. Here we address the basis on which active sites typically present on oxide catalysts distinguish between almost identical bonds in reactants and selective oxidation products [26]. Generally selectivity in oxidation catalysis involves activation of the reactant through rupture of a C-H bond (the k l route in scheme 1), whereas diminishing selectivity is associated with rupture of any bond in the selective oxidation product (the k2 route in scheme 1). As a means of validating this hypothesis the upper selectivity limit, attained at a fixed conversion, in all 14 reactions used in this study was plotted against the function:

D~ where D~

(reactant)- D~

or C-C (product)

(reactant) is the bond dissociation enthalpy of the weakest C-H bond in the

substrate and D~

or C-C (product) is the bond dissociation enthalpy of the weakest bond

in the selective oxidation product.

1101

1

1cu3

.It 2

9e e ~ - " ~

.....

9 84 ,

80 9

A

oo

z=

11

40

2 20

t -70

-120

t

0

i

-20

13 ~ 1 4 ....l

30

D*Hc.I.I r,,,,~to,t- D ' H e . ,

130

80

0~ c-c ~..a,,.-t ( k J I m o l )

Figure 3 : Selectivity in product versus D~ C-H reactant" D~ C-H or C-C product at 30% conversion. 1, Ethylbenzene to Styrene. 2, 1-Butene to Butadiene. 3, Acrolein to Acrylic Acid. 4, Ethane to Ethylene. 5, n-Butane to Maleic Anhydride. 6, Propene to Acrolein. 7, Methanol to Formaldehyde. 8, Ethanol to Acetaldehyde. 9, Propane to Propene. 10, n-Butane to Butenes. 11, Propane to Acrolein. 12, Methane to Ethane. 13, Ethane to Acetaldehyde. 14, Methane to Formaldehyde [ 1]. The observed correlation (Figure 3) shows that there is a clear relationship between limiting selectivities and the nature of the weakest C-H or C-C bonds in the reactants and products. Active sites in oxide catalysts are capable of activating target bonds in the reactant that are up to 30-40 kJ mole 1 weaker than similar bonds in the selective oxidation product. When bigger differences in bond energies arise drastic reductions in selectivities result because the discriminating capacity of the active site has been exceeded. Scheme 2 presents a sequence of reactions starting with propane and leading to propene and acrolein. In the figure the percentage cited under each arrow is the limiting selectivity that has been reported in the literature for the reaction in question at 30% conversion. These data were generated on the same basis as those used for figures 2 and 3 above. The numbers above the arrows are the values which apply for the function: D~ (reactant) - D~

or C-C (product). Scheme 2

I [ [ ~C~C~C

I I I I

41 ldmole-1 >

14 ldmole-1

700/o --i--\

/

55 kJmole-1 40%

C~C~C~O

I

l

1102 This coherent reaction network clearly demonstrates the importance of the 30-40 kJ mole -] selectivity limit. When it is exceeded, as is the case with propane oxidation to acrolein, selectivity declines drastically. Similarly the accummulated data for propane and propene ammoxidation [27,28] to acrylonitrile indicate selectivities at 30% conversion of 50% and 85% respectively. These data are consistent with the 41 kJ mole .] difference in bond enthalpies shown in scheme 2 for propane and propene. Having established the discriminating limits of typical active sites in selective oxidation catalysts the second question to be addressed is the factors which determine the ratio of kl to k3 in scheme 1. The same methodology was followed here except that selectivities at 10% conversion were taken as a measure o f k l to k3. The basic approach is to assume that the kl route is favoured by activation or rupture of a C-H bond in the reactant whereas the k3 route is favoured through activation or rupture of a C-C bond. This part of the analysis will be restricted to the oxidative dehydrogenation of ethane, propane and n-butane. The corresponding selectivity conversion plots are presented in figure 4.

ethane ---> ethylene

propane-->

propene

100 9

60

9

9

".j

"

9 9 9

,,9 , . . "e

40

t

Oo 9 o

*,

9

"

9

sel 20

9

9

"#

"

40 : ' ~ r . ' , ' ~ . . .

"

""

sel 20 ~#~" : ee

0 I

I

I

I

t

0

20

40

60

8O

100

conversion (%)

0

.

o~

9

I

~

I

t

20

40

60

80

100

conversion (%)

n-butane ---> butenes

80 :':" 60

9 ., "

40

" .~;:..--

"" . " ' . sel

20

0

.

:~ e .

9t .

0

.

..

-.

**

Figure 4: Selectivity-conversion plots for

9

I

I

I

20

40

60

conversion (%)

100

ethane [29-38], propane [39-56] and n-butane [57-64] oxidative dehydrogenation.

1103 Generally C-C bonds in alkanes are much weaker than the corresponding C-H bonds, as shown in Table 1. Clearly then active sites do not distinguish between these bonds on a bond strength basis only and in view of the structure of these alkanes it is reasonable to suggest that steric factors must play a role here, since C-C bonds are generally more difficult to accommodate within an active site than C-H bonds.

Table 1: Bond Dissociation Enthalpies in Alkanes [26]

DC-HkJ molel(C-H3)

Alkane

Dc.c kJ mole "1

Ethane Propane n-Butane *(The value in brackets

376 420 368 417 360 (343)* 403 refers to the central C-C bond in n-butane)

DC-HkJ molel(C-H2) n/a 401 390

A more reasonable correlation emerges with the selectivity at 0-10% conversion when we examine the ratio of C-C to C-H bonds in a particular alkane, as shown in figure5.

100 --I-90

"-

Selectivity at 0% conversion Selectivity at 5% conversion Selectivity at 10% conversion

80 Sel 70 0

i

I

i

0.1

0.2

0.3

0.4

Ratio C-C to C-H Bonds

Figure 5: Influence of the ratio of C-C to C-H bonds in ethane, propane and n-butane on the selectivity in oxidative dehydrogenation at 0, 5 and 10 % conversion.

1104

3. CONCLUSIONS There was a clear upper limit in terms of selectivity-conversion beyond which experimental studies have not advanced for many selective oxidation reactions. These limits have been achieved through detailed catalyst design and reactor optimization. This work shows that active sites on oxidation and ammoxidation catalysts are capable of selectively activating, typically, a C-H bond in a reactant, rather than a similar C-H or C-C bond in the product provided that the bond dissociation enthalpy of the weakest bond in the product is no more than 30-40 kJ mole~ weaker than the bond dissociation enthalpy of the weakest bond in the reactant. When these limits are exceeded selectivity falls drastically. This work also indicates that primary activation of alkanes is through C-H bonds although the corresponding C-C bonds are much weaker. Cleavage of a C-C bond in the primary activation step leads directly to carbon oxide formation, but this step is less favoured because stedc factors make it diitictflt for the C-C bonds to be accommodated at the active site.

REFERENCES .

2. .

4. 5.

9

10 ll 12 13 14 15 16

C. Batiot and B.I( Hodnett, Appl. Catal. A, 137 (1996) 179. J.M. Thomas and W.J. Thomas, Principles and Practice of Heterogeneous Catalysis, Ed. VCH Weinheim, 1997. D. Voet and J.G. Voet, Biochemistry, J. Wiley & Sons Inc., 1995. B. Schiott and I~A. Jorgensen, Catal. Today, 16 (1993) 79. ILK. Grasselli in Surface Properties and Catalysis by Non-Metals, (Ed. Bonnelle, J.P. et al), D.Reider Publishing Company, Dordrecht, (1983) 273. M. Turco, G. Bagnasko, P. Ciambi, A. La Ginestra and G. Russo, Stud. in Surf. Sci. Catal, 55 (1990) 327. F.M. Bautista, J.M. Campelo, A. Garcia, D. Luna, J.M. Marinas and 1LA. Quiros, Stud. in Surs Sci. Catal, 82 (1994) 759. V.G. Roguleva, M.A. Nikiphorova, N.G. Maksimov and A.G. Anshits, Catal. Today, 13 (1992) 219. ICD. Campbell, Catal. Today, 13 (1992)245. G.J. Tjatjopoulos and I.A. Vasalos, Catal. Today, 13 (1992) 361. J.S.J. Hargreaves, G.J. Hutchings, IEW. Joyner and C.J. Kiely, Catal. Today, 13 (1992) 401. S.J. K o ~ J.A. Roos, J.M. Diphorn, l~H.J. Veehoi~ J.G. Van Ommen and J.1LH. Ross, Catal. Today, 4 (1989) 279. G.J. Hutchings, M.S. Scurrell and J.1L Woodhouse, Catal. Today, 4 (1989)371. C. Chevalier, P. Ramirez, M. Ceruso, A. Choplin and J.M. Basset, CataL Today, 4(1989) 433. A. Kiennemann, 1L Kieffer, A. Kaddouri, P. Poix and J.L. gehspringer, Stud. in Surf Sci. Catal., 55 (1990) 365. A.A. Kadughim~ O.V. Krylov, S.E. Plate, Y.P. Tulenin, V.A. Selezaev, A.V. Bolrov and Y.M. Kimelfeld, Stud. in Surf. Sci. Catal., 55 (1990) 447.

1105 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51

P. Kovacheva, N. Davidova and A.H. Weiss, Stud. in Surf. Sci. fatal., 82 (1994) 403. S.T. Brandao, L. Lietti, P.L. Villa, S. Rossini, A. Santucci, IL Millini~ O. Forlani and D. Sanfilippo, Stud. in Surf. Sci. fatal., 82 (1994) 443. T. LeVan, M. Che, M. Kermarec, C. Louis and J.M. Tatibouet, fatal. Lea., 6 (1990) 395. Choudhary, AICHE J., 37 (1991) 915. F.P. Larkins and M.IL Nordin, J.Catal., 130 (1991) 147. Yang, Bull. Soc. Chim Belg., 100 (1991) 5. Hinson, J.Chem. Soc. Chem_ Commun., (1991) 1430. Kiwi, J. Phys. Chem_, 96 (1992) 3, 1344. K o ~ Thesis Twente (1990). Handbook of Chemistry and Physics 61 st edition, R.C. Weast editor, 1980-1981 A. Corma, J.M. Lopez Nieto, N. Paredes and M. Perez, Appl. fatal., 97 (1993) 159. IL Burch and E.M. Crabb, Appl. fatal., 100 (1993) 111. Y.-C. Kim~ W. Ueda and Y. Moro-oka, Appl. fatal., 70 (1991) 189. ILK. Grasselli and J.L. Callahan, J. fatal., 14 (1969) 93. K. Seshan, H.M. Swaan, ILH.H. Smits, J.G. Van Ommen and J.ILH. Ross, Stud. in Surf. Sci. fatal., 55 (1990) 505. J. Barrault and L. Magaud, Stud. in Surf. Sci. fatal., 82 (1994) 305. I. Matsuura and N. Kimura, Stud. in Surf. Sci. fatal., 82 (1994) 271. B. Grzylowska, P. Mekss, IL Grabowski, K. Wcislo, Y. Barbaux and L. Gengembre, Stud. in Surf. Sci. fatal., 82 (1994) 151. J.G. Eon, P.G. Pries de Oliveira, F. Lefebre and J.C. Volta, Stud. in Surf. Sci. fatal., 82 (1994) 83. P.G. Pries de Oliveira, J.G. Eon and J.C. Volta, J. fatal., 137 (1992) 257. L. Magaud, Thesis Poitiers (1994). H.H. Kung, US Patent No. 4 777 319 (1988). IL Burch and R. Swarnakar, Appl. fatal., 70 (1991) 129. IL Burch and E.M. Crabb, Appl. fatal., 97 (1993) 49. S.S. Hong and J.B. Moffat, Appl. fatal., 109 (1994) 117. J.C. Vedrine, J.C. Le Bars, J.C. Vedrine and A. Auroux, Appl. fatal., 88 (1992) 179. A. Erdohelgi and F. Solymosi, Appl. fatal., 39 (1988) Lll. G.C. Colorio, B. Bonnetot, J.C. Vedrine and A. Auroux, Stud. in Surf. Sci. fatal., 82 (1994) 143. S. Bordoni, F. CasteUani, F. Cavani, F. Trifiro and M.P. Kulkarni, Stud. in Surf. Sci. fatal., 82 (1994) 93. J.G. Mc Carty, A.B. Mc Ewen and M.A. Quinlan, Stud. in Surf. Sci. fatal., 82 (1994) 405. M. Merzauki, B. Taouk, L. Monceaux, E. Bordes and P. Courtine, Stud. in Surf. Sci. fatal., 72 (1992)165. P.M. Michalakos, M.C. Kung, I. Jahan and H.H. Kung, J. fatal., 140 (1993) 226. E. Morales and J.H. Ltmsford, J. fatal., 118 (1989) 255. L. Mendelovici and J.H. Lunsford, J. fatal., 94 (1985) 37. T. Hayakawa, A.G. Andersen, H. Orita, M. Shimizu and K. Takehira, fatal. Lett., 16 (1992) 373.

1106

52 53 54 55 56 57 58 59 60 61 62 63 64

O. Desponds, ILL. Keiski and G.A. Somorjai, Catal. Lett., 19 (1993) 17. L. Mendelovici, Chem_ Lett., (1982) 1469. S. Bordoni and F. Trifiro, J. Che~L Soc. Faraday Trans., 90 (1994)19, 2981. H.M. Swaan, Thesis Twente (1992). A. Kaddouri, European Congress (1993). M.A. Chaar, D. Patti, M.C. Kung and H.H. Kung, J. Catal., 105 (1987) 483. M.A. Chaar, D. Patti and H.H. Kung, J. Catal., 109 (1988) 463. D. Patel, P.J. Andersen and H.H. Kung, J. Catal., 125 (1990) 132. L. Owens and H.H. Kung, J. Catal., 148 (1994) 587. R.G. Rizayev, R.M. Talyshinskii, J.M. Seifitllayeva, E.M. Guseinova,Y.A.Panteleyeva and E.A. Mamedov, Stud. in Surf. Sci. Catal., 82 (1994) 125. A. Corma, J.M. Lopez Nieto, N. Paredes, A. Dejoz and I. Vasquez, Stud. In Surf.Sci.Catal., 82 (1994) 113. H.H. Kung, US Patent No. 4 777 319 (1988). D. Bhattacharyya, S.~ Bej and M.S. Rao, Appl. Catal., 87 (1992) 29.

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

1107

A novel computer-aided technique for the development of catalysts for propane ammoxidation to acrylonitrile X.-Q. Wu, Q.-X. Zhang, Q.-L. Dai, Z.-Y. Hou and D.-W. Lu Chemical Engineering Department, Zhejiang University Hangzhou 310027, People's Republic of China

In this paper, a new computer-aided technique was presented, with which the experimental procedure of developing catalysts is scheduled sequentially. In each sequential step the neural networks model and multi-objective optimization are used to determine optimal design for the next experiment. The sequential method proved very efficient in developing catalysts for propane ammoxidation to acrylonitrile. And the yield of acrylonitrile corresponding to the best catalyst was up to 58.9%.

1. INTRODUCTION The work of developing multicomponent catalysts, such as catalysts for selective oxidation of hydrocarbon, is very complicated, which often takes much time and spending. Conventionally, developing this kind of catalysts has been a repetitious trial-and-error procedure under the directions of qualitative analyses and experiences. To a multicomponent catalyst, many factors, e.g., preparation conditions, active components, structure and reaction conditions will influence its catalytic performance. This influence behaves generally in a synergistic way. The conventional method or trial-and-error procedure can not disclose and describe such a complicated influential relationship, so the efficiency ofdeveloping catalysts is low. In order to improve the developmental method, researchers have introduced computer-aided means, e.g., artificial intelligence, optimization and graphics, etc., and the usefulness of the means has been proved in some cases [ 1-3]. The process for direct acrylonitrile (ACN) synthesis from propane ammoxidation: C3H8 + NH3 + 202 > CH2CHCN+4H20 AH= - 151.13 kcal/mol (1) with its commercial potentiality, has been widely investigated in the recent years. Various catalytic systems have been reported for the reaction, yet the selectivity and yield of acrylonitrile were unsatisfactory. Moreover, it seems that V-Sb-A1 mixed oxides represent the most promising candidate components for the catalysts because of their ability to selectively convert an alkane to an unsaturated nitrile [4-6].

This work is partially supported by the National Natural Science Foundation of China.

1108 In the paper, we presented a new computer-aided technique, which takes the catalyst development as an iterative procedure, and each iteration includes four steps, i.e. distributing experimental points, carrying out experiments, modeling the catalytic relationship and forecasting the optimal design. The technique was applied to developing catalysts for propane ammoxidation to acrylonitrile.

2. METHODOLOGY

The key to improve the method of developing catalysts is to set up some quantitative catalytic relationships, with which one can make the developmental procedure become a combinational one of qualitative analyses, quantitative predictions and experiments. For some simple catalytic systems, without any pre-determined experiment, quantum chemistry can be used to estimate catalytic properties quantitatively. Unfortunately, up to now it is difficult to apply this method to the multicomponent catalytic systems. A certain amount of experiments is necessary to develop the multicomponent catalysts. But, how to decrease the amount of experiments noticeably is a matter to which researchers have paid great attention, it is also the problem to be solved in this paper. Hence, we proposed a novel computer-aided technique, by which the procedure of developing catalysts is transferred to an iterative or sequential one. 2.1. Principle of the technique In the technique, the orthogonal design, neural networks(NN) and multi-objective optimization are adopted to construct the iterative procedure. The principle and steps of the technique are detailed below:

1). Adopting orthogonal design to distribute a number of exploratory (or original) points evenly within the whole experimental space; 2). Performing the experiment at each located point and getting the data; 3). Using all current data (including new and old data) to compose the sample set for NN training and training NN to obtain the catalytic relationship model; 4). Carrying out multi-objective optimization based on the trained NN model to forecast optima (if exist several local optima); 5). Calculating the iterative precision with the predicted and tested results; if the precision unsatisfied, taking all optima as new experimental points, then turn to step 2; 6). If the precision satisfied, giving out the total optimum and stop the iterative procedure. Here the iterative precision is given as: = lira,-r

.ll/llr

.ll

(2)

where Yopt and Yoxp are the results of catalytic performance obtained by the optimization and experiment respectively. The diagram of the iterative procedure is shown in Figure 1. According to the strategy of the iterative procedure, it can be expected that, with the iteration

1109 going on, the precision of the NN model and optimization will be raised gradually. Especially, the precision in the sensitive regions will be raised more rapidly than others. 2.2. Distribution of the original experimental points It is important to determine the experimental space, since the prediction ability of the presented technique is limited to the space. If some factors are not included in the space, their effects will not exist in the relationship models. For a catalytic system, essential influential factors can be selected based onthe theoretical and empirical knowledge as well as the results of literature. Then the experimental space can be determined reasonably. In order to have the NN model in the first iteration be capable of describing the catalytic relationship approximately, the original points must be distributed evenly in the experimental space. For this purpose, the orthogonal design or other method should be used to schedule the original points. At these experimental points the catalysts are prepared and evaluated.

BEGIN)

I Distributing the original experimental points ]

I Performing experiment at each located point I

1

Modeling the catalytic relationships via neural networks with all current data

Forecasting each local optimum via multi-objective optimization

~'

Yes

Taking each local optimum as new experimental points during the next iteration

C_fixdng out the final optimal design

T (END)

Figure 1. The diagram of the iterative procedure in the proposed computer-aided technique for developing catalysts.

1110 2.3. Modeling of the catalytic relationship via neural networks It is very useful to extract the valuable information, both qualitative knowledge and quantitative catalytic model, from all current experimental data. Depending on the information, the experiment in the next iteration can be designed elaborately. In this way the iterative procedure can be led to a hopeful direction. As we known, the catalytic performance of a multicomponent catalyst will be affected by many factors. The correspondent theoretical model describing the catalytic relationship between performance and factors can hardly be built up. Regression method is often used to correlate low nonlinear multivafiable relationships, but for a high nonlinear or strongly synergistic relationship, its limitation arising t~om the determination of the structure of the regression model becomes very serious. However, the neural network model is a useful tool for correlating such a sophisticated relationship between the outputs data (results) and the inputs data (factors) [7]. Neural networks have been successfully used for a number of chemistry applications including correlation of structure-activity or structure-spectrum [8-9], estimation of acid strength of mixed oxides [ 10] and product distribution [ 11]. Although a broad range of NN architectures and learning paradigms are available [12], the back propagation algorithm for multilayer feedforward networks [13] is the most popular approach for engineering and chemistry applications. The multilayer feedforward NN shown in Figure 2 is composed of many interconnected processing units or neurons organized in successive layers. The network is called fully connected because each neuron distributes its output to each neuron in the next layer. The first or input layer is composed of fan-out units vx46ch do not perform any computation but simply distribute their input to all neurons in the next layer. The last layer is the output layer. Between the input and the output layers there can be several hidden layers but only one will be considered here for simplicity. It is assumed that all the hidden and output layer neurons are identical although other choices are possible.

11

02

I1

W~

O/

m 9

J

Is input layer

"X l,N2 )f~

m.. ,ffay r'"

inputs

outputs

output layer

"

Figure 2. Three-layer feedforward neural network.

Figure 3. Neuron i in layerj.

A neuron in the hidden or output layer can be represented as in Figure 3. The neurons perform the following computations. The/th neuron in layerj receives N inputs {I1,'",/N} from layer j-lwith connection strengths or associated weights { W~/,.-., W~ }. The neuron first ~omputes the weighted sum of the N inputs:

1111

S/'= ~ W Jklo, +b J,

(3)

k=l

where b / is a bias term, The bias term need not appear explicitly because it can be interpreted as a weight associated with a constant input of one. The output of the neuron is a nonlinear function of the sum in (3)" O[ = f (S/ ) (4) where f denotes the activation function. The activation function is often chosen to be a sigmoidal functionf(x)= 1/(1+ e -~ ). NN training consists of finding a parameter vector Wthat minimizes the mean square output error:

: ( w-) = T , _1_

'

If ;- f ,ll

(5)

where Wis comprised of all the elements w j M is the number of the data included in the training sample set. Y~ denotes the actual output, and Y~ is the corresponding network prediction. Using a gradient descent method, a solution of W is obtained with J( W ) reaching to its minimum. In this work, the input variables of the input layerX = (x~,x2,...,x,) r are catalytic influential factors, i.e., components of a catalyst and reaction conditions. The output variables of the output layerY=(y~,y2,...,yr,) r are results relating to the catalytic performance, such as catalytic activity and selectivity. Based on the sample set including all the current data, the NN is trained. Here the trained NN is written as: Y = F( X , W)

(6)

where F denotes the NN model function of the catalytic relationship. 2.4. Prediction of the optimal design via multi-objective o p t i m i z a t i o n

With the help of the trained NN model, the effects of factors on the catalytic performance can be simulated numerically. On the other hand, the optimal design for influential factors to gain the best catalytic performance can be also predicted by using the NN model. This will be realized by solving the following optimization problem with multi-objective: Opt{yl,Y2,''',ym} I7 = F ( X , W )

t

s.t.

(7)

X ~R

where R is the feasible region for X . To solve the problem (7), we should transfer it from a multi-objective problem to a single-objective one. A case of the transformation is as below:

M~

1 ~ J= m-E. (Y,- Y7)2 ;=

? = F (2 ,O7 ) s.t.

XeR

(8)

1112 where Y7 is the ideal value of the ith sub-objective. A number of conventional optimization techniques, complex method for instance, should be used with respect to solving the single objective problem The result is an effective (or Pareto) solution of the multi-objective problem and just the optimal design for the experimental procedure.

3. EXPERIMENTAL Depending on the sequential method mentioned above, in our work of developing catalysts for propane ammoxidation to acrylonitrile, we also focused efforts on the catalytic active system of V-Sb-AI mixed oxides. The work was divided into three parts, i.e., optimizations of the composition of promoters and supporters, the composition of main components as well as reaction conditions.

3.1. Catalyst preparation In a stirred flask equipped for heating under reflux, the solid of N]-I4VO3 was dissolved in hot distilled water, to the hot solution the solid of Sb203 was added The slun~ was maintained under reflux for 18 hours. Then the solution was added to a prepared soft and homogeneous A1203-SiO2 gel, and followed by the addition of (]~]A)5I-Is(H2(WO4)6), SnO2, (]N[I-I4)H2PO4, CrO3, (NH4)6Mo7024"4H20, etc.. With stirring the resulting slun~ was evaporated. The thick material was dried in an oven at 114~ overnight, then ground and dried in air at 350~ for 5 hours. After cooled, the dried material was pelletized then screened and the 30 to 60 mesh particle size was collected, the screened material was finally heat treated at 610~ in air for 3 hours.

3.2. Catalyst evaluation The catalyst evaluation mentioned in section 4.1 and 4.2 was carried out in a tubular stainless steel fixed bed reactor. The gaseous feed components were metered through mass flow controllers into the top of the reactor at atmospheric pressure. Steam was introduced by air through a water bath. Air was taken as the source of oxygen. For the evaluation of catalysts, the reaction temperature was 500~ the molar feed ratios was C3HdNH3/Oz/H20 = 1.0/2.2/2.1/3.0, and the reaction contacting time is 14.0 h'gC3Hdgcat The catalyst was activated by introducing air at 610~ for two hours and NH3 at 5300C for 30 minutes. The analyses of the products was performed with on-line GC. The catalyst used in the expeiiment of optimizing reaction conditions is the best one obtained in the work of optimizing the main components. In order to observe the effect of each condition on the catalytic performance, the ranges of conditions were broadened deliberately.

4. RESULTS AND DISCUSSION

4.1. Optimization of the composition of promoters We selected the contents of promoters, such as P, K, Cr, Mo, the content of V-Sb and the weight-ratio of A1203/SIO2 in the supporter as six influential factors. Originally, 19 sample points were distributed, at which the catalysts were prepared and evaluated, thus all of the data were collected to compose the original training sample set. Among the data, the ranges of the catalytic performance were, the conversion of propane Xp 5 3 . 3 % - 82.1%, the selectivity of acrylonitrile SACS2 3 . 5 % - 50.5% and the acrylonitrile yield YAcs 14.3% ~ 34.9%.

1113 Based on the sample set, a NN configured in 6-20-12-2, that is six units (six influential factors) in the input layer, twenty units in the first hidden layer, twelve units in the second hidden layer and two units (the conversion of propane Xp and selectivity of acrylonitrile SACN)in the output, layer, was trained. The typical error curves of NN self-learning is shown in Figure 4. The fitting ability of the trained NN was tested with some data outside the sample set. We can see from Figure 4, up to 130000 times of the NN learning, the overfitting (or overlearning) problem is forthcoming. At this point the NN learning process should be finished. 0.09 9

~

T

:3 TestingPatteras I 0.07 0.06 . . . .

~li_

0

l~r~tlllUlg ratL~lns I 0.08

_d -~--'--~

,

~

1 0"04 o.o ~.~.~,j, 0.02 O.Ol o 50000 100000 150000 200000 LearningTimes

Figure 4. Root-mean-square error curves for NN training.

Yacn% 60 50 40 30 20 10 0 ori.

[] opt. m Exp.

1st

2nd

3rd

Figure 5. The best ACN yields in each iteration for optimi~ng promoters.

Then the sequential steps mentioned in section 2 were repeated till the precision was satisfied, and the catalytic results corresponding to the optimal composition were XP 80.0%, SAC~53.7% and YAC~43.0%. For this case the iteration was undertaken thrice, in each one four local optimal composition were predicted and then tested experimentally. The best acrylonitrile yields (including both results of calculation and experiments) in each iteration are shown in Figure 5.

4.2. Optimization of the composition of main components Based on the best composition of promoters, the composition of main components was optimized. The contents of V, Sb, W, Sn and supporter were taken as five influences, and XP, SAC~ also as the outcomes. 20 Original sample points were distributed, at which the catalysts were prepared and evaluated. The ranges of the results were, Xv 11.6%~88.2%, SAC~ 15.0% 44.4% and YAC~ 2.4% - 28.9%. Accordingly, A 5-20-12-2 network was adopted in the sequential iterative procedure. The results of the optimization and evaluation in the first and second iteration are listed in Table 1 and 2. Table 1 Results of the optimization and evaluation in the first iteration Optimization (%) Catalysts Evaluation (%) No. XP SACN XP SACN 96.987 58.041 98.216 50.984 C-11 94.424 57.046 96.039 50.651 C-12 82.344 65.999 76.096 58.131 C-13 C-14 86.560 57.960 80.986 45.598 ,..

Relative errors (%) XP

SACN

1.254 1.682 -8.211 6.883

-13.840 -12.626 -18.595 -21.329

1114 From Table 1, we can find that the highest YACNof evaluations is 50.1%, better than the best result in the work of optimi~g promoters. But the relative errors of the acrylonitrile selectivity between the optimization and evaluation are big. Therefore, the data of these four points were added to the original ~rnple set for training the NN model. Thus the trained NN model became more accurate, so did the optimization. This fact can be seen from Table 2. Table 2 Results of the Catalysts No. C-21 C-22 C-23 C-24

optimization and evaluation in the second iteration Optimization (%) Evaluation (%) Relative errors (%) XP XP SACN XP SACN SACN 95.085 67.635 93.696 59.762 -1.461 -11.64 81.246 64.126 80.348 66.722 -1.105 +4.048 94.206 66.431 83.760 67.647 -11.09 +1.830 86.016 64.889 84.994 62.882 -1.188 -3.093

The results in the second iteration showed that the relative errors between optimization and evaluation were less than those in the first iteration, and YACNreached to 55.0%. After performing the third iteration, the best catalyst of the optimization was similar to the best one tested in the 2nd iteration. It meant that the catalyst better than the best one obtained did not exist in the experimental space. So the iterative procedure Was finished, and the results relating to the best composition of main components were Xp 93.7%, SAcN 59.7% and YAcN 56.0%. The best YACNin each iteration is shown in Figure 6.

4.3. Optimization of reaction conditions It is knportant to consider the effects of reaction conditions, such as reaction temperature, ratios of ammonia to propane(N/C), oxygen to propane (O/C) and steam to propane (H/C), on the catalytic performance. To investigate and optimize these effects, we used a 4-12-7-2 network to correlate the relationship between the catalytic performance and conditions. Yacn% 70 60 50 40 30 20 lO l, 0 ori.

/

~Opt. mExp.

Yacn% 60

iji!i - i i lil

40

[!ii!' i"i-iiili"i] iiii"i 1st

2nd

3rd

Figure 6. The best YACN in each iteration for optimi~ng main components.

OOpt. IExp.

30 10 0 off.

1st

2nd

3rd

Figure 7. The best YACN in each iteration for optimi~g reaction conditions.

The original sample set consisted of the data of 16 experimental points. As above, the iterative procedure was undertaken again to search for the best reaction condition. Six experimental points were tested in three following iterations. Only twice was the procedure

1115 repeated, the iterative precision was satisfied. The final best results were: Xp 85.2%, SACN 69.2% and YAcs 58.9%, which corresponded to the conditions of reaction temperature 524.8~ and C3Hs/NH3/Oz/I-I20 = 1.0/3.0/4.4/6.0. At these molar feed ratios, a lower concentration of propane in the feed gas led to a lower space time yield. However, the results exhibited the synergistic effects of reaction conditions on the catalytic performance. By adding some desired constraints on the influential variables in problem (7), the iterative procedure will make out more reasonable reaction conditions for the commercial meaning. Furthermore, the best ACN yields in each iteration are shown in Figure 7. 4.4. Discussion

In three parts of work to optimize the composition of promoters and main components, as well as the reaction conditions, totally, 81 experimental points were tested. The best ACN yields of each part are shown in Figure 8. The best YAcs was raised sharply from initial 34.9% to final 58.9%. This fact demonstrates the proposed technique is very effective for developing catalysts.

60

Yacn% m

50 40 30 20 10 0

m 1 iiiiiiiiI

m m iiiiiiiil !iiiiiii!i

I iiiiiiii/iiiiiiil/iiiiiiii , m Init.

Prom.

Main

Cond.

Figure 8. The best ACN yields in each iterative procedure for optimizing promoters, main components and conditions. Otherwise, there still existed some problems to be considered in our work. First, the method of distributing the original experimental points should be capable of making the distribution reflect the tendency of catalysis sufficiently with the least points. The more the points is distributed, generally the more evident the tendency is, but the larger the amount of experimental work is. On the other hand, a small data sample set is not capable of disclosing the catalytic tendency adequately, then result in an overfitting problem in NN learning. Therefore it is an interesting question for the technique how to avoid the dilemma. As a successive research of this work, we are having a try by using uniform design to take the place of the orthogonal design. Secondly, for the development of multicomponent catalysts, the influential factors are too many to be included wholly in an iterative procedure. As we done in this work, dividing a whole catalytic system into three parts, some synergistic effects of the factors might disappear. Possibly this difficulty may be overcome by combining the theoretical or heuristic knowledge and the modeling of NN. Under the guidance of heuristic analyses, the experimental space will be determined strategically.

1116 5. CONCLUSION To develop a multicomponent catalyst, the information, the quantitative relationship between the catalytic performance and influences particularly, extracted from the data is very useful. It is necessary and crucial to work out an interactive strategy for controlling the whole process from data acquiring, correlating the data to forecasting the optimal design. For this sake, in the paper, a new computer-aided technique was proposed and then used to develop catalysts for propane ammoxidation to acrylonitrile. Under a certain reaction condition, an optimal catalyst with high-performance of Xp 85.2%, SACN 69.2% and YACN58.9% was obtained. This example demonstrated that the technique can raise the efficiency of developing catalysts obviously.

REFERENCES

1. 2. 3. 4. 5. 6. 7. 8. 9. 10. 11. 12. 13.

H.C. Foley, E. E. Lowenthal and X.-D. Hu, Computer Aided Innovation of New Materials II, Elsevier Science Publishers B.V., (1993) 1101. R.A. Van Santen, Chem. Eng. Sci., 45 (1990) 2001. E.R. Becker and C. J. Pereira (editors), Computer-aided Design of Catalysts, Marcel Dekker, Inc. 1993. G. Centi, R. K. Grasselli and F. Tfifiro, Catal. Today, 13 (1992) 661. R. Nilsson, T. Lindblad and A. Anderson, J. Catal., 148 (1994) 501. A.T. Guttmann, R. I~ Grasselli and J. F. Brazdil, US Patent, No. 4 746 641, (1988). X.-Q. Wu, B.-Y. Li and D.-W. Lu, Proceedings of the 7th National Conference on Chemical Engineering, Beijing, China, (1994) 1136. J. Zupan and J. Gasteiger, Anal. Chim Acta., 248 (1991) 1. P.C. Jurs, CICSJBulletin, 11(5)(1991) 2. S. Kito, T. Hattori and Y. Murakami, Ind. Eng. Chem. Res., 31 (1992) 979. S. Kito, T. Hattori and Y. Murakami, Appl. Catal. A: General, 114 (1994) L173. B. Widrow and M. A. Lehr, Proc. IEEE, 78 (1990) 1415. D. E. Rumelhart and J. L. McCleUand, Parallel Distributed Processing: Explorations in the Micro structure of Cognition, MIT Press, Cambridge, MA, 1986.

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

1117

Catalysts by Rational Design: Prediction and Confirmation of the Properties of the Co/Ce/Br liquid-phase autoxidation Catalyst Based on the Kinetic Similarity to the Co/Mn/Br Catalyst Rob K.Gipe a and Walt Partenheimerb,c a7081 93rd Ave. SE, Mercer Island, WA 98040, USA b Central Research and Development, DuPont Experimental Station, Wilmington DE 19803-0262, USA CWork performed at the Amoco Research Center, Naperville, IL 1. ABSTRACT The kinetics of m-chloroperbenzoic acid (MCPBA) with Co(II)/Mn(II)/Br(-I) and Co(II)/Ce(III)/Br(-I) mixtures in acetic acid is reported. MCPBA first reacts with Co(II) to give Co(Ill), then Co(III) oxidizes Ce(III) to Ce(IV), which is followed by Ce(IV) oxidizing Br(-I) to dibromine, see figure 1. This is similar to the observed reactions of MCPBA with the Co(II)/Mn(II)/Br(-I) combination. We therefore predicted that the Co/Ce/Br combination would be a similar autoxidation catalyst to Co/Mn/B ,i.e.,there is a synergistic interaction when either Mn or Ce is added to a Co/Br catalyst as well as a lowering of the rate of carbon dioxide formation. This was subsequently observed. The reaction of toluene with dioxygen, catalyzed by Co/Br, Mn/Br, Ce/Br, Co/Mn/Br, and Co/Ce/Br mixtures is reported. Problems encountered in determining a synergistic interaction are delineated and include 1) the co-oxidation of a number of intermediates, 2) having different compositions of intermediates at different times, and 3) that fact that a substantial amount of the initially added bromide exists as benzylic bromide. Hammett studies of the Co/Mn/Br and Co/Ce/Br catalysts give similar rho values of-1.2 and -1.4, respectively, indicative that both catalysts operate via a similar mechanism. 2. INTRODUCTION The discovery of metal/bromide catalysts in 1954 [1] evolved into one of the major industrial processes producing oxygenates from hydrocarbons using dioxygen as the primary oxidant: metal/bromide hydrocarbon + 0 2 ;, oxygenate + H20 ( 1) HOAc The primary example is the production of terephthalic acid from p-xylene which is used in the production of polyethylene(terephthalate). A recent review describes the oxygenation of at least 251 hydrogens to give 279 products using at least 35 different catalyst combinations. There are a number of reports which include cerium as one of the elements

1118 of the catalyst package, see table 1. Addition of cerium to a Co/Br catalyst has been reported to increase the rate of conversion of 2-methylnaphthalene [3]. This information led one of the authors to extend our kinetic studies to include interactions of cerium with cobalt and bromide. This kinetic study led to the prediction that 1) cerium would be synergistic with Co/Br in the same way that manganese is, and 2) reduction in the formation the wasteful by-products, carbon dioxide and monoxide, would occur.

MCPB

~2"

Co(HI

Mn(HI

Br CH3 CHa

MCBA

Co(ID~

MCPBA-~

Co(

- Mn(ll) ~

~

Br"

O CHs ~

Ce(IV)

2

Br

CHI

MCBA ~-~

Co(H)~,/

~

Ce(III)

Br

~3

CHa

Figure 1. Most favorable kinetic pathways when MCPBA is added to Co/Mn/Br and Co/Ce/Br mixtures in 10% water/acetic acid 3. EXPERIMENTAL

3.1 Kinetic studies These were determined as previously reported [4]. 3.2 Autoxidation of toluene Measurements were made in a glass cylindrical reactor as previously described [5]. Initial conditions were 1.09M toluene in 100.0 ml anhydrous acetic acid at 98-100~ The concentrations of metals in the Co/Br, Mn/Br, Ce/Br catalysts is 0.0200M and 0.0100M for each metal in the Co/Mn/Br and Co/Ce/Br catalysts. The bromide/metals ratio is always 1.00 mol/mol. Initial salts used were cobalt(II) and manganese(H) acetate tetrahydrates, cerium(m) acetate hydrate, and sodium bromide. Rates of oxygen uptake and carbon monoxide and carbon dioxide formation were calculated from the GC data knowing the flow rate of gases through the reactor and the composition of air.

1119 Table 1 Reported Metal/Bromide Catalyzed Autoxidations containing Cerium Catalyst a Co/Ce/Br 68 b Co/Ce/Br 76 c Co/Ce/Br d Co/Ce/Br

reagent/product tetralin/acetoxytetralin 2,6-dimethylnaphthalene/2,6-dicarboxynaphthalene p-cresyl acetate/4-acetoxybenzoic acid

solvent/yld Ac 20 HOAc Ac20

90

decahydronaphthalene/mixture of acetates and ketones Ac20 2,6-diisopropylnaphthalene/2,6-dicarboxynaphthalene o-toluic acid/o-phthalic acid 4,4-dimethylbiphenyl/4,4'-dicarboxybiphenyl 2-methylnaphthalene/2-naphthoic acid 2-methylnaphthalene/2-naphthoic acid benzylacetate/benzaldiacetate HOAc 45 toluene/benzylacetate Ac20 15

e f g h i j k

Co/Ce/Br Co/Ce/Br Co/Ce/Br Co/Ce/Br Co/Ce/Br Co/Ce/Zn/Br Co/Ce/Zn/Br

1 m n o p q

p-nitrotoluene/p-nitrobenzoic acid Co/Ce/Zr/Br 2,6-diisopropylnaphthalene/2,6-dicarboxynaphthalene Co/Mn/Ce/Br 1,1 -bis(3-methylphenyl )ethane/1,1 -bis(3,3'-benzoic)ethane Co/Mn/Ce/Br o-toluic acid/o-phthalic acid Co/Mn/Ce/Br Co/MoJI'i/Ce/Br 1,2,4-trimethylbenzene/ 1,2,4-tricarboxybenzene 2,6diisopropylnaphthalene/2,6-dicarboxynaphthalene Co/Mn/Ce/Br

a. Jap. Pat. 56118042 (1981) b.T. Maki and Y. Asahi, Jap Pat 61210052 c. K.Yazu,M.Saito, K.Ukegawa,T. Nakayama, S. Tadao, T. Suzuki, M. Ono, J. Miki,N.Takei,K.Tate,Nippon Kagaku Kaishi, 1,(1991)92-96. d. K.Yazu,T.Wakabayashi,T. Nakayama, Chem Lett, (1986) 1409-1412. e. S.Hayashi,T.Matsuda,A.Sasawaka, Ger Pat DE 3707876 A1, (1987). f. T. Nakayama, E. Nakamura, K. Koguchi,Nippon Kagaku Kaishi, 4, (1982), 650 g.H. Mami et al, Jap Pat 310846, (1988). h. Y. Ogawa, T. Yamada, US 93,339,202(1993) i. Y. Ogawa, T. Yamada, Aromatikkusu,46(3/4),80-3 (1994)(Japan)(CA121(3):34977f) j. Jap. Pat. 56104845 (1981) k. Jap. Pat. 56104846 (1981) 1. Jap. Pat. 59080637 (1984) m. P.A.Sanchez,D.A.Young,G.E.Kuhimann.W. Partenheimer W P Schammel, US Pat 4,950,786, (1990). n. H. Mami et al, Jap Pat 310846, (1988). K.Yazu,T.Nakayama, Nippon Kagaku Kaishi,3 (1988)304-310.(CA 109(21): 189633s) o. Jap. Pat. 47042639 (1972) to K. Nakaoka and S. Wakamatsu p. C. Fumagalli, L. Capitanio, G. Stefani,EP513835, (1992). (CA118(8):60707a) q.T. Yamada, K. Sugiura, Y. Doko, K. Maeda, R. Minami,Y. Nagao, JP4330039

1120 3.3 Hammett studies These were performed in the same reactor using identical procedures as those described above. Two or three reagents were simultaneously oxidized and the first order rate constants of their disappearance were calculated. Each experiment was repeated twice and the averaged rate constants were used in figure 2. M-xylene, at a concentration of 1.00 M, was used as the kinetic standard in each experiment, i.e., the rates of all the other reagents were compared to it. The concentration of all the other reagents were 0.150M. ..................................... ~-.~5

~.. D-OCH3

4-

p-t-t)utyf~t~

m

E -0.5

1

0.5

~m CH3 (

0 ~5

m-Cl "%

n.-C'.t~9~H.':I

2 Co(III)a + MCBA (2) HOAc/H20

t 1/2 =0.016 sec.

1121 In contrast to Co, Ce(III) acetate reacts quite slowly with MCPBA. The reaction of 0.0029M MCPBA with 0.0250M cerium(m) acetate results in a colorless solution slowly turning yellow at room temperature. The yellow color is due to an increase in absorbance in the 350-400 nm region due to the formation of cerium(IV): 23 ~ 2 Ce(III) + MCPBA

;, 2 Ce(IV) + MCBA

(3)

HOAc/H20 t 1/2 = 190min When a solution of Ce(III) (0.005M) and Co(II) (0.005M) is added to a solution of 0.0580 M MCPBA, the UV-VIS indicates the presence of only Ce(IV) and Co(II) after 4 sec (the time of mixing). Since MCPBA does not react that quickly with Ce(III), rxn 3, but does react quickly with Co(II), rxn 2 , we surmised that the MCPBA first reacted with the Co(II) to form Co(III)a and then Co(III)a reacted with Ce(III). This was confirmed using the stopped-flow apparatus. At 600nm, there was an initial increase in absorbance due to formation of Co(III)a which was followed by a decrease due to the reaction of Co(III)a with Ce(III): 23 ~ Co(III)a + Ce(III)

:~, Co(II) + Ce(IV)

(4)

HOAc/H20 t 1/2 = 0.042 sec The chemistry of cerium(Ill) acetate is similar to that of Mn(II) acetate in that both do not react as fast with MCPBA as Co does, but do react quickly with Co(III)a.

Ce(IV) (0.00125M) was prepared by mixing MCPBA, Co(II), and reacting the mixture with 0.0125M KBr:

Ce(III), and then

23 ~ Ce(IV) + Br-

- .......

- > Ce(III) + 1/2 Br 2

(5)

HOAc/H20 t 1/2 = 38 min. The rate of reaction 5 is only 0.73 that observed for Mn(III) + Br- when corrected for the slightly different concentrations of the metals and bromide and assuming a first order dependence of Ce(III) and Br-, see table 2. Co/Ce/Br are illustrated in figure 1.

The kinetic similarity of Co/Mn/Br and

There are at least three different forms of cobalt(III) in acetic acid which have been dubbed Co(III)a, Co(III)s and Co(III)c by Jones [6]. Co(III)a reacts further to form

1122 Co(HI)i, a change easily observed in the visible region, and this has a half-life of 1.7 min at 30 ~ C. Co(III)a, 'cobalt(Ill) active', is a more reactive form than Co(HI)i, 'cobalt(llI) stable'. Table 2 illustrates that Co(III)i reacts 2-3 orders of magnitude more slowly than Co(III)a toward Mn(II) or Ce(III) (when corrected for the differences in concentration). The spectrum of Ce(IV) does not change thereafter for 20 minutes suggestive that there are not different forms of Ce(IV) as there are for Co(l/I). Table 2 Rate Constants for Selected Reactions for MCPBA, Co(g), Co(III), Ce(III),Ce(IV), Mn(II),Mn(III), and Bromide in 10% Water/Acetic acid a 9

Reactants

~

Products

k,s- 1b

temp.C comments

MCPBA + Co(II

MCBA + Co(l/I0 66(2)

30

MCPBA + Ce(III) MCPBA + Mn(II) MCPBA + KBr

MCBA + Ce(IV) 6.2(.2)x10 -5 MCBA + Mn(III) 0.017 MCBA + KBr 3 0.08

23 23 23

(c) (d)

MCPBA Co(III)a Co(III)s + Ce(III) Co(III)s + Mn(II) Co(III)a + Ce(III) Co(III)a + Mn(II) Ce(IV) + KBr

MCBA 2x10 -6 Co(III)s 0.0070(.0002) Co(II) + Ce(IV) 0.00067 Co(II) + Mn(III) 0.0015 Co(II) + Ce(IV) 16.3(1.6) Co(II) + Mn(III) 6.6(0.1) Ce(III) + KBr 3 0.00031 (0.1)

25 25 23 23 23 23 23

(e)

Mn(III) + NaBr

Mn(II) + NaBr 3 0.0066(.0004)

23

(f)(g) (f)

(h)

a. [MCPBA]o=0.0005 M and all others 0.0100M unless otherwise stated. All data have been measured in our labs except the stopped flow measurements which were performed at Purdue University together with Bill Schleper in Dale Margarum's lab. b. Standard deviation, in parenthesis(), based on at least three independent measurements. c. Initial concentrations are Ce(I11)=0.0025 M, MCPBA=0.0029M d. Autocatalytic reaction, rate refers to fast part of S curve e. Rate of thermal decomposition reported by C.F. Hendriks, H.C.A. van Beek, and R.M. Heertjes, Ind. Eng. Chem. Prod. Res. Dev., 18 (1979) 38 for perbenzoic acid. A number of other peracids give approximately the same rates. f. Initial concentrations are Co(III)s=0.00025, Ce(III) or Mn(II)=0.0025M g. Using a sample of Co(Ill) prepared via ozone. h. Initial concentrations are Ce(IV)=0.00125M, KBr=0.0125M.

4.2 The Synergistic Interactions in Co/Ce/Br and its Similarity to Co/Mn/Br The synergistic interaction for the Co/Mn/Br catalyst has been previously reported by Ravens [9] based on the replacement of some of the cobalt by manganese in a Co/Br catalyst and by us [ 10] based on the fact that the sum of the activities of Mn/Br and Co/Br catalysts is less than the Co/Mn/Br catalyst.

1123 We will define the synergy factor (SF) as the rate of reaction (R) in the presence the catalyst with components X1/X2/X3... divided by sum of the rates of reaction of the individual components X1,X2,X3... (or sums of these components). For Co/Mn/Br and Co/Ce/Br catalysts, respectively, we have: SF = RCo/Mn/Br/(RCo+RMn+RBr+RCo/Br+RMn/B r)

(6)

SF = RCo/Ce/Br/(RCo+RCe+RBr+RCo/Br+RCe/B r)

(7)

A SF value > 1.00 indicates a synergistic interaction and a value < 1.00 indicates an antagonistic interaction while a value of 1.00 would indicate the absence of synergy. Cobalt(H) acetate itself is an autoxidation catalyst [11 ], as is manganese(H) acetate [ 11 ], and as are bromide compounds [12]. Under the mild conditions employed in these experiments (approximately 1 atmosphere of air and 100~ Co catalyzed oxidation of toluene has an induction period of about 1 hr, the rate of oxygen uptake is roughly 0.3 ml O2/min and hence this will be subsequently ignored (RCo =0.0). We do not observe oxygen uptake with manganese(H) acetate, cerium(Ill) acetate or sodium bromide catalysts with the conditions employed in these experiments, hence their rates are zero (RMn=RBr=Rce=0.0). Values of SR > 1.0 are always observed for both the Co/Mn/Br and Co/Ce/Br catalysts. We also find a higher degree of synergy in the Co/Mn/Br catalyst than the Co/Ce/Br one, i.e., SFCo/Mn/Br > SFCo/Ce/Br. For example from the rates of oxygen uptake we have: Table 3 Results from the Autoxidation of Toluene using Co/Br,Mn/Br, Ce/Br, Co/Mn/Br and Co/Ce/Br Catalysts rate oxygen uptake,ml O2/min time, hr Co/Br Mn/Br Ce/Br Co/Mn/Br Co/Ce/Br SFCo/Mn/Br SFCo/Ce/Br 1.00 2.00 3.00 4.00 5.00

1.52 1.21 1.02 1.12 1.01

1.20 1.12 0.78 0.60 0.51

0.69 0.66 0.65 0.35 0.32

7.20 7.18 7.49 3.86 3.02

4.26 3.43 3.04 3.33 3.63

2.64 3.08 4.16 2.24 1.99

1.93 1.83 1.82 2.26 2.73

The SF can also be calculated in other ways. For example, the yield of benzoic acid for Co/Br, Mn/Br, Ce/Br, Co/Mn/Br, and Co/Ce/Br at 5.0 hr is 4.6, 1.4,0.0, 40.4 and 20.0 mol %. This gives a SFCo/Mn/Br =6.7 and SFCo/Ce/Br = 4.3. Figure 3 gives the activity of the catalysts, based on the conversion of toluene, where Co/Mn/Br > Co/Ce/Br > Co/Br > Mn/Br > Ce/Br. Synergy factors can be calculated at a given value of conversion. From figure 3 at 3.00 hr one obtains SFCo/Mn/Br =1.38 and SFCo/Ce/Br = 1.21. Using

1124 the data on figure 3, one obtains the first order rate constant for toluene disappearance as 1.57, 1.43, 0.86, 5.60 and 3.22 (xl03,s -1) for Co/Br, Mn/Br, Ce/Br,Co/Mn/Br, and Co/Ce/Br catalysts, respectively. From these rate constants one obtains SFCo/Mn/Br = 1.86 and SFCo/Ce/Br = 1.32. The latter are probably the most meaningful synergy factors. The SF values on table 3 should be measured at conditions where the concentration of the oxidant and catalyst are identical. This is very difficult to do for the following three reasons: 20 90

8O 7O

~r

6o

0 so W ,-

-r

Co/Br

-a-

Mn/Br

j.~r"

Ce/Br

/~

=

40

'0 9 ,.m

>m

~ "0

--x- Co/Mn/Br

/X------- N

18

M

\

14 12

,,,o .c

C o / C e / B r .-- x j

8 20

.

Co/Br

Ibe~

[]

"

A

Mn/Br

m 2

10

0

0 0

2

4

time,hr

Figure 3. Activity of Selected Catalysts

0

20

40

\

C /Br

•

Co/Mn/Br

=

C o )/ C e / B r 60

80

Conversion,%

Figure 4. Benzaldehyde Formation for Selected Catalysts 1. One is not observing only the autoxidation of toluene in these experiments but rather the co-oxidation of toluene with benzaldehyde, benzyl alcohol, benzyl acetate, and the benzyl bromide. Benzaldehyde, benzyl alcohol, and benzyl acetate are all more reactive than toluene [13]. 2. We do not know if the distribution of these intermediates will remain constant for a given conversion of toluene with a given catalyst. Figure 4 indicates otherwise. Unfortunately, the concentrations of the benzyl alcohol and acetate were not measured. The benzaldehyde yield, at a given value of conversion, varies considerably from catalyst to catalyst. 3. The amount of catalytically active bromide is different for each due to the variation of the concentration of (z-bromotoluene during the experiments. (z-Bromotoluene is an inactive form of bromine in these reactions [ 14]. The appropriate manner to express the yield of benzylic bromides is on the initial sodium bromide added rather than on the initial toluene basis since sodium bromide is the limiting reagent. The benzylic bromide yields, on a sodium bromide basis, range from 22% for the Mn/Br catalyst to 93 for the Co/Mn/Br catalyst, see figure 5. Thus the effective bromide concentration varies from 7 to 78%

1125 Interestingly, the more active catalysts contain the highest amount of inactive bromide, i.e., the highest yield of o~-bromotoluene, see figure 5 ! 4.3. The Steady State Oxidation States of the Metals and Reduction in the Formation of Carbon Oxides (CO, CO2).

Observation of the colors during the metal/bromide oxygenation of toluene are consistent with very low steady state concentrations of the metals in their higher oxidation states. Previously we had estimated based on UV-VIS studies that only 0.6% of the cobalt in a Co/Mn/Br catalyzed oxidation of p-xylene consisted of Co(III) [10]. In the absence of 100 90 80

_~ r

9

70 60

4O

r

20

X

9 Co/Br

Xlx'X'

9 "0

A& A= A

5O O I-J IZl

x

9 9 9

.

9Mn/Br 9

=A

9Ce/Br

m

a

9

9 Co/Br

[]

&e

~2

9M n / B r

3O

x Co/Mn/Br

9Ce/Br

' v

9

x

Co/Ce/Br

X

X

x Co/Mn/Br

10

x 0

20

40

Co/Ce/Br

60

80

Converslon,%

Figure 5. Formation of o~-Br-Toluene For Selected Catalysts

o 100

~" 0

20

40

60

80

converslon,%

Figure 6. Formation of Carbon Dioxide for Selected Catalysts

bromide, the UV-VIS are consistent with substantial amounts of Co(III) and Mn(III). The observed colors of the oxygenations of toluene, during the first hour, are: Catalyst Observed During Rxn Colors of Metal Acetates in Acetic acid Co green Co(II), pink Mn brown Co(III), green Co/Br pink Ce(III), light yellow Mn/Br colorless Ce(IV), colorless Ce/Br colorless Mn(II), colorless Co/Ce/Br pink Mn(III), brown Co/Mn/Br pink The replacement of some of the cobalt by manganese reduces the steady state concentration of Co(III) even lower during the catalytic reaction which results in a lower rate of carbon dioxide formation [10]. This is also observed when some of the cerium is replaced by cobalt i.e. in a Co/Ce/Br catalyst, see figure 6.

1126 4.4 Hammett Studies of Co/Mn/Br and Co/Ce/Br catalysts Hammett studies of metal/bromide catalysts have been previously reported by Kamiya [15] and ourselves [10]. The results for a Co/Ce/Br catalyst are given on figure 2. Six different reagents are reported using a range of sigma+ substituent constants from 0.79 (p-nitrotoluene) to -0.78 (p-methoxytoluene). The rho value of the Hammett equation (the slope of figure 2) is -1.4(0.12) which is similar to that of-1.2(0.11) [17] determined for a Co/Mn/Br catalyst. The negative value of the rho values indicate a build-up of positive charge in the aromatic ring in the transition state [ 16]. In cobalt only oxidations, which utilize Co(III) as the oxidant, the rho value is -1.8, consistent with a radical cation mechanism [ 17]. The lesser values of the metal/bromide rho values reported here, and by Kamiya [ 15] indicate a different transition state than a radical cation. The value of rho for the benzylic abstraction of hydrogen atoms by the bromide atom of-1.82 (in acetic acid) [18] is also not consistent with the values of-1.2 to -1.4 reported here as is benzylic abstraction by benzyl peroxy radicals in toluene which give a value of-0.63 [19]. The transition state may therefore involve bromine atom abstraction but with the bromide partially or wholly bonded to the manganese(II) metal as originally suggested by Kamiya [ 15] and supported by the authors [2,4,10]. REFERENCES

1. R. Landau and A. Saffer, Chem. Eng. Prog., 48,(1968)20. 2. W. Partenheimer, Catalysis Today, 231 (1995). 3. Y. Agawa, T. Yamada, Aromatikkusu 46(1994)80 (Japan), CA121(3):34977fS. 4. T. Oyama, and J. W. Hightower, "Catalytic Selective Oxidation", Amer. Chem. Soc., 1993, chapter 7 by W. Partenheimer and R.K. Gipe. 5. W. Partenheimer, J. Mol. Catal., 67(1991)35. 6. Jones, G.H., J. Chem. Research (S) (1981) 228-229. 7. Jones G.H.,J. Chem. Soc., Chem. Commun.,(1979)536;Jones, G . H . , J. Chem. Research (M), (1981)2801 ;Jones, G.H.J. Chem. Research (S) (1982)207. 8. There are at least three different forms of cobalt(Ill) in acetic acid. Following Jones notation, we have labeled these Co(III)a, Co(III)i, and Co(III)c. 9. D.A.S. Ravens, J. Chem. Soc.,55(1959) 1768. 10. D.W. Blackburn, "Catalysis of Organic Reactions", Marcel Dekker,Inc., 1994, chapter 14, by Walt Partenheimer. 11. S. Carra, E. Santacesaria, Catal. Rev.-Sci.Eng., 22(1980)75. 12. J.E. Mclntyre and D.A.S. Ravens, J. Chem. Soc., (1961) 4082. 13. R.A. Sheldon and J.K. Kochi, "Metal-Catalyzed Oxidations of Organic Compounds", Academic Press, New York, N.Y. 1981, p 23. 14. K. Sakota, Y. Kamiya, and N. Ohta, Bull. Chem. Soc. Japan, 41 (1968)641. A.S. Hay, and H.S. Blanchard, Can J. Chem.,41(1965)1306. See also reference 2. 15. Y. Kamiya, J. Catalysis,44(1974)480. 16. T.H. Lowry, K.S.Richardson, "Mechanism and Theory in Organic Chemistry, Harper and Row, 3rd Edition, page 143. 17.C.F. Hendriks, H.C.A. Beck, P.M. Heerjes, Ind. Eng. Chem. Prod. Res. Dev.,17(1978)266. The latter report a value of-1.9 but did not account for the number of equivalent hydrogen atoms. We have taken the values from this reference

1127 and plotted them vs. sigma + values to obtain value of- 1.8. 18. J.R. Gilmore, J.M. Mellor, Chem. Comm., (1970),507. 19. G.A. Russell, J. Amer. Chem. Soc.,78(1956), 1047.

This Page Intentionally Left Blank

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 1997 Elsevier Science B.V.

1129

T h e K i n e t i c s of the Partial O x i d a t i o n of M e t h a n e to F o r m a l d e h y d e over a S i l i c a - S u p p o r t e d V a n a d i a Catalyst A.W. Sexton and B.K. Hodnett Department of Chemical and Environmental Sciences, University of Limerick, IRELAND. ABSTRACT A kinetic study has been carded out on the partial oxidation of methane to formaldehyde over a silica-supported vanadia Catalyst. The results indicate that oxygen was adsorbed on the catalyst and took part in the reaction in an Eley-Rideal or Mars-van Krevelen manner. The nature of the interaction with the catalyst was dependent on whether the reaction took place in methane rich (PcH4 - 80 kPa) or lean (PcH4 = 4 kPa) conditions. A reaction mechanism for the partial oxidation of methane to formaldehyde is proposed, which is consistent with the data reported here. Methanol oxidation experiments over this catalyst suggested that it was not an intermediate under the conditions employed during this study. INTRODUCTION In recent years there has been much interest in the conversion of methane to value added products, such as ethane/ethylene [1], methanol [2], formaldehyde [3-5] and synthesis gas [6]. Many studies have been carded out on the partial oxidation of methane to formaldehyde over silica [7], and over molybdena [8,9], vanadia [5,10] supported on silica, or FeNbB-O [11]; with nitrous oxide [7-9] or oxygen [7,10] as the oxidant.

Table 1 Published Kinetic Findings on the Partial Oxidation of Methane to Formaldehyde with Oxygen Catalyst

Rate Law

SiO 2 MoO3/SiO 2

kpo20pCH41

V205/SiO 2

kpo2~

FeNbB-O * - Refers to methane conversion

HCHO Ea / kJ mol "l 250* (507- 527~ 117* (527- 597~ 189 227 256

Ref. [7] [14] [15] [11]

Several authors have presented data on the activation energy of the reaction of methane to formaldehyde with N20 [7-9,12,13] or 02 [7,14,15] as oxidant. Those using oxygen are

1130 summarised in Table 1. Rate laws, measured under various conditions, have also been proposed in the literature [7-9,12-15]. Here the results of a new kinetic study of the partial oxidation of methane to formaldehyde, using air/oxygen are presented, along with a proposed mechanistic scheme consistent with the data.

EXPERIMENTAL The catalyst studied was 1 wt% vanadium (as metal) supported on Cab-O-Sil (M5), hereafter referred to as 1V-cabosil. The catalyst was prepared by the wet impregnation method described earlier [5]. Catalytic testing was carried out in a fixed bed quartz microreactor with on-line analysis of the reaction products, as described previously [16]. The catalyst reached steady state within minutes throughout these experiments. Products were analysed every 30 minutes, for 2 hours at each temperature. The influences of W/F, methane and oxygen partial pressures in methane rich (up to 80 kPa CH4 in the feed) and methane lean (up to 4 kPa CH4 in feed) conditions were examined at 550 and 600 ~ Negligible reaction was observed at lower temperatures. The reactant partial pressures were varied, keeping the total flow rate at 25 ml min ~, by adding helium ballast to the system. Oxygen (02) was the chosen oxidant. The oxygen partial pressure was varied in methane rich (Pcm = 81 kPa) and methane lean ( P c H 4 -- 20 kPa) conditions. The CH4 partial pressure was varied in the range 0-81 kPa (keeping the oxygen partial pressure at 20 kPa). The effects of W/F (catalyst mass used/reactant gas flow rate) were assessed, using air as the oxidant. This was done in methane rich conditions, using 0.1 g 1V-cabosil while varying the feed gas flow rate from 6.25 - 100 ml min 1. Methanol oxidation experiments were carried out in order to determine if methanol was an intermediate in the production of formaldehyde from methane. To this end a methanol saturator was placed upstream of the reactor. The saturator was submerged in an ice/acetone bath (at -16 to - 20 ~ keeping the saturated methanol partial pressure at 5 kPa. This was approximately equivalent to the total carbon containing products generated during standard reaction conditions. The gas feed stream to the saturator consisted of 81 kPa helium and 20 kPa air. The flow rate was varied from 6.25 - 100 ml min "~. RESULTS In order to determine the catalyst stability two experiments were carried out. Firstly, the catalyst was tested in 10 ~ intervals, from 450-600 ~ The catalyst temperature was than lowered in 10 ~ intervals back down to 450 ~ No significant change in activity or product distribution was noted on the downward cycle. Secondly, the catalyst was tested for 50 hours in methane rich conditions. Again with no significant change in activity or selectivity. Hence, the catalyst was deemed to be stable under the experimental conditions employed here. If the methane-air system was being used in a kinetically controlled regime the observed reaction rates would be independent of reaction mixture delivery rate to the catalyst bed. The methane conversion rate increased up to 25 ml min ~ at 550 ~ and with almost all flow rates

1131 at 600 ~ These data indicate that at 600 ~ the reaction is under diffusion control. However, at 550 ~ the reaction was seen to be operating under kinetic control, over the range of flow 25-100 ml min ]. Figure 1 illustrates the selectivity-conversion relationship between the various products obtained at 550 and 600 ~ during these experiments. It can be clearly be seen that HCHO was a primary product. Indeed as conversion increased HCHO selectivity decreased, with an analogous increase in CO selectivity. Further increase in CH4 conversion lead to the onset of increased CO2 selectivity, indicating a sequential reaction from methane to products as follows: C H 4 '-)

HCHO 41, CO r CO2 100

100 80

80 -

~60

.~60.>, 40

.v..~

40

CO

r~20

~20

--

0

1 : 2 Conversion7 %

3

1.5

i

i

v"

2.0 2.5 Conversion / %

Figure 1. Selectivity as a Function of Methane Conversion at (a) 550 ~ 0.1g 1V-cabosil, 81 kPa C H 4 and 20 kPa air

3.0

and (b) 600 ~

The Arrhenius plot for methane activation and formaldehyde production are shown in Figure 2. As can be seen from the plot shown in Figure 2, two distinct regions were noted in the Arrhenius plot for methane activation. There was a linear region which corresponded to a lower temperature range of 490-550 ~ with a second from 550-600 ~ The activation energy for methane conversion, in the temperature range 490-550 ~ was 323 kJ mol l. The activation energy for methane activation decreased towards zero over the temperature range 550-600 ~ indicating that the reaction was limited by diffusion at the highest temperatures studied. Over the same temperature range (490-550 ~ the activation energy for formaldehyde was 242 kJ mol l. These values are consistent with those previously reported by Otsuka and Hatano [ 11] (see Table 1). The plots for methane activation and formaldehyde formation were noted to be coincident at the low range of temperatures examined (480-500 ~ while some divergence between the two sets of data was noted at higher temperatures. This was due to the fact that formaldehyde selectivity was high at low temperatures, while sequential reactions became more important as the temperature was increased. The influences of reactant partial pressures on formation rates of the various products are illustrated in Figure 3.

1132

7

1.1

1.2 1.3 1/T x 10-3 / K "1

1.4

Figure 2. Arrhenius Plot for Temperature Range 480-600 ~ W / F = 8 . 9 6 x 10 "2 k g r a i n m o l "1, 81 k P a CH4 a n d 2 0

kPa air 3.0

a

HCHO+CO

CO

3.0

L~ 2.0

9

1.0

,-

'~

~

.,~

2.0-

The maximum rate ever observed, during this study, for C O 2 production was 300 ~tmol kg 1 s~ while the maxima for formaldehyde and carbon monoxide were 1300 ~tmol kg ~ s~ and 3400 ~tmol kg 1 s "1 respectively. Hence, combination of the rates for HCHO and CO was the important factor in determining the overall shapes of the total product formation curves and CO2 will not be discussed further here. The rate of evolution of formaldehyde as a function of methane partial pressure (Pen4) in the range 20-81 kPa is shown in Figure 3a. Formaldehyde and carbon monoxide rates increased fairly linearly with increasing Pen4. In methane lean (Pcm = 4 kPa) conditions overall reaction rates were 10 times lower, but very little CO or CO2 was produced (Figure 3c). The formaldehyde production rate decreased slightly for Po2 in the range 20-

~

0

20

40

60

80

100

A

CO

1.0 0.0

0.0

t []

0

'

I

5

'

I

'

10

I

15

'

t j

20

Po2 / kPa

PCH4 / k P a

0.30 "7

% 0.20 O

Figure 3. Product Formation Rates as a Function of Reactant Partial Pressures (a) 20 kPa 02 (b) 81 kPa CH 4 (c) 4 kPa CH4, T = 550 ~ W/F = 8.96 x 10.2kg min mol~.

~0.10 0.00

II

0

i

B I

25

,--

i

B

i

,--

50 75 Po2 / kPa

,

m

100

1133 81 kPa and fell more rapidly above 81 kPa. This suggested that the rate determining step, in methane lean conditions, may have involved the contact of methane with the catalyst active site, which was inhibited by the large excess of oxygen in the system. For methane rich feeds, formaldehyde and carbon monoxide production increased with increasing oxygen partial pressure (Figure 3b). An attempt was made to ascertain the reaction orders for each of the three products, namely HCHO, CO and CO2, with respect to both methane and oxygen partial pressure in methane rich conditions and with respect to oxygen in methane lean conditions. The reactions were found to be neither first nor second order with respect to either of the reactants. The difficulty in determining reaction orders may have been due to the fact that in methane rich conditions almost total oxygen conversion was achieved, resulting in a considerable gradient in oxygen concentration in the catalyst bed. The data were then subjected to the test plots for Langmuir-Hinshelwood, Eley Rideal and Mars-van Krevelen mechanisms, allowing for both associative and dissociative desorption cases. Very poor correlation coefficients were - 0.0015 obtained with all Langmuir-Hinshelwood o reaction test plots. Figure 4 illustrates a sample test plot for a Mars-van Krevelen 0.0010 (MvK) reaction mechanism, with an 550~ associative oxygen adsorption step. It was difficult to distinguish between 0.0005 o associative and dissociative oxygen adsorption with MvK reactions. It is, 0.0000 ' I ' I ' however, clear that oxygen adsorption was an 0.00 0.10 0.201 0.30 important factor in the reaction in methane 1/POE / kPa rich conditions 9 In methane rich feed conditions the data slightly favoured a model Figure 4. Sample Mars-van Krevelen Test which entailed dissociative oxygen adsorption Plot for Associative Oxygen Adsorption at as the slow step in Methane Rich Conditions For methane rich conditions the data W/F = 8.96 x 10-2kg min mol~., 81 kPa CH4 indicates that dissociative adsorption of oxygen was an important factor, based on the test plots for an Eley-Rideal mechanism. Conversely, associative oxygen adsorption appeared to be important in methane lean conditions. There was little conclusive evidence, in the form of these correlation coefficients, to chose between a Mars-van Krevelen or an Eley-Rideal type mechanism. However, our observations indicate that the most appropriate forms of the rate equations for HCHO production are the following:

6oo

METHANE RICH CONDITIONS

9 __ RHCHO kPcH4

{

PO2 I+KPo2 0.5

po 2

METHANE LEAN CONDITIONS" RHCHO = kPeH 4 1+ Kpo2

(Eqn. 1)

}

(Eqn.2)

1134

100

80

~,

~

60

-

40

r~

20

CO

o

'

30

,,~

40

50

'

,

'

60

,

'

70

I

'

,

80

90

'

100

Conv./% Figure 5. Selectivity as a Function of Methanol Conversion at 550 ~

0.1 g 1V-cabosil, 5 kPa CH3OH,21 kPa Air in feed

The role of methanol as a possible intermediate is examined in Figure 5. Figure 5 shows that methanol is readily converted to HCHO in the standard reaction conditions over 1V-cabosil. However, a number of factors indicate that methanol is not an intermediate in the reaction. Firstly, methanol was never detected under TAP reaction conditions [4,15]. Secondly, McCarthy has indicated that the rate constants for a large number of fundamental reactions of methane and its derivatives with surface oxygen species may be calculated using the extended version of the Arrhenius equation indicated in Eqn. 3. (Eqn. 3)

.

t

McCarthy has reported rate constants for the following reactions [19]" CH4 +O(surO #- CH 3 + OH(surf)

(Rxn. 1)

CH3OH + O(surOd-- CH3OH + OH(surf)

(Rxn. 2)

The values for the extended Arrhenius equation are given here [19]" ko(cH4)= 1.73 x 1012mol cm s k0(cmo~= 1.35 x 1012mol cm s n(CH4) = 0

n(CmOH) = 0

Ea(CH4) = 63.73 kJ mol 1

E~cmor0 = 46.99 kJ mol ]

These data predict more methanol than formaldehyde in the product stream during methane oxidation; this clearly is not the case. DISCUSSION

The results indicate only sequential reactions occur in the partial oxidation of methane to formaldehyde, with no evidence for the existence of parallel reaction pathways to COx. An Eley-Rideal/Mars-van Krevelen type of mechanism was found for the partial oxidation of methane to formaldehyde. The differences in the rate equations were due to differences in the amounts of oxygen present in methane rich and lean conditions. Methanol was not an intermediate in the reaction. Methanol oxidation experiments indicated that methanol was oxidised sequentially to formaldehyde, carbon monoxide and hence to carbon

1135 dioxide. However, no evidence was observed to suggest that any methanol was present during the partial oxidation of methane over the catalyst. Results from the kinetic experiments for the partial oxidation of methane indicated that a complex series of fundamental reactions leading to formaldehyde took place. Two different kinetic expressions have been derived to describe the partial oxidation of methane in rich and lean feed conditions. An attempt is made here to elucidate why this might be so. During the partial oxidation of methane to formaldehyde, methane reacted with some form of active oxygen at the catalyst surface. In methane lean experiments the reaction mixture contained up to 97 kPa O2 (96% of the total reactants in the system). In methane rich conditions the maximum amount of O2 present was 4 kPa (4% of the total). In each instance oxygen had to absorb at the catalyst surface before methane oxidation could occur. Since the chance of encountering another oxygen molecule before reaction with surface electrons, was much lower in methane rich feed conditions the rate of reaction was governed by Eqn. 1 (i.e. dissociation of oxygen occurred as indicated in Rxn. 3a). Conversely, in methane lean conditions oxygen arrived at the catalyst surface faster than electrons could be delivered, hence, the associative adsorption of oxygen was favoured (Rxn. 3b); Eqn. 2 describes the system behaviour under these conditions. This rationale affords an explanation as to why two different rate equations were found to apply in methane rich and lean feed experiments. These data are consistent with the following reaction scheme:

O2(g) + 2e~ur0 # 2Oiads )

(Rxn. 3a)

O2(g) + e(surf) e O2(ads)

(Rxn. 3b)

2CH4(g) + O (ads) ~ CH ~ + OH iads)

(Rxn. 4a)

CH4(g ) + O~ds) ~ CH~ + OHiads )

(Rxn. 4b)

2CH ~ 4- O (Latt) @ CH 3O" + 2e~ur0

(Rxn. 5)

CH30"~ HCHO + H"

(Rxn. 6)

Proposed Reaction Scheme for Formaldehyde Production The evidence to date suggests that methane does not chemisorb on the catalyst surface [4,15]. The partial oxidation of methane has been studied in a temporal analysis of products (TAP) reactor, over 1V-cabosil. A feature of the TAP system is that comparisons of residence times of various components in the reactor, and hence on the catalyst surface, can be made. Kartheuser has shown that methane and an inert gas with molecular mass - 16 g mol ~ had the same residence times in the TAP reactor, over 1V-cabosil, implying that methane did not adsorb on the catalyst surface [15]. This is consistent with the Eley-Rideal and Mars-van Krevelen mechanisms. Hence, methane from the gas phase (Rxn. 4) reacted with surface oxygen to form CH 3" radicals. It must be noted that both forms of Rxn. 4 describe the conversion of methane, but under different conditions. Reaction 4a predominated in methane lean cases, while Rxn. 4b was more relevant to methane rich conditions. Radicals generated

1136 via Rxn. 4 could further react with the lattice oxygen anions to form CH30"radicals (Rxn. 5). These radicals could subsequently decompose giving formaldehyde (Rxn. 6). ACKNOWLEDGEMENT

This work was supported in part by the EU Joule programme, contract number JOUF-0044-

c(yr) REFERENCES

[1] M. Makri, Y. Jiang, I.V. Yentakakis and G.Y. Vaneyas, Studies in Surface Science and Catalysis, Elsevier, Amsterdam, 101 (1996) 387 [2] K. Katja, X.M. Song, and M. Huuska, Catal. Today, 21 (1994) 513 [3] F. Arena, F. Frusteri, D. Miceli, A. Parmaliana and N. Giordano, Catal. Today, 21 (1994) 505 [4] B. Kartheuser, B.K. Hodnett, H. Zanthoffand M. Baems, Catal. Letters, 21 (1993) 209 [5] M. Kennedy, A. Sexton, B. Kartheuser, E. Mac Giolla Coda, J.B. McMonagle and B.K. Hodnett, Catal. Today, 13 (1992) 447 [6] K. Heitnes, S. Lindberg, O.A. Rokstad and A. Holmen, Catal. Today, 21 (1994) 471 [7] S. Kasztelan and J.B. Moffat, J. Chem. Soc., Chem. Commun., (1987) 1663 [8] H.-F. Liu, R.-S. Liu, K.Y. Liew, R.E. Johnson and J.H. Lunsford, Jr. Am. Chem. Soc., 106 (1984)4117 [9] M.M. Khan and G.A. Somorjai, J. Catal., 91 (1985) 263 [10] N.D. Spencer and C.J. Pereira, J. Catal., 116 (1989) 399 [11] K. Otsuka and M. Hatano in "Methane Conversion by Oxidative Processes Fundamental and Engineering Aspects", E.E. Wolf (Ed.), Van Nostrand Reinhold, New York (1992) [12] M.D. Kennedy, Ph.D. Thesis, University of Limerick, Ireland (1992) [13] K.J. Zhen, M.M. Khan C.H. Mak, K.B. Lewis and G.A. Somorjai, J. Catal., 94 (1985) 501 [ 14] N.D. Spencer and C.J. Pereira, AIChE J., 33 (1987) 1808 [ 15] B.J. Kartheuser, Ph.D. Thesis, University of Limerick, Ireland (1993) [16] E. Mac Giolla Coda, M. Kennedy, J.B. McMonagle and B.K. Hodnett, Catal. Today, 6 (1990) 559 [17] G.C. Bond, "Heterogeneous Catalysis - Principles and Applications", Oxford Chemical Series, Oxford University Press, Oxford (1974) [ 18] A.W. Sexton, Ph.D. Thesis, University of Limerick, Ireland (1995) [19] J.G. McCarthy in "Methane Conversion by Oxidative Processes - Fundamental and Engineering Aspects", E.E. Wolf (Ed.), Van Nostrand Reinhold, New York (1992)

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 1997 Elsevier Science B.V.

1137

Catalytic Destruction of Volatile Organic Compounds on Platinum/Zeolite A. O'Malley and B.I( Hodnett Dept. of Chemical and Environmental Sciences, University of Limerick, IRELAND.

ABSTRACT A range of platinum (Pt) exchanged zeolites were tested for the catalytic destruction of volatile organic compounds (VOC's) in air. The range of zeolites included ZSM-5 and [3zeolite in their acidic forms and the results are compared with conventional Pt/AI203, Pt/TiO2 and Pt/SiO2 catalysts. Catalysts typically contained 0.5-6.5wt% platinum and dispersions, as measured by hydrogen chemisorption, were in the range 39-65%. Toluene and ethylbenzene, prime indoor sources of VOC's, originating from paints, adhesives and combustion products were fed to the reactor in the concentration range 200-5000ppm with excess oxygen. Typically in our reaction conditions, 100% conversion of toluene or ethylbenzene could be achieved over Pt/~-zeolite and Pt/ZSM-5 below 175~ 1.1NTRODUCTION Volatile Organic Compounds, VOC's, are an important class of air pollutants, commonly found in the atmosphere at ground level in all urban and industrial centres. Strictly speaking, the term volatile organic compounds refers to those organic compounds which are present in the atmosphere as gases, but under normal conditions of temperature and pressure would be liquids or solids. Those VOC's which are present in the atmosphere as a result of human activities, arise mainly from motor vehicle exhausts, evaporation of petrol vapours from motor cars, solvent usage, industrial processes, oil refining, petrol storage and distribution, landfilled wastes, food manufacture and agriculture. Natural biogenic processes also give rise to substantial ambient concentrations of organic compounds and include emissions from plants, trees, animals, natural forest fires and anaerobic processes in bogs and marshes [ 1]. Emission inventories are now available for VOC's for most European countries. European emissions of low molecular mass volatile organic compounds from human activities amounted to about 24 million tonnes/year in 1989. This total is comparable with levels of sulphur dioxide and nitrogen oxides (as NO2), each of which are of the order of 20 million tonnes/year for Europe [ 1]. The US Clean Air Act of 1990 calls for a 90% reduction over the next nine years in emissions of toxic chemicals, 70% of which are volatile organic compounds [2]. A protocol, signed by most of the European countries, calls for a 30% cut in VOC emissions by 1999 relative to the 1988 levels [3]. For VOC destruction, catalytic oxidation is one of the most important air pollution control techniques. In this process VOC's are oxidised over a catalyst at temperatures much lower than those required for thermal oxidation. The temperature for 100% conversion of a given

1138 VOC to CO2 and H20 varies depending on the nature and concentration of the VOC and the catalyst. Typical literature reports include a 5%COO catalyst on a carbonaceous support which totally oxidised 600ppm toluene at 210~ and 1300ppm toluene at 250~ [4]. Pt/monolith completely oxidised 500ppm toluene at 230~ [5]. Zeolites also showed a good performance for the complete oxidation of VOC's at low temperature. For example 140ppm ethene was totally oxidised at 200~ using ZSM-5 and 150~ using Cu/ZSM-5 [6]. In this work two aspects of the destruction of VOC's are investigated. The first relates to the performance of a new series of platinum exchanged zeolites and the second elucidates the structural factors which determine the ease of oxidation of two VOC's, ethylbenzene and toluene. 2.EXPERIMENTAL

2.1 Catalyst Preparation A series of Pt/A1203 catalysts containing 0.5-6.5wt% Pt were prepared by wet impregnation. Briefly, the required amount of chloroplatinic acid, H2PtCI6.6H20, (Aldrich), was dissolved in excess ethanol and added to pre-calcined Al203 (supplied by Rhone Poulenc, particle size 212-500nm). The prepared mixture was stirred constantly for 5 hours. The excess liquid was removed under vacuum at 80~ in a rotary evaporator. The resulting solid was dried in air at 100~ for 12 hours and calcined at 450~ for 4 hours [7]. 2wt% Pt supported on TiO2 (Aldrich) and 2wt% Pt supported on SiO2 (Carbosil) were also prepared by wet impregnation. ~Itte range of Pt/zeolite catalysts were prepared by ion exchange [8]. The zeolites, NH4+/ZSM-5 and NH4+/[3-zeolite (supplied by P.Q.), were converted to their acidic form by heating to 427~ in air. The required amount of Pt(NH3)4C12.H20 precursor (Johnson Matthey) was added to a suspension of H20 and the zeolite. This was stirred at reflux for 6 hours. The sample was filtered and washed three times with hot water, then dried for 12 hours at 120~ in air and calcined for 4 hours at 450~ 2.2 Characterisation The platinum contents of the prepared catalysts were measured by Atomic Absorption Spectroscopy using a Varian Spectra AA 400 plus. In preparation for analysis, all samples were dissolved in aqueous HF and only measured platinum loadings will be referred to below. Surface area measurements at 77K were performed using a Micromeritics Gemini lIl 2375 surface area analyser with nitrogen as the adsorbing gas. All samples were outgassed at 200 ~ C before analysis. Thermogravimetric analyses were carried out on a Stanton Redcrofi TG770 thermobalance with a Bausch and Lomb Omnigraphic 2000 XY Recorder. Approximately 8mg of catalyst was heated at a rate of 10~ rain1 in a flow of 30 mls rainl air from 20~ to 900~ All samples were crushed to a fine powder and analysed by X-ray diflffaction from 20 = 10% 70 ~ using a Philips Diffractometer with nickel filtered CuKa radiation (~,=1.5406A~ Platinum dispersion was measured by hydrogen chemisorption. A dynamic pulse technique was used similar to that reported by Heck et al. Briefly, 100mg of catalyst was introduced into a quartz glass vertical reactor and kept in position by 2 plugs of quartz wool. A pulse of

1139 hydrogen (1%H2 in argon) was injected into an argon stream and hydrogen uptake was monitored using a TCD. 2.3 Catalyst Testing The catalytic destruction of toluene and ethylbenzene was carried out using a continuous flow system All gases (helium, 99.99% and 29.7% oxygen in helium) were supplied by BOC gases and the flow rates were controlled by mass flow controllers (Tylan General model FC2900). Toluene was introduced into the vapour phase by passing helium through a saturator. The saturator consisted of a stainless steel tube filled with molecular sieve (4A) soaked in the organic liquid. The partial pressure of toluene in the helium stream could be controlled by adjusting the saturator temperature or by diluting the toluene/helium stream leaving the saturator with oxygen/helium from the mass flow controllers. The reactant stream consisted of 5000ppm toluene or 1000ppm ethylbenzene, 12vo1% O2 and the remainder helium, at a total flow of 50ml/min, unless otherwise stated. In all experiments 0. l g catalyst was tested using a quartz glass vertical reactor tube. The catalyst was kept in position by 2 plugs of quartz wool. The feed stream entered a series of 3-way valves which allowed the reactor to be bypassed or not as required. A K type thermocouple was inserted into the catalyst bed. The temperature of the furnace was controlled by a Eurotherm Before testing, the catalyst was pre-treated for 90 minutes in a stream of helium or hydrogen at 30ml/min at 400~ Atter pre-treatment the catalyst was cooled to the desired temperature and the reaction was started by introducing the reactant stream Product gases were analysed by on-line gas chromatography (Varian model 3400 GC) using a porapak T column and a Thermal Conductivity Detector. 3. RESULTS AND DISCUSSION An investigation into catalyst deactivation was carried out under typical reaction conditions. Briefly, the reactant stream containing 5000ppm toluene, was passed over 2.5wt% Pt/[3-zeolite at 156~ Under these conditions toluene conversion was 30% and no loss of activity occurred over a 7 hour period, in spite of the fact that the used catalyst was found by thermogravimetry to contain 10wt% coke. Figure 1 illustrates the conversion of toluene with reaction temperature for a series of platinum supported alumina catalysts. A1203 presents little activity over the temperature range studied. At 400~ a conversion of 20% was observed. Impregnating the support with platinum resulted in a significant increase in catalyst activity. A 2wt% Pt/AI203 converted 100% toluene at 250~ Increasing the platinum content decreased the temperature for 100% conversion. 100% conversion oftohene was achieved over 6.5wt% Pt/A1203 at 220~

1140

=

100-

0

9 6.5wt% PtJAI203

1,=.

W

|>

5.5wt% Pt/AI203 9 4.5wt% Pt/AI203 x 2wt% Pt/AI203

75-

O~ o .~ @

D

v

|

50-

0

AI203

25-

m

0

I-

loo

3oo Temperature

400

(oC)

Figure 1. Conversion of toluene with reaction temperature using the indicated Pt/A1203 catalysts. (PTol,=e=5000ppm, Po2=12vol%, W/F=0.06g s/ml). A similar trend was observed when ZSM-5 was used as support for various platinum contents as shown in figure 2. ZSM-5 showed no activity below 225~ and reached 25% conversion at 400~ Exchanging the ZSM-5 with 0.5wt% platinum resulted in 100% conversion of toluene at 227~ and a further improvement in activity could be achieved by increasing the platinum content to 2wt%. c 0 Itl

100

v

75-

[] 2wt% Pt/ZSM-5 9 1.9wt% Pt/ZSM-5 x 0.5wt% Pt/ZSM-5

50-

o ZSM-5

.m

L.

r0 e,,

25m

0 I--

0 100

1,.,1,~1,

I

200

v

31~0

4( )0

T em perature (oC)

Figure 2. Conversion of toluene with reaction temperature using the indicated Pt/ZSM-5 catalysts. (PTol,=e=5000ppm, Po~= 12vo1%, W/F=0.06g s/ml). A slightly different pattern was observed with the series of Pt/13-zeolites as seen l~om figure 3. 13-zeolite exhibited minimal activity even at 400~ Total conversion of toluene at 170~ was achieved by exchanging 0.5wt% Pt into the zeolite. Increasing the platinum content presented no further improvements in activity.

1141 tO W

100

o m

z~ 2.3wt% Pt/13-zeolite 9 2wt% Pt/13-zeolite x 0.5wt% Pt/13-zeolite

75

> C 0

o 13-zeolite

50

e~

25 o p.

OI 100

o 2()0

p 300

Temperature

~

400

(oC)

Figure 3. Conversion of toluene with reaction temperature using the indicated Pt/13-zeolite catalysts. (Pxol,=e=5000ppm, Po~= 12vo1%, W/F=0.06g s/ml). Figures 4(a) and (b) summarise the effect of platinum content on the temperature for complete conversion and for 50% conversion of toluene, respectively. Increasing the platinum content of alumina results in a decline in the temperature required for complete conversion (figure 4(a)). This is also true for platinum exchanged ZSM-5. However, varying the platinum content on 13-zeolite does not si~ificantly alter the temperature for complete conversion. A similar dependency on platinum content was observed for the temperature required for 50% conversion. (a) (b) 250

250

225-

2250 o o

~

tO

o Pt/ZSM-5

200-

x

~'~ 200-

9 Pt/AI203 PUp-zeolite

1-

175

150

o

i

~

~

~

Pt (wt%)

~

6

~

o Pt/ZSM-5

9PvA~o3

175150

o

X PtJl3-zeoUte

i

~,

3

~,

~

6

~

Pt ( w t % )

Figure 4. The effect of platinum content on the temperature for (a) complete conversion and (b) 50% conversion of 5000ppm toluene, respectively. (Po~=12vol%, W/F=0.06g s/ml). Table 1 compares the platinum dispersion of a 6wt% Pt/Al203 and 2wt% Pt/[3-zeolite before and after reaction. Also included is the coke content of the t~esh and used catalysts. The 6.5wt% Pt/Al203 presented a platinum dispersion of 39% while the 2wt% Pt/13-zeolite was significantly lower at 21%. Following reaction, this dispersion remained unchanged, however 10wt% coke was detected on 6.5wt% Pt/Al203 and 12wt% coke on 2wt% Pt/13zeolite.

1142

Catalyst

Dispersion (%)

6.5wt% Pt/A1203(flesh) 6.5wt% Pt/A1203(used) 2wt% Pt/13-zeolite (flesh) 2wt% Pt/~-zeolite (used)

39 39 21 21

Table 1

Coke (wt%) 0 10 0 12

Two aspects of these results will be discussed here, namely the low platinum dispersion and the presence of si,,~ificant amounts of coke on alumina and zeolite catalysts after testing. The latter finding is probably due to a combination of test temperatures below 250~ and the propensity for toluene to form coke. In any event, coke formation did not seem to inhibit the oxidation reaction to any significant extent. The low platinum dispersions, confirmed by XRD measurements, indicate that the active species may reside outside the zeolite micropores on Pt/[3-zeolite. This phenomenon was also illustrated by Smimitios et al [10]. The role of the zeolite support then needs to be addressed if its function is not to supply the dispersing medium for the platinum phase. Figure 5 compares various platinum supported catalysts for converting 1000ppm ethylbenzene. Platinum supported on [3-zeolite, TiO2 and SiO2 present similar activities, completely converting ethylbenzene at 170~ 173~ and 179~ respectively. However 2wt% Pt/AI203 required a temperature of 233~ for complete oxidation. The supports alone showed no si~ificant activity in the temperature range tested. 100

N

9 =

O

,=

"-

,.r

c

M.I

O O

"9

9 2wt%

Pt/TiO2

9 2wt% Pt/SiO2

75 50

Pt/AI203

o

2wt%

x

0.5wt% Pt/~-zeolite

25 0 100

200

Tem perature

3()0

400

(oc)

Figure 5. Conversion of ethylbenzene with reaction temperature using the indicated catalysts. (PF~hylb~e= 1000ppm, Po~= 12vo1%, W/F=0.06g s/ml). Figure 6 presents a comparison of ethylbenzene and toluene conversions over 2wt% Pt/A1203.

1143 100 C 0 W t....

75-

q)

50-

C 0

25-

(.1

0 100

150

200

2,50

300

Temperature (oC) Figure 6. Conversion of(o) 5000ppm ethylbenzene and (o) 5000ppm toluene with temperature using 0.1g 2wt% Pt/AI203. (Po2=12vol%, W/F=0.06g s/ml). This data indicates that ethylbenzene is more easily destroyed by oxidation than toluene. This finding is consistent with many literature studies which indicate variable destructibility depending upon substrate structure. Here it is postulated that destructibility is related to the bond dissociation enthalpy of the weakest C-H bond in the substrate. For ethylbenzene this correspond to the methylene C-H bond on the ethyl group with a bond dissociation of 332.3kJ/mol. For toluene the weakest C-H bond is on the methyl group with a value of 368.2kJ/mol [11]. The hypothesis is that the slow step in the VOC destruction is the generation with the aid of a catalytically active site of a radical l~agment according to H

I R2-- CmR1

I R;

R2m(2mR1 + H"

I

R3

Thereafter total oxidation proceeds via gas phase radical chemistry. The conversiontemperature profile for a given substrate depends upon the activity of the catalyst and the inherent destructibility of the substrate. Here we propose that the lower the bond dissociation enthalpy of the weakest bond the more readily the substrate can be activated by the catalyst. At the basis of this hypothesis is the suggestion that active sites in total oxidation catalysts distinguish between C-H bonds on a bond strength basis only. REFERENCES 1. 1LG. Derwent in Volatile Organic Compounds in the Atmosphere, Ed. 1LE. Hester and R.M. Harrison, Env. Sci. Tech. (1995). 2. J.N Armor in Environmental Catalysis, Ed. J.N. Armor, ASC Symposium Series 552, (1994) 298. 3. Chemistry and Industry, (Dec 2,1991) 855. 4. M.T. Vandersall, S.G. Maroldo, W.H. Brendley, Jr.I~ Jurczyk and 1LS. Drago in Environmental Catalysis, Ed. J.N. Armor, ASC S)anposium Series 552, (1994) 339.

1144

5. B.L. Mojet, M.J. Kappers and J.T. Miller, Stud. Surf. Sci. Catal. 10 (1996) 1165. 6. L.M. Parker and J.E. Patterson in Environmental Catalysis, Ed. J.N. Armor, ACS Symposium Series 552, (1994) 301. 7. J.F. Le Page in Applied Heterogeneous Catalysis: Design, Manufacture and Use of Solid Catalyst, Ed. Technip(1987). 8. P.K Alm, S.Nishiyama, S. Tsuruya and M. Masai, Appl. Catal. A:General, 101 (1993) 207. 9. 1LM. Heck and 1L Farrauto in Catalytic Air Pollution Control, Ed. Van Nostrand Reinhold (1995). 10. P.G. Smirnitios and E. Ruckenstein, Appl. Catal. A:General, 117 (1994) 75. 11. C.Batiot and B.KHodnett, Appl. Catal. A:General, 137 (1996)179.

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 1997 Elsevier Science B.V.

1145

High temperature propane oxidation to reducing gas over promoted Ni/MgO catalysts. Role of impregnation condition and promoter on properties of catalysts M.V. Stankovi~ and N.N. Jovanovi6* Institute of Chemistry, Technology and Metallurgy, Center for Catalysis and Chemical Engineering, Njegogeva 12, 11000 Beograd, Yugoslavia

ABSTRACT The selectivity of variously promoted magnesia supported nickel catalyst, prepared under different impregnation conditions, has been investigated for the reaction of propane oxidation by air to CO and He. Promoted Ni/MgO catalysts have been prepared by multiple successive impregnation to give catalysts between 1 and 5 wt% nickel loadings. The catalysts samples were prepared by soaking of the MgO support in aqueous solutions of nickel nitrate and the corresponding nitrates of the metals used as promoters, followed by drying, calcination and reduction. Study of the oxygen chemisorption on the catalyst samples showed that the catalyst obtained from impregnation solution of the lowest concentration have the smallest Ni crystallites, the Ni crystallite size increases with the number of impregnation steps, and that the effectiveness of promoters to produce small Ni crystallites decreases in the order AI203>MgO>CaO. From the reaction study at temperatures in the range from 690 to 950~ performed over promoted Ni/MgO catalysts, it has been observed that selectivity towards CO+H2 formation decrease in the presence of promoter used in the order AI203>MgO>CaO, and that Ni loading in the catalyst appears to be arround 3 wt% to achieve the highest selectivity to CO and 1-/2 under experimental conditions studied. A correlation between the mean Ni crystallite size and the selectivity to CO and 1-/2 formation was established: the smaller Ni crystallites were more selective than the larger ones. 1. I N T R O D U C T I O N Reducing gas used for the treatment of metals in industry can be produced from propane in the presence of air or Oz over a supported metal catalyst with high selectivity towards CO+He formation. This process is performed at temperatures higher than 800~ There are several different reactions for high temperature propane oxidation in the presence of air [1]. The main reaction which occurs during propane oxidation with Oz (12.3% C3Hs in air) is the reaction of carbon monoxide and hydrogen formation. The theoretical

*The work was partly supported by the Serbian Ministry of Sciences and Technology.

1146

conditions for the formation of CO and 1-12from C3Hs are shown by equation: C3Hs + 3/2 02 ~ 3CO + 41-12 Aft=-227 kJ/mol

(1)

Propane can react with 02 to form complete combustion products (4% C3Hs in air): C3Hs + 502 ~ 3CO2 + 4H20

Aft=-2043 kJ/mol

(2)

Finally, C3H8 can crack to form CI-I4 and coke: C3Hs ~ 2CH4 + C

A//'=-47 kJ/mol

(3)

It is generally accepted that the overall process folows the reverse water-gas shift reaction: CO + H20 r CO2 + H2 A/-/'= -41 kJ/mol (4) As a result of these reactions a mixture of CO+H2+CO2+CI-I4+H20 is obtained which complies with thermodynamic predictions, and tend to effect complete equilibrium among all the components of the product gas. Conversions close to the equilibrium values can be achieved with considerable ease over supported Ni catalysts. To favour propane oxidation according to reaction (1) a selective catalytic material must be used. For practical purposes, nickel is usually impregnated on a suitable porous support which provides thermal stability at working temperatures [2]. But selectivity of a catalyst may depend on various other factors like composition, concentration of active component, physical and structural parameters. The effect of these parameters on the behaviour in propane oxidation of the Ni supported on mullite has been studied in our previous papers [3,4]. With this objective the present work was undertaken to investigate the effects of some parameters of the impregnation process on the selectivity of promoted Ni/MgO catalysts for the reaction of propane oxidation by air to CO and H2. The properties of catalysts relevant for their selectivity such as the Ni-loading in produced catalyst samples, the nickel surface area and mean Ni-crystallite size as well the pore size distribution of catalyst samples are presented. 2. EXPERIMENTAL 2.1. Preparation of catalyst samples Various promoted Ni/MgO catalysts have been prepared by multiple successive impregnation of the MgO support. The support used, which was calcined at 1450~ had the following characteristics: macroporous spherical granules having diameter of about 20 mm, BET surface area of 0.25 m2g 1 and specific pore volume of 0.145 cm3ga. Aqueous solutions of Ni-nitrate and nitrate of the corresponding promoter, with atomic ratio 10:1, respectively, were used. All the impregnation steps were performed at 25~ with the ratio of solution volume to support mass of 3 cmaga, and the solution/support contact time of 30 min. After each impregnation the catalyst precursors were subsequently dried at ll0~ for lh and calcined at 400~ for 2h to convert nitrate salts onto oxides.

1147 2.2. Physico-chemical characterization of samples The content of nickel in catalyst samples was determined by a standard chemical analysis, using dimethylglyoxime. The content of promoters in catalyst samples were determined by atomic absorption spectroscopy (Varian AA 775 series). The specific surface area, SB~r, of the support and of the prepared samples was evaluated by the BET method from the nitrogen adsorption isotherms determined at -196~ in a high vacuum apparatus. The specific pore volume, Vp, and the pore size distribution for the support as well as the catalyst samples were determined by mercury porosimetry (Carlo Erba, Porosimeter Model 2000). Surface area of Ni, SNi, and the mean Ni crystallite size, dNi, of the prepared catalyst samples were calculated from the oxygen chemisorption data, which were obtained by the pulse chromatographic method [5]. All the catalyst samples were previosly reduced by hydrogen at 300~ for 2 h and finished at 450~ for 1 h. The Ni surface area was calculated assuming that the chernisorption of oxygen onto the supported nickel correspond to one on the pure nickel metal. The Ni surface area was calculated assuming a chemisorption stoichiometry O/Ni~un=l and surface nickel atom average area of 0.067 nm 2. The mean Ni crystallite size was derived according to the relation: 5" 103 dNi=~ nm (5) ~Ni " SNi

where YNi is specific density of nickel, g cm 3, and S~ is surface area of nickel, m z gNi-1. 2.3. Catalyst selectivity tests The high temperature propane oxidation by air, catalysed by prepared samples, was studied in the temperature range from 690- 950~ in a flow fixed-bed quartz reactor on-line connected with an analytical system. In all the catalytic run almost equally amount of catalysts of about 40 g, and the catalyst fraction granulated from 2 to 3 mm were used. Propane and air mixture with a volume ratio of 1:7.14, were passed over the catalyst at a gas hourly spave velocity (GHSV) of 300 h a, and at atmospheric pressure. Before the catalytic tests the catalyst samples were carefully reduced in situ with propane and air mixture, at a volume ratio of 1:9.6, respectively. Analysis of C3H8, 02, CO, H2, CO 2 and CH 4 in gas mixture in the inlet as well as in the outlet of the reactor was performed using a Perkin Elmer gas chromatograph (columns: 4 m x 3 mm I.D. 60/80 Poropak Q, and a 2 m x 3 mm I.D. 60/80 Molecular Sieve 5A, both at 150~ A calibration mixture (Messer-Griesheim) was used as the reference in quantitative analysis of the product samples. The water content in the reaction products was determined by on-line connected Prolabo-hygrometer. Mass balance accurate to + 1% were obtained for all the analyses. The conversion of propane was defined as moles of propane converted per mole of propane introduced according to: n ~ - ner X= (6) o ner

1148 In relation (6) n~ denotes the molar flow-rate of propane to the reactor and nw represents the molar flow-rate of propane at the outlet of the reactor. The selectivity for each product is calculated by equation: ui "~ d n i / d t

SF = 100

%

(7)

E Di"1 dni/dt

where ui is stoichiometdc coefficient of product "i", ni is amount of the product 'T', and the quantity dni / dt is called the rate of formation of the product 'T' [6].

3. RESULTS AND DISCUSSION 3.1. Physico-chemical properties Some properties of the prepared magnesia supported nickel catalyst samples are summarized in Table 1. Table 1 List of samples studied Sample Ni loading

Promoter

SBL-r

Vp

SNi

dNi

No.

Designation*

wt%

wt%

m 2 g-1

cm 3 g-1 m e gNi"1

nm

0 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16

M NiAI/M-1-2 NiAI/M-1-3 NiAI/M-1-4 NiAI/M-1-5 NiAI/M-1-6 NiAI/M-2-1 NiAI/M-2-2 NiAI/M-2-3 NiAI/M-2-4 NiAI/M-3-1 NiAI/M-3-2 NiAI/M-3-3 NiAI/M-3-4 NiAI/M NiMg/M NiCa/M

0 1.10 1.62 2.13 2.63 3.13 1.13 2.12 3.04 3.87 1.79 3.19 4.15 4.95 3.31 3.24 3.32

0 0.10 0.14 0.19 0.23 0.27 0.10 0.18 0.26 0.34 0.16 0.28 0.36 0.43 0.29 ** 0.30

0.25 0.6 0.8 1.0 1.2 1.3 1.1 1.2 2.1 2.2 1.5 1.6 1.7 1.9 1.4 1.4 1.3

0.145 0.138 0.134 0.130 0.124 0.116 0.136 0.134 0.128 0.117 0.132 0.126 0.113 0.103 0.130 0.125 0.136

0

0 15.1 12.8 12.6 11.6 11.2 9.0 8.7 8.4 8.1 6.3 5.7 5.5 5.2 15.7 11.1 7.7

37 44 45 48 50 62 65 67 69 89 100 103 107 36 51 73

*Designation of the samples: The first chemical symbol behind Ni indicates the used promoter, the letter M refers the MgO-support; the f'n~t number denotes the nickel concentration in the impregnation solution and second number refers the number of successive impregnation steps. ** no determined.

1149

The Ni-loading for samples listed in Table 1, depending from Ni concentration in the impregnation solution and the impregnation steps used, varies in the range from 1.10 wt% to 4.95 wt%. Comparing data in Table 1 for the catalyst samples obtained by equal impregnation steps, but with various Ni concentration in impregnation solution, exhibit that the Ni-loading increased roghly proportional to the Ni concentration in the solution. The achieved Ni-loading within the pores of MgO-support is in the range from 66 to 54 %, in relation to the theoretical Ni-loading, which is calculated by taking into account the Ni concentration in impregnation solution and assuming the pore volume impregnation mechanism in absence of Ni2*-ions adsorption on the magnesia support surface area. Comparing data in Table 1, it is obvious that increasing the number of impregnation steps raises the Ni-loading in the catalyst samples, reduces the pore volume and increases the BET surface area. Compared with the MgO support, an increase in the BET-surface area and a decrease in the pore volume of the obtained catalyst samples can be explained by the porosity of nickel oxide deposit within the support. From data presented in Table 1, can be observed that the Ni surface area, SNi, and the mean Ni crystallite size are affected by both the Ni concentration in impregnation solution and the number of impregnation steps. Increasing the number of impregnation steps with constant Ni concentration in solution (samples No. 1 to 5; No. 6 to 9 and No. 10 to 13) decreases the Ni surface area and increases the mean Ni crystallite size in the produced catalyst samples. It can be explained by growth of the Ni crystallites as a result of deposit of Niz+ from impregnation solution on the Ni crystallites formed in the previous impregnation steps. Raising the Ni concentration in impregnation solution (see samples obtained by equal impregnation steps in Table 1) decreases the Ni surface area and increases the mean Ni crystallite size in the obtained magnesia supported nickel catalyts. Comparing data in Table 1 at the approximately constant Ni-loading in the produced A1203- promoted Ni/MgO catalysts of 3.11_+0.07 wt% (the samples No 5, 8 and 11) it is obvious that the sample No. 5 produced by 6step impregnation in the solution of 1 mol Ni2* dm 3 has the most developed Ni surface area of 11.2 m 2 gNi"1, and the smallest mean Ni crystallite size of 50 nm, howewer the sample No. 11 obtained by 2-step impregnation in solution of 3 mol Ni 2+ dm 3 have the smallest Ni surface area of 5.7 m 2 gN(,1 and the largest mean Ni crystallite size of 100 nm. The influence of promoter used on the Ni surface area and the mean Ni crystallite size in the promoted Ni/MgO catalyst samples can be observed from the values collected in Table 1 (samples No.14 to 16). These samples were prepared by soaking the magnesia support in the aqueous solutions concentration of 0.75 Mnickel nitrate and 0.075 M nitrates of corresponding promoter by required impregnation steps to provide the desired Ni-loading of about 3 wt%. In the presence of the promoter used, the Ni surface area decreases and the mean Ni crystallite size increases according to the following order: AI203>MgO>CaO. Taking into account that the particle radius of A1203 is 6.3 nm, of MgO is 20.8 nm and of CaO is 65.6 nm [7], this nonreducible promoters effect on the mean Ni crystallite size is in a good agreement with an increasing particle size of the promoter used.

1150

The derivative of the cumulative pore volume curves with respect to pore diameter for MgO support and obtained AlzO3-promoted Ni/MgO catalyst samples are depicted in Figure 1. The magnesia support has a porous structure with prevalent pore diameters in the range of macro pores approx. 8000 nm. The diagrams in Figure 1 for three series of prepared catalyst samples show that the peak amplitude in the pore diameter range of about 8000 nm decreases by increasing the number of impregnation steps at constant Ni concentration in impregnation solution, i.e with the increasing of the achieved Niloading. The pore size distribution curves for catalyst samples show a significant broadening to the pores with smaller diameters. It is interesting to note that the observed changes in the pore structure of catalyst samples are more intense for the samples prepared from impregnation solution with greater Ni z+ concentration than for the ones prepared from the diluted solution.

~

Figure 1. Derivative of the cumulative pore volume curves with respect to pore diameter for the support and catalyst samples. The number of curves corresponds to the number of samples in Table 1.

3.2. Catalyst selectivity The conversion for propane oxidation in air at temperatures in the range studied for the catalyst samples listed in Table 1 was total. CO, H2, CH4, CO2 and H20 were the only detectable products on all catalysts. Results on the study of high-temperature propane oxidation by air, catalysed b y prepared catalyst samples, are demonstrated in Figures 2 and 3. The diagrams in Figure 2 present the selectivity for products formation of main reaction in propane oxidation by air on prepared catalyst samples listed in Table 1. These results demonstrated that the selectivity for the samples prepared by multiple impregnation in 1 M Ni2+ solution increase with increasing of the Ni-loading

1151

up to 3.13 wt%. The selectivity as a function of temperature on the samples obtained by 2-step and 3-step impregnation changes rapidly up to final investigated temperature, whereas the selectivity on the samples obtained by 5-step and 6-step impregnation at the temperature of 790~ reaches about 94%, and at higher temperatures change is small. Comparing the selectivity for the samples obtained by multiple impregnation with 2 M Ni 2+ solution it can be observed that the maximum selectivity is achieved on the catalyst with 3,04 wt% Ni, i.e for the catalyst with the medium Ni-content.

1 O0

7" +

1 O0

Owl

7" +

90

90

B

B B

~, 8o

B" so ti1

~ if)

~

70

~6o

m

8

NiAI/M-2-2 NiAVM-2-3 NiAI/IVI-2-4

-e- , NiA,VM-;1-6 '

700

800

'

Temperature,~

~o

' 1 000

1 O0

90

7§

~,oo

~

900 TemperabJre,~ l

960

I O0

+

70

B 2t B'

/

"

7oo

,

"

.

l

ooo

,

1 000

,

90

8o

I1)

~

oo

70

~ o3

-3 -2

7(1

// -B- NiA~-3-3

N;o~

~

700

NiAaA-3-4 800

900

Temper,al~r'e, oC

9 - o - NiAVM ---m- NiMg~

NiCB/M

Iooo

%0

860 8~o 960 ~o 1ooo Temperature, o C

Figure 2. Dependence of selectivity of CO + H2 formation in propane oxidation on the temperature for the catalyst samples listed in Table 1.

1152 1 00

25,

20

(%1 "r 90 +

8

.B

B

B" ao

~ O3

10

70

~0

--e- NiAVM-I-6 ' 7~0 ' o~0 ' 9~0

Temperatur'e,~

"

20

I

- e - NiAVIVI-1-6 -49- NiAI/IVl-2-3 A NiAI/IVl-3-2

1

I

1ooo

~o

- e - NiAVIVl-1-6 -B- NiAI~-2-3 NiAI/lVl-3-2

03

~,[o. 7ao. o6o. o~o.1 000 Temperat;ure,~

1ooo

25

!

" i 10

mo ~oo Temperab.re,~

7m

20

- e - NiAI/IVI-1-6 -49- NiAVIVl-2-3 --A- NiAI/IVl-3-2

~1o (D (D

~]0

700

800

900

Tempera~re,~

1 000

Figure 3. Dependence of the selectivity of products formation in propane oxidation on temperature for the AlzO3-promoted Ni/MgO catalyst with Ni-loading of ~3 wt% and different mean Ni crystallite size. The samples correspond to those in Table 1.

On the catalyst samples obtained by immersion of magnesia support in solution of highest Ni-concentration, and studied in the propane oxidation, the best selectivity was exhibited on the sample prepared by 2-step immersion and with lowest Ni-loading of 3.19 wt%. On the other two samples, having greater Ni-loading, exhibit lower selectivity at all investigated temperatures. The selectivity towards CO+H2 formation in the propane oxidation on catalyst samples obtained by impregnation with more concentrated solutions was lower than

1153 the one observed on samples prepared from 1 M N i z+ solution, even in the case when the samples had greater Ni-loading. The diagrams in Figure 2 also show the effect of the promoter used on the selectivity of CO+H2 formation in propane oxidation by air. The best selectivity at all temperatures was obtained on the Al203-promoted catalyst. The selectivity for the products formation of the main reaction (1) in the presence of investigated promoter decreases in order to AI203>MgO>CaO, which is in the correlation with the previously mentioned effect of these promoters on the Ni surface area in magnesia supported nickel catalysts. The selectivity of each product formation in propane oxidation by air on the samples with Ni-loading of about 3 wt% is presented in Figure 3. These samples are chosen from each examined catalyst series ( see Figure 2) due to highest selectivity expressed towards products formation of the main reaction. The results show that the principal reaction in propane oxidation by air on these samples is reaction (1) and the other reactions occur only partially. The best selectivity is obtained on the catalyst prepared by 6-step imregnation in the least concentrated Ni2+-solution. From these results, it can be concluded that on this sample, in a rather small extent, all mentioned reactions occur: propane combustion (2), propane cracking (3) and water-gas shift (4). On this catalyst sample reactions (2) and (4) were more favoured than reaction (3), which occurs at temperatures below 900~ only. It is already observed that the selectivity to products formation of main reaction is significantly smaller on the catalyst sample prepared by immersion in solution of the highest Ni 2+ concentration. The obvious differences in selectivity among these three catalyst samples could be correlated with the Ni surface area in these samples (Table 1). Thus the selectivity to formation of products of main reaction is remarkably higher on the smaller Ni crystallites, having the greater surface area, than on the larger ones. 4. CONCLUSION Promoted Ni/MgO catalyst for the propane oxidation by air to produce reducing gas (CO and H2) is obtained by multiple successive impregnation of the low area magnesia support. The Ni-loading in the prepared catalyst is increasing with both concentration of Ni z+ in impregnation solution and the number of impregnation steps. The catalyst sample with the Ni-crystallite having the smaller size is obtained by the multiple impregnation in solution of the lowest Ni concentration. In the presence of the promoter used, the mean Ni crystallite size in the promoted magnesia supported nickel catalyst increases according to the following order: AIzO3> MgO>CaO. The selectivity to the main products formation on the catalyst in the presence of investigated promoter decreases in order to AIzO3> MgO>CaO. The selectivity to formation of main products is connected with Ni-crystallite size, i.e. on the smaller Ni-crystallite the selectivity is higher than on the larger ones.

1154 CO and H2 are formed with greater than 95% selectivity in the propane oxidation at temperatures higher than 800~ on the Al203-promoted catalyst prepared by 6-step impregnation of magnesia support in the 1 M NiZ§ REFERENCES 1. M. Huff, P.M. Tomiainen and L.D. Schmidt, Catal. Today, 21 (1994) 113. 2. W.J.M. Vermeiren, E.Blomsma and P.A. Jacobs, Catal. Today, 13 (1992) 427. 3. N.N. Jovanovi~ and M.V. Stankovi~, Appl. Catal., 30 (1987) 3. 4. N.N. Jovanovi~, M.V. Stankovi~ and G.A. Lomi~, in A.Andreev r al.(editors), Heterogeneous Catalysis, Proceedings 8th International Symposium, Vama, Institute of Catalysis, Bulgarian Academy of Science, Sofia, Bulgaria, 1996, Part 2, 553. 5. N.E. Buyanova, A.P. Kamaukhov L.M. Kefr I.D. Rather and O.N.Chemyavskaya, Kinet. Katal. 8 (1967) 868. 6. R.L. Burwell, Pure Appl. Chem., 46 (1976) 71. 7. V.V.Veselov, T.A. Levanyuk, P.S. Pilipenko and N.T. Meshenko, in Nauchnye osnovy kataliticheskoi konversii uglevodorodov, Akademiya Nauk Ukrainskoi SSR, Naukova dumka, Kiev, 1977, 84 ( in Russian).

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

1155

Structural sensitivity of the oxidation reactions catalyzed by dispersed transition metal oxides: role of defect structure. V. A. Sadykov, S. F. Tikhov, S.V. T sybulya, G. N. Kryukova, S. A. Veniaminov, V. N. Kolomiichuk, N. N. Bulgakov, L. A. Isupova, E. A. Paukshtis, V. I. Zaikovskii, G. N. Kustova, L. B. Burgina. Boreskov Institute of Catalysis SD RAN, Novosibirsk, 630090, pr. Lavrentieva, 5, Russia

1. INTRODUCTION Structural sensitivity of the catalytic reactions is one of the most important problems in heterogeneous catalysis [ 1,2]. It has been rather thoroughly studied for metals, while for oxides, especially for dispersed ones, situation is far less clear due to inherent complexity of studies of their bulk and surface atomic structure. In last years, successful development of such methods as HREM and STM along with the infrared spectroscopy of test molecules has formed a sound bases for elucidating this problem in the case of oxides. In the work presented, the results of the systematic studies of the bulk/surface defect structure of the oxides of copper, iron, cobalt, chromium, manganese as related to structural sensitivity of the reactions of carbon monoxide and hydrocarbons oxidation are considered. 2. EXPERIMENTAL

2.1. Catalyst preparation As starting materials, nitrates, hydroxocarbonates, (oxo)hydroxides, oxalates, ammonium oxalatoferriate and ammonium dichromate of"pure for analysis" or "specially pure" grades were used. Oxides were usually prepared by precursors thermal decomposition at 300-400 ~ with a subsequent annealing in air or in He flow in the range of temperatures up to 1100 ~ Some details of the preparation procedures are given in [3-7]. Only samples which were found to be surface pure by SIMS and XPS were used in studies. To change defect structure of oxides in soft conditions, mechanical activation in the high-powered mill EI-2x150 was applied [8-9].

2.2. Catalysts characterisation Electron microscopy. Samples were examined in JEM-100 CX, JEM-200C and JEM-400 C microscopes. Specimens were deposited onto a carbon film supported on a copper grid [3-5]. XPD patterns were obtained with a URD-6 diffractometer (Germany) using Cu I~ radiation. The Polycrystall program was used to determine the structural parameters [6, 8]. lR-spectra of the lattice modes of oxides were obtained using M-80 spectrometer [8]. Relative densities of extended defects were evaluated by X-Ray Small Angle Scattering method (SAXS) using Cu I ~ radiation with a nickel filter and an amplitude analyzer [4,8]. Mossbauer spectra of 57Fe (including experiments with oxides doped with this ion tracer) were acquired using an NF-640 spectrometer in the temperature range 298-4.2 K [7, 10].

1156

Surface coordinatively unsaturated centers were studied by the infrared spectroscopy of adsorbed test molecules (CO, NO) using IFS-113V Bruker spectrometer [11, 12].

Heats of oxygen adsorption and amounts of a weakly bound oxygen were determined using a high temperature Calvet microcalorimeter, TPD and electrochemical method [ 13]. Catalytic properties in the reactions of carbon monoxide oxidation (all oxides) and butene oxidative dehydrogenation (iron oxides) were studied using a microreactor with the vibrofluidized bed of catalysts and pulse/flow kinetic installation [4]. Catalytic activities were characterized by the reaction rate W (molec. CO/m2s) in differential conditions and first-order rate constant K (dm 3 butene (STP)/m2.s.atm), respectively.

3. RESULTS AND DISCUSSION 3.1. Genesis of defect structure

Types of defects and their relative stabifity. Main types of defects are given in the Table 1. Table 1. Main types of defects found in dispersed transition metal oxides Oxide system

Predominant types of defects

CuO

(001 ) and (100) twins; screw dislocations along , microstrains, misfit dislocation network at CuO/Cu20 interface, grain boundaries. ~-Fe203 Cation vacancies and interstitials; (0001) twins and stacking faults; (1120), (10~2) and (11~3) twins; {0001 } screw dislocations, grain boundaries, surface steps, surface spinel precipitates. Co304 Cation vacancies and interstitials, (111) twins and stacking faults, grain boundaries, microstrains, misfit dislocation network at C0304/C00 interface MnO2 Dislocations and (100) stacking faults; intergrowth of e and 13phases. Fe304 -y-Fe203 Cations vacancies and superstructure; (110) stacking faults and twins CoO Clusters of point defects; (110) twins; surface steps, dislocations, spinel microinclusions, planar---defec--t-S---stabilizedbyi________mpurit:_____!_e. .......................... s: As dependent upon the genesis, three types of systems could be distinguished: 1. Low-temperature systems (calcination temperatures 400-500 ~ where anion residues or excess oxygen remain in the lattice, and such point defects as cation vacancies (up to several % of the regular sites occupancy) are generated due to electroneutrality requirement [5-7]. Among extended defects, twins, bulk or near-surface stacking faults and diclocations confined to the most densely packed planes predominate, suggesting their generation by the precursor phase topotactic transformation. Estimations of the extended defects density from the data of XPD combined with TEM and Mossbauer spectroscopy [3, 4, 6, 7, 10, 14] have shown that it varies from a 101 to a 10.5 per unit cell, reaching a maximum for samples prepared via thermal decomposition of oxohydroxides or nitrates. The density of the bulk planar defects was the lowest when carbonates, oxalates or ammonium dichromate were used as starting compounds [ 15, 16]. Since anion sublattice of such precursors is completely rearranged in the course of decomposition, this fact could be explained by the absence of any topotaxy in this case. The same was true when aging of the amorphous hydroxides occurs in the acid solutions, so that

1157 crystallization of oxide particles proceeds via dissolution -precipitation route. In this case, nearly perfect single crystal platelets of oxides (i.e. hematite) were obtained [ 12]. In addition to topotaxy, relative occurence of the extended defects seems to depend upon microimpurities, which are capable to effect their excess energy and, hence, stability. In such a way, residual hydroxyls are expected to diminish the effective charge of the anion sublattice, thus reducing electrostatic repulsion in the faulted regions and decreasing extended defects energies [8]. Similarly, some extended defects can appear due to oxygen deficiency, and crystallographic shear structures in ruffle and related oxides are well-known examples. For samples of Cr203 prepared by thermal decomposition of Cr (III) nitrate solution, we have revealed a new type of extended defects generated by the excess of oxygen. In this case, high oxidative potential due to nitrate anions decomposition, favors transition of up to 50% of chromium cations into 4+ state. As a result, oxide particles appear which are composed of slabs of ot-Cr203 (structural type of corundum) and CrO2 (structural type of ruffle) with thickness ca 200 A and the most developed surface faces parallel to the (101)R or (100)R planes. These slabs are stacked in rather thick (up to 1500 A) particles with the [211]R axes of symmetry (Fig.la), which have a great number of surface steps.

Fig. 1. A typical image of particle (a), HREM picture (b) and microdiffraction (c) for lowtemperature Cr203 samples prepared from nitrate. According to HREM data (Fig. l b), stacking is not coherent and possibly occurs along various directions favourable for epitaxy of these phases. Moreover, in the direction, a superstructure with period ca 13 A was observed, that is rather close to ot-Cr203 cell length in the direction. Microdiffraction for such particles is very complex due to splitting of reflexes but remains of a single crystal type (Fig. 1c). After air annealing at 500 ~ CrO2 phase disappears without changing particles' shape and thickness. Particle's sizes in the (11~0) direction estimated from analysis of the halfwidths of the (11"~0) and (222[0) XPD peaks remain nearly the same as in the two-phase particles (ca 300 A). It means that (11~0) type stacking faults are formed with densities up to 0.1/unit cell. 2. Middle-temperature systems (up to 900 ~ Usually, increase of calcination temperature leads to annealing of the genetic planar defects or their reconstruction [3,4,7]. Simultaneously, for dispersed oxides, recfistallization proceeds, and extended defects of a new types are generated by sintefing, collapse of fine intraparticle pores or by reversible phase transition [6,7 ]. As a result, in this temperature region the overall density of the extended defects tends to increase (Fig. 2). Stability of extended defects was found to be greatly enhanced by microimpufities segregation in their vicinity revealed by EDX spectral analysis and Mossbauer spectroscopy of the 57Fe ion tracer [6,7,14]. As follows from the analysis of the extended defects structure [ 17 - 19], in their vicinity some cations

1158 are shifted into interstitial positions forbidden in the ideal structure. For Co304 samples with high density of the (111) stacking faults, up to 5% of cobalt cations were found to be transferred from the regular tetrahedrons into neighboring empty octahedra. For defect hematite samples, XPD data imply presence of up to 1-2 % of cations in the interstitial positions. In the IR spectra of the lattice modes of defect oxides, additional absorption bands appear. Thus, for hematite, absorption band at 570 - 580 cm1 typical to Fe in Td emerges [20].

400 O It% t~

O

O OO tt3

200

0 573

1

873

!

1173 T, K

Fig. 2. Integral density of extended defects estimated from SAXS data (1) and relative intensity of I R " defect" band (2) versus annealing temperature for hematite samples.

Fig. 3. A typical image of CoO particle with a surface layer ofCo304.

As follows from Fig. 2, the relative intensity of a such "defect" band indeed correlates with the integral density of bulk extended defects which dominate in the middle-temperature region. For Co304 with high density of defects in the (111) plane, in the 400-500 cm1 region absorbance is observed due to fragments with a local CoO structure, which is predicted for such defects. 3. High-temperature systems (ca 1000-1100 ~ where density of the bulk defects falls (Fig. 2). For some oxide systems, a great number of surface extended defects was detected. Thus, for (z - Fe203 , prismatic faces, which are atomically flat at moderate temperatures of calcination [21 ], in the high-temperature region undergo reconstruction forming steps [6]. Microimpurities segregation at the surface generates also patches with a local superparamagnetic spinel structure [ 10]. A great number of misfit dislocations are generated in particles of oxides after phase transition (Co304 - CoO, CuO-Cu20) [4,7]. 4. Soft ways to change defect structure. Mechanochemical activation was found to be very efficient in changing defect structure [8, 9, 15, 16]. For dispersed oxides, mechanochemical activation was found to generate a great number of point defects thus changing oxides stoichiometry. For such oxides as CuO, Co304, ot - Fe203, interaction of point defects with extended defects was found to "wash out" the latter via climb mechanism [8, 14, 15]. Another efficient route to change defect structure of the surface/near-surface layer is a reversible hydroxylation/dehydroxylation or oxidation/reduction of the oxides at moderate temperatures [22]. As a result, a great number of near-surface extended defects were generated. Fig. 3 demonstrates a typical moir6 picture arising due to oxidation of the surface layer of CoO particle into Co304, which shows a well-developed misfit dislocations network.

3.2. Structure of surface defect centers and bonding strength of reagents. IR spectroscopy of adsorbed CO~NO combined with isotope dilution experiments in the adlayer revealed both isolated and clustered coordinatively unsaturated cations on the surface.

1159 Discrimination between these centers was based on the next points [ 11,12]: 1) Clustered centers have a lower values of Vco and VNo due to a ligand effect; 2) Vibrations of isotopically identical CO (NO) molecules adsorbed on clustered centers are dynamically coupled, hence, clustered centers can be distinguished by the adlayer isotope dilution experiments. Integral coverages less than 5-10 % of monolayer used in our experiments proved that these centers are defects and not regular sites. 3) clustered centers are easily reduced by CO even at room temperatures giving rise to absorption bands typical to Me+/Me~ centers or subcarbonyls. A typical spectral picture illustrating isotope dilution experiment for Cr203 is shown in Fig. 4. 0,7

!

I

I

O-1 W.10

o-2

-16

D-3

B

I

K.10 5

15

30

l0

20 l0

I

2200 2150 Wavenumbers,

I

2100 cm

-I

Fig. 4. IR spectra in the carbonyl region for reduced Cr203. A- 12CO, B-13CO+12CO (7:1) mixture. 298 K.

460

600

860

10'00 T ~

Fig. 5. Rate constants of butene oxidation (1,2) and rates of catalytic CO oxidation (3) versus temperature of hematite samples calcination. Bu:O2 = 1:10 (1) or 1:1 (2).

Table 2 lists band positions for corresponding carbonyl complexes, heats of desorption estimated from the activation energies of desorption, and isotope shift values. In general, results obtained for NO as a test molecule and not shown here for the sake of brevity, correlate well with those obtained by using CO [ 12]. In most cases, for oxides pretreated in oxidizing conditions, no bands which could be assigned to complexes with the regular surface cations (Co 3+, Fe 3+, Cr 3+' Mn 4+, Cu 2+) were observed even at liquid nitrogen temperature [8, 11, 12, 14-16, 22]. It implies that cations in the regular positions of the most developed faces are effectively shielded by the oxygen anions. Isolated coordinatively unsaturated cations with decreased effective charge appear after weak reduction at moderate temperatures and/or vacuum pretreatment. The density of clustered centers was found to correlate with the density of bulk or nearsurface extended defects [8, 12, 22]. Hence, clustered centers can be assigned to surface steps including those at outlets of the bulk extended defects. Density of such centers estimated by using known values of the absorption coefficients is not higher than several percent of monolayer, broadly varying for different samples of the same oxide system [ 12, 14, 22]. Atomic models of surface centers and heats of oxygen adsorption. For main types of the most developed surface faces of oxides studied here, models of their atomic arrangement based upon minimization of the surface energy in the framework of the semiempirical Interacting Bonds Method in a slab approximation were proposed [4, 13, 23].

1160

Table 2. Characteristics of the oxides surface centers by infrared spectroscopy of adsorbed CO Oxide

Carbonyl band position, cm 1 2140 - 2150 2110 - 2120 2170 2120-2150 1980-2080

CuO CoOx

FezO3

Band assignment Cu+-CO isolated Cu + - CO clustered CoZ+-CO isolated Co1+(2+1 -CO clustered Co ~ - CO Fe z+ - CO isolated Fe z+ - CO clustered Fe 1+ - CO clustered Fe ~ - CO Cr 3+ - CO isolated Cr 3+clustered

2190 - 2200 2170 2100-2140 2060 2200 2170

CrzO3

AH of CO desorption, kcal/mol 20-25 15-17 19-20 5-7 25-30

Dynamic shift, cm 1 0 10-20 -

4-5 20 20 > 20 4-5 10-15

0 15-20 0 18

In Table 3, heats of oxygen adsorption calculated by this method are compared with the experimental values. Regular centers on the most developed densely packed planes are mainly covered by tightly bound bridged (Mz O) oxygen forms. On-top (M-O) oxygen forms are located at surface cations on more open planes where steric hindrances for MzO form exist. On-top forms also appear on the densely packed planes at surface steps including those at intersection of the bulk extended defects with the surface. MzO form can be converted into MO form by creating cation vacancy. As can be seen from the Table 3, the most weakly bound MO forms appear at cluster centers in the vicinity of a surface step comprised of surface cations in regular positions and subsurface interstitial cations. Formation of the Me-Me bonds between these cations during oxygen removal ("breathing bonds") decreases energy of the Me-O bond rupture. Estimated heats of oxygen adsorption agree rather good with the experimental values. For the majority of oxide systems, rather low amount of a weakly bound oxygen (usually, not higher than 10% of monolayer) was found. As dependent upon preparation conditions, for samples of the same phase it varies in a wide range, that agrees well with its assignment to defect centers. The situation is quite different for broadly nonstoichiometric Mn304+xand y-FezO3 spinel oxides, which have a rather big (up to 1 monolayer) amount of a moderately bound oxygen. This oxygen is assigned to MO form located at cations having neighboring cation vacancy. Table 3. Calculated and experimental values of the enthalpy of oxygen adsorption (kcal/mol). Oxide

AH calculated Regular centers

........

u/5 . . . . . . . . . . . . . . . . . . . .

Co304

.

.

.

.

.

.

.

.

.

.

.

.

.

.

.

.

.

..................

MzO: 130; MO (Oh): 40-50; MO (Td) 960-70. MO: 40-70; M20:130

Fe203 .

.

.

.

.

.

.

.

.

.

.

.

.

.

.

.

.

.

.

.

AH experimental

Defect centers

............

Centers at outlet of(110) stacking fault: 15.

Centers at outlet of(0001) fault: 18-30

.......

15; 40; 120.

......20; 37;___!__20.

1161

3.3. Structural sensitivity versus flexibility of the oxides surfaces Structure sensitivity. For most systems studied, for the same oxide phase, specific catalytic activity in the reaction of CO oxidation measured in conditions when effect of the reaction media on defect structure could be neglected, broadly varies depending upon genesis (Table 4). For iron oxides, rate constant of the butene oxidative dehydrogenation at 300 C varies from ca 0.1.10 -4 to 30 * 10.4 dm 3 Bu/m 2.s.atm, and catalytic activities in both reactions were found to correlate (Fig. 5). It means that for such oxide systems we have a clear case of structural sensitivity. For two spinel systems with a broad bulk nonstoichiometry - Mn304+x and y-FezO3, pretreated in oxidizing conditions, specific catalytic activity of various samples is rather close. In this case, apparent structural insensitivity clearly correlates with a big amount of a moderately bound oxygen (vide supra). For all other oxide systems, attempts to find any correlation between activity and the density of surface/bulk point defects were unsuccessful. Instead, specific activity was found to correlate with densities of the surface clustered defects. As an example, Fig. 6a shows such trend for the case when clustered defects concentration is characterized by normalized optical density of carbonyl complexes absorption band. Similarly, Fig. 6b illustrates correlation of specific activity with the amount of weakly bound oxygen located on clustered defect centers. In the latter case, correlation splits into three branches corresponding to various types of dominating extended defects for samples annealed in various temperature ranges [6]. For low- and middle-temperature samples, rather good correlations with the density of bulk extended defects were also observed [4, 22, 24]. Hence, structural sensitivity for the systems considered here is explained by a broad variation of the density of active centers located at surface extended defects including those at outlets of bulk defects. Flexibility of the bMk/surface structure and reaction media effect. For such systems as manganese oxides, copper oxides, spinel iron oxides Fe304-y-Fe203 [4, 5, 24, 25 ], reaction media effect at enhanced temperatures (up to 400 ~ ) and at Prolonged (up to 103 h) exposures in reaction mixtures was found to remove all initial differences in the phase composition and defect structure. All extended defects were washed out due to interaction with a flux of point defects created by reaction media. As a result, a constant level of the catalytic activity was achieved for these oxide systems demonstrating apparent structural insensitivity of the reaction of CO oxidation. Hence, in this case, great flexibility of the oxide bulk structure allows to reach the same true steady state of the catalyst. "~

"

lgw 17

/

/

.

9

20'

16 /

~ 10 ~i

15

~ P'S "103 Fig. 6a

I/z ~22

Coverage, %mon. Fig. 6b

Fig. 6. Correlations of catalytic activity of hematite samples with surface normalized intensity of the carbonyl band at 2110 cm 1 (a) and surface coverage by a weakly bound oxygen (b). 1 - low-temperature, 2 -middle -temperature and 3- high-temperature samples, respectively.

1162 In general, restrictions on flexibility are imposed by a low excess energy of extended defects, high lattice energy and a narrow homogeneity range, which hinder restructuring at typical temperatures of catalysis [24]. Such oxides as C o 3 0 4 , ot-Fe203 and ~-CrzO3 meet these criteria, and their steady-state catalytic activities vary nearly in the same limits as the initial ones. However, for these oxides as well as for all other oxides with clustered surface defects, a limited or partial flexibility of the surface layer was revealed [12, 13, 14, 22, 24]. Table 4. A scale of the rate of catalysis (W) variation in CO oxidation on dispersed oxides. Oxide

CuO

Co304

c~-Fe203 Fe304-

.............................................................................................................................................

W initial

MnOz

Mn203

Mn304

ot-Cr203

................................................................................................................................................................

4 "1016+ 5.10 TM + 1.1015 + 3"1016+ 5.1018; 3.1019; 1.1018; 2.1018; 185 ~ 25~ 227 ~ 227 ~

6"1017+ 3'1018; 140 ~

1.5"1017+ 1.4"1018; 140 ~

4'1016+ 2 "10151"1017; 8.5* 1016; 140 ~ 185 ~

W

steady- 4,1016, state 185 ~

2,1016 + 6,1014+ 6.1016+ 1.1017, 5"1017; 3"1017; 9.1016" 227~ 140 ~ ..........

227 ~ ~

.........

~ :

......

2.1017, 6.1016+ 4.1015 + 227~ 1"1017" 9"1016;

227 ~ ~ ~ .

.

227 ~ .

.

.

.__::

.

.

.

.

:::.

.

.

.

185 ~ .

.

.

.

.

.

~

The essence of this phenomenon is that properties of defect clustered centers and kinetic features of the oxidation reactions depend upon stoichiometry of the surface layer. For oxides studied, surface reduction is a topochemical type process and proceeds via spreading of the reduced zone from the extended surface defects accompanied by a cations redistribution between the regular and the interstitial positions. Reoxidation as well as hydroxylation/carbonization causes shrinkage of this zone. Table 5. Effect of oxides pretreatment on the kinetic features of low-temperature CO oxidation. ..............................................................................................

Sample

W1 a

ot-Cr203

w

E2 b

n CO 3 a b

a

b

2* 1014 3'1016

14

5-8

2* 1016 1" 10 TM

10

0

ic6;iS

.........

n 02 4 a b

W(CO) a b

0.5

0.2

0

0

1

2.4

1

0.5

0

0.5

1

100

75 ~ Co304

25 ~ a and b- pretreatment for 1 h at 400 ~ in O2 and He, respectively. For ct-Cr203, before pretreatment in He, 50 % of oxygen monolayer was removed by CO reduction at 300 ~ rate of catalysis (molec. CO/m2s); pulse regime, a steady state after 60 min contact with reaction mixture (1% CO + 1% O2 in He ) at a given temperature. 2 _ activation energy (kcal/mol); 3 and 4 _ reaction orders; 5 _ ratio of the rates of CO2 evolution in pulses of reaction mixture and 1% CO in He, respectively. 1

_

1163 The concept of a partial flexibility allowed to explain why activation energy of CO oxidation catalyzed by oxides of this type was higher when determined at a steady-state of the surface as compared with that found for a constant state of the surface (in pulse experiments) [25]. The main idea is that a temperature increase favors the surface reduction and, hence, increases the number of clustered defect centers. In some cases, variation of the properties of clustered defect centers with their degree of reduction was found to affect mechanism of the catalytic reaction of CO oxidation (transition from the oxidation-reduction scheme to a Langmuir-Hinshelwood (L-H) type). An example of this kind is given in Table 5 for Co304 and (z-Cr203. All typical features of the L-H type mechanism (high activity, low activation energies, fractional reaction orders, higher rate of CO2 evolution in pulses of CO+O2 as compared with that in CO pulses) are observed for reduced samples. 4. CONCLUSIONS Structural sensitivity manifestation for reactions of catalytic oxidation on transition metal oxides depends upon atomic structure of the surface planes, types and densities of the surface/bulk defects and structure flexibility. ACKNOWLEDGMENTS

The research described in this publication was made possible in part by Grants No RPVOOO and RPV300 from the International Science Foundation and the Russian Ministry of Science. REFERENCES

1. M. Boudart. J. Mol. Catal., 30 (1985) 27. 2. G.A. Somorjai. In" Annual Review of Physical Chemistry"', Vol. 45, p. 721. Annual Reviews, Palo Alto, CA, 1994. 3. G.N. Kryukova, V. I. Zaikovskii, V. A. Sadykov, S. F. Tikhov, V.V. Popovskii and N. N. Bulgakov. J. Solid State Chem., 74 (1988) 191. 4. V.A. Sadykov, S. F. Tikhov, G. N. Kryukova, V. V. Popovskii, N. N. Bulgakov and V. N. Kolomiichuk. J. Solid State Chem., 74 (1988) 200. 5. G.N. Kryukova, A. L. Chuvilin and V. A. Sadykov. J. Solid State Chem., 89 (1990) 208. 6. G.N. Kryukova, S. V. Tsybulya, L. P. Solovyeva, V. A. Sadykov, G. S. Litvak and M. P. Andrianova. Materials Sci. Eng. A, 149 (1991) 121. 7. V.I. Kuznetsov, V. A. Sadykov, V. A. Razdobarov and A. G. Klimenko. J. Solid State Chem., 104 (1993) 412 8. V.A. Sadykov, L. A. Isupova, S.V. Tsybulya, S.V. Cherepanova, G. S. Litvak, E. B. Burgina, G. N. Kustova ,V. N. Kolomiichuk, V. P. Ivanov, E. A. Paukshtis, A.V. Golovin and E. G. Avvakumov. J. Solid State Chem., 123 (1996) 191. 9. L.A. Isupova, V. A. Sadykov, I. A. Pauli, O.V. Andryushkova, V. A. Poluboyarov, G. S. Litvak, G. N. Kryukova, E. B. Burgina, L. P. Solovyeva and V. N. Kolomiichuk. In: "Proceedings, Int. Sem. on Mechanochemistry and Mechanical Activation, S. Petersburg, Russia, 1995.

1164

10. V. I. Kuznetsov, V. A. Sadykov, M. T. Protasova and G. S. Litvak. Izv. SO AN SSSR, Set. Khim. Nauk, 2 (1990) 112. 11. Yu. A. Lokhov, M. N. Bredikhin, S. F. Tikhov, V.A. Sadykov and A. G. Zhirnyagin. Mend. Commun., 1 (1992) 10. 12. S. F. Tikhov, V. A. Sadykov, V. A. Razdobarov, and G.N. Kryukova. Mend. Commun., 1 (1994) 69. 13. V. A. Razdobarov, V. A. Sadykov, S. A. Veniaminov, N. N. Bulgakov, V. V. Popovskii, G. N. Kryukova and S. F. Tikhov. React. Kinet. Catal. Lett., 37 (1988) 109. 14. S. F. Tikhov, V. A. Sadykov, G. N. Kryukova, E. A. Paukshtis, V. V. Popovskii, T. G. Starostina, V. F. Anufrienko, V. A. Razdobarov, N. N. Bulgakov and A.V. Kalinkin. J. Catal. 134 (1992) 506 15. L. A. Isupova, V. Yu. Alexandrov, V. V. Popovskii, E. M. Moroz, G. S. Litvak and G. N. Kryukova. Izv. SO AN SSSR, Set. Khim. Nauk, 1 (1989) 39. 16. L. A. Isupova, V. Yu. Aleksandrov, V.V. Popovskii, V. A. Balashov, A. A. Davydov, A. A. Budneva and G. N. Kryukova. React. Kinet. Catal. Lett., 31 (1986) 195. 17. P. Veyssiere, J. Rabier and J. Grilhe. Phys. Stat. Sol. (a), 31 (1975) 605. 18. Ph. R. Kenway. J. Am. Ceram. Soc., 77 (1994) 349. 19. L. A. Bursill, and R. L. Withers. Phil. Mag. A, 40 (1979) 213. 20. G. N. Kustova, E. B. Burgina, V. A. Sadykov and S. G. Poryvayev. Phys. Chem. Minerals, 18(1992),379. 21. G. N. Kryukova, A. L. Chuvilin and V. A. Sadykov. In: Materials Research Soc. Meeting Symp. Series, v. 295, p. 179-182, Materials Research Society, Pittsburgh, PA, 1993 22. V. A. Sadykov, Yu. A. Lokhov, S. F. Tikhov, G. N. Kryukova, M. N. Bredikhin, V. V. Popovskii, N. N. Bulgakov, L. P. Solovyeva, Ii P. Olenkova and A. V. Golovin. In: "Proceedings, 6th Intern. Symp. Heterogen. Catal.", Sofia, Bulgary, 1987, p. 359. 23. N. N. Bulgakov and V. A. Sadykov. React. Kinet. Catal. Lett., 58 (1996) 397. 24. V. A. Sadykov, S. F. Tikhov and V. A. Razdobarov. In: Unsteady-state Processes in Catalysis. Proc. Int. Conf., Novosibirsk, 1990. (Yu.Sh. Matros, Ed.). VSP, Utrecht, The Netherland, p. 407. 25. V. A. Sadykov and S. F. Tikhov. J. Catal., 165 (1997) 279.

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

1165

Oxidation of cyclohexane using polymer bound Ru(III). complexes as catalysts Jacob John, Mahesh K. Dalai and R. N. Ram* Chemistry Department, Faculty of Science, M. S. University ofBaroda, Baroda-390 002, India. Polymer bound Ru (III) complexes were synthesised by sequential attachment of chloromethyl group, 1,2-diaminopropane (DAP) as a ligand and metal chloride to styrene-divinyl benzene copolymer with 8% and 15% cross-linking. Synthesised catalysts were characterised by different techniques such as FTIR, reflectance UV'Vis spectroscopy, SEM, TGA, ESR, NMRand ESCA. Various physico chemical properties such as moisture content, bulk density, surface area and swelling behaviour in different solvents were also studied. The corresponding homogeneous complex [RuDAPCI2] CI was also synthesised. Catalytic activity of these catalysts was tested for oxidation of cyclohexane by varying the temperature of the system as well as concentration of substrate and catalyst. Values of energy of activation and entropy of activation have been evaluated from the kinetic data. A probable reaction mechanism has been proposed. I. INTRODUCTION The oxidation of hydrocarbons is an extremely important commercial reaction for functionalizing hydrocarbons to yield products that are either important in themselves or are intermediate enroute to other chemicals [ 1]. In fine chemicals, because of stringent ecological standards, more emphasis is given to oxidation by molecular oxygen or hydrogen peroxide in preference to non-.environmental friendly metal oxides. The catalytic activity of transition metal ions was reported to be low as such in homogeneous system while an enhanced activity was observed when they were heterogenized by supporting on to a solid support [ 1]. Research for viable polymer supported catalysts for laboratory and industrial oxidation reactions has received recent scientific interest [2-4]. The main problem is the leaching of the metal ion from the surface of the support when it is immobilized by the use of the monodentate or non-chelating ligands. We have reported catalytic activity of different heterogenized chelated metal complex catalysts for various hydrogenation reactions [5-10]. Covalently attached polymer bound multidentate amines could bevaluable starting material to synthesise polymer bound chelates and macrocycles [ 11]. Present study deals with synthesising the polymer bound Ru (III) complexes using 1,2-diaminopropane (DAP) as a ligand and oxidation of cyclohexane under mild reaction conditions using above catalysts. 2. EXPERIMENTAL 2.1 Materials and Equipments Styrene, divinyl-benzene (DVB), dioxane, methanol and cyclohexane were purified according to published methods [12] Styrene-divinylbenzene copolymer (XAD-8) was procured from Fluka

1166 AG, Switzerland. 1,2-dichloroethane and 1,2-diaminopropane were distilled before use. Aluminium chloride was purified by sublimation. RuCI3.3H20 (Lobachemie, Bombay) was used as received. Elemental analyses and TGA were carried out in our laboratory on a Coleman Analyser and a Shimadzu Thermal Analyser DT-30, respectively. The surface area of the support as well as the catalyst was measured using a Carlo Erba Strumentzion 1800. Swelling studies of the catalysts were carried out using polar and nonpolar solvents atconstant temperature. The detailed procedure has been described earlier [ 13]. UV-Vis reflectance spectra of the solid samples were recorded on a shimadzu UV-240 spectrophotometer with reference to nonabsorbing BaSO 4 as a standard and liquid sample in methanol. FTIR and NMR were recorded on a Perkin Elmer R-32 instrument. EPR spectra were scanned on a Breaker ESP 300k band spectrometer using a 100 KHz field modulation on powdered samples at 298 K in N 2 atmosphere. Scanning electron micrographs were recorded on a Jeol SJM T-300. ESCA was recorded on VG model ESCA-3 Mark (II) U.K. with A I K and MgK~ as the radiation sources.

2.2 Synthesis of the Catalysts Styrene-divinyl benzene copolymer with 15% crosslinking was synthesised by the suspension polymerization technique using benzoyl peroxide as an initiator [5]. After polymerization the beads were washed with distilled water, water ethanol (1:1) mixture and ethanol. It was finally, soxhlet extracted with ethanol benzene (1:1) mixture. In a separate series of experiment commercially available styrene divinyl benzene copolymer (XAD-8) was taken as a polymer support. Polymer beads were chloromethylated with HCI, para formaldehyde and acetic anhydride using 1,2 -dichloroethane as a solvent using AICI3 as a catalyst [ 14]. Chloromethylated XAD-8 copolymer was modified by introducing 1,2-diaminopropane.as ligand. The method is reported earlier in detail [5]. The functionalized polymer beads were kept in contact with an ethanolir solution of RuCI3.3H20 (0.4% w/v) for 7 days. There was a change in the colour of the supematent liquid from dark orange to light orange after complexation and beads turned to light grey indicating formation of metal complex on to the polymer matrix. The metal content was determined by refluxing the polymer supported catalyst with cone. HCI (AR) for 24 hrs and then estimating the metal concentration in the solution by spectrophotometfic method after complexation with a nitroso-R salt [ 15]. Synthesised catalysts were named as NPML where N=percent crosslinking, P = styrene-divinyl benzene copolymer, M = Metal (Ru) and L=Ligand (DAP). The following catalysts were prepared. Catalyst B = 15PRu(III)DAP Catalyst F = 8PRu(III)DAP 2.3 Kinetics of Oxidation The kinetics of oxidation of cyclohexane was studied at atmospheric pressure by measuring oxygen uptake using a glass manomatric apparatus. The initial rate was calculated from the slope of the plot of oxygen uptake at various interval of time. The detailed procedure and experimental set up are described earlier [ 12]. The products were analyzed by gas chromatograph .. Oxidation of cyclohexane produced two products, cyclohexanone (1) and cyclohexanol (El). Catalyst B gave 11.2% I and 20.4% Ill whereas catalyst F produced 12.3% I and 19.8% H.

1167 3. RESULTS AND DISCUSSION 3.1 Characterization of Catalysts Physicochemical properties of the supported catalysts are given in Tables 1-3. A decrease in the surface area was observed after loading the metal ions. This is in accordance with the earlier results [5, 9 and 12]. This might be due to blocking of pores of the polymer support after introduction of the ligand and the metal ions. The change in morphology of the polymer support after ligand and metal ion introduction was observed by SEM (Fig. 1) Successful functionalization of the polymer was confirmed by elemental analyses at different stages of preparation of the catalyst. Catalyst F (8% cross linked) was found to have more swellability compared to catalyst B (15% cross-linked). A decrease in swelling was observed as the nature of the solvent was changed from polar to non-polar. Methanol was found to be suitable swelling agent and employed for oxidation reaction because of better swellability with the catalyst and miscibility with the substrate. The UV-Visible reflectance spectra showed d-d transitions at 210 and 500 nmmi.'ght be due to Ru(III). ESCA studies of the polymer bound ruthenium catalysts gave peaks due to Ru (3d 3/2), Ru (3p 3/2) N (1 S), CI (2p 3/2) and C (1 S) for ruthenium - DAP, indicating +3 oxidation state of the metal ion. The unbound complex [Ru DAP C12]CI was characterized by ESR. The gr gH and gov values were found to be 1.897, 2.504 and 2.360 respectively indicating thereby the presence ofgu in low spin +3 oxidation state. The NMR spectrum of the ligand 1,2-DAP showed 2 peaks at 1.00 and 2.42 due to amino and methylene protons respectively. NMR of the unbound complex [Ru DAP C12] CI showed a shiR in peak by 0.2 due to methylene group while amino proton appeared in the same region with multiple splitting which indicates different electronic environment ofligand after complexation. The formation of metal complex on the surface of the polymer was confirmed by FTIR. The various infrared frequencies were assigned as shown in Table 4. TG analysis indicates that thermal stability of the polymer support increases on increasing cross linking of the polymer however no significant change .was observed when metal complex was loaded on the support. Initial weight loss observed might be due to moisture content. It was concluded that catalysts could be used safely upto 100~ (Fig. 2). A probable structure of the catalyst has been proposed based on the spectroscopic data (Scheme 1).

Table 1 Physicalproperties of the supported catalyst Physical property Pore volume (cm 3g-l) Surface Area (m~g-l) Apparent bulk density (g m3) Moisture content (wt%)

Catalyst B 0.398 (0.425) 26.842 (27.932) 0.50 1.28

Catalyst A 0.569 (0.525) 149.60 (160.00) 0.39 0.38

(Values for the respective polymer support are indicated in the parenthesis)

1168 Table 2 Elemental analysis at different stages of preparation of.catalyst B and F x C

87.21 84.28

Y

z

H

CI

C

H

N

C

7.35 7.04

8.9 4.8

84.21 83.33

7.40 6.98

2.59 1.77

83.92 82.88

H

N

7.36 6.77

2.53 1.53

Ru

3.25 x 10.2 1.57 x 10.3

x = after chloromethylation; y = after ligand introduction; z = after complex formation. Table 3 Swelling studies of the supported catalysts Swelling (tool %) Solvent

Catalyst B

Water Methanol Ethanol Dioxane DMF Acetone THF Benzene n-Heptane

Catalyst F

0.692 0.611 0.402 0.198 0.137 0.110 0.098 0.086 0.040

0.962 0.610 0.601 0.497 0.442 0.420 0.411 0.308 0.095

Table 4 IR frequencies (cm -~) Catalyst

Ru-CI

Ru-N

B

225

300

C-N 1072

3407 1598

1170

F

240

314

1069

3437 1633

1272

/CH3 Ci~ CI mm,-R . t t ~ CIj ~NH:Scheme 1

N-H

-CH2CI

1169

Fig. 1 Scanning electron micrographs of (a) P (S-DVB) with 15% cross-linking (b) Catalyst B

1170

3.2 Oxidation reactions The kinetics of oxidation of cyclohexane for polymer bound catalysts B and F were investigated. The stirring of the reaction mixture was maintained at an optimised rate (700 rpm) throughout the experiment to minimise diffusion. The reaction was carried out in a kinetic regime at atmospheric pressure and in the temperature range of 30-45~ The influence of various parameters on the rate of oxidation was studied (Tables 5-7). 3.2.1 Influence of Cyciohexane concentration The effect of substrate concentration on the rate of oxidation was determined in the range of 5.94 x 103 to 23.00 x 10-3 mol 1~ at 35~ and 1 atm. pressure at a constant catalyst concentration of 3.22 x 10.5 mol 1-! of Ru(III) for catalyst B and 1.55 x 10.6 mol 1-~ of Ru (III) for catalyst E It was observed that the rate of oxidation increases linearly with respect to substrate concentration. The order of reaction calculated from the linear plot~; of log (initial rate) vs. log [cyclohexene] was found to be fractional for both catalysts. 3.2.2 Influence of catalyst concentration The effect of catalyst concentration on the rate of oxidation was studied in the range of 1.61 x 10.5 to 6.43 x 10.5 mol 1-! of Ru (III) for catalyst B and 0.776 x 10.6 to 3.100 x 1 0 .6 mol 11 of RU(III) for catalyst F at constant substrate concentration of 11.80 x 10-3 mol 1!, 35~ and 1 atm pressure (Tables 5 and 6). The order of reaction calculated from the plots of log (initial rate) vs. log [Catalyst] was found to be fractional with respect to catalyst concentration for both the catalysts. This may be due to non accessibility of catalytic sites, as well as steric hinderance because of complex nature of the catalyst [5]. 3.2.3 Influence of temperature Catalytic oxidation was studied over a range of 30-45~ at a fixed catalyst concentration of 3.22 x 10.5 mol 1-~ ofRu (III) for catalyst B and 1.55 x 10.6 mol 1! ofRu (III) for catalyst F, at 35~ 1 atm pressure and a substrate concentration of 11.80 x 10.3 mol 1-!. An increase in the rate with temperature was observed. The values for energy of activation calculated from the slope of the plot of log (initial rate) vs. 1/T (Fig. 3) were found to be 5.86 and 9.36 Kcal mol t for catalyst B and F respectively; corresponding entropy of activation was found to be-59.91 and -45.44 eu. 3.3 Oxidation of Cyciohexane using unbound complex [Ru DAP CI2]CI Oxidation of cyclohexane was also studied using homogeneous complex of Ru(IIl) with 1,2 DAP under similar condition (Table 8). However for convenience same quantity of catalyst could not be used as the same quantity of unbound catalyst gave immeasurable.oxygen uptake. Inspite of using larger amount of Ru (III), a lower reaction rate was observed as compared tO polymer supported catalyst. Effect of various parameters such as concentration of substrate and catalyst, temperature, amount and nature of solvent is seen and the results are summerised in Table 8. The energy of activation was found to be 7.04 Kcal moP. 4. RATE EQUATION The reaction mechanism for the oxidation of olefins by metal ions / complexes in homogeneous medium is studied widely and the formation of peroxo and oxo complexes was suggested to be responsible for the transfer of oxygen to the substrate. Vaska et al., have reported formation of peroxo complexes when dioxygen is bound covalently to the metal centre [ 16]. The formation of oxo complex and the transfer of oxygen via this route has been suggested by Taqui Khan et al.; in the oxidation of olefins catalysed by Ru(III) complex in homogeneous medium [ 17]. On the basis

1171 of experimental results as well as evidence from the literature, the following mechanism and rate equation are proposed. /O--O Ru (III) complex +02 ~ Ru(IV) R~uu(IV) m > Ru(V) = O Ru (V) = O + substrate

> product + Ru(III) complex

Keeping the amount of 02 constant, the rate law may be written as 9 R' = k [Catalyst] [Cyclohexane] Thus on increasing the amount of the catalyst as well as the concentration of cyclohexane, an increase in the rate is observed (Tables 5 and 6).

Table 5 Summary of the kinetics of oxidation ofcyclohexane by polymer bound catalyst B in 20ml methanol at 1 atm pressure. [Ru(III)] [Cyclohexane] Temp Rate of Reaction (mol 1l) 105 (mol 1~) 103 (~ (ml min-I) 102 3.22

5.94 11.80 14.80 17.80 23.00

35

2.58 2.66 3.05 3.24 3.65

3.22 4.01 4.83 6.43

11.80

.35

2.66 3.85 4.25 4.64

3.22

11.80

30 35 40 45

2.52 2.66 3.50 3.93

5. CONCLUSION Polymer bound Ru(III)-DAP complexes were found to be stable upto 100~ These catalysts were found to be effective for oxidation of cyclohexane under mild operating conditions. The rate of reaction was studied by varying different parameters and the order of reaction with respect to [catalyst] as well as [substrate] was found to be fractional for both the catalysts. This

1172 Table 6 Summary of the kinetics of oxidation of cyclohexane by polymer bound catalyst F in 20ml methanol at 1 atm pressure [Ru(III)] (mol I-') 10 6

[Cyclohexane] (mol 1") 103

Temp (~

Rate of Reaction (ml min-') 102

1.55

5.94 11.80 14.80 17.80 23.00

35

2.50 2.98 3.72 4.10 4.50

0.776

11.80

35

2.77 2.98 3.78 4.40 4.60

1.55

11.80

30 35 40 45

2.79 2.98 3.56 4.61

might be due to non-availability of active sites as well as steric hinderance. The entropy of activation calculated from the kinetic data was found to be -59.91 and -45.44 eu for catalysts B and F respectively which indicates loss of freedom due to fixation of catalyst molecules on the polymer matrix. The activity of these catalysts for the oxidation.of cyclohexane was observed to be higher than their homogeneous counterparts. A probable reaction mechanism is also proposed.

C

.'xt/1

7I-

__o !

-1.4

v

o

2O 200

Temp.(C)

400

Fig. 2 DTA-TG curves of (P) P(S-DVB) with 15% cross-linking (C) catalyst B

15

3.20

I/Txl6

Fig. 3 Arrhenius plots for catalysts B and F

3.30

1173 Table 7 Summary of the kinetics of oxidation of cyclohexane by homogeneous complex [Ru DAP C12]C1 [Ru] (mol 1!) 103

[Cyclohexane] (mol 1!) 103

Temp. (~

Rate of reaction (ml min-~) 102

1.60

5.94 11.80 14.80 17.80 23.00

35

2.01 3.16 3.34 3.50 3.98

0.643 0.803 0.960 1.280 1.600

11.80

35

1.90 1.99 2.08 2.18 2.34

1.60

11.80

30 35 40 45

2.81 3.38 3.98 4.84

Acknowledgement Authors would like to thank Prof. A.C. Shah, head, Chemistry department and R & D, IPCL for facilities as also to UGC, New Delhi for financial support to one of us (JJ).

REFERENCES 1. 2. 3. 4. 5.

F.R. Hartley, Supported Metal Complexes, Reidel, Dordrecht, 1985. M. M. Miller and D. C. Sherrington, J. Catal., 152 (1995) 368. D. C. Sherrington and H. G. Tang, J. Catal, .142 (1993) 540. W. Derong and S.Licai, Chem. Eng. Sci. (1992) 3673. Jacob John, M. K. Dalai, D. R. Patel and R. N. Ram, J. Macromol. Sc. Pure Appl. Chem. A.34 (3) (1997) 409. 6. M. K. Dalai and R. N. Ram, European Polym. J. 1997, (In Press). 7. D. R. Patel, M. K. Dalai and R. N. Ram, J. Mol. Catal. A : Chemical, 109 (1996) 141. 8. J. John and R. N. Ram, Polym. Intl., 34 (1994) 369. 9. J.N. Shah, D. T. Gokak and R. N. Ram, J. Mol. Catal., 60(1.990) 141, 10.D.T. Gokak, B.V. Kamath and R.N. Ram, React. Polym. 10 (1989) 37. ll.R.S. Drago and J. Gaul, J. Am. Chem. Sot., 102 (3) (1980) 1036. 12.B.S. Furniss, A. J. Hannaford, V. Rogers, P. W. G. Smith and A. R. Tatchell (Eds) Vogel's Textbook of Practical Organic Chemistry, 4th Ed., ELBS and Longmann, London, 1978.

1174 13. D. T. Gokak, B. V. Kamath and R. N. Ram, J. Appl. Polym. Sci., 35 (1988) 1523. 14. J. D. Spivack and S. Vauey, U. S. Pat. 3 281 505 (1966). 15. A. K. Singh, M. Katyal and R. P. Singh, J. Ind. Chem. Soc., 53 (1970) 691. 16. L. Vaska, Acct. Chem. Res., 9 (.1976) 275. 17. M.M. Taqui Khan, Ch. Sreelatha, S.A. Mirza, G. Ramchandraiah and S.H.R. Abdi, Inorganica Chimica Acta, 154 (1988) 103.

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

1175

Photoimmobilized catalysis for low-temperature oxidation of olefins L.V. Lyashenko, V.M. Belousov, E.V. Kashuba L.V.Pisarzhevskij Institute of Physical Chemistry, National Academy of Sciences, Ukraine, 252028, Kyiv, Prospekt Nauki, 31

Heterogeneous metal complex catalysts, synthesized by a new photoimmobilization method, for low-temperature oxidation of olefins C3-C4 to partial oxidation products are proposed. The process selectivity is about 100 %.

1. INTRODUCTION The catalytic oxidation of light olefins on complex oxide catalysts is known to proceed at temperatures above 630 K. Promising catalysts for the low temperature oxidation of hydrocarbons can be heterogeneous metal complexes whose efficiency is largely determined by the method of their preparation. We have developed a new nontraditional method to synthesize heterogeneous metal complex catalysts, which distinguishes by using UV-irradiation directly in the process of complex immobilization on the support[ 1-3]. Here we present the results of studying the oxidation of propene to acrolein and isobutene to methacrolein on metal complex catalysts synthesized through photochemical reactions. We have established that the application of photoimmobilized metal complex catalysts decreases the reaction temperature from 630 K to 300-320 K at selectivity about 100 %.

2. EXPERIMENTAL

2.1 Preparation of catalysts. The catalysts were prepared by the immobilization of metal ions on the support surface during metal photoreduction under UV irradiation from the salt solutions in proton solvents in the presence of benzophenone. The reaction was carried out in a molybdenum glass vessel using mercury lamp under stirring for 2 h. Large pore silica gel with a surface area of 380 mE/g and particle diameter of 0.05 mm was used as a support. After settling the precipitate was washed out by decantation with isopropanol and dried at 365-370 K in flowing argon. The catalysts thus prepared will be referred to as photoimmobilized ones (Me-Ph). For comparison, non-irradiated immobilized catalysts were synthesized under similar conditions (Me-Iml) as well as ones from solution in CC14 (Me-Im2).

1176 We also prepared supported catalysts by the impregnation of silica gel in isopropanol with metal salt, followed by its decomposition under heating (Me-Sup). The heterogeneous metal complex catalysts containing Cu, Mo, W, V, Ti, Ag were examined. 2.2 Experimental technique. Catalytic activity was measured in a flow reactor. Reaction mixture to test catalysts was 1,22 vol.% of hydrocarbon in air. The oxidation products were analysed chromatographically. To gain information on metal surface compounds of both photoimmobilized and immobilized samples their diffuse reflectance spectra in IR, visible and UV region were measured using spectrophotometer "Hitachi-340". ESR spectra of catalysts also were examined using radiospectrometer "Varian-E-9".

3. EXPERIMENTAL RESULTS AND DISCUSSION 3.1. Isobutene oxidation. The data obtained on isobutene oxidation are represented in the Table 1, Figures 1 and 2. To compare the catalytic activities of the samples synthesized by various methods and containing different amounts of metal, we calculated the atomic catalytic activity (ACA), i.e. the rate of methacrolein formation (mol/s) per one g of metal. Table 1. Oxidation of isobutene to methacrolein on metal complex catalysts prepared by various methods Catalyst Temper.of Selectivity and ACA reaction yield of meth106 start,K acrolein, % mol/s g Me Cu-Ph 290 100 25 4.2 Cu-Iml 438 34 21 0.7 Cu-Im2 453 32 11 0.3 Cu-Sup 433 40 12 0.6 W-Ph W-Iml W-Im2 W-Sup

305 310 400 490

100 90 35 0

40 13 1 0

5.2 3.6 0.4

Mo-Ph Mo-Iml Mo-Im2

293 293 330

100 100 80

28 11 4

5.5 1.6 0.1

V-Ph V-Im2 V-Sup

295 346

100 12 74 2 inactive till up 510 K

1.4 0.1

1177

100

20

Q-

2

o

50 ~O r-~ G) ,H .H rH o 0 cO 4-~ r

5

'~ r-I g)

323

373

/.23

2O

/.73 T(K)

523

573

100

0

0,1

!

15 10

50

323

373

/.23

/.73 T(K)

523

573

Figure 1. Temperature dependence of methacrolein (1) and C02 (2) yield on Cu-Ph (a) and Cu-Iml and Cu-Sup (b) catalysts. It can be seen that photoimmobilized catalysts are quiteadvantageous as compared to other ones. A specific feature of isobutene oxidation on photoimmobilized catalysts obtained consists in the fact that this process takes place almost at room temperature. For example, on Cu-Ph reaction starts at 290 K, with increasing temperature the amount of methacrolein formed rises, attains its maximum value at 355-365 K and then at 430 K it decreases. The maximum degree of isobutene conversion is 25-30 %. On Cu-Iml and Cu-Sup isobutene oxidation starts at 420-450 K (Figure 1). Partial oxidation products (methacrolein, acetaldehyde, acetone - -~ 0.05 vol. %) are formed with maximal rate at 520-550 K. Above 570 K isobutene is oxidized to CO2 and H20. The similar results are obtained over Ti-containing catalysts. In the presence of Ti-Ph isobutene is oxidized to methacrolein in the range of 330-420 K and to COz and H20 at 430-520 K. The Ti-Iml, Ti-Im2 and Ti-Sup are inactive at low temperatures, producing CO2 and H20 at the temperatures above 520 K.

1178

A/ t0 3o

mol/s

gMe

t 20

2

10

,3 z/oo T,/4 Figure 2 Temperature dependence of methacrolein formation rate on Mo-Ph (1), Mo-Iml (2) and Mo-Im2 (3) catalysts. The ACA value for Mo-containing photoimmobilized samples (Mo-Ph) is the same order of magnitude as that for the catalysts immobilized from isopropanol (Mo-Iml). However the methacrolein yield relative to the initial isobutene on Mo-Ph is much higher than on Mo-Iml. ACA of the catalysts, synthesized in isopropanol (Mo-Iml) is higher than that for the catalysts prepared from CC14 (Mo-Im2). On Mo-Sup samples isobutene undergoes no oxidation at temperatures of up to 490 K, starting from 510 K a cracking process takes place. The oxidation of isobutene on W- and V- containing catalysts proceeds in the same way as on Mo-containing samples. Another peculiarity of isobutene oxidation on Cu, Mo, W, V, Ti catalysts is the high selectivity towards methacrolein formation (100 % selectivity is observed at 300-390 K). In the region of maximal yield ofmethacrolein there is no even trace amounts of CO2. In some cases (Mo-Ph, W-Ph) acetone and the cracking products are observed at temperatures above 390 K.

3.2. Propene oxidation to acrolein The oxidation of propene on photoimmobilized Cu-Ph, Mo-Ph and W-Ph catalysts occurs at a measurable rate already at room temperatures. The selectivity to the partial oxidation product, acrolein, proved to be 100 %. The rate of formation of the reaction products on the above catalysts against temperature is represented in Figure 3. The maximal rate of the acrolein formation is observed on Cu-Ph, Mo-Ph and W-Ph at 345, 350 and 340 K respectively. Above these temperatures the photoimmobilized catalysts lose their activity.

1179

[ -w-.to 400

SO

~ ~, f,

300

,,

,|,,

35O

Figure 3. Temperature dependence of acrolein formation rate on Cu-Ph (1), Mo-Ph (2) and WPh (3) catalysts. Within 310-430 K the rate of propene oxidation on Me-Im2 and Me-Sup is less by order of magnitude as compared to that on Me-Ph. At constant temperature the activity of the photoimmobilized catalysts decreases gradually. Temperature rise from 300 to 360 K leads to a new increase in the rate of partial oxidation production. Time decrease in the activity of these catalysts at constant temperature is described by the equation W=A lnt, which is typical for processes with a variable number of active sites. Decrease in the activity of Me-Ph catalysts after attaining the low-temperature maximum can be ascribed to reoxidation of photoreducted metal ions. 3.3. Other oxidation reactions The supported Ag-containing compounds are known to be typical catalysts for production of the ethene oxide from ethene. It was interesting to examine Ag-Ph catalysts in this reaction and to compare them with Ag-Im2 and Ag-Sup. It was established that ethene oxide, CO and CO2 were formed on Ag-Sup at temperature range of 370-410 K. Such the products were formed at the presence of Ag-Im2, although the reaction took place at more high temperatures (450-500 K). The major product of partial oxidation on Ag-Ph was butene-2,3 oxide. It has been formed begining from 340 K with 100 % selectivity and ethene conversion being of 30-35 % at 408-415 K. The ethene oxide was detected in trace amounts.

1180

/

/ I

/

doo gOo

..

' mJ

r /,7/,5 r / L

9

|

i

Figure 4. The diffuse reflectance spectra of W-containing catalysts 91,1"- W-Ph, 2,2"- W-Im2, 3 -WO3. Mo-and W-containing systems were examined in the reaction of n-butane oxidation. Mo-Ph and W-Ph catalysts were found to be active in this process. There were two temperature ranges for n-butane oxidation - the low temperature one at 310-410 K and the other at the temperature higher than 520 K. In the high temperature range the reaction products were acetaldehyde, CO2 and HEO. At low temperatures only one product of oxidation was formed. By the methods of chromatography, field mass-spectrometry, IR and NMR-13C spectroscopy it was determined this product to be an ethene oxide. The Mo-Iml, Mo-Im2, Mo-Sup were active only at high temperature range, to produce the products of deep oxidation only.

3.4. The nature of the active centers of heterogeneous metal complex catalysts. The diffuse reflectance spectra in visible and UV regions indicate, that support surface contains metal as separate complexes or their associates in all systems studied: the absorption edge for the immobilized samples is shii~ed towards the short-wave region in the spectrum as compared to that of solid oxides (for example, Figure 4). The near IR spectrum of the samples of Me-Im2 exhibits two bands at 1.37 and 1.90 I~ that are ascribed to vibrations of the OH-group in H20 molecule. For the Mo, V, W catalysts immobilized from isopropanol, besides these bands, the spectrum also exhibits 1.42, 1.681.801~ bands attributed to C-H vibrations of the -CH3 and _CH groups and a 1.32 !~ band

1181 ascribed to vibrations of the OH-group in the alcohol molecule. Hence it can be suggested that the active compound in these samples is present in the form of separate complexes, whose coordination sphere involves water and alcohol molecules. Since these complexes are not washed out by solvent, one may suggest that they are strongly bonded to the support surface. We have also examined ESR spectra of the catalysts. These spectra indicate the presence of W 5+ and Mo 5+ions in the systems studied. ESR signal for the samples synthesized from CC14is negligible. The data obtained permit to state that the site of the low-temperature active center is W 5+ and Mo 5+, involved in the WO 3+ and MoO 3+ oxo ions. The intensity of ESR signal correlates qualitatively with the activity of W- and Mo-containing catalysts in the sequence: Me-Ph > Me-Iml > Me-Im2. ESR spectra of V-containing systems show a SFS structure, characteristic ofvanadyl ions in an octahedral environment. Chemical analysis indicated the presence of Cu + ions on Cu-Ph catalysts and nearly complete absence in immobilized and supported samples. Thus we may conclude that the low-temperature catalyst surface contains metal ions (electronic and ESR spectral data) in the form of complexes whose coordination sphere involves water and solvent molecules (IR spectral and elemental quantitative analysis). On the base of spectroscopy studying the composition of surface compounds may be offered as

CH3 Si ~

O ~

(OU)x

C ~

O...Me (n-1)+/

CH3

~

(UO)y

(ROH)z

The correlation between spectra of the solution from which the immobilization was carried out and diffuse reflectance spectra of the catalysts shows first forming of the reduced metal complex under UV-irradiation followed by its immobilization on the silica surface.

4. CONCLUSION By the photoimmobilization of Cu, V, Mo, W, Ti and Ag complex ions on silica gel the systems including the metal ions in a low oxidation degree were synthesized. Their coordination sphere was consisted of hydroxy ions, water and isopropanol molecules. Such systems possess unique catalytic properties. They catalyse the hydrocarbon selective oxidation at sufficiently low temperatures: isobutene at 300-320 K, propene at 310-350 K, n-butane at 330-350 K. The significant selectivity towards partial oxidation products (about 100 %) is observed at the quite high hydrocarbon conversion catalytic ( 30 - 40%). A special feature of photoimmobilized systems is their ability to catalyse the reaction in unexpected direction as compared to immobilized and supported catalysts of the same chemical nature. Two new catalytic reactions, proceeding over the photoimmobilized catalysts at a low temperature, have been discovered. One of them is the n-butane oxidation to ethene oxide over

1182 Mo-Ph and W-Ph catalysts and another one is the formation of butene-2,3 from ethene over Ag-Ph systems. To our mind, such properties of photoimmobilized systems can to be explained by two factors: the first one is the presence of metal ions in a partially reduced state and the second deals with the presence of organic ligands in the surface complexes. Due to the first one the reaction can be performed at a quite low temperature while the second favors the high selectivity of the process. Photoimmobilized catalysts have been recently shown to be active and of a steady state action in hydrogenation and ammonolysis reactions.

REFERENCES

1. E.V.Kashuba, L.V.Lyashenko and V.M.Belousov, Kinetikai kataliz. 30 (1989) 474. 2. V.MBelousov, E.V.Kashuba and L.V.Lyashenko, Ukr.khim.zur. 57 (1991) 287 3. L.V.Lyashenko, V.M.Belousov and E.V.Kashuba, Teor. exper.khim. 33 (1997) in press.

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

1183

Selective oxidation catalysis over heteropoly acid supported on polymer In Kyu Song a, Jong Koog Lee, Gyo Ik Park and Wha Young Lee* aDepartment of Industrial Chemistry, Kangnung National University, Kangnung, Kangwondo 210-702, Korea Department of Chemical Engineering, Seoul National University, Shinlim-dong, Kwanak-ku, Seoul 151-742, Korea* Membrane-like heteropoly acid-blended polymer film catalysts were prepared using a common solvent (or mixed solvents) and they were tested as fixed-bed catalysts for the ethanol conversion reaction in a continuous flow reactor. It was found that heteropoly acid catalyst was finely and uniformly distributed through the polymer matrix. All the film catalysts showed the higher selectivity to acetaldehyde than the bulk solid catalyst. Conversion and selectivity over the film catalysts were also affected by the nature of solvent and polymer. Microporosity of the film catalyst was controlled by the phase separation method. The microporous film catalyst could be regarded as a highly dispersed heteropoly acid catalyst supported on polymer matrix. The film catalysts were characterized by IR, TPD, SEM, EDX, DSC, and ESCA. 1. INTRODUCTION

Heteropoly acids (HPAs) are inorganic acids as well as oxidizing agents [1-3]. They are highly soluble in polar solvents such as water, alcohols, and amines, but some HPAs are insoluble in non-polar chemicals such as benzene and olefins [4-6]. The solubility of HPAs in turn is closely related to their ability to adsorb reactants. Polar substances readily penetrate into the bulk of HPAs to form a pseudo-liquid phase [7], whereas non-polar chemicals are mostly adsorbed on the surface of HPAs [8]. Owing to these characteri.stics, HPAs have been widely investigated and have been used in a commercial process producing methacrylic acid [9-11]. It is well known [12-17] that the acid and redox properties of HPA can be controlled in a systematic way by replacing the protons with metal cations and/or by changing the heteroatom or the framework transition-metal atoms. Novel catalysis of HPA has been also modified by combining HPA with ion exchange resins [18] or conjugated polymers [19]. Another method for the modification of novel catalysis of HPA is to blend HPA with polymer using solvents to form a membrane-like film catalyst [2022]. Dispersity and redox property of HPA can be easily controlled by this method to meet the need for low temperature oxidation reactions. Another advantage of the film catalyst is that porosity can be also controlled by the membrane preparation technique [23] as described in this work.

1184

Membrane-like HPA-blended polymer film catalysts were prepared using organic solvents and they were tested as fixed-bed catalysts for the ethanol conversion reaction in a continuous flow reactor. The effect of solvent and polymer on the catalytic activity of the film catalyst was examined. The porosity of the film catalyst was also controlled by the phase separation method.

2. EXPERIMENTAL 2.1. Preparation of the film catalyst H3PMo12040 (PMo, Aldrich) was purified and calcined at 300oc for the precise quantification. Polysulfone (PSF, Udel-1700 from Union Carbide), polyethersulfone (PES, Victrex 5200P from ICI), and polyphenylene oxide (PPO, poly-2,6-dimethyl1,4-phenylene oxide from Aldrich) were used as blending polymers. Dimethylformamide (DMF), and mixture of methanol (M) and chloroform (C) were used as blending solvents. A homogeneous PMo-polymer-solvent solution was prepared at room temperature. This solution was casted on a glass plate with a constant thickness in ambient condition (56% relative humidity) to form a membrane-like film catalyst. The homogeneous solution was also casted under different relative humidity (RH) in order to control the porosity of the film catalyst by the phase separation method [23]. All the dried film catalysts were thermally treated at 150oc before reaction and characterization. Composition and preparion method of the film catalysts will be described in Section 3 in detail. 2.2. Reaction and characterization Vapor-phase ethanol conversion reaction was carried out in a continuous flow reactor. The film catalyst was cut into small pieces (2 mm x 2 mm) and used as a fixed-bed catalyst. All the film catalysts were treated at 150oc for 1 hr by passing air (5 cc/min) before the reaction. Ethanol was preheated for vaporization and fed to the reactor together with air as both a carrier gas and an oxygen source. Net amount of PMo was 50 mg in each run. The products under steady state condition were analyzed with an on-line GC (HP 5890 II). Conversions and selectivities were calculated on the basis of carbon balance. The film catalysts were characterized by IR (Midac Co. M2000), TPD, SEM (Jeol JMS-35), EDX (Philips PV-9900), DSC (TA Instruments TA200), and ESCA (Perkin-Elmer PHI 581). All the film catalysts were thermally stable during the reaction because the reaction was carried out at temperatures below the glass transition temperature (Tg) of corresponding polymers. Reaction conditions will be described in detail for each run. 3. RESULTS AND DISCUSSION 3.1. Preparation and characterization of PMo-PSF-DMF A homogeneous PMo-PSF-DMF solution was successfully obtained by dissolving both PMo and PSF in a common solvent of DMF. The viscous and greenish PMo(4.76 wt%)-PSF(23.81 wt%)-DMF(71.43 wt%) solution was casted on a glass plate with a constant thickness at 56% RH to form a membrane-like film, and subsequently it was dried for 4-5 hrs at the same condition. The thickness of the prepared PMo-PSF-DMF film catalyst was 0.017 mm.

1185

The film catalysts comprising PMo and each polymer material show two different thermal behaviors. Tg of a polymer increases when it forms a physicochemical blending with PMo, and, on the other hand, Tg of some polymers decreases when it forms a physical blending. Fig. 1 shows the thermal behavior of PMo-PSF-DMF and PSF-DMF (PMo-free film). Tg of PSF-DMF and PMo-PSF-DMF were found to be 187oc and 174oc, respectively. This result means that PMo in PMo-PSF-DMF acts as an impurity and that the blending of PMo with PSF is physical. The physical blending was also confirmed from IR measurements. Typical bands of the Keggin structure of PMo in the PMo-PSF-DMF film was not changed. Fig. 2 shows the crosssectional SEM micrograph of PMo-PSF-DMF film. No visible evidence representing PMo was found in the PMo-PSF-DMF film and there was no distinctive difference between PSF-DMF and PMo-PSF-DMF. This indicates that PMo was not recrystallized into the large particles but was finely distributed as fine particles invisible in the SEM in the PMo-PSF-DMF. The uniform distribution of PMo in the PMo-PSF-DMF film was also confirmed by EDX analysis as shown in Fig. 3.

3o

,T

~ .8,7~

(9

*

0

&

100

I

I

200

Temperature (~

I

,

300

Figure 1. Thermal behavior of (a) PMo-PSF-DMF and (b) PSF-DMF. In order to investigate any interaction between PMo and PSF, the oxidation state of molybdenum in bulk PMo and in PMo-PSF-DMF film was measured by ESCA. The spectrum could be fitted with only one doublet corresponding to Mo 3d3/2 and Mo 3d5/2. The binding energies of Mo 3d3/2 and Mo 3d5/2 in both catalysts were 235.3 eV and 232.1 eV, respectively. It was confirmed that there was only one type of molybdenum (VI)in both PMo and PMo-PSF-DMF. A DMF-TPD experiment on bulk PMo revealed that DMF desorption started at 150oc and reached the maximum at 270oc and 337oc. These temperatures are higher than Tg of PMo-PSF-DMF film. This fact suggests that DMF (organic base) is chemically adsorbed on the acid sites of PMo in the PMo-PSF-DMF film and affects the acidic function of the film catalyst. It is summarized that highly dispersed PMo catalyst supported on PSF was obtained by blending these two materials. Although the oxidation state of molybdenum of PMo in the PMo-PSF-DMF was not changed, the acidic catalytic activity of the film catalyst might be affected by DMF that was adsorbed on the acid sites of PMo.

1186

Figure 2. Cross-sectional SEM micrograph of PMo-PSF-DMF (xl,000).

Figure 3. EDX image of PMo-PSF-DMF film by mapping on only molybdenum.

3.2. Catalytic activity of PMo-PSF-DMF

Table 1 shows the catalytic activity of bulk PMo and PMo-PSF-DMF film catalyst at 170oc. Acetaldehyde is formed by oxidation reaction while ethylene and diethylether are formed by acid-catalyzed reaction over PMo catalyst [24]. As shown in Table 1, PMo-PSF-DMF film catalyst shows the higher ethanol conversion than bulk PMo. The PMo-PSF-DMF shows remarkably enhanced yield and selectivity for acetaldehyde, but it shows obviously decreased yield and selectivity for ethylene and diethylether compared to the bulk PMo. The oxidation activity of the film catalyst was about 10 times higher than that of bulk PMo. It was believed that the enhanced oxidation activity of the film catalyst was due to the fine distribution of PMo through PSF matrix whereas the reduction of an acidic activity was due to DMF which was strongly adsorbed on the acid sites of PMo. The DMF effect was also confirmed by the catalytic activity of PMo-DMF in Table 1. PMo-DMF showed the suppressed acidic catalytic activity compared to the bulk PMo. Above results imply that the film

1187

catalyst can be applied to the low temperature oxidation reactions to obtain high yield and selectivity for oxidation product by enhancing PMo dispersion and by suppressing the acid-catalyzed reaction. Table 1 Catalvtic activitv of PMo-PSF-DMF film catalvst at 170oc Catalyst Amount of EtOH converted to each product Ethanol (xl05 mole/g-PMo-hr) & (% selectivity) conversion(%) CH~CHO C;~H4 C2H5OC2H5 Bulk PMo 2.60(16.4) 2.41 (15.2) 12.00(68.4) 2.8 PMo-DMF a) 4.16(64.7) 0.82(12.7) 1.50(22.6) 1.0 pMo-PSF-DMF b) 26.70(80.9) 9.87(3.0) 5.30(16.1) 5.3 W/F=32.43 g-PMo-hr/EtOH-mole, air=5 cc/min, film thickness=0.017 mm, a)PMo recrystallized from DMF, b)PMo(4.76 wt%)-PSF(23.81 wt%)-DMF(71.43 wt%) solution was casted and dried at 56% RH

3.3. Preparation and characterization of PMo-polymer-MC A new method for the preparation of PMo-imbedded polymer film catalyst using mixed solvents was successfully developed toward the modification of novel catalysis of HPA. HPA and polymer can be easily blended if both materials are soluble in a common solvent as in the case of PMo-PSF-DMF. Though HPA and polymer are not soluble in the same solvent, if a solvent dissolving HPA and another solvent dissolving polymer are miscible, HPA and polymer can be blended using the solvent mixture. Methanol was used for PMo and chloroform was used for polymer. A homogeneous PMo(1.22 wt%)-polymer (6.90 wt%)-methanol(4.41 wt%)chloroform(87.47 wt%)solution was casted on a glass plate at 56% RH, and subsequently it was dried for 4-5 hrs at the same condition. The thickness of the prepared film catalyst was 0.017 mm. PSF, PES and PPO were used as blending polymers. PMo-free polymer films were also prepared at the same condition for comparison. The PMo-blended polymer film catalyst was denoted as follow ; for example, PMo-blended PSF film catalyst prepared using methanol (M)-chloroform (C) mixture was denoted as PMo-PSF-MC. Fig. 4 shows the SEM micrograph of bulk PMo and PMo-PPO-MC film catalyst. The bulk PMo was large cluster with diameters of 10-100 llm. All the film catalysts retained greenish color implying the fine distribution of PMo through each polymer matrix. No visible evidence for PMo in PMo-PSF-MC and in PMo-PES-MC was found in the SEM images as in the case of PMo-PSF-DMF. This indicates that PMo is uniformly and finely distributed through PSF and PES matrix. On the other hand, PMo in PMo-PPO-MC exists as agglomerates having diameters of 1 I~m or less, although PMo dispersion is much improved compared to the bulk PMo. Thermal behavior of the films and film catalysts are shown in Fig. 5. Tg of PESMC was 236oc whereas that of PMo-PES-MC was 219oc. Tg of PMo-PSF-MC was not detected from a room temperature to 350oc. However, considering that the physical state of PMo-PSF-MC was changed after the reaction over 170oc and became fragile, Tg of PSF was supposed to be lowered after blending with PMo. The decreased Tg of PSF and PES after blending with PMo means that there is n o

1188

interaction or no chemical bonding between PMo and polymer in the film catalyst, and PMo acts as an impurity for each polymer as in the case of PMo-PSF-DMF. The increased Tg of PPO after blending with PMo suggests that there is a certain interaction between PMo and PPO and that PMo is not an impurity for PPO. Chemical state of PMo-PPO-MC is not clear, but it can be presumed that PMo-PPOMC may show different catalytic activity from PMo-PSF-MC and PMo-PES-MC.

Figure 4. SEM micrograph of (a) bulk PMo (x480) and (b) PMo-PPO-MC (x3,000).

A

,~

(d)

o

m

14.

"

(b).

~

~

236~

185oC

m

r ..r

221~

0

!

I

100

I

I

200

I

,Y'-"-

300

Temperature (~ Figure 5. Thermal behavior of (a) PPO-MC, (b) PMo-PPO-MC, (c) PSF-MC, (d) PMo-PSF-MC, (e) PES-MC, and (f) PMo-PES-MC.

1189

3.4. Catalytic activity of PMo-polymer-MC

Ethanol conversion and product selectivity over the film catalysts are listed in Table 2. All the film catalysts show the higher ethanol conversion than bulk PMo. The enhanced conversion over the film catalyst is believed to be due to enhanced surface area of PMo. The conversion is in the follwoing order; PMo-PSF-MC > PMoPES-MC > PMo-PPO-MC > PMo. PMo-PPO-MC shows the smallest conversion among three film catalysts as expected in SEM images of Fig. 4. This may be partly resulted from partial agglomeration of PMo throughout PPO. Bulk PMo and PMo-MC show similar conversion and selectivity. This means that the mixed solvent has no influence on the catalytic activity of PMo unlike DMF of the PMo-PSF-DMF. It also suggests that the main reason for the enhanced activity of the film catalyst is not the effect of mixed solvent but the enhanced PMo dispersion upon blending. The enhancement of ethanol conversion over the film catalyst contributes to the increase of acetaldehyde yield via oxidation reaction over highly dispersed PMo throughout the polymer support. The selectivity to the oxidation reaction over PMo-PPO-MC is three times or more to other two film catalysts and that to dehydration reaction is 50% or less. Lower surface area of PMo-PPO-MC may be responsible for the lower activity but different selecivity can not be explained by the different surface area. The increased Tg of PPO after blending with PMo suggests that there is some interaction (like chemical bonding) between PMo and PPO. The interaction between two materials can contribute to the inhibition of acidic activity of PMo-PPO-MC. Table 2 Catalytic activity over PMo-polymer-MC at 170oc Selectivity(%) Catalyst Conv e rsion (%) CH3CHO C2H4 C2H~OC~H 5 Bulk PMo 6.9 12.8 8.4 78.8 PMo.MC a) 7.4 10.5 8.6 80.9 PMo-PSF-MCb) 39.5 20.0 16.1 63.9 PMo_PES_MC b) 33.7 9.0 22.4 58.6 p Mo_PPO_MC b) 13.4 59.4 9.8 30..8 W/F=169.1 g-PMo-hr/EtOH-mole, air=5 cc/min, film thickness=0.017 mm, a)PMo recrystallized from methanol-chloroform mixture, b)PMo(1.22 wt%)-polymer (6.90 wt%)-methanol(4.41 wt%)-chloroform(87.47 wt%) solution was casted and dried at 56% RH In order to confirm the individual effect of polymer materials on the catalytic activity, the perm-selectivities of O2/ethanol through the film catalysts were measured at 80oc where the extent of reaction was negligible. As shown in Table 3, the ratio of O2/ethanol in permeation side is smaller than that in feed side. This means that the 02 permeability is smaller than the ethanol permeability and that the permeation rate of 02 is the rate-determining step. The lowest ratio of O2/ethanol through PMo-PES-MC may be responsible for the lowest acetaldehyde selectivity over PMo-PES-MC. O2/ethanol ratio in permeation side and acetaldehyde selcetivity showed the same trend in the following order as shown in Table 2 and Table 3 ; PMo-PPO-MC > PMo-PSF-MC > PMo-PES-MC.

1190

Table 3 Permeability ratio of 02/ethanol through the film catalyst at 80oc Catalyst Pressure (atm) , Permeability ratio of O2/ethanol Feed side Permeation side PMo-PSF-MC a) 0.9 1.04 0.57 PMo-PES-MC a) 0.9 1.04 0.11 PMo-PPO-MC a) 1.1 1.04 0.85 film thickness=0.017 mm, permeation area=17.65 cm 2, a)PMo(1.22 wt%)-polymer (6.90 wt%)-methanol(4.41 wt%)-chloroform(87.47 wt%)solution was casted and dried at 56% RH

3.5. Porosity control of PMo-PSF-DMF Another advantage of PMo-PSF-DMF is that porosity of the film catalyst can be controlled by the membrane preparation technique. The homogeneous PMo(4.76 wt %)-PSF(23.81 wt %)-DMF(71.43 wt%) solution was used for the preparation of microporous film catalyst by the phase separation method. Water vapor was used as a non-solvent for PSF. Phase separation rate was controlled by modulating water vapor concentration (RH). RH might affect DMF evaporation rate and phase separation rate. Fig. 6 shows the cross-sectional SEM micrograph of the film catalysts which were prepared at different condition. PMo-PSF-DMF(56V) was prepared by casting the solution at 56 % RH and by drying it in vacuum. PMo-PSF-DMF(85) was prepared by casting and drying the solution at 85 % RH. PMo-PSF-DMF(56V) had no micropores because all water vapor and DMF were evacuated as shown in Fig. 6. PMo-PSF-DMF(56V) was a non-porous film and it showed no measurable microporous properties. On the other hand, PMo-PSF-DMF(85) had well-developed micropores and honey comb type cells with an average pore diameter of 0.25 l~m.

Figure 6. Cross-sectional SEM micrograph of (a) PMo-PSF-DMF(56V) (x4,000) and (b) PMo-PSF-DMF(85)(x2,000).

1191

Its total pore area was about 25 m2/g. The blending of PMo with PSF is a physical one and PMo-PSF-DMF film can be regarded as a finely distributed PMo catalyst supported on PSF as described in Section 3.1. Most of PMo in the PMo-PSFDMF(85) film catalyst is presumed to exist on/near the surface of pore wall as an encapsulated and physisorbed state after the phase separation process. It is proposed that each of micropore and honey comb type cell acts as a micro-reactor having well distributed PMo on/near the wall. 4.

CONCLUSIONS

Membrane-like HPA-blended polymer film catalysts were prepared using organic solvents (a common solvent or mixed solvents) and they were tested as fixed-bed catalysts for the ethanol conversion reaction in a continuous flow reactor. It was found that PMo was finely and uniformly distributed through the polymer matrix. All the film catalysts showed the higher acetaldehyde yield and selectivity than the bulk solid PMo. Conversion and selectivity over the film catalyst were also affected by the nature of solvent and polymer. A microporous film catalyst was also successfully prepared by the phase separation method with the modulation of RH. The film catalyst could be regarded as a highly dispersed heteropoly acid catalyst supported on polymer. It was concluded that the film catalyst could be applied to the low temperature oxidation reactions to obtain high yield and selectivity for oxidation product by enhancing catalyst dispersion and by suppressing the acid-catalyzed reaction. REFERENCES

1. N. Mizuno and M. Misono, Chem. Lett., (1984) 669. 2. M. Ai, Appi. Catal., 71 (1981) 88. 3. I. V. Kozhevnikov, Catal. Rev.-Sci. Eng., 37 (1995) 311. 4. M. Misono, Mater. Chem. Phys., 17 (1987) 103. 5. J. B. Moffat, J. Mol. Catal., 52 (1989) 169. 6. T. Okuhara, T. Nishimura and M. Misono, Chem. Lett., (1995) 155. 7. M. Misono, K. Sakata, Y. Yoneda and W. Y. Lee, in T. Seiyama and K. Tanabe (Eds.), New Horizons in Catalysis, Proc. 7th Int. Cong. Catal., Tokyo, 30 June4 July 1980 (Stud. Surf. Sci. Catal., Vol 7B), Kodansha-Elsevier, 1980, p. 1047. 8. M. Misono, Catal. Rev. -Sci. Eng., 29 (1987) 269. 9. M. Ai, J. Catal., 116 (1989) 23. 10. N. Mizuno, T. Watanabe and M. Misono, Bull. Chem. Soc. Jpn., 64 (1991) 243. 11. H. Mori, N. Mizuno and M. Misono, J. Catal., 131 (1990) 133. 12. H. C. Kim, S. H. Moon and W. Y. Lee, Chem. Lett., (1992) 1987. 13. M. Ai, Appi. Catal., 4 (1982) 245. 14. N. Mizuno and M. Misono, J. Mol. Catal., 86 (1994) 319. 15. S. S. Hong and J. B. Moffat, Appl. Catal., 109 (1994) 117. 16. C. L. Hill and C. M. McCartha, Coor. Chem. Rev., 143 (1995)407. 17. T. Okuhara, N. Mizuno and M. Misono, Adv. Catal., 41 (1996) 113. 18. K. Nomiya, H. Murasaki and M. Miwa, Polyhedrons, 5 (1986) 1031. 19. A. Pron, Synth. Met., 46 (1992) 277. 20. J. K. Lee, I. K. Song, W. Y. Lee and J. J. Kim, J. Mol. Catal., 104 (1996) 311.

1192

21. 22. 23. 24.

I. K. Song, S. K. Shin and W. Y. Lee, J. Catal., 144 (1993) 348. I. K. Song, J. K. Lee and W. Y. Lee, Appl. Catal., 119 (1994) 107. Y. B. Kim, MS Thesis, Seoul National University, Seoul, Korea (1996). T. Okuhara, A. Kasai, N. Hayakawa, Y. Yoneda and M. Misono, J. Catal., 83 (1983) 121.

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 1997 Elsevier Science B.V.

A s t u d y of V2Os-KaSO4-SiO~. c a t a l y s t s o x i d a t i o n of t o l u e n e to b e n z a l d e h y d e

1193

for

catalytic

vapor-phase

A. O. Rocha Jr a, A. L. Chagas a, L. S. V. S. Sufi~ ", M. F. S. Lopes a and J. A. F. R. Pereira b aEscola Polit~cnica - Universidade Federal da Bahia - Rua Aristides Novis, 2, 20 andar, Federaqio, 40210-630 - Salvador, Bahia, Brasil bFaculdade de Engenharia Quimica - U N I C A M P - C.P. 6066 - 13081-970 Campinas, Silo Paulo, Brasil The vapor phase catalytic oxidation of toluene to benzaldehyde has been studied over V90~-I~SO4-SiO9 catalysts in an isothermal differential reactor. The experiments were carried out at atmospheric pressure, temperatures from 410oC to 470~ and the modified spatial time (W/FTo) ranging from 0 to 180 g cat/mol toluene/h. The experimental tests showed the best performance for the catalyst obtained by co-precipitation. These results may be due to a crystalline phase identified in the process of catalyst characterization. Reaction kinetics was determined using the Mars and van Krevelen model. 1. INTRODUCTION The gas phase partial oxidation of toluene to benzaldehyde is an industrially important reaction due to the fact that benzaldehyde is a common intermediate in a wide variety of chemical reaction processes. Almost haft of benzaldehyde world production is employed in the synthesis of food additives (mainly flavoring). Although the partial oxidation of aromatic hydrocarbons is widely treated in the literature, the oxidation of toluene to benzaldehyde presents few information, thus requiring systematic studies (1). The catalysts which have presented the most suitable characteristics for this oxidation are the metal oxides and metal oxides mixtures of transition elements of the V and VI groups, and the literature reports information related to the formulation, preparation and evaluation of the catalysts (2 - 6), although very few data have been published related to the reaction kinetics. Gunduz and Akpolat (5) present experimental kinetic data of gas phase oxidation of toluene to benzaldehyde over V205 catalysts. Their results are based on the redox model and are restricted to the temperature of 430oC. Also, it is not found in the literature enough data which allow to analyze the activity and behavior of V20~ catalysts based only on their physical characteristics.

1194 2. E X P E R I M E N T A L

2.1. Preparation of catalysts In order to obtain a better understanding of V205 - K2SO4 - Si02 catalysts behavior, in a wide range of operational conditions, a systematic study of two catalysts was carried out. One of them was prepared using the co-precipitation technique and the other obtained through impregnation on a commercial support. They are denominated respectively VK-COPREC catalyst (coprecipitation) and VK-SUPPORT catalyst (impregnation). The VK-SUPPORT was prepared through wet impregnation of the active phase on 20 g of commercial silica (SHELL $980) using an oxalic acid solution of NHaVO3 (7.8940 g) and K2SO4 (6.5919 g). The process of impregnation was carried out in a rotoevaporator at 80oC for 8 hours. Afterwards the impregnated support was dried in an oven at 95 - 100oC during 4 hours. The VK-COPREC catalyst was obtained through co-precipitation using Ray and Mukherjee (6) methodology. Silica gel was initially prepared from the addition of sulfuric acid to a solution of sodium silicate (22.35 g). The prepared gel was washed and dissolved in a solution of potassium hydroxide (20.55 g) to originate potassium silicate, and to this solution diluted H2SO4 was added. The precipitate, after being dried for 10 hours in an oven, was impregnated with a NH4VO~ solution in a roto-evaporator for 30 hours, and then dried out at 95 - 100oC in an oven.

2.2. Experimental Apparatus All the catalysts were tested in an experimental set-up operated at atmospheric pressure, as shown on the flow diagram (figure 1). A stream of air, after being filtered and dried, was divided in two. One of the streams was saturated by bubbling into toluene, and the other, consisting of pure air, was adjusted and mixed to the first one to obtain a stream with the required reactor inlet concentration. The reaction takes place in a stainless steel differential reactor maintained in an oven with controlled temperature. Samples of the exit gas of the reactor were fed to a gas chromatograph with a 10'x 1/8" packed chromatographic column of 10% OV-101 in Chromosorb W type AW-DMCS 80/100, which permitted an excellent separation between toluene and benzaldehyde. The calibration methodology was similar to the experimental procedures using streams of known composition of toluene or benzaldehyde bypassing the reactor. For the benzaldehyde calibration, nitrogen was used instead of air to avoid any possible oxidation of the benzaldehyde to benzoic acid. Catalyst particles size -35 +48 Tyler mesh were used in all tests. Porosity was measured using a mercury porosimeter. A 0.1356 ~m pore mean diameter was determined. The Satterfield and Sherwood (7) methodology was used to verify that reaction occurs without any diffusional limitation (internal or external). The effective diffusivity was estimated from the porosity measurements and binary ~sion coefficient and pore tortuosity published in the literature, leading to an estimated value of 10 1 for the generalized Thiele Modulus based on the reaction rate. The effectiveness factor was then considered as 1.0.

1195 The external diffusional effects were evaluated from the generalized JD factor, being (Pwb - Pws) _= 10-Satin, where Pwb and Pws are the partial pressure of toluene in the gas phase and at the catalyst surface, respectively. Thus, Pws iS about 0.4% less than Pwb, SO it was assumed that PT~ _--Pws. In all the experiments 0.3 g of catalyst were used, and the spatial time varied in a range of 0 to 180 gcat.h/mol. Toluene concentration in the reactor feed stream was 0.5% molar, and the reactor was operated at temperatures ranging from 410~ to 470~ The catalyst was oxidized and activated within the reactor.

T B - Termostatic Bath C - Compressor G C - Gas Cromatograph C S - Silicagel Column C F - Filter T C - Temperature Controller F O - Oven I R - Integrator/Registrator T I - Temperature Indicator M - Manometer R E - Reactor R1, R2 - Rotameter SB - Saturator filled with benzaidehyde ST - Saturator filled with toluene T1 to T 5 - Thermocouple CS

F

Figure 1 - Experimental apparatus

3. R E S U L T S AND D I S C U S S I O N The experimental results for catalyst evaluation show that the VK-SUPPORT (V205 and I~SO4 in a silica support) presents low activity and is extremely unstable, the activity decaying very rapidly with time. On the other hand, the VK-COPREC (V20~ - K2SO4 - SiO2 obtained through co-precipitation) shows to be active, presenting a selectivity of almost 100% at 450~ Also, this catalyst is very stable for long periods of reactor operation. Figure 2 presents the selectivity of toluene to benzaldehyde as a function of the reactor temperature for W / Fwo = 110 g cat.h/tool, for the VK-COPREC catalyst.

1196 100 90 A

o~

80 70 60

o~

50 40

rl)

30 20

W/FTo= 111 gcat h/mol

10 0 410

I

I

I

I

I

420

430

440

450

460

470

Reaction Temperature (~ Figure 2 - Effect of t e m p e r a t u r e on the selectivity

During the experimental runs it was shown that the above selectivity is dependent on the reduction level of the catalyst. These data agree with the published results of Trimm and Irshad (2). It was found that the selectivity of toluene to benzaldehyde increases with reaction time, while the toluene conversion shows to be practically constant. These results indicate that a given ratio of V +~ / V +4, after the process of oxidation must be attained, so that satisfactory values of toluene to benzaldehyde conversion are obtained.

3.1. M o d e l i n g of the r e a c t i o n The reaction kinetics was analyzed using the Mars and van Krevelen (8) redox model. C 7 H 8 + 0 2 -~ C 7 H 6 0 + H 2 0 (1) The model equation which shows a better fitting to the experimental data presents two parameters kl (the constant rate of toluene oxidation) and k2 (the constant rate of catalyst oxidation). The reaction rate is second order dependent on toluene partial pressure (PT) and first order dependent on oxygen partial pressure (Po), leading to the following equation

rT =

kl k2 P~ PO

(2)

kl P~ + k2 PO The model parameters were adjusted through the Marquardt method (9), for each temperature, and from the values of kl and k2 the Arrhenius equation constants were determined. Their values are A1 - 7,29 x l0 s mol/g h atm 2 A2 - 7,65 x 105 mol/g h atm

E.1 - 22950 cal/mol Eal - 16021 cal/mol

1197

where A, and A2 are the pre-exponential factors and Eal and Ea2 the values of the activation energy, for the Arrhenius equation. On figure 3 a comparison between the experimental data of the reaction rate and the calculated values from equation (2) is presented.

2,5oE-o3

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r

r

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0,00E+00 0,00E+00 5,00E-04 l OOE-03 1,50E-03 2,00E-03 2,50E-03

r e x p (mol/g cat.h)

Figure 3 - Comparison between the experimental data and the calculated values.

3.2. Physical characteristics of catalyst The Electron Scanning Microscope analysis of the VK-SUPPORT and VKCOPREC catalysts show a significant difference between the two used catalysts, although both catalysts present similar textures when they are new. For the VK-SUPPORT catalyst it was found that a substantial part of the impregnated phase had been lost in the used catalyst compared to the freshly impregnated catalyst, as may be clearly seen on figures 4 and 5. The VKCOPREC catalyst shows a very different behavior (figures 6 and 7). While the fresh catalyst shows a similar texture to the VK-SUPPORT, the used VKCOPREC catalyst shows the formation of a crystalline phase, with needle shape crystals, suggesting that this new phase composed of a mixed oxide, is responsible for the high selectivity and stability of this catalyst. The X-ray diffraction tests permit to identify various phases that are present in the two catalysts, both fresh and after use. The X-ray diffraction results on figures 8 to 11 show that the VK-COPREC catalyst presents a better definition of a crystalline structure than the VK-SUPPORT catalyst. The diffraction results show also that the silica, in both catalysts, presents an amorphous structure.

1198

9~ ' : ' : ~ ' : - : : ' ~ : - : : ~ : ' ' . . : + ' . w '.',.," '^ .~.~ -' ', ' ~ .~.[~.~f~. ~.".~'~. "..-,~.'4:/.::":-:.:.74.:-:.:.~:.:~:..-:.:~.'.:.:+:.~.u~,.:.r " .: ~ ,. .~:..~.~.+: +:.:.: :. :.:.:+.::...:.: : :~4:+:~+. :...::.:.~.~.~....q.:.:~;,.-~::

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The K2SO~ phase is identified in the VK-SUPPORT catalyst, while in the VKCOPREC catalyst peaks relative to the I~Na(SO4)2 phase are present. This result is coherent with the fact that the silica gel is obtained through co-precipitation from sodium silicate, and even with an efficient washing it is practically impossible to eliminate all the sodium (10). The X-ray diffraction tests also show the presence of the K4V207 phase in the impregnated catalyst. In the coprecipitate catalyst the phases where vanadium is present do not show good definition. It seems that various phases are present both in the form of vanadium oxides as in the form of potassium vanadates. In the used coprecipitate catalyst, as shown on figure 11, a well defined peak, corresponding to a value of 9.26 in 20, is probably due to the sodium vanadate lines. The results show that the high selectivity and stability of the VK-COPREC catalyst may be due to the presence of phases containing sodium vanadates and sulfates.

1199

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1500

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2o Figure 11 - X-ray diffraction p a t t e r n s of the used coprecipitate catalyst.

1200 The thermo-gravimetry analysis (TGA), figures 12 and 13, indicate that the VK-SUPPORT catalyst presents a marked peak which corresponds to the loss of water through hydroxyls at temperatures ranging from 160~ to 300~ while for the VK-COPREC a uniform distribution of terminal hydroxyls is detected, indicating a better dispersion of the active phase in the same range of temperature.

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1201 4. CONCLUSIONS The results shown in the present work permit to conclude that the V~OsK2SO4-SiO~ catalyst obtained through coprecipitation shows a good activity and selectivity for the catalytic oxidation of toluene to benzaldehyde, and that the conversion to benzaldehyde depends on the V§ § ratio. The results also show that the proposed redox model fits with accuracy to the experimental data. The results of scanning electron microscopy, the X-ray diffraction and the thermogravimetry (TGA) analysis suggest that the high selectivity and stability of the VK-COPREC catalyst are due to the needle shape crystalline phase formed after the catalyst activation and experimental runs are carried out. This catalyst also shows a better dispersion of the active phase than the supported catalyst. REFERENCES

1. Haber, New Developments in Selective Oxidation by Heterogeneous Catalysis, Amsterdam, Elsevier, 1992. 2. D.L. Trimm and M. Irshad, Journal of Catalysis, 18 (1970) 142. 3. K. Papadatos and K.A. Shelstad, Journal of Catalysis, 28 (1973) 116. 4. A J., Van Hengstum, J.G. Van Ommen, H. Bosch and P.J. Gellings, Applied Catalysis, 8 (1983) 369. 5. G. Giinduz and O. Akpolat, Ind. Eng. Chem. Res., 29 (1990) 45. 6. S.K. Ray and P.N. Mukherjee, Indian Journal of Technology, 21 (1983) 137. 7. Satterfield and T.K. Sherwood, The Role of Diffusion in Catalysis, AddisonWesley Publishing Company, Inc., 1963. Mars and D.W. van Krevelen, Chem. Eng. Sci. Suppl.,3 (1954) 41. 9. Marquardt, Soc. Ind. Appl. Math. J., 11 (1963). 10.A.B. Stiles, Catalyst Supports and Supported Catalysts, Butterworth Publishers, United States of America, 1987. o

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3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 1997 Elsevier Science B.V.

1203

O2/H2 oxidation of hydrocarbons on the catalysts prepared, from Pd(ll) complexes with heteropolytungstates N.I.Kuznetsova, L.I.Kuznetsova, S.V.Koscheev, E.B.Burgina

L.G.Detusheva,

V.A.Likholobov,

M.A.Fedotov,

Boreskov Institute of Catalysis, Novosibirsk 630090, Russia Palladium(II) complexes with heteropolyanions PWllO397"and PW90349-, and originally synthesized bimetallic Pd(II)-Fe(III) complexes with PW90349"were used for the preparation of SiO2 supported catalysts of hydrocarbons oxidation. The composition of the complexes in aqueous solution was characterized by 31p NMR and IR spectroscopy. The supported samples prepared from Pd(II) complexes with the PWllO397"anion exhibited a considerable activity in a liquid-phase oxidation of benzene and cyclohexane with a gas mixture of O2/H2. H2 pretreatment of the catalysts gave rise to increasing the yield of oxigen containing organic products. It was witnessed by XPS and IR studies that heteropolytungstate principally retained its structure and a part of Pd(II) ions became reduced to Pd(0) after supporting and H2 treating the samples at a temperature of 100~ Decomposition of the heteropolytungstate proceeded at a temperature of 450~ resulting in a loss of catalytic properties of the sample. The samples prepared from Pd(lI) complexes with another heteropolytungstate PW90349" anion showed poor catalytic activity in oxidation of hydrocarbons. By contrast, bimetal Pd(II)Fe(III) complexes with the same anion gave active catalysts after supporting and H2 treatment. The specific interaction of palladium species with PWllO397"or Fe(III) in the complexes with heteropolytungstates determine the catalytic properties of the supported samples.

1. Introduction

Hydrocarbon oxidations on platinum metal catalysts readily proceed in a liquid phase in the presence of co-reductant, most simply, dihydrogen. Under the action of co-reductant the active oxygen species are generated from dioxygen. The properties of active oxygen are close to the properties of hydroperoxides, therefore, the two type systems operate in a similar way. With O2/H2 mixture the oxidative transformation of aromatics [1-4], cyclohexene [5-6] and saturated hydrocarbons [7-8] have been studied. It should be noted that some of the systems are effective enough to generate interest of industrial chemistry. The complex composition of the systems using O2/H2 oxidant gives evidence that pure metal catalysts do not operate effectively without promoter, which can be halide ions and/or transition metal oxides. Our idea was to bring into contact platinum metal ions with probable promoter by means of synthesis of metal ion containing heteropolytungstates. As revealed in our previous study, palladium complexes with heteropolytungstate anion PWllO397 do catalyze oxidation of benzene with O2/H2 mixture. Water soluble palladium complexes operate in a two-phase liquid medium. This paper describes preparation, characterization and catalytic performance of SiOa supported palladium complexes with heteropolytungstates.

1204

2. Experimental 2.1. Preparation of Pd(II) complexes with heteropolytungstates in aqueous solutions.. Palladium(II) complexes with the PWI10397" anion, PWllPd and PWllPd2, were prepared as previously described [9] through the interaction of Pd(H20)42§ and NaTPWI iO39 in aqueous solution. Resulting solution contained 0.03 M PWllO397" at a Pd(II) to PWIIO397" ratio of 1 and 2. An alternative procedure was started from Pd(OAc)2 (0.15 mmol in acetone solution) that was mixed with equimolar amount of NaTPWIIO39 in a water. Acetone was evaporated with air flow at room temperature and an aqueous solution was added by dilute HC104 to pH 2 and PWllO397" concentration of 0.03 M. The resulting compound was denoted as PWlIPd(OAe). Palladium(II) complexes with the PW90349 anion, (PWg)2Pd3 and PWgPd3, were prepared from Pd(H20)42+ in sulfuric acid solution, that was added by Na2CO3 concentrated solution to pH 2, and a solid salt NasHPW903424H20 [ 10] was immediately introduced under stirring. A ratio of reagents Pd(II) to PW90349" was 1.5 or 3. After 1-2 minutes dissolving was completed and pH was readjusted to 3.7. Resulting solution contained 0.02 M PW90349. Palladium(II)-iron(III) complexes with the PW90349- anion, (PWg)2PdFe2, (PWg)2Pd2Fe and PWgPd2Fe, were prepared through mixing stoichiometric amounts of Fe2(SO4)3 in a water with Pd(H20)42+ in sulfuric acid. The solution was added by Na2CO3 to pH 2 and a solid salt NasHPW903424H20 was immediately introduced. A ratio of reagents Pd(II) to PW90349" was 0.5, 1 or 2. After 1-2 minutes dissolving was completed and pH was readjusted to 3.7. Resulting solution contained 0.02 M PW90349". 3~p NMR characteristics of the aqueous solutions prepared are given in Table 1. Complexes with heteropolytungstate anions were precipitated through adding 10-fold excess of (C4H9)4NNO3 or CsNO3. From elemental analysis of appropriate tetrabutylammonium salts, the atomic ratio of P:W:Pd:Fe were calculated. Table 1.31p NMR characteristics of complexes in aqueous solution (chemical shift 8 and integral intensity of the peak referred to the total 31p (in brackets)) and composition of the corresponding tetrabutylammonium salts. Complex

8, p.p.m. (% referred to E 31]~))

P:W:Pd:Fe

PW1 iPd PWIIPd2 PWllPd (OAc) (PW9)zPd3 PW9Pd3 (P W9)2Pd2Fe (PW9)2PdFe2 PW9PdzFe

-12.8(55); -12.8(60); -12.8(55); - 12.5(90) - 12.5(75) -45.8(30); -37.8(82); -37.0(12);

1.0:11.3:1.1 1.0:10.6:1.9 no data 2.0:18.4:2.3 1.0:10:2.7 1.0:10.1:0.8:0.5 1.0:8.1:0.4:0.9 1.0:8.0:1.3:0.9

-13.2(45) -13.2(40) -13.2(45)

-66.2(55) -66.0(12) -45.5(18); -63.9(38); -78.8(22)

1205

2.2. Synthesis of Si02 supported samples. Solid samples were prepared through the impregnation of SiO2 (powder with surface area of 200 m 2 g-l) with aqueous solutions of Pd(II) or Pd(II)-Fe(III) complexes with heteropolytungstate anions. Moisture was evaporated to dryness with air flow at room temperature; whereupon the samples were dried at 100~ for 1 hour. Reference samples Pd(OAc)2/SiO2 and Pd(OAc)z/Na7PWllO39/SiO2 were prepared by means of standard impregnation procedure. Acetone solution of palladium acetate and aqueous solution of heteropolytungstate salt were used. After evaporation of the solvent at room temperature, the samples were dried at 100~ The second sample was prepared by means of consequent impregnating and drying procedures. When used, the calcination was performed in air at 400 ~ or 450~ -. Hydrogen treatment of the samples was made in a gas flow (Hz:N2=l:l) at a temperature from room to 200~ for 1 hour.

2.3. Characterization. 31p NMR spectra were registered on a Bruker MSL-400 spectrometer at room temperature. The resonance frequency of 161.98 MHz and accumulation frequency of 0.05 Hz were used, chemical shifts (8) were measured with respect to an external aqueous solution of 85% HaPO4. 0.1 M solution of Na7PWllOa9 has been used to estimate the concentration of heteropolytungstate. IR spectra were recorded in KBr on Specord IR-75 (TBA-salts) or on BOMEM spectrometer (supported samples). CO adsorption measurements were performed at ambient temperature in a pulls mode with hydrogen as a carrier gas. The samples were pretreated in hydrogen flow at 100~ X-ray photoelectron spectroscopy (XPS) data were registered on ES-2401 instrument using low power (100 wt) Mg I ~ radiation source. The spectra were obtained at room temperature, the binding energy was corrected against one of Si 2p (103.5 eV) as a standard.

2. 4. Catalysts testing. Catalytic reactions were carried out in a glass thermostated flask. To observe gas consumption, the flask was connected with a gas burette. The flask was loaded with a portion of catalyst, solvent (CH3CN) and substrate (C6H6 or C6H12), the system was blown with a gas mixture of O2/H2 and sealed. Suspension was subjected to intensive stirring. After 1 hour the reaction was stopped and the organic products were analyzed by GC (Tswet-500 instrument with FID and a column of 0.4% 1nitroanilineantraquinone on carbon black) and GC-MS (LKB model 2091). 3. Results.

3.1. Characteristics of Pd(II) and Pd(II)-Fe(III) complexes with heteropolytungstates. The composition of aqueous solutions denoted as PWI1Pd and PWllPd2 (Table 1) was recently studied [9]. When the molar amounts of Pd(II) and PWI10397" being equal, the solution contained two type complexes of identical composition: PWllPd and PWllPd-OPdWzlP. In the case of PWI~Pd2 solution, overstoichiometric palladium formed bi- and

1206 polynuclear fragments bonded to heteropolytungstate. Both solutions included appreciable amounts of Na2SO4. The compound PWllPd(OAc) was obtained in a solution free of sulfate ions. From alp NMR data the complex was very similar to PWI1Pd, and heteropolytungstate anions were involved into binding with Pd(II). Heterpolytungstate PW90349" anion possesses three vacancies in the structure, being capable to form compounds of the type (~-1,2,3-PW9V30406 [ 11]. Two-charged M ions tended to give complexes of a composition (PW9Oa4)2M3, where M may be Pd(II) ions [12]. To follow this direction, in this study two complexes of the different average composition were prepared under variation of a ratio of reagents: Pd(H20)42+ and NasHPW9034. Both solutions exhibited one of the same principal signal in alp NMR 6=-12.5 ppm, that was different from 6--12.8 ppm for PWIIPd. IR spectra of both PW9Pd3 and (PW9)EPd3 TBA-salts comprised the typical frequency bands of Keggin-type metal-substituted heterpolytungstates (Fig.l): Vas P04 about 1070 cm "1, Vas W--O 950 cm "1, Vas W-O-W(M) about 880-700 cm -] [13]. The distinction of the composition of PW9Pd3 (or 1 (PW9)2Pd3) from PWI]Pd TBA-salts suggested stability of the parent PW9 fragment. When Fe(III) being added concurrently with Pd(II), the formation of bimetal Pd(II)Fe(III) complexes with PW90349 was observed. The bimetal complexes provided several broad 31p NMR peaks (Table 1) distinguished from the signals of both Pd(II) and Fe(III) complexes with PW90349.(Fe(II) complex with PW90349" prepared according to the known procedure [13] and oxidized with dioxygen to Fe(III), gave alp NMR signal at 6=286 ppm) Broadening and considerable negative chemical shift of 31p signals reflected cp~ the specific interaction of heteropolyanion with paramagnetic Fe(III) ions. No changes of 1200 1000 800 600 principal IR bands of TBA-salt were produced Wavenumber, cm-I by Fe(III) ions (Fig.l). Three solutions were m

i

i

i

i

- - - i

....

~ -

i

i

Fig.1. IR spectra of TBA-salts of the obtained which contained Pd(II) and Fe(III) PWqPd3 (1) and PWqPdEFe (2) complexes, ions totally involved into bimetal complex with PW90349 anions. The average composition of the complexes is given in Table 1.

3.2. Characteristics of supported samples. Because of a strong absorption of SiO2 and Na2SO4 in the IR region around 1100 cm "1, the characteristic frequencies of SiO2 supported heteropolytungstates in a region of 1000-700 cm "1 were considered. The principal IR bands of the supported and dried at 100~ samples were identical to those of the solid Cs-salt of PWI]Pd (compare spectra 1 and 2 in Fig.2). Negligible transformation of the spectrum occurred after the low temperature hydrogen

1207 treatment of P W I I P d / S i O 2 that indicated the parent heteropolytungstate anion to keep structure after low temperature reduction of the sample. The hydrogen treatment at the temperature of 100~ gave rise to appearing a small shoulder with v 978 c m "1 (spectruna 3) that could be explained by the reversible transformation of the heteropolytungstate anion [ 14], and was developed with increasing the temperature of treatment. Calcination of the samples in air brought about the transformation of the parent PWllO397 anion to PW120403 at 400~ (v 984 cm -1 on spectrum 4) and the complete destruction of heteropolytungstate to WO3 at 450~ (broad band around 850 c m -1 o n spectrum 5). Issuing from IR spectra, the conditions of stability of the supported heteropolytungstate were consistent with the data of XPS. As shown in Fig.3, spectra 1-2 for W 4f, the electronic state of tungsten corresponding to W 6+ [ 15], was kept after the hydrogen treatment. However, sintering of the sample prior to reduction resulted in some broadening W 6+ signal without distinct splitting the 4f7/2and 4t"5/2 lines (compare spectra 2 and 3), that indicated the destruction of the regular structure of anion. XPS spectra gave also interesting information on the palladium species on the surface. The peak at E=336.5 eV in the spectrum of the sample before reduction (1 for Pd 3d in Fig.3) was attributed to Pd 2+ [16]. The appearance of the Pd ~ species after reduction was evidenced by shifting this peak to 334.8 eV and by admixing the signal of reduced Pd species at about 340 eV (spectrum 2). The similar spectrum belonged to the F., ~ -7. sample passed sintering prior to the hydrogen treatment (spectrum 3). The clearly appeared in spectrum 3 band at 331 eV was ascribed to palladium hydride. Distinctive features of spectrum 3 was more symmetrical shape of the peaks at 334.8 eV and around 340 eV that is explained by lower contribution of oxidized state. Thus, the hydrogen treatment of the supported samples at 100~ resulted in reduction of a part of Pd 2+ to Pd ~which proceeded more completely after sintering. A part of accessible palladium species on the surface of the Pd containing supported samples was estimated by means of CO adsorption measurements; the data are depicted in Fig.4. ~ooo 9b0 a6o 700 Wavenumber, cm-1 So high dispersity of palladium (if the conventional stoichiometry of absorption of CO:Pd=I is accepted) is not surprising for the reference samples Fig.2. IR spectra of Cs+-salt of prepared from acetone solution, but is unusual for P W l l P d (1); untreated PW11Pd(OAc)/SiO2, because a water solution was used in PWllPd/SiO2 (2); PWllPd/SiO2 this case. This is, probably, resulted from the interaction after H2 treatment at 100~ (3) of Pd(II) with the heteropolytungstate in a parent and after sintering at 400~ (4), solution. All the other samples containing NaESO4 w e r e 450~ (5). t... o rs~

1208 characterized by smaller fraction of accessible P d , probably, because of blocking effect of inert salt. W 4f 60

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,

,

,

,

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,

,

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,

,

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,

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,

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,

,

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,

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,

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.

.

.

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.

.

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~

.

.

.

.

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45

cBa3H, 3B

,

320

330 3Hepm~[

340 CBII3H,

350 3B

Fig.3. XPS spectra for W 4fand Pd 3d of PWllPd/SiO2" untreated (1), H2 treated at 100~ (2), sintered at 450~ and H2 treated at 100~ (3).

3.3. Catalysis. The catalytic properties of the supported samples were tested in oxidation of cyclohexane and benzene with a mixture of O2/H2 gases at a temperature of 20-40~ . Cyclohexanol and cyclohexanone were obtained from cyclohexane and phenol with admixtures of cyclohexanol and hydroquinone (no more than 2% mol. of each) was obtained from benzene. A suspension of unreduced samples in CH3CN consumed the O2/H2 mixture. Upon reaction with O2/H2 gases and hydrocarbon substrates the samples were subjected to reduction reflected by a change of a color from yellow to gray. The yield of organic products is depicted in Fig.4 (upper diagram). The samples containing Pd(II) complexes with PWllO397" showed substantial catalytic activity, whereupon, PWIIPd/SiO2 producing higher yield of organic products than PW11Pd(OAc)/SiO2. The PWIIPd/SiO2 catalyst became inactive after sintering at 450~ All the other samples were inert in unreduced state, except for small activity of one of bimetal Pd(II)-Fe(III) sample.

1209 A

O

E

ElO

II

0

phenol

2 2

3

T-4

5

6

7

8

9

10

11

3

4

5

6

7

8

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11

a_

~

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~

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% Pd COads/Pd

1

2

ta

o_

0

~

EL

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n

I::L

EL

a_

0.54 0.31

~

o

=

"~ EL

1.08 0.54 0.54 0.60 0.54 0.57 0.60 0.38 0.19 0.8 0.23 0.80 0.14 0.78 0.75 0.25 0.27

Fig.4. The yield of the products of hydrocarbons oxidation in the presence of SiO2 supported Pd(II) complexes with heteropolytungstates: (A) before H2 treatment, (B) after H2 treatment at 100~ (CO adsorption was measured after the treatment). Conditions: 0.03g of catalyst suspended in lml of CH3CN, 0.1ml of benzene or cyclohexane, O2/H2=1/2,1 hour. Preliminary gas-phase hydrogen treatment of the samples gave rise to increasing the consumption of gases and, in most cases, the yield of the organic products (Fig.4, down diagram). Reduction of the PWIlPd/SiO2 and PWI1Pd2/SiO2 samples improved their activity in oxidation. Increasing the temperature of reduction from room to 100~ resulted in increasing the yield of organic products, but further increase of the temperature had an adverse effect on the yield (Fig.5). On the contrary, activity of the PW11Pd(OAc)/SiO2 catalyst was lower after than before gas-phase reduction, and lower than activity of the analogous PWllPd/SiO3 sample, even though the former one was characterized by more accessible Pd surface. Reference samples, which did not contained palladium complexes with heteropolytungstate, remained inactive in oxidation even after reduction. Analogous behavior was exhibited by the PWllPd/SiO2 sample sintered at 450~ it has only negligible activity in reduced state. These date suggest the promotion effect of the PWllO397anion, the presence of which is necessary to create the activity of palladium species in oxidation. The specific structure of the heteropolytungstate appeared to be a crucial factor since destruction of the anion resulted

1210 in a considerable loss of activity in oxidation. On the other hand, the reference Pd(OAc)2sample did not catalyze oxidation even though it contained the necessary promoter. The reason is the method of preparation providing no chemical bonding between Pd(II) and heteropolytungstate.

NaPWllO39/SiO2

9 phenol 9 cyclohexanol

,.. 16 14 12

~

" ~

r,

1

_

_ _

anone

22 20

.,

10 8

0

-1

50

1~

150

2~

~c

16 14

20

25

30

35

40

T~

Fig.5. The yield of the products v s temperature of H2 pretreatment of the PWllPd/SiO2 catalyst. Fig.6. The yield of organic products v s temperature of reaction in the presence of P W 1 1 P d / S i O 2 , HE pretreated at 100~ Conditions: catalyst 0.03g, CH3CN lml, benzene or cyclohexane 0.1ml, O2/H2=1/2, 1 hour. Another complex with the PWI10397" anion containing acetate anion was available in water solution but appeared unstable in supported state. Destruction upon drying and the following reduction at 100~ made the supported complex less active catalyst than analogous PWIIPd/SiO2 sample. Hydrogen treatment of the samples both in a gas phase and in liquid resulted in reduction of a part of Pd(II) to Pd (0), as detected by XPS. After reduction Pd(0) species retained location near by the heteropolytungstate anion, which can participate the catalytic process in this position. As seen from Fig.6, the most favorable temperature of oxidation catalyzed by reduced PWI1Pd/SiO2 sample, was 30~ At the temperature of 30~ the PWllPd/SiO2 catalyst produced 9 mol/g-at Pd h of phenol and 23 mol/g-at Pd h of cyclohexanol plus cyclohexanone. About 20% of a O2/H2 mixture consumed was spent to oxidation of organics, and the rest formed a water. Two solutions (PW9)EPd3 and PW9Pd3 was shown by 31p NMR to contain one of the same complex. In the second solution the overstoichiometric Pd(II) ions can be connected to heteropolyanion in the form of bi- and polynuclear fragments. Both (PW9)EPd3/SiO2 and PW9Pd3/SiO2 samples possessed low catalytic activity in oxidation even in reduced state, that indicated insufficient promoting ability of the PW90349" heteropolyanion. On the contrary, two of three samples containing bimetal Pd(II)-Fe(III) complexes with the PW90349" anion developed appreciable activity after reduction. The Fe(III) ions rather than heteropolytungstate are responsible for the promotion effect in this case. Bimetal complexes with heteropolytungstate included Pd(II) and Fe(III) ions in different combinations, some of them were advantageous to create the catalytic activity of palladium. The detailed structure of

1211 active in catalysis Pd(II)-Fe(III) complexes is supposed to be the subject of the following study.

4. Conclusion. The results of this study have shown that prepared from Pd(II) complexes with the PWllO397 anion, SiO2 supported samples behave in catalysis of hydrocarbons oxidation in a similar way as their water soluble precursors, but exhibited improved performance because of the more advantageous composition of the catalytic system. In the present conditions, the catalytic activity of palladium species is developed only in the contact with a proper promoter. The effective contact of palladium with the second component is provided when complexes with the heteropolytungstates are applied for the synthesis of the catalysts. The method is appeared powerful in both situations: when heteropolytungstate anion is a promoter by itself and when it serves as a matrix to assemble together Pd(II) with promoting Fe(III) ions. The present study obviously demonstrates the significance of the promotion effect in oxidation with a O2/H2. There is deficient in direct information on the mechanism of this phenomenon. The role of promoter can be speculatively proposed as (i) facilitating oxygen transfer from activator (palladium species) to organic substrate and/or (ii) stabilizing the proper palladium oxidation state. In the latter case oxidized palladium is suggested to participate dioxygen activation. Acknowledgment The study was supported by Russian Foundation for Basic Research, Grant No. 95-0308754a. References 1. A.ltoh, Y.Kuroda, T.Kitano, G.Zhi-Hu, A.Kunai, K.Sasaki, J. Mol. Catal. 69 (1991 ) 215. 2. Jap. Appl. No5-4935, 1993. 3. T.Kitano, Y.Kuroda, M.Mori, S.Ito, K.Sasaki, J. Amer. Chem. Soc. Perkin Prans. 2 (1993) 981. 4. T.Miyake, M.Hamada, Y.Sasaki, M.Oguri, Appl. Catal. 131 (1995) 33. 5. N.I.Kuznetsova, A.S.Lisitsyn, V.A.Likholobov, React. Kinet. Catal. Lett. 38 (1989) 205. 6. N.I.Kuznetsova, A.S.Lisitsyn, A.I.Boronin, V.A.Likholobov, in" B.Delmon and J.T.Yates (Ed.), Stud. Surf. Sci. Catal. Vol.55 (1990) p.89. 7. US Pat. No5.409.876, 1995. 8. S.B.Kim, K.W.Jun, S.B.Kim, K.W.Lee, Chem. Lett. 7 (1995) 535. 9. N.I.Kuznetsova, L.G.Detusheva, L.I.Kuznetsova, M.A.Fedotov, V.A.Likholobov, J. Mol. Catal. A" Chemical (1996) in press. 10 R.Contant, J.-M.Fruchart, J.-P.Ciabrini, M.Fournier, Inorg. Chem., 16 (1977) 2916. 11. P.J.Domaille, J. Amer. Chem. Soc. 106 (1984) 7677. 12. W.H.Knoth, P.J.Domaille, R.L.Harlow, Inorg. Chem. 25 (1986) 1577. 13. C.Rocchiccioli-Deltcheff, R.Thouvenot, R.Franck, Spectrochimica Acta, A 32 (1976) 587. 14. R.I.Maksimovskaya, Kinet. i kataliz (Rus.), 36 (1995) 910. 15. Handbook of x-Ray Photoelectron Spectroscopy. Ed: C.D.Wagner, W.M.Riggs, L.E.Davis et.al. Perkin-Elmer Corporation, Physical Electronics Division, Eden Prairie, Minnesota, 1979. 16. Analysis of surface by Auger and X-ray photoelectron spectroscopy. D.Brigs, P.M.Sikh, Ed., Moskow, Mir, 1987, p.598.

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3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

1213

O x i d a t i o n of n - p e n t a n e to phthalic or maleic a n h y d r i d e : The role of t h e VPO c a t a l y s t s t r u c t u r a l disorder Z. Sobalikl, S. Gonzalez 2, P. Ruiz3 and. B. Delmon3 1 j. Heyrovsky Institute of Physical Chemistry, Academy of Sciences of the Czech Republic, Dolejskova 3, 182 23 Prague 8, Czech Republic 2 On leave from Instituto de Quimica. Universidad de Salamanca, Salamanca, Espana. 3 Unitd de catalyse et chimie des matdriaux divisds, Universitd catholique de Louvain, Place Croix du Sud, 2 bte 17, B-1348, Louvain-la-Neuve, Belgium.

ABSTRACT Vanadium phosphate catalysts are obtained from precursors prepared by a two-step synthesis. In the first step, a VOPO4-mixed isobutanol-water intercalate was obtained by precipitation from a solution containing vanadyl isobutoxide, H3P04 and carefi~y adjusted water content (precursor A). In the second step, precursor B was formed by reflux of precursor A in (i) an inert (n-octane) or (ii) reductive (isobutanol) medium. By such a procedure, precursors and catalysts (with PN atomic ratio equal to 1.05) displaying widely different structural defects (XRD, IR) were prepared. Catalysts were tested in the oxidation of n-pentane into maleic (AM) and phthalic (PA) anhydrides. Formation of PA demands a highly ordered structure, while AM could be formed on a highly defective VPO catalyst. INTRODUffHON The oxidation of n-pentane to phthalic anhydride over vanadium phosphate catalyst (VPO) is an interesting reaction that triggers question on the mechanism of selective oxidation reactions. In addition to this fundamental aspect, the reaction might perhaps be of practical usefulness. It would indeed be industrially attractive to replace o-xylene by n-pentane as a feedstock for phthalic anhydride (PA). The corresponding reaction mechanism necessarily involves a very complex set of active sites arranged according the a special surface topology. Such a topology could ensure the exact timing of the consecutive steps, including abstraction of 10 H-atoms, insertion of 3 oxygen atoms and cyclisation of 2 molecules of n-pentane to form PA, and at the same time providing an easy desorption of the final anhydride. Surface properties of a VPO catalyst and the dynamics of surface adsorbed species on the selectivity to PA or maleic anhydrides (MA) and the practical usefulness of this reaction was evaluated by Centi [1]. He showed that, for

1214 economic reasons, a more efficient use of n-pentane is necessary, namely an increase of selectivity of the process to PA must be achieved for future industrial applications. At the present stage of research, parameters controlling the PA/MA formation ratio have only been partially identified and much progress in this field is necessary for future developments. The objective of this paper is to study the influence of the conditions during VPO preparation on the properties of the final precursor and to correlate the solid state properties of these precursors with the catalytic activity in npentane oxidation and in particular in the selectivity to PA. Our attention focused mainly on the extent of structural defects observed in precursors and catalysts. In order to prepare VPO catalysts with varying extent of structt~ral disorder but with constant V/P ratio, the catalyst precursors were prepared via a VOPO4 mixed alcohol/water intercalate, containing various amounts of water, and further converted into the final precursor in an inert (n-octane) or reductive (isobutanol) medium. These two parameters, i.e. water content in the preparation mixture and the character of the medium during transformation into the VPO catalyst precursor, is shown to play a decisive role in the control of the structural characteristics of the final catalyst. Precursors were characterised by elementary analysis, XRD, FTIR, SBET, porosity and XPS methods. The role of structural disorder on the catalytic activity was evaluated using the following classification of structural defects: (i) long range defects showing destruction of links between vanadyl phosphate layers both in XRD and IR spectra; (ii) short range defects exhibiting distortion inside the vanadyl phosphate sheets shown by broadening of the respective IR bands. EXPERIMENTAL a) Materials: V205 (Janssen Chimica, purity > 99.9%), isobutanol (Aldrich, by Karl Fisher titration the content of water was found to be 0.1%), H3PO4 (Fluka, purity > 99%), benzene (Aldrich, purity > 99%), and n-octane (Aldrich, purity > 99%) were used as obtained for V P O preparation. Water was purified on Milli-Q plus system, SiC (particle size 0.3 ram) provided by Prolabo was used as obtained.

b) Vanadyl isobutoxide (V-ISOB) was prepared by refluxing 40g of V2Os in 500 ml isobutanol with addition of 50 g benzene with water continuously removed in a form of water-benzene azeotrop and separated in a Dean-Stark apparatus. Typically, about 6 ml water were isolated by dissolving about 30 g V205. ARer refluxing for about 40 hours, the rest of the unreacted V205 was isolated by filtration of the mixture by a membrane filter (Schelicker and Schuele 100) and the solvent was evaporated to 50 cc at lower pressure for about 16 hours. The vanadium content found in the colourless liquid products was over 95% of the theoretical value expected for pure vanadyl isobutoxide. c) Catalysts precursors preparation. The simplified presentation of the preparation procedure explaining the notification of the precursors during the preparation steps is shown in Fig. 1.

1215 Table I. Amount of water used in the preparation of Precursors A and their identification. order (1) 0 1 2 3 4 5 6 7 number mole (2) 0.2 0.65 1.1 1.9 2.4 3.1 4.2 5.3

(H20N)

ratio (1) this order number was used to denote the individual A, B and C samples. (2) mole of water added to vanadium ratio in the preparation solution of precursor A.

Precursors

A were prepared by addition of H s P O 4 to vanadyl isobutoxideAsobutanol/water mixture. The P N ratio of the mixture was 1.0. The water content, expressed as HsO/V mole ratio in the mixture, was adjusted in individual preparations to a value between 0.2 and 5.3. The mixture was vigorous stirred for 48 hours. The solid isolated by filtration was dried under vacuum at 50 ~ for 2 hours. The notification of the individual samples prepared and differing in the water/vanadium ratio is shown in Table I. Precursors Bo were prepared by refluxing 5.5 g of a precursor A in 100 ml of n-octane for 20 hours. The solid was isolated by evaporation of the solvent under low pressure and dried under vacuum at 80 ~ for 18 hours, filtered and dried at 80 ~ for 18 hours. Precursors Bi were prepared by refluxing 5.5 g of precursor A in 40 ml of isobutanol for 16 hours, filtered and dried under vacuum at 80 ~ for 18 hours. Catalysts C were prepared by equilibration of a precursor B at the npentane/air mixture up to 400 ~ for 20 hours. The notification of an individual Precursor Bo or Bi and a sample C is related to a parent precursor A (see also Table I ). Characterisation methods. Vanadium and phosphorus contents were determined by an ICP analysis (inductive coupled plasma atomic absorption) a f a r dissolution in 0.1 M nitric acid, carbon by measuring the omount of COs produced by total oxidation using Coulomat 702 Stroelheim. The BET specific surface area was obtained using ASPAP 2000 (Micromeritics) by nitrogen adsorption at- 196 ~ after degassing samples at 125 ~ Powder X-ray diffraction analysis was obtained in a high resolution X-ray diffractometer Siemens D-500 XRD using CuKa radiation. For XRD spectra interpretation data available for VOPO4 hydrates and alkoholates i.e., for VOPO4.2 H20 data identified by McMurdie [2], taking the most intensive peak at about 20 =11.8 ~ and indexed as (001), for VOP04. H20 data presented by Ladwing [3] taking the prominent band at 20 =12 ~ indexed as (001), an with the second most intensive line (002) at 2 o =28 ~ and for VOPO4-aliphatic alcohol intercalates as presented by Benes et al. [4]. The identification of XRD reflections of VOHPO4.1/2 H20 and (VO)2P207 were taken according to Bordes et al [5]. IR spectra were recorded with a Bruker IFS 88 spectrometer using KBr technique. Observed infrared bands were interpreted using assignment presented by Ladwig [3] for VOP04 hydrates and by Busca [6] for VOHP04.1/2 H20 and (VO)2P207. XPS analysis was performed with an SSX-100 model photoelectron spectrometer (FISONS) using monochromatized aluminium anode. The binding energy (BE) values were calculated with respect to C ls (BE of C-CH

1216

fixed at 284.8 eV). The fitting routine was used supposing Gaussian/Lorentzian ratio 85/15, the P/V, C ~ and C.~ ratio using sensitivity factors provided by manufacturer and with SiO~ as an external reference. The BE of V2p~ of 518.1 eV and 517.4 eV were used for V5+ and V4+, respectively [7].

Q

O

.

mixing RT

48 hours

_

VOPO4-water--isobutanol intercalate

A(0) to A(4)

n'~

.~-\

r~ ] ~

i

A[0] to A[7]

~risobutanol eflux

R1

/R

2

~ i aBi solution~

~ i a Bi solid - ~

-, Bi 3) /

R 1 - solid state reduction in inert media (n-octane) R2 - reduction "via solution" in isobutanol media R3 - reduction "via solid" in isobutanol media Figure 1. Scheme of samples preparation.

Catalytic tests were made in a U-tube reactor using 0.2 g of VPO d) catalyst. The reaction stream (total flow 30 ml/min) was : n-pentane 0.6 %

1217 vol., oxygen 20 %vol., helium 5 %vol, balance nitrogen. The inlet and the outlet gas streams were analysed by an on-line gas chromatograph equipped with a TCD detector on TENAX and Porapak Q columnA (Altech Associates, Inc.), running under temperature program from 70 to 240 ~ The outlet of the reactor prior to analysis was kept at about 150 ~ to prevent condensation of products. Conversion of n-pentane is defined as the number of n-pentane converted by the number of n-pentane feed (in % tool.). Selectivities to maleic and phthalic anhydrides were expressed as the fraction of n-pentane converted into AM or PA (in % tool.). The data were measured under conditions where the molar conversion of n-pentane was always equal or lower than 15%. 3. RESULTS 3,L P r e p a r a t i o n of catalyst precursor~ Precursors A were prepared in a form of yellow (A[0]) to dark green (A[7]) solid and identified as VOP04 - mixed isobutanol/water intercalates with the amount of isobutanol per VOP04 molecule varying from 1.6 to 0.05, and displaying the basal spacing decrease from about 14.6 /k to about 7.4 ~, respectively.

Precursors B were prepared by various preparation routes and thus the reduction step of V5+ to V4+ was realised in different reaction conditions: (i) By reduction in an inert media (Precursorsr Bo) the colour of the suspension turned due to V5+ reduction under reflux from yellow to black after 3 hours of reflux, and so the isobutanol intercalated in the Precursor A served as a reduction agent. The IR spectra of samples display a broad unresolved band between ca 800 to 1400 cm -1, with a m a ~ m ~ m at about 1085 cm -1 and several very diffuse low intensive bands at the 400 - 700 cm -1 region, tentatively assigned to a very defective structure of a vanadyl phosphate. This was supported by XRD results showing amorphous product, displaying only a low intensity XRD reflection at the 20 = 28.5 ~ (ii) Substantial differences were observed during preparation in the reductive medium (Precursors Bi) using precursors A low (samples A[0] - A[3]), or high (A[4] -A[7]) content of water and accordingly the ~ m p l e s prepared in the reductive media are further divided in the following subgroups: (ii-a) Precursors Bi[0] - Bi[3]. These precursors were formed in the form of a blue solid by slow precipitation from a dark green solution obtained during reflux. Accordingly these are referred to as precursors Bi "via solution". (H-b) Precursors Bi[4] -Bi[7]. No apparent dissolution of the solid was observed by starting from precursors A with a high water content. Instead, the colour of the dispersed solid turned during the first two hours of reflux from yellowgreen into pale blue and then does not change during the rest of the time of the reflux and these samples are further referred in the text as precursors Bi prepared "via solid". Comparing both types of Bi samples it could be summarised that the via solution prepared samples differ in the IR and XRD spectra (see Figure 2 and 3, respectively) in the following structural features : i) there is suppression of Ill bands assigned to POH (~p POH at 1133 cm-1 and v P-(OH) at 933 cm-1. ii) there is a very broad and low intensity of the group of IR bands belonging to the O-P-O structure (80PO at 413 482, 534 and 550 cm-1);

1218 iii) A missing or badly developed IR band at 686 cm -1 assigned to the coordinated water molecules, accompanied by a parallel shift in the position of the maximum of the 8 H20 band. (iv) they display only one broad band in the XRD spectra with 20 of about 30.4 ~ assigned to (202) plane of VOI-IP04. 1/2 H20. (v) the missing line at about 20 = 15.5 ~ indicate a very defective structure with a high disorder along the (010) plane.

g l]

l-

v V9

1800

t(100

1400

1200

1000

w w e n u n d w ~ om"t

Figure 2.

FTIR spectra of precursors Bi

800

600

400

1219

w

m

m

It

20

s

40

Figure 3 X R D spectra of precursors Bi 3,2. Charaetet4zation of s ~ u e t t u ~ defecfs of the catalyst precursors An attempt to characterise the level of the structural order of the layered structures of the vanadyl phosphates is given considering that these defects correspond generally to two categories, i.e. long range or short range defects. The idea is to distinguish roughly between the disorder effecfing interlayer connections of the layers of the cryst~ lattice Gong range defects) and the local disorder around the vanadium atoms (short range defects), see also e.g. [8] and [9] using similar concept. The general features of these two categories used were as follows: i) long range defects showing destruction of links between vanadyl phosphate layers both in XRD and IR spectra. In this category we consider defects indicated by broadening and eventually missing of the XED bands and destruction of links between vanadyl phosphate layers, as suggested by the missing of bands of Sip P-OH and v P-OH at 1135 and 927 cm -1, respectively, and of ~ I-I20 at 686 cm -1. ii) short range defects exhibiting distortion inside the sheets of the structure and manifested by broadening of the IR bands of the vanadyl phosphate molecule. In this category are considered broad or not distinguished other IR bands of vanadyl phosphate molecules at the 1400 - 400 cm -1 region, suggesting defects inside the V O P 0 4 sheets, and so a distortion of the structural order on a molecular level.

1220

The general characteristics of the s~mples prepared by various preparation routes were characterised as follows: Precursors B~[1]-[4] display high extent of defects, both of long range and short range character. Precursors B~ prepared "via solution" (Bt[1] -[3]) contain mostly long range defects, but tl~e short range order is mostly retained. Precursors B~_ prepared "via solid" (Bi[4]-[7]), displayed both short and long range order. By comparing these structural characteristics with the amount of the organic species present in precursors B, and expressed as the (C/V) value, the general trend between the efficiency of elimination of fragments of the reducing agent during formation of precursor from precursor A, and 'the extent of their structural defects is also suggested.

Catalysts Samples C were identified as (VO)2PsO7 with various extent of defects. Only samples prepared via solid displayed some extent of both the short and long range order. Most significantly the IR bands at 790 and 743 cm-1 associated with links between layers were reliable detected only in ~mples via solid and were missing in samples prepared from precursors B via solution. 3.4. Catalytic activity Results of catalytic activity tests of catalysts C together with some structural characteristics are presented in Table II. Table II. Characterisation of the preparation procedure of the VPO catalysts and their activity, in n-pentane oxidation. structural PrecursorB preparation SSET n-pentane SMA.% SPA.% defects (I) route m2/g conversion % evaluation cs) ,

13o(2) Bo(3) t3o(4) Bi(2)

octane 20.0 10.5 S+L 17.5 none octane 24.4 9.9 S+L 18.7 none octane 12.9 11.1 S+L 19.0 none isobutanol, 14.9 11.7 S 5.6 none via solution Bi(3) isobutanol, 16.9 12.8 S 5.6 none via solution B~(4) isobutanol, 26.3 15.6 0 24 4.8 via solid B~(6) isobutanol, 23.7 12.7 0 31 4.6 v/a sol/d B~(7) isobutanol, 19.0 12.4 0 26 9.6 via solid (I) precursor B used for catalytic test. (2) structural defects evaluation: (S+L) - both short and long range defective structure; S - short range defects, long range ordered structure; O - both short and long range ordered structure.

1221

Clearly two regions are observed: one in which both short and long range order are present, and the region in which no or only short range order is observed. M A was the only product of selective oxidation over catalysts produced from Bo or Bi by via solution preparation route, i.e. starting from precursor A with low content of water. Catalysts obtained by the via solid route produced by n-pentane oxidation both M A and PA, and the P A / M A ratio was the highest for samples with the m a x i m u m of water used in the synthesis of precursor. 4. DISCUSSION

By comparing the selectivity of the different catalysts for P A and M A formation in n-pentane oxidation over V P O catalysts, we concluded that P A formation displayed a different sensitivity to the nature of the structural defects of the catalyst. Accordingly, products of selective oxidation of n-pentane over V P O catalysts could be controlled by varying the conditions of the precursor preparation, and formation of M A and P A is directly related to the structural defects of the prepared V P O catalysts. In our previous report [10], it was shown that the formation of maleic and phthalic anhydrides could be controlled by varying the experimental conditions of the precursor preparation. In that work such effect was produced by changing the time provided for developing the V O P 0 4 mixed water/isobutanol intercalate (Precursor A of the present notation). The preparation of VOHPO4.1/2 H20 via full development of the VOPO4-mixed water/isobutanol intercalate favoured the formation of a precursor which promotes production of PA. Only M A and no P A was observed over the catalyst prepared directly from solution (i.e.to some extent similar to samples Bill] -[3] of the present work). Selectivity was correlated with the structural order of the samples. Nevertheless, due to some specific features of the Precursor A formation under conditions used in the previous work [10], these samples also displayed differences in both the X P S surface atomic P/V ratio and the bulk P N value, thus making less convincing the arg~_~ments for the role of structural order of the V P O catalysts in the selectivityto P A or MA. On the contrary, in the present work, samples show the same and nearly stoichiometric P/V bulk atomic ratio. The procedure of the development of the VOHPO4.1/2 H20 (Precursor B)was then controlled principally by the amount of water added during preparation of precursors. The prepared somples then differed by the amount of intercalated organic material as shown by the changes of the C N values. This results in a distortion of the layered structure of the precursor and consequently in the V P O catalyst (XRD and FTIR). Our results confirm the fact the selectivity in the formation of M A and P A is directly related to the structural characteristic of the prepared V P O catalysts. CONCLUSIONS

Using the evidence of the changes in the M A and PA selectivity and the characterisation of the nature of the structural disorder, we suggest that M A could be formed on a highly defective V P O catalyst,but that both the short and long range structural orders are necessary for P A formation by oxidation of npentane over V P O catalysts. It could be speculated that a high structural order is necessary to create a tridimensional surface topology around the

1222 complex active sites with an optimal distribution of the individual active functions to provide for the concerted process of the PA formation. ACKNOWLEDGMENTS The stay of Mrs S.R.G. C a r r ~ was supported by the Direcci6n General de Investigaci6n Cientffica DGICYT (Programa FPU), Ministerio de Educaci6n y Ciencia de Espafia, which is gratefully acknowledged. The Service de Programmation de la Politique Scientifique (Belgium) is gratefuly acknowledged for its Concerted Action grant, especially for the support of Ing. Z. Soba!ik and Dr. P. Ruiz. The authors also thank Mr. M. Genet for his help in the XPS analysis. R~'~CI~ 1. G. Centi, J. T. Gleaves, G. Golinelli and F. Trifiro, III European Workshop Meeting "New Developments in Selective Oxidation" (P.Ruiz and B. Delmon, Eds), Stud. Surf. Sci. Catal., 72 (1992) 231. 2. H. McMurdie, Powder Diffraction 1 (1986) 98. 3. G. Ladwig, Z. Anorg. Allg. Chem., 338 (1865) 266. 4. L. Benes, J. Votinsky, J. Kalousova, and J. Klikorka, Inorg. Chim. Acta, 114(1986)47. 5. E. Bordes, P. Courtine, and J.W. Johnson, J. Solid State Chem., 55 (1984) 270. 6. G. Busca, F. Cavani, G. Centi, and F. Trifiro, J. Catal., 99, (1986) 400. 7. S. R. Carraz~in, C. Peres, J.P. Bernard, M. Ruwet, P. Ruiz, and B. Delmon, J. Catal., 158 (1996) 452. 8. S. Albonetti, F. Cavani, F. Trifiro, P. Venturoli, G. Calestani, M. Lopez Granados, and J.L.G. Fierro, J. Catal. 160 (1996) 52. 9. L.M. Cornaglia, C. Caspani, and E.A. Lombardo, Appl. Cats]. 74 (1991) 15. 10. Z. Sobalik, S. Gonzalez and P. Ruiz, Preparation of Catalysts VI. Scientific bases for the Preparation of Heterogeneous Catalysts, Stud. Surf. Sci. Catal. 91(1995)727.

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 1997 Elsevier Science B.V.

1223

E l e c t r o c h e m i c a l O x i d a t i o n of P r o p e n e U s i n g a M e m b r a n e R e a c t o r w i t h Solid E l e c t r o l y t e S. Hamakawa, T. Hayakawa, K Suzuki, M. Shimizu and K. Takehira National Institute of Materials and Chemical Research, 1-1 Higashi, Tsukuba, Ibaraki 305, J a p a n

An electrochemical reactor, using ceria-based solid electrolyte coated with yttria-stabilized zirconia (YSZ), and gold and silver as the anode and cathode, respectively, has been employed for the selective oxidation of propene at 450~ On applying a direct current to this system, acrylaldehyde was formed at the gold anode, and its formation rate increased with increasing direct current. Selectivities to acrylaldehyde, CO and C02, based on converted propene, were 13.4, 25.6 and 61%, respectively, when the YSZI SDC was used as an electrolyte membrane. It is found t h a t selectivity to acrylaldehyde in this study was higher than that (ScHo=8.5%) obtained by using SDC alone as an solid electrolyte. This result suggests t h a t YSZ coating on the ceria-based solid electrolyte leads to inhibit complete oxidation of propene at the surface of cerium oxide. From the dependence of the selectivity to oxidation products on the thickness of YSZ, it is considered t h a t the selective oxidation of propene occurred at the Au-YSZ-gas triple phase boundary by the oxygen species pumped electrochemically through the ceria-based solid electrolyte and the YSZ.

1. I N T R O D U C T I O N In the oxidation chemistry, activation of molecular oxygen can be achieved by its reduction, i.e., 02 -* 0 2 ~ 0 --* 0 2 . The transient intermediate oxygen species, e.g. 0 2 , 0 , would play an important role as an active species for oxidation of hydrocarbons. We have reported that a fuel cell type reactor with an oxide ionic conductor, e.g. yttria-stabilized zirconia (YSZ), for the selective oxidation of hydrocarbons, as shown in Figure 1, can serve as an 'oxygen pump' [1 ], by which the oxygen flux transferred across the YSZ can be controlled by the electric potential externally applied between the anode and cathode [2-4]. Four electron reduction of molecular oxygen, 02, to oxide ion, 02. , at the cathode and oppositely four electron oxidation of oxide ion to molecular oxygen takes place at the anode under oxygen pumping conditions. Therefore, it is likely t h a t the

1224 active oxygen species, e.g. 02, O-, can be produced on the anode surface during oxygen pumping conditions. In addition, an electrochemical membrane reactor system has a lot of advantages to CnHm the hydrocarbon oxidation; (1) the rate of oxidation can be controlled 02 externally by varying the electric current or potential, (2) the CO, H2 contribution of the gas-phase CnHm.2 I "~~electrode reaction by 02 can be reduced, (3) CnHm O 1 ~ (catalyst) the system is safe from the view oxide ion conductor point of possible explosion and (4) a fuel cell reactor for energy Figure 1. Principle of the hydrocarbon generation can be constructed. oxidation using an electrochemical Thus, the application of membrane reactor. electrochemical cell using the YSZ for the oxidation of ethene [5], ethylebenzene [6] and methane [7-9] over metal or metal oxide catalysts as the anode has been studied. Furthermore, we have demonstrated that a ceria-based solid electrolyte (CeO2)0.s(Ln01.5)o.2 (Ln=Sm, Gd, Y) is an attractive candidates for use in an electrochemical reactor for the selective oxidation ofpropene at relatively low temperatures [10]. This is due to the high oxide ionic conductivity at the low temperature of 350~ compared with t h a t of YSZ. However, the complete oxidation of propene to carbon dioxide took place easily at the surface of CeO2, so that the selectivity to acrylaldehyde was lower than that using YSZ. Recently, K. Eguchi and H. Arai et al. reported t h a t a fuel cell with SDC coated with YSZ prepared by the ion plating method attained higher power density than that with YSZ alone at 800~ [11]. This is due to the achievement of constructing from a dense film of YSZ on the SDC to suppress the reduction by the fuel gas. In this paper, we report the investigation of a ceriabased solid electrolyte coated with YSZ to inhibit the complete oxidation of propene at the surface of Ce02 as well as the selectivity to partial oxidation to acrylaldehyde. 2. E X P E R I M E N T A L

The apparatus for partial oxidation of propene was constructed as shown in Figure 2. Gd and Sm doped ceria disk, (CeO2)0.s(Gd01.5)0.2 (GDC) and (CeO2)o.s(Sm01.5)0.2 (SDC) (0.8 mm thickness and 13 mm diameter, Anan-kasei Co.), respectively, were used as an electrolyte membrane. YSZ thin film was prepared on the disk by the spin-coating method with the YSZ slurry. The obtained specimen was sintered at 1000~ for 1 hr in air. The thickness of YSZ was controlled by the rotation rate of spin coating apparatus. A thin film of gold

1225

Au wire

5% C3H6--~~--~

t~[ll~ I

Quartz tube ~ ~ ] 1 ~

I[[[[ Ill I /Aluminatube 1111 ~ ]

. [

Glass seals

~111

::-

u lam 0

Sample (0.8 mm)~---I [i

I

~

---[1!~1--~ Thersnocoupm

1111

o lOO% o 2 - - - ~ - 1

Figure 2. Schematic diagram of the electrochemical membrane reactor. (1.0 ~m thick) was prepared as the anode on the YSZ film by a vacuum vaporization method. Porous silver cathode (2-3 ~m thick) was then prepared by painting silver paste on the other face of the disk, followed by annealing at 550~ for 1 h. The two electrodes (projected area; 0.5 cm 2) were connected withgold wires to an electrical circuit in order to control the oxygen transfer flux from the Ag cathode to the Au anode through the electrolyte. The disk was mounted between two vertical alumina tubes and sealed by low melting point glass. The reactor was placed in the electrical furnace to control the temperature of the electrochemical cell system. A gaseous mixture of propene (5%), nitrogen (5%)and helium (90%) was passed (1.5 l.h 1) over the anode at 450~ for testing the catalytic activity of the various electrolyte systems. Oxygen gas (1.8 l.h 1) was introduced into the cathode compartment. The products in the effluent gas were determined by gas chromatography using a thermal-conductivity detector and nitrogen as an internal standard. Porapak Q (GL. Science), PEG 6000 (GL. Science) and molecular sieve 13X (GL. Science) columns were used for analyzing and estimating the rates of hydrocarbons, oxygenated compounds and inorganic gases, respectively. The oxygen mass balance was estimated from the comparison between the electrochemical oxygen supply and the oxygen consumption calculated from the amounts of oxygenated compounds. The conductivity of the electrolyte membrane was measured by the ac impedance method [12].

1226 3. R E S U L T S AND D I S C U S S I O N 3.1. Characterization of the c e r i a - b a s e d solid electrolyte c o a t e d with YSZ SEM observations of the surface and the cross-section of the Gd doped ceria coated with YSZ (YSZ I GDC) show t h a t the thin YSZ film is porous and the YSZ particles are fairly uniform in size. The thickness of YSZ film was about 20 ~m. The conductivity of this composite membrane was 4 . 3 9 x 1 0 4 S'cm "1 at 500~ under air atmosphere, which was almost same order as t h a t of Sm doped ceria coated with YSZ (YSZ I SDC) (1.51 X 10 .4 S'cml). However, either value was lower than t h a t of GDC alone (4.03 x 10 -3 S.cm -1) or SDC alone (4.66 x 10 .3 S.cm-1). This may be due to the low conductivity of YSZ film. 3.2. Partial oxidation of p r o p e n e Propene oxidation was carried out by using an electrochemical reactor constructed from a Sm doped ceria electrolyte coated with YSZ (YSZ [ SDC) as a membrane. In a blank t e s t where nitrogen gas alone was passed over the Au anode instead of the reaction gas at 450~ it was confirmed t h a t the oxygen pumping was well controlled by the applied current, i.e., the amount of oxygen gas evolved at the anode coincided well the value calculated from the electric current by using Faraday's law. When the propene mixed gas was introduced into the anode space at 450~ the reaction cell gave a stable voltage of 480 mV (EMF) under open circuit condition. Applying the electric current to the reaction cell, propene oxidation took place, and the rate of oxidation increased with ~. 6increasing electric current (i.e., .~ YSZISDC increasing oxygen flux). No -~ evolution of molecular oxygen was ~k 5 observed in this case. Namely, -~ 4cch all oxygen species pumped ~ electrochemically through the o 3YSZI SDC membrane was CO consumed by the propene o 29 J A oxidation over the Au anode. ~ c~4o Acrylaldehyde, carbon monoxide and carbon dioxide were observed ~ o as the main products under the ~ 0 T" '~ ! 0 0.5 1.0 1.5 closed circuit conditions. The O formation rates of oxidation Current / mA products are plotted as a function of the current in Figure 3. The Figure 3. Partial oxidation ofpropene using formation rate of acrylaldehyde the YSZISDC electrochemical membrane increased linearly with increasing reactor. the electric current. This result

1227 indicates t h a t acrylaldehyde formation proceeds by the reaction of propene with the oxygen species pumped electrochemically through the YSZISDC membrane from the cathode to the anode. When the electric current was 1.1 mA, i.e., 10.2 ~mol~r of oxygen flux, the formation rates of acrylaldehyde, carbon monoxide and carbon dioxide were 0.33, 1.9 and 4.53 ~tmol~r, respectively, and the oxygen consumption rate of 9.02 ~tmol/hr was calculated by assuming water as another oxidation product. Mass balance calculated between oxygen evolution and consumption of propene oxidation was about 90%. 3.3. Effect of the YSZ c o a t i n g on the s e l e c t i v i t y to oxidation p r o d u c t s We have studied the partial oxidation of hydrocarbons with electrochemical reactor using an oxide ionic conductor, e.g. YSZ, SDC, etc. [2-4, 10]. In these studies, it was found t h a t a ceria-based solid electrolyte is useful for the propene oxidation to acrylaldehyde at relatively low temperature of 350~ [10]. In this case, however, the acrylaldehyde selectivity was lower than t h a t obtained with the electrochemical reactor constructed from YSZ. This may be due to the high activity ofceria surface for the complete oxidation of hydrocarbons [13,14]. The effect of the YSZ coating on the ceria-based solid electrolyte was shown in Table 1. When the YSZI SDC membrane was used as the solid electrolyte, selectivities to acrylaldehyde (ScHo) carbon monoxide (Sco) and carbon dioxide (Sco2) based on converted propene was 13.4%, 25.6% and 61%, respectively. Here, it should be emphasized that the selectivity to acrylaldehyde increased with YSZ coating compared with t h a t (ScHo =8.5 %) obtained by using SDC alone as a solid electrolyte. In addition, it was found t h a t carbon monoxide formation was observed in the present study, although its formation was not detected in the case of SDC alone. The same phenomena were observed, when the Gd doped Table 1.

Selectivity to oxidation products in the propene oxidation at 450~ Selectivity / %

Solid electrolyte

Thickness of YSZ

CO

C3H40

C02

SDC GDC

--- % --- %

8.5 % 12.0 %

91.5 % 88.0 %

YSZ/SDC

25.6 %

13.4 %

61.0 %

14.6~m

YSZ/GDC

20.7 %

18.2 %

61.1%

20.0~m

YSZ

10%

40.0%

50.0 %

1228

ceria coated with YSZ (YSZIGDC) was used as a electrolyte membrane for the electrochemical oxidation of propene. These results suggest that YSZ coating on the ceria-based solid electrolyte lead to inhibit complete oxidation of propene at the surface of SDC or GDC. However, taking into account the fact t h a t the acrylaldehyde selectivity in the either case of using YSZI SDC or YSZ I GDC was still lower than t h a t obtained with YSZ alone, it is likely that the surface of ceria was not completely covered by YSZ could not be obtained in this study. 3.4. R e a c t i o n site for t h e a c r y l a l d e h y d e p r o d u c t i o n

Generally, in the electrochemical membrane reactor, e.g. the AuIYSZIAg system, molecular oxygen is reduced to oxide ion by four-electron transfer at the Ag cathode, incorporated into and transferred though the YSZ as oxide ion, and then reoxidized to molecular oxygen by four-electron transfer at the Au anode. In addition, it is considered that oxygen evolution at Au anode occurs at the triple phase boundary between Au electrode-YSZ-gas phase [15,16]. Accordingly, we consider t h a t hydrocarbon, e.g. propene, reacts with the oxygen species appeared at the triple phase boundary through the YSZ and forms oxygenated compounds, e.g. acrylaldehyde. Furthermore, it is likely that the oxygen species pumped electrochemically at the triple phase is active for the partial oxidation of hydrocarbons, since the gold has no ability to dissociatively active molecular oxygen [17] and YSZ has no activity for the partial oxidation of hydrocarbons [4]. In order to clarify the reaction site of the partial oxidation of propene using the ceria-based solid electrolyte coated with YSZ as a membrane, we have studied the dependence of the selectivities to oxygenates on the thickness of YSZ. When the Sm doped ceria coated with YSZ (YSZISDC), each selectivity of the oxidation products did not dependent on the thickness of YSZ, as shown in Figure 4.

I 9.~

75 C02

0

O o 9~

50

op,.~

o o

25

4,

9 ~,4

n

l

-

CO C3H40

oP.,l

~

O~ 0

!

10

20

30

40

50

Thickness of YSZ / grn Figure 4. Dependence of the selectivity to oxidation products on the thickness of YSZ.

1229 This result indicates that the main reaction site is not inside the YSZ film, but its surface. Therefore, it is likely that the selective oxidation of propene occurred at the Au-YSZ-gas triple phase boundary by the oxygen species pumped electrochemically through the SDC and then the YSZ, as shown in Figure 5. However, it was found that acrylaldehyde selectivity in the case of using YSZISDC or YSZIGDC was still lower than that obtained with an electrochemical reactor constructed from YSZ alone, as shown in Table 1. Accordingly, a possible occurring of the propene oxidation at the surface of SDC and/or at the triple phase boundary between Au-SDC-gas cannot be denied in this study.

CO

CH2=CHCHO

C3H6

YSZ

Figure 5. Mechanism of the selective oxidation ofpropene using an electrochemical membrane reactor.

4. CONCLUSIONS An electrochemical cell system with ceria-based solid electrolyte coated with YSZ prepared by the spin coating method showed higher selectivity to acrylaldehyde than that with ceria-based solid electrolyte alone. This may be due to the fact that a film of YSZ on the ceria-based solid electrolyte to suppress the complete oxidation of propene. When the YSZISDC disk was used as an electrolyte membrane, selectivity of the oxidation products did not depend on the thickness of YSZ. This indicates that the selective oxidation of propene occurred at the Au-YSZ-gas triple phase bouridary by the oxygen species pumped electrochemically through the ceria-based solid electrolyte and the YSZ.

REFERENCES 1. 2. 3. 4. 5.

E.C. Subbarao, 'Solid Electrolytes and Their Applications,' Plenum Press, New York, 1980. T. Tsunoda, T. Hayakawa, K. Sato, T. Kameyama, K. Fukuda and K. Takehira, J. Chem. Soc., Faraday Trans., 91 (1995) 1111. A.P.E. York, S. Hamakawa, T. Hayakawa, T. Tsunoda, K. Sato, and K. Takehira, J. Chem. Soc., Faraday Trans., 92 (1996) 3579. S. Hamakawa, K. Sato, T. Hayakawa, A. P. E. York, T. Tsunoda, K. Suzuki, M. Shimizu and K. Takehira, J. Electrochem. Soc. 144, (1997) 1. M. Stoukides and C. G. Vayenas, J. Catal., 70 (1981) 137.

1230 6. 7. 8.

9. 10. 11. 12. 13. 14. 15. 16. 17.

J.N. Michaelis and C. G. Vayenas, J. Catal., 85 (1984) 477. D. Eng and M. Stoukides, Catal. Rev. -Sci. Eng., 33 (1991) 375. T. Hayakawa, K. Sato, T. Tsunoda, K. Suzuki, M: Shimizu and K. Takehira, J.Chem. Soc., Chem. Commun., (1994) 1743; K. Sato, J. Nakamura, T. Uchijima, T. Hayakawa, S. Hamakawa, T. Tsunoda and K. Takehira, J. Chem. Soc., Faraday Trans., 91 (1995) 1655. K. Otsuka, S. Yokoyama and A. Morikawa, Chem. Lett., (1985) 319. S. Hamakawa, T. Hayakawa, H. Yasuda, K. Suzuki, M. Shimizu and K. Takehira, J. Electrochem. Soc., 143 (1996) 1264. T. Inoue, T. Setoguchi, K. Eguchi and H. Arai, Solid State Ionics, 35 (1989) 285. T. Yajima, H. Iwahara and H. Uchida, Solid State Ionics, 47 (1991) 117. J.M. Deboy and R. F. Hicks, Ind. Eng. Chem. Res., 27 (1988) 1577. T. Hattori, J. Inoko and Y. Murakami, J. Catal., 42 (1976) 60. H. Yanagida, R. J. Brook and F. A. KrSger, J. Electrochem. Soc., 117 (1970) 593. J. Sasaki, J. Mizusaki, S. Yamauchi and K. Fueki, Bull. Chem. Soc. Jpn., 54 (1981) 1688. G.I. Golodets, 'Heterogeneous Catalytic Reactions Involving Molecular Oxygen,'Elsevier, New York, 1983.

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 9 1997 Elsevier Science B.V. All rights reserved.

Vanadium pentoxide catalytic partial oxidation of 1-butene

membrane

1231

reactor

for

Sangjin Moon, Tayoon Kim, Seungdoo Park a, Jihoon J u n g b, Sukin Hong Dept. of Chemical Eng. Korea University, Seoul, Korea SKI petrochemical research center, Suwon-si, Korea b Dept.of Chemical Eng. Kyonggi University, Suwon-si, Korea a

ABSTRACT The vanadium pentoxide cata~tic membrane reactor was prepared by coating its sol inside the Vycor ~ tube membrane. After heat treatment of the prepared membrane, the [010] planes of vanadium pentoxide layer were grown largely which contributes to partial oxidation reaction of l-butene to maleic anhydride. The partial oxidation of l-butene to maleic anhydride was carried out in the catalytic membrane reactor. The maximum selectivity of 95% was obtained at 350~ when the surface velocity was 500cm/h. And at this condition, oxygen permeability was almost four times higher than the reaction had not occured.

1. I N T R O D U C T I O N One of the most industrially important reactions using vanadium pentoxide(V205) catalyst is the partial oxidation of l-butene to maleic anhydride [1]. Partial oxidation reactions are inherently unselective and often make by-products of little or no value. Oxygen-rich feeds result in low product selectivities and high hydrocarbon conversions [2]. Because partial oxidation and total oxidation always proceed competitively, the selectivity of maleic anhydride from l-butene is low. Though fixed bed reactors or fludized bed reactors have been used for partial oxidation for the past 30 years, the selectivity of maleic anhydride has not been obtained higher than 69% [3]. Some attempts have been reported on a new type of reactor to overcome the above limit. This is a membrane reactor which offers some advantages. A membrane reactor plays a

1232 dual role both catalyst and membrane, which is commonly regarded as a barrier capable of being selectively permeated by some components of a mixture or, at least, of changing the composition of a fluid stream that flows through it due to a certain driving force such as a pressure, concentration, or electric potential gradient.[4] First, one of the products is separated by the membrane and chemical equilibrium is shifted. As a result, the conversion is increased. This type of membrane can be used for water-gas shift reaction, dehydrogenation and so on [5]. Second, the formation of by-products is suppressed due to presence of a membrane between the two reactants, so that selectivity is improved. This second type can be applied for the partial oxidation a n d hydrogenation reaction that produces a lot of valuable by-products during the reaction, but lacks in study compared to the first type of membrane reactor. The aim of this research was to prepare the V20~-coated catalytic membrane reactor using sol-gel method, and to investigate reaction charateristics of that through the partial oxidation of 1-butene to maleic anhydride. And also, we investigated the enhancement of oxygen permeability through V2Os-coated catalytic membrane during partial oxidation reaction occurs.

2. EXPERIMENTAL 2.1 Preparation of a V205-coated catalytic m e m b r a n e by sol-gel method V205 sol was made by melting its powder at 900~ and then pouring it into deionized water. After that the mixture was stirred by magnetic stirrer and then it was filtrated. This procedure uses the property that amorphous V205 dissolves easily in water [6]. As a support for the membrane reactor, Vycor | glass with quartz tube connected to either end was used. The length and the diameter of Vycor are 100mm and 7 ram, respectively. In order to coat inside the Vycor with V20~ sol one end of the Vycor was connected to syringe and the other end of it was put into V205 sol. After the sol was sucked in, it was allowed to stay within the Vycor for a certain time to form thin layers on the inside surface of the Vycor, and then the remaining sol was pushed out. The thickness of the thin layers was controlled by the length of time the V205 sol remaining in the Vycor, and by the viscosity of sol. The coated Vycor was dried at room temperature for 24 hours and, in order to crystallize the V205, the heat t r e a t m e n t was carried out by raising the furnace temperature by 1 ~ per minute and keeping it at 450~ for an 0hour. And the V205 membrane was also prepared by dip coating the V205 sol in YSZ disk to characterize the membrane surface and its cross section.

1233

2.2 Reaction of 1-butene and p e r m e a t i o n of oxygen through the membrane The VzO5-coated catalytic membrane connected to the quartz tube was fixed within the 1/2 inch stainless tube by using CAJON ~ ultra-torr fitting. Both nitrogen as the transport gas and 1-butene were passed through the catalytic membrane under the application of oxygen to the outside of the catalytic membrane. The oxygen permeated through the V2Os-coated catalytic membrane and reacted with 1-butene on the coated thin layer. At this time the concentration of the mixture of 1-butene and nitrogen was sustained at 1%, and the surface velocities (volumetric flow rate/surface area of reactor) of 1-butene and nitrogen was set to 500cm/h and 200cm/h, respectively. Maleic anhydride and unreacted compounds were analyzed by on-line gas Chromatography. The permeabilities of oxygen through the membrane were calculated by measuring the decreasing oxygen pressure using a pressure transducer connected to a recorder. To investigate the reaction effects on permeabilities of oxygen, the pure nitrogen was fed into inside the Ve05-coated catalytic membrane, and uncoated Vycor was also used under the same reaction conditions as above. Fig.1. shows the reaction apparatus.

6

1

2

3

.....

1 . 0 2 2. butene 3. N2 4. MFC 5. furnace 6. pressure transducer 7. A/D converter 8. thermocouple 9. membrane

Fig.1. The a p p a r a t u s of catalytic membrane reactor system.

2.3 Characterization Both the specific area and the pore size distribution were measured by adsorbing nitrogen while the existence of macro pores was confirmed by mercury porosimeter. The crystal structure of V205 according to the crystallization temperature and time was investigated by X-ray diffractometer. The morphology of VzO5 sol was observed by TEM and the thickness of the thin layer coated on the Vycor was measured by

1234 SEM. To analyze the unreacted 1-butene and maleic anhydride in membrane reactor, gas chromatography equipped with FID Tenax-GC | column was used. Also to analyze the CO and CO2, chromatography equipped with TCD and molecular sieve 5A Porapak Q column was used.

the and gas and

3. R E S U L T 3.1 S u r f a c e

area

and

pore

diameter

The specific surface area of vanadium pentoxide coated membrane was 1.1 me/g, and the average pore diameter of it, measured by nitrogen adsorption, was 30A. Macro pores identified by mercury intrusion did not exist, and the porosity was approximately 9%. From results of isotherm adsortion-desorption curve,as shown in Fig.2, V2Oscoated membrane has slit-shaped pore structure. Hence the pores mostly exist between the V205 layers, and the surface of those are nearly non-porous.

13. !-" 03 "~

4O

40

35

35 D

25

E := o

20-

~

10

o -o

5

>

30

30

15

0

-

0.0

25

L

20

L O

15

c~

10

l

I i 1 , I , I , 0.2 0.4 0.6 0.8 1.0

~ 0.0

l , l 0.2 0.4

i 0.6

0.8

Relative Pressure (P/P o)

Relative Pressure (PIP o)

(a)

(b)

Fig.2 Adsorption and desorption curve of V 2Os-coated layer with nitrogen (a) before and (b) after calcination, respectively ( [-];adsorption

9

; desorption )

1.0

1235

3.2 X-Ray Diffraction Analysis The X-ray diffraction reasults as a function of the heat treatment of coated membrane are shown Fig.3. The V20~-coated membrane prior to calcination was an amorphous structure without any characteristic peak. When the membrane was heated to 200~ the crystallization was slightly developed, and at 300~ and 400~ the crystallization continued to occur. But at temperature over 400~ the crystal never grew.

[0101

10

20

(a)

30

40

10

20

10

I

40

[010] [1011

20

30 (b)

[010] [2001

[101] [400]

[1011

[4001

I I 30

(c)

[4001

[200]

40

10

20

I

30

40

(d)

Fig.3. The XRD pattern of vanadium pentoxide layer with calcination conditions. (a) before calcination (b) 25"C-200~ (c) 25~ (d) 25"0-450"C

Therefore the heat treatment temperature shoul be higher than 400~ The characteristic peak of vanadium pentoxide occured at 28=15.4 ~ 20.3 ~, 22 ~, 25.6 ~ 26.2 ~ and 31 ~ Each peak shows [200], [010], [110], [210], [101], [400] plane. In the case of the membrane heated up to 200~ [200] plane, [010] plane, [110] plane and [400] plane started to occur but the rate of crystal growth is low. The membrane that was heated upto 300~ shows only the great growth of [010] plane while the [210] plane and [010] plane grew slightly. The membrane which was heated at 450~ showed the great growth in [200] plane, [010] plane, [110] plane and [400] plane, but the crystalization rates on other planes were negligibly small. Of the four crystalline structure, the active sites of oxidation reaction are [010] plane at 20.3 ~ and [101]

1236 plane at 26.2 ~, where [010] plane is the main active site of partial oxidation [13]. Therefore, the higher the rates of intensity of [010] plane and [101] plane are, the more favorably the partial oxidation occurs. The intensity rates are expressed by the morphological factor which is defined by 11OlOl/111o11. In the case of the vanadium pentoxide coated membrane, this factor is ten times higher than those of the ordinary vanadium pentoxide based catalysts. 3.3 P a r t i a l O x i d a t i o n R e a c t i o n The conversion in the membrane reactor was comparatively low. The maximum conversion was only 22% and 33%, respectively, when surface velocity was 500 cm/hr and 200 cm/hr. On the other hand, the selectivity of maleic anhydride in the membrane reactor was very high as shown in Fig.4. When the surface velocity was 200 cm/hr and 500 cm/hr, the maximum selectivity was 80% and 95% at 350~ respectively.

100

100

~

80

:~ 60

~

6o

40

g~ 40

20

2o 00

200

300

400

oo

Temperature ('(3) Fig.4. S e l e c t i v i t y of m a l e i c a n h y d r i d e over V205-coated membrane reactor ( 0 2 pressure at 15 psi) 9 Z~

surface 9 velocity 500 cm/h and 200 cm/h,respectively

200

300

400

Temperature ( ~ ) Fig.5. Selectivity of maleic a n h y d r i d e with 0 2 pressure over V205-coated m e m b r a n e reactor (surface v e l o c i t y 200cm/h) 9B •

" 02 pressure 15, 30 and 45 ,respectively.

The total oxidation occurs when 1-butene reacts to the gas phase oxygen, where the partial oxidation occurs when 1-butene reacts to the lattice oxygen. In the membrane reactor, 1-butene and gas phase oxygen are separated by the membrane, and 1-butene and oxygen in the gas phase cannot react directly, and therefore the total oxidation do not occur. On the other hand 1-butene adsorbed on the V20~ membrane reacted only with the lattice oxygen in the VeO~ membrane, and the desired lattice oxygen was supplied from the other side of V205 membrane. Therefore the partial oxidation occured only at the membrane reactor and the selectivity of maleic anhydride was ncreased.

1237 The conversion in the membrane reactor was very low compared to that in the fixed bed reactor because the amount of V205 catalyst coated on the membrane reactor was only 0.01cc. This restricts the reaction to a small area. The conversion at the surface velocity of 500 cm/hr was lower t h a n that at 200 cm/hr, because the amount of oxygen through V205 membrane wall do not changed with surface velocity of 1-butene. However, the selectivity of the former case was higher in the reverse because the contact time between 1-butene and catalyst is reduced. When contact time is long, a portion of maleic anhydride reacts with the catalyst again, and tends to be converted into CO2 and other by-products and, therefore, lowers the selectivity of maleic anhidride. Because butadiene and acetic acid are produced under 250~ and a portion of lattice oxygen is used for total oxidation above 400~ the selectivities of maleic anhydride under 250~ and above 400~ are relative low. In order to increase the conversion in the membrane reactor, the membrane reactor should be made in hollow fibers, or the unreacted 1-butene should be recycled. 1-butene can be condensed easily because it has a high liquefaction temperature. If the membrane process through recycling is developed, both the selectivity and the yield could be increased. When the oxygen partial pressure applied from the outside was changed to 15, 30 and 45 psi, the effect of oxygen pressure was observed. As shown in Fig.5, the selectivity decreases as the oxygen partial pressure increases. Most notably when the pressure increases from 15 psi to 30 psi the decrease in the selectivity was slight, but when pressure was raised to 45 psi the selectivity showed a substantial drop. This is due to the fact that, the permeation rate of oxygen through membrane was too low as the oxygen pressure was low. The oxygen was supplied only by the adsorption of oxygen on the V20~ membrane. When the pressure was above a certain value, the oxygen was permeated through the membrane and reacted directly with 1-butene, which causes the total oxidation. The permeabilities of oxygen through the Vycor were constant with the pressure difference and decreased with the square root of the temperature (figure not included). So the main mechanism for gas permeation is Knudsen diffusion. However, the permeabilities of oxygen through the V20~-coated catalytic membrane were deviated from Knudsen diffusion. As the partial oxidation proceeded, the permeabilities of oxygen through the V2Os-coated catalytic membrane were increased slightly with temperature. But at the temperature of maximum selectivity obtained, the permeability of oxygen was enhanced almost four times higher than that of the V20~-coated catalytic membrane that did not participated in the reaction. (see, to Fig.6 when the feed gas was nitrogen only.) The enhanced permeability of oxygen through V2Os-coated catalytic membrane at 350~ was most likely due

1238 to the partial oxidation reaction based on Redox mechanism of 1-butene. But further studies must be conducted for a more precise explanation regarding the enhancement of permeability. 12 o 10

g

m

6

•

~ 4 9

E o.

2

_

V2Os-coated Vycor

B

I

o

0

1oo

200

I

I

300

400

Temperature (~ Fig.6 P e r m e a b i l i t y of 0 2 with t e m p e r a t u r e at p r e s s u r e d i f f e r e n c e 15psi ( feed gas

9 /X I I , 1-butene and N2 mixture

OO

,N2

)

4. CONCLUSION The V2Os-coated catalytic membrane was applied to the partial oxidation reaction to produce maleic anhydride from 1-butene. The maximum selectivity was 95% at 350~ and the permeability of oxygen through the V2Os-coated catalytic membrane was enhanced. More careful studies are needed about this permeability enhancement.

REFERENCES 1. 2. 3. 4. 5. 6.

G. A. F. G. A. J.

Centi and F. Trifiro, Chem. Rev., 88(1988) 55 Tonkovich et al., Chem. Eng. Sci., 51(1996) 789 Cavani, Ind. Eng. Pro. Res. Dev., 22(1983) 565 Saracco and Vito Specchia, Catal. Rev. Sci. Eng. 36(1994) 305 Champagnie et al., Chem. Eng. Sci., 45(1990) 2423 Bullot et al., J. Non-Cryst. Solids, 68(1984) 123

3rd World Congress on Oxidation Catalysis R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons (Editors) 1997 Elsevier Science B.V.

Peroxidase a n o n

1239

of phenol by catalase immobilized on carbon materials

E.Horozova, N.Dimeheva and Z.Jordanova Department of Physical Chemistry, University of Plovdiv; 24, Tsar Assen St, Plovdiv - 4000, Bulgaria SUMMARY The peroxidase activity of immobilized catalase on the oxidation of phenol has been studied. The immobilization was carded out from eatalase solutions with pH < 3,5 on two kinds of soot differing in the average size of the particles building them up. The effect of the initial concentration of phenol on the rate of its peroxidase oxidation by catalase immobilized on the soot of finer-grained structure has been studied. The relationships obtained are described by the equation of Michaelis-Menten. The kinetic parameters (the constant of Miehaelis - Km, the maximum reaction rate - V, the rate constant - k and the activation energy - E a of the process were calculated. It was found that catalase adsorbed on the soot of larger globular particles does not take part in the peroxidase oxidation of phenol. INTRODUCTION In bioeatalytic systems, eatalase is mainly used in immobilized state. High activity of immobilized catalase was achieved on its sorption immobilization on cellulose [6], on silica gel modified with fatty acids or phospholipids [7] as well as on activated carbon fibres and fabrics/tissues/[8]. Biocatalytie activity of catalase immobilized on cellulose was also studied in nonaqueous solvents [9,10]. In [9] it was found that unlike the enzyme dissolved in waterdimethylformamide medium, on the oxidation of o-dianisidine in the presence of dimethylformamide the immobilized eatalase does not show any peroxidase activity. It was used [10] for working out an organic-phase amperometric biosensor by immobilizing the enzyme in a polymeric film on a glass-carbon surface. The effect of the polyacrylamide pad for immobilization on the thermodynamic activity of immobilized eatalase was studied in [ 11]. Catalase is also used in co-immobilization with other enzymes such as glucose oxidase [12,13], lactate dehydrogenase [14], peroxidase [15], for creating enzyme membranes needed on the determination of the respective substrates. To study the possibility for peroxidase activity of catalase initiated by the process of its immobilization on carbon materials is he goal of the present work. The peroxidase function of eatalase in solutions was well studied on the oxidation of ethyl alcohol, phenol [ 16,17] and arornatie amines [18]. However, there is no data about peroxidase activity of catalase in immobilized state. Such data is greatly needed because eatalase monomers are stable in acid solutions in which the use ofperoxidase is impossible due to its deactivation.

1240 MATERIALS AND METHODS The catalase used in this work was EC 1.11.1.6 from Penicillium chrysogenum 245 (Biovet -Bulgaria), M, = 244 000. -! The specific activity of the enzyme is 1000 U x m g . The reagents for the solutions: Na2HPO4.12H=O, citric acid, phenol and H202, all with analytical grade qualification "pure for analysis". The solutions were prepared with bidistilled water. Carbon materials: soot "NORIT" with fine-grained structure - the average size of particles being 5 x 104 - 45 x 104A, and "PM-100" built up of larger globular particles with an average size of 21 x 104 - 340x 104A. The adsorption of catalase on both types of soot was performed by an adsorption method under static conditions from 1 ml solution of catalase with concentration of the enzyme C=lx 10"4M in citrate buffer (pH =3.02) per 10 mg soot in 24 hours. The amount of catalase was determined spectrophotometrically by the decrease in the concentration of the catalase in the solution after the adsorption. The spectrophotometric measurements were carried out on a Specord UV VIS (Carl Zeiss, Jena, Germany). The amount of the catalase in the solution was determined on the basis of a calibration graph (Fig. l-b) for the maximum at ~,~ = 280nm. The value of the extinction coefficient was ees0 = 0.9x 1051.mol'~.crff1. The peroxidase activity of catalase in solution and in immobilized state on both types of soot was studied on the oxidation of phenol. The kinetics of the enzyme reaction was monitored spectrophotometricaUy by the decrease in the concentration of the substrate at = 270 nm A

- -

.

I

.

C"

.

"

"-

III

II

1

A

II

0.7

0.7

0.6

0.6

0.5

0.5

0.4

0.4

0.3

0.3

0.2

0.2

O.l

0.1

.

280

300

~,, nm

.

.

.

.

: . .

__

: -

:-.r_-

::

~.-.~

-.:

.-

b

/.

260

.

.

i

I

..

. J.

2

. . . ..

t .

3

.

i

4

_.

.

._

1

. .

5

.l

6

..

l_

?

C x 10', mol.l ~ Figure 1. Catalase absorption spectra (a) and a calibration graph (b) for determination of the catalase concentration in the solution; a: catalase concentration in citrate buffer (pH = 3.02) l x l 0 "4M; 1 cm cell.

1241 RESULT AND DISCUSSION On the adsorption of catalase from solutions with a concentration of C = 1x lO4M in citrate buffer with pH = 3.02, a dissociation of the enzyme to subunits takes place [2,18]. No matter how the catalase dissociation is initiated, the dissociation process is always accompanied by a certain loss of catalase activity and optic changes - a shift towards the shortwave region [ 19]. The catalase tetramer dissociation to monomers [ 17] runs according to scheme E4 r 4E. For dissolved catalase and for eatalase immobilized on soot "NORIT" the dependence of the initial oxidation rate of phenol on the initial concentration of the substrate was studied (Fig.2), where A is the absorbanee corresponding to the current concentration of the substrate (phenol). The obtained relationships are described by the equation of Michaelis-Menten :

v..~iSl v = K. +lSl

(])

where V is the rate of the enzymatic reaction, and Vm~- the maximum rate, i.e. the rate at which the ~ e - s u b s t r a t e complex react; K ~ - Michaelis constant, equal to such a concentration of the substrate at which the rate of the enzymatic reaction is half of its maximum; [S] - the concentration of the substrate. A transformation of these relationships by the method of Lineweaver - Burk brought about to equation:

1

K= 1

v-

v..~ s + v..--~-

1

(2)

used to built up a graph in coordinates 1/V - 1/[S] (Fig.3). The obtained straight line intersects K. the X and Y axes. The slope gives ( V=u )' the intercept of the Y-axis ( 1/Vm~ ), the intercept of the X-axis (-1/K~). The values obtained for K~ and V=,~ are given in Table 1.

Table 1. Kinetic parameters ofper0xidase oxidation ofphen01 by eatalase Catalase

Kin, M

V~,

l~t 285 K

E,, kJ.mol q

,S "1

303 K

In solution 5.48x104 1.43 0.9x10 "4 1.3xlO "4 14.56 C=1.36x lO'SM Immobilized on "NORIT" 2.73• .3 16.70 7.3x10 "4 9.5x10 "4 10.95 -~ ~ 0 . 3 7 m g . . . . . . ................... . ........................................ Immobilized DOES NOT SHOW PEROXIDASE ACTIVITY on "PM- 100"

1242 ~ :

A 0.7

=~-

_

.

.

.

.

.

.

.

.

.

.

.

.

.

:

.

-~

.~

......

V-I

1

V-~