11th International Congress on Catalysis - 40th Anniversary

Studies in Surface Science and Catalysis 101 11TH INTERNATIONAL CONGRESS O N CATALYSIS 40TH ANNIVERSARY

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Studies in SurfaceScience and Catalysis Advisory Editors: B. Delrnon and J.T. Yates Vol. 101

IITH INTERNATIONAL CONGRESS ON CATALYSIS = 40TH ANNIVERSARY PART A Proceedingsof the 1I t h ICC, Baltimore, MD, USA, June 30 -July 5,1996 Editors

Joe W. Hightower Department of Chemical Engheering, Rice University, Houston, TX 77251- 1892,USA

W. NicholasDelgass School of Chemical Engineering, Purdue University, West Lafayette, IN 47907,USA

Enriquelglesia Alexis T. Bell Department of Chemical Engineering, University of California, Berkeley, CA 94720-9989,USA

1996 ELSEVIER Amsterdam - Lausanne- New York -Oxford

-Shannon -Tokyo

ELSEVIER SCIENCE B.V. Sara Burgerhartstraat 25 PO. Box 21 1,1000 AE Amsterdam, The Netherlands

ISBN 0-444-81947-9

0 1996 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, PO. 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 referredtothecopyright owner, Elsevier Science B.V., unlessotherwisespecified. 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

V

TABLE OF CONTENTS Part A Preface J.W. Hightower

xix

Plenary Lectures P- 1

P-2

P-3

P-4

P-5

Driving Forcesfor Innovation in Applied Catalysis I.E. Maxwell

1

Constrained Geometry and Other Single Site Metallocene Polyolefin Catalysts: A Revolution in Ole@ Polymerization J.C. Stevens

11

Characterizationand Chemical Design of Oxide Surfaces Y . Iwasawa

21

Photocatalysis:State of the Art and Perspectives K.I. Zamaraev

35

Towards Molecular Design of Solid Catalysts 51

A. Baiker

40th Anniversary Lectures

R- 1

R-2

A Retrospective View of Advances in Heterogeneous Catalysis: 1956-1996,Science R.L. Bunvell, Jr.

63

A Retrospective View of Advances in Heterogeneous Catalysis: 1956-1996 Technology H . Heinemann 69

Catalysis on Carbides, Nitrides, and Sulfides A- 1

An Example of Novel Basic Catalysts: The Aluminophosphate Oxynitrides or “AIPONs” A. Massinon, E. Gueguen, R. Conanec, R. Marchand, Y. Laurent, and P. Grange

The figures before the articles indicate the numbers used during the Conference

77

vi A-2

A-3

A- 5

A-6

Reaction Kinetics of the Hydrodenitrogenation of Decahydroquinoline over NiMo(P)/AI,O, Catalysts M. Jim and R. Prins

87

Effect of Spillover Hydrogen on Amorphous Hydrocracking Catalysts A.M. Stumbo, P. Grange, and B. Delmon

97

Hydrodesulfirization of Benzothiophene Catalyzed by Molybdenum Surfde Cluster Encapsulated into 5eolites M. Taniguchi, S. Yasuda, Y. Ishii, T. Murata, M. Hidai, and T. Tatsumi

107

Role of Adsorbed Hydrogen Species on Ruthenium and Molybdenum Suljdes. Characterization by Inelastic Neutron Scattering, Thermoanalysis Methods and Model Reactions M. Lacroix, H. Jobic, C. Dumonteil, P. Afanasiev, M. Breysse, and S. Kasztelan

117

General Papers A-7

A-8

A-9

A-10

Characterization of a Zeolite Membranefor Catalytic Membrane Reactor Application A. Giroir-Fendler, J. Peureux, H. Mouanega, and J.-A. Dalmon

127

Catalytic Reduction of SO, Stored in So, Transfer Catalysts - A TemperatureProgrammed Reaction Study G. Kim and M.V. Juskelis

137

Effect of Tunnel Structures of BaTi,O, and Na,Ti,O,, on PhotocatalyticActivity and Photoexcited Charge Separation M . Kohno, S . Ogura, K. Sato, and Y. Inoue

143

Organochromium Complexes in Homogeneous Olejn Polymerization G. Bhandari, J.L. Kersten, R.R. Kucharczyk, P.A. White, Y. Liang, and K.H. Theopold

153

Synthesis of Fine Chemicals A-1 1

A-I2

A-I3

A- 14

Selective Catalytic Oxidation with Air of Glycerol and Oxygenated Derivatives on Platinum Metals P. Fordham, R. Garcia, M. Besson, and P. Gallezot

161

Selective Methylation of Catechol: Catalyst Development and Characterisation L. Kiwi-Minsker, S. Porchet, P. Moeckli, R. Doepper, and A. Renken

171

Selective Oxidation with Copper Complexes incorporated in Molecular Sieves R. Raja and P. Ratnasamy

181

Discovering the Role of Au and KOAc in the Catalysis of VinylAcetate Synthesis W.D. Provine, P.L. Mills, and J.J. Lerou

191

vii A-15

A-16

A-17 A-18

A-19 A-20

A-2 1 A-22

A-23

Formation of Citraconic Anhydride by Vapor-Phase Decarboxy-Condensation of Pyruvic Acid M. Ai and K. Ohdan

201

Heterogeneous Enantioselective Dehydration of Butan-2-01 S . Feast, D. Bethell, P.C.B. Page, M.R.H. Siddiqui, D.J. Willock, G.J. Hutchings, F. King, and C.H. Rochester

21 1

Enantioselective Hydrogenation Catalysed by Palladium T.J. Hall, P. Johnston, W.A.H. Vermeer, S.R. Watson, and P.B. Wells

22 1

Stereochemical Studies of the Enantio-diferentiating Hydrogenation of Various Prochiral Ketones over Tartaric Acid-Mod$ed Nickel Catalyst T. Sugimura, T. Osawa, S. Nakagawa, T. Harada, and A. Tai

23 1

Enantio-diflerentiation over Heterogeneous Catalysts. The Shielding Eflect Model J.L. Margitfalvi, M. Hegedus, and E. Tfirst

24 1

Racemization of (IS)-(-)-exo-2,4-Dideuteroapopineneover Pd: Evidence for an Intramolecular I,3-Deuterium Shij G.V.Smith, B. Rihter, A. Zsigmond, F. Notheisz, and M. Bartok

25 1

Fe/MgO Catalysts for the Selective Hydrogenation of Nitriles G. Bond and F.S. Stone

257

Selective Synthesis of Ethylenediaminefrom Ethanolamine over Modified H-Mordenite Catalyst K.Segawa, S . Mizuno, M. Sugiura, and S. Nakata

267

Competitive Reaction Pathways in Propane Ammoxidation over V-Sb-Oxide Catalysts: An IR and Flow Reactor Study G. Centi and F. Marchi

277

Research on Model Catalysts

A-25

A-26

A-27

A-28

‘Seeing’the Active Site in Catalysis. STM and Molecular Beam Studies of Surface Reactions M. Bowker

287

Catalytic Formation of Carbon-Carbon Bonds in Ultrahigh Vacuum: Cyclotrimerization of Alkynes on Reduced TiO, Surfaces K.G.Pierce, V.S.Lusvardi, and M.A. Barteau

297

A New Approach to Understanding the Rochow Process: Synthesis of Methylchlorosilanesfrom CH, + C1 Monolayers on Cu,Si in Vacuum D.-H. Sun, A.B. Gurevich, L.J. Kaufman, B.E. Bent, A.P. Wright, and B.M. Naasz

307

Ruthenium as Catalystfor Ammonia Synthesis M. Muhler, F . Rosowski, 0.Hinrichsen, A. Homung, and G. Ertl

317

viii A-29

A-3 0

Surface-Structure-DependentReaction Pathways of Methyl Groups on Ni(lO0) and Ni(ll1) Surfaces R.B. Hall, M. Castro, C.M. Kim, and C.A. Mims

327

Normalization by oxygen Uptake of the Rates of Oxidative Dehydrogenation of Methanol and Ethanol S . Tanabe, H.E. Davis, Jr., D. Wei, and R.S. Weber

337

Environmental Catalysis A-3 1

A-32

A-33

A-34

Some Evidences of a Bifunctional Mechanismfor the Reduction of NO on Pd Based Catalysts H. Praliaud, A. Lemaire, J. Massardier, M. Prigent, and G. Mabilon

345

Characterization of Pd-based Automotive Catalysts R.W. McCabe and R.K. Usmen

355

Process Development for the Selective Hydrogenolysis of CCIF, (CFC-12) into CHZF2 (HFC-32) A. Wiersma, E.J.A.X. van de Sandt, M. Makkee, H. van Bekkum, and J.A. Moulijn

369

Catalytic Fluorination over Chromium Oxides. Preparation of Hydrofluorocarbons S. Brunet, B. Boussand, and J. Barrault

3 79

Oxidation of Methane A-35

A-36

A-37

Oxidative Coupling of Methane to Ethylene with 85% Yield in a Gas Recycle Electrocatalytic or Catalytic Reactor-Separator M. Makri, Y . Jiang, I.V. Yentekakis, and C.G. Vayenas

3 87

The Active w g e n for Direct Oxidation of Methane to Methanol in the Presence of Hydrogen Y . Wang and K. Otsuka

397

Photocatalytic Production of Methanol and Hydrogenfrom Methane and Water C.E. Taylor, R.P. Noceti, J.R. D’Este, and D.V. Martello

407

Oxidation Catalysis A-38

A-39

A-40

Role of A- and B-Cations in Catalytic Property of Substituted Hexaaluminate (ABAI,,O,, a) for High Temperature Combustion K. Eguchi, H . Inoue, K. Sekizawa, and H. Arai

417

AdMetal Oxidesfor Low Temperature CO Oxidation G. Srinivas, J . Wright, C.-S. Bai, and R. Cook

427

Photocatalytic Oxidation of Air Contaminants by Chlorine (CL!) or Hydroxyl (OH) Radicals or Holes @+): Mechanistic Correlations 0.d’Hennezel and D.F. Ollis

43 5

ix

A-4 1

A-42

A-43

A-44

A-45

A-46

A-4

A-24

Partial Oxidation of Methane to Synthesis Gas over Ru/TiO, Catalysts Y. Boucouvalas, Z.L. Zhang, A.M. Efstathiou, and X.E. Verykios

443

The Oxidative Transformation of Methane over the Nickel-Based Catalysts Modified by Alkali Metal Oxide and Rare Earth Metal Oxide Q. Miao, G. Xiong, S. Sheng, W. Cui, and X. Guo

453

Design of Stable Catalystsfor Methane-Carbon Dioxide Reforming J.A. Lercher, J.H. Bitter, W. Hally, W. Niessen, and K. Seshan

463

Comparison of Perovskite and Hexaaluminate-type Catalysts for CO/H,-fieled Gas Turbine Combustors C . Cristiani, G. Groppi, P. Forzatti, E. Tronconi, G. Busca, and M. Daturi

473

Hydrocarbon Activation and Oxidation on Transition Metal Mixed Oxides. Ft-IR and Flow Reactor Studies E. Finocchio, R.J. Willey, G. Ramis, G. Busca, and V. Lorenzelli

483

Biomimetic Oxidation on Fe Complexes in Zeolites (3.1. Panov, V.I. Sobolev, K.A. Dubkov, and A S . Kharitonov

493

A Molecular Approach to Synergy Generation in Co-Mo Binary Sulfide Catalysi‘s for Hydrodesulfirization Y. Okamoto and H. Katsuyama

503

Electrochemical Promotion of NO Reduction by CO and by Propene A. Palermo, M.S. Tikhov,N.C. Filkin, R.M. Lambert, I.V. Yentekakis, and C.G. Vayenas

5 13

Catalysis on Acids and Bases B- 1

B-2

B-3

B-4

B-5

B-6

Promotion of Molecular Hydrogen on Solid Acid Cracking Activity T . Shishido, T. Nagase, K. Higo, J. Tsuji, and H. Hattori

523

Selective Isomerization of Alkanes on Supported Tungsten Oxide Acids E. Iglesia, D.G. Barton, S.L. Soled, S. Miseo, J.E. Baumgartner, W.E. Gates, G.A. Fuentes, and G.D. Meitzner

533

Tungsta and Platinum-Tungsta Supported on Zirconia Catalysts for Alkane Isomerization G. Larsen, E. Lotero, and R.D. Parra

543

n-Butane Isomerization on Ni-promoted Sulfated Zirconia Catalysts W.E. Alvarez, H. Liu, E.A. Garcia, E.H. Rueda, A.J. Rouco, and D.E. Resasco

553

Rare Earth Modified Silica-Aluminas as Supportsfor Bifunctional Catalysis S.L. Soled, G. McVicker, S. Miseo, W. Gates, and J. Baumgartner

563

What NMR Has Told Us about Solid Acidity J.F. Haw and J.B. Nicholas

573

X

B-7

B-8

Novel Microporous Solid “Superacids”:CsjY3J W,,O,, (2sxs3) T. Okuhara, T. Nishimura, and M. Misono

581

Comparison of the Reactivities of H,P W,,O,, and HJi W,,O,, and their K’,NH,’ and Cs’ Salts in Liquid Phase Isobutane/Butene Allylation N. Essayem, S. Kieger, G. Coudurier, and J.C. Vedrine

59 1

B-9

Coupling of Alcohols to Ethers: the Dominance of the Surface SN2Reaction Pathway K. Klier, Q. Sun, O.C. Feeley, M. Johansson, and R.G. Herman 601

B-10

Characterization of Two Different Framework Titanium Sites and Quantijication of Extra-famework Species in TS-I Silicalites L. Le Noc, D. Trong On, S. Solomykina, B. Echchahed, F. Beland, C. Cartier dit Moulin, and L. Bonneviot

61 1

Environmental Catalysis

B-11

B-12 B-13 B-14

B-15 B-16

B-17

B-18

B-19

Injluence of Suljiu Dioxide on the Selective Catalytic Reduction of NO by Decane on Cu Catalysts F. Figueras, B. Coq, G. Mabilon, M. Prigent, and D. Tachon

62 1

Copt Clusters in NaMordenites as Catalystsfor SCR of No, L. Gutierrez, A. Ribotta, A. Boix, and J. Petunchi

63 1

Decomposition of Nitrous Oxide over ZSM-5 Catalysts F. Kapteijn, G. Mul, G. Marbin, J. Rodriguez-Mirasol, and J.A. Moulijn

64 1

On the Role of Free Radicals NO, and 0, in the Selective Catalytic Reduction (SCR) of No, with CH, over CoZSM-5 and HZSM-5 Zeolites D.B. Lukyanov, J.L. d’Itri, G. Sill, and W.K. Hall

65 1

An Infared Study of NO Reduction by CH,Over Co-ZSM-5 A.W. Aylor, L.J. Lobree, J.A. Reimer, and A.T. Bell

66 1

Precious Metal Loaded In/H-ZSM-5for Reduction of Nitric Oxide with Methane in the Presence of Water Vapor M. Ogura and E. Kikuchi

67 1

Interfacial RhOJCeO, Sites as Locationsfor Low Temperature N,O Dissociation J. Cunningham, J.N. Hickey, R. Cataluna, J.-C. Conesa, J. Soria, and A. Martinez-Arias

68 1

The Activity of VOJZrO, for the Selective Catalytic Reduction of NO V . Indovina, M. Occhiuzzi, P. Ciambelli, D. Sannino, G. Ghiotti, and F. Prinetto

69 1

Selective Reduction of NO, by Propene over Au/y-Al,O, Catalysts M.C. Kung, J.-H. Lee, A. Chu-Kung, and H.H. Kung

70 1

x1

B-20

The Role of Surface-Generated Gas-Phase Methyl Radicals in the Reduction of NO by CH, over a Sr/La,O, Catalyst S. Xie, T.H. Ballinger, M.P. Rosynek, and J.H. Lunsford

71 1

Part B Catalysis on Zeolites and Microporous Solids B-2 1

B-22

B-23

B-24

B-25

B-26

Isomerization and Hydrocracking of Decane and Heptadecane on Cubic and Hexagonal Faujasite Zeolites and Their Intergrowth Structures E.J.P. Feijen, J.A. Martens, and P.A. Jacobs

72 1

Solid State Ion Exchange of Alkali Metal Cations into Zeolite Y: Physicochemical Characterization and Catalytic Tests J. Weitkamp, S. Emst, M. Hunger, T. Roser, S. Huber, U.A. Schubert, P. Thomasson, and H. Knozinger

73 1

Role of Bronsted and Lewis Acidity in the Conversion of n-Pentane on Dealuminated H-Y, H-Mordenite and HZSM-5 V . Gruver, Y . Hong, A.G. Panov, and J.J. Fripiat

74 1

The Montmorillonite Catalyzed Production of Phenol and Acetone from Cumene Hydroperoxide W .A. de Groot, E.L.J. Coenen, B.F.M. Kuster, and G.B. Marin

75 1

Activation of Alcohols and Ketones by Surface Hydroxyls of Strong Solid Acids L. Kubelkova, J . Kotrla, J. Floriin, T. Bolom, J. Fraissard, L. Heeribout, and C. Doremieux-Morin

761

Heterogeneous Catalystsfor the Direct, Halide-fiee Carbonylation of Methanol B. Ellis, M.J. Howard, R.W. Joyner, K.N. Reddy, M.B. Padley, and W.J. Smith

77 1

B-27

Activation of Ethane on Modified ZSM-5 Zeolites Studied under Transient Conditions 78 1 A. Hagen, O.P. Keipert, and F. Roessner

B-28

Imaging of n-Hexane in Zeolites by Positron Emission Profiling (PEP) R.A. van Santen, B.G. Anderson, R.H. Cunningham, A.V.G. Mangnus, J. van Grondelle, and L.J. van IJzendoorn

79 1

Ethylene Dimerization in Nickel Containing MCM-41 and AIMCM-41 Studied by Electron Spin Resonance and Gas Chromatography M. Hartmann, A. Poppl, and L. Kevan

801

Alkane Partial Oxidation with Iron N,N'-bis(2-Pyridinecarboxamide)Complexes Encaged in Zeolite Y P.P. bops-Gerrits, M. L'abbe, W.H. Leung, A.-M. Van Bavel, G. Langouche, I. Bruynseraede, and P.A. Jacobs

81 1

B-29

B-30

xii Catalyst Characterization B-3 1

B-32

B-33

B-34

B-35 B-36

B-37

B-3 8

B-39

B-40

B-4 1

B-42

Physicochemical Characterization and Catalytic Properties of Solid Superacids Based on Surfated Zirconia Modified with Supported Platinum L.M. Kustov, T.V. Vasina, A.V. Ivanov, O.V. Masloboishchikova, E.G. Khelkovskaya-Sergeeva, and P. Zeuthen

82 1

Brensted Acid Strength of Solids Studied by 'HNMR: Estabhhing the Scale; Infruence of Lewis Acid Sites L. Heeribout, V. Semmer, P. Batamack, C. Doremieux-Morin, R. Vincent, and J. Fraissard

83 1

The Development of Strong Acidity by Non-Framework Aluminium in H-USY Determined by A1 XAFS Spectroscopy D.C. Koningsberger and J.T. Miller

84 1

Mesoporous MCM-41 Aluminosilicates as Model Silica-Alumina Catalysts: Spectroscopic Characterization of the Acidity F. Di Renzo, B. Chiche, F. Fajula, S. Viale, and E. Gmone

85 1

Elucidating the Nature of the Cobalt Centres in CoAPO-I8 Acid Catalysts L. Marchese, G. Martra, N. Damilano, S. Coluccia, and J.M. Thomas

86 1

Highly Dispersed Titanium Oxide on Silica: Preparation, Characterizationby XAFS, and Photocatalysis S . Yoshida, S. Takenaka, T. Tanaka, H. Hirano, and H. Hayashi

87 I

An Advance in Raman Studies of Catalysts: UltravioletResonance Raman Spectroscopy C. Li and P.C. Stair

88 1

The Dynamics of Surface-Catalyzed Reactions Studied by Infrared Chemiluminescence of the CO and CO, Products K. Watanabe, H . Uetsuka, H. Ohnuma, and K. Kunimori

89 1

Electron Availability and the Surface Fermi Level Local Density of States. An Alternative Way to See CatalyticActivity of Metals Y.Y. Tong, A.J. Renouprez, G.A. Martin, and J.J. van der Klink

90 1

A New Characterization Method for Adsorbed Hydrogen on Supported Pt Particles K. Asakura, T. Kubota, N. Ichikuni, and Y. Iwasawa

91 1

Hydrogen Chemisorption and Mobility on RdSiO,, WRdSiO, and Ru-Ag/SiO, R.L. Narayan, N . Savargaonkar, M. Pruski, and T.S. King

92 1

The Interaction of Rh with Intrinsic Defects of Reduced RwCeO, Catalysts: A Comparative XPS/UPS and 'H-NMR Study A. Pfau, J. Sanz, K.D. Schierbaum, W. Gopel, J.P. Belzunegui, and J.M. Rojo

93 1

...

XU1

B-43

In Situ Characterizationof the VanadiumSilicalite Catalyst (73-2) and its Photocatalytic Reactivity M. Anpo, S.G. Zhang, and H. Yamashita

94 1

Catalyst Deactivation B-44

Promotion and Deactivation of V,O,/TiO, SCR Catalysts by SO, at Low Temperature W.S. Kijlstra, N.J. Komen, A. Andreini, E.K. Poels, and A. Bliek 95 1

B-45

Structural Aspects of Activation and Deactivation of Cobalt Catalysts in Hydrogenation of Carbon Dioxide B. Klingenberg, G. Frohlich, F. Grellner, U. Kestel, G. Meyer, M. VoB, D. Bor,gnann, G. Wedler, J. Lojewska, T. Lojewski, and R. Dziembaj

96 1

Nature of H-Mordenite Deactivation Phenomena During SCR of No, M . Lezcano, A. Ribotta, E. Miro, E. Lombardo, J. Petunchi, C. Moreaux, and J.M. Dereppe

97 1

B-46

Oxidation Catalysis

c-1 c-2

c-3

c-4

c-5

C-6

c-I C-8

Eflects of Cs and V on Heteropolyacid Catalysts in Methacrolein Oxidation L.M. DeuBer, J.W. Gaube, F.-G. Martin, and H. Hibst

98 1

The Oxidation of C, Molecules on VanadylPyrophosphate Catalysts V.V. Guliants, J.B. Benziger, and S . Sundaresan

99 1

Direct Oxidation of Isobutane into Methacrylic Acid over Cs, Ni, and V-substitutedH,PMo,,O,, Heteropoly Compounds N . Mizuno, W. Han, T. Kudo, and M. Iwamoto

1001

Role of Water on the Properties of FePO Catalysts used for the Oxidative Dehydrogenation of Isobutyric Acid J.M.M. Millet, M. Forissier, D. Rouzies, P. Bonnet, and J.C. Vedrine

1011

Gas-Phase 0, Oxidation of Alkylaromatics with Ag'-Doped CVD Fe/Mo/DBH J.S. Yo0 and C. Choi-Feng

1021

Mechanistic Approach of the Oxidative Dehydrogenation of Propane over VMgO Catalysts by in situ Spectroscopic and Kinetic Techniques A. Pantazidis and C. Mirodatos

1029

Mechanochemistry in Preparation and Modification of Vanadium Catalysts V.A. Zazhigalov, J. Haber, J. Stoch, L.V. Bogutskaya, and I.V. Bacherikova

1039

Oxidative Dehydrogenation of Propane on Rare Earth Vanadates. Influence of the Presence of CO, in the Feed B. Zhaorigetu, R. Kieffer, and J.-P. Hindermann

1049

xiv C-9

C- 10

Structure Sensitivity of Oscillating Partial Oxidation of Propene on Widely Dispersed Platinum Catalysts M. Kobayashi, T. Kanno, H. Takeda, and S. Fujisak

1059

Catalytic Oxidation of Propane to Aclylic Acid with Molecular Oxygen Activated over Reduced Heteropolymolybdates W. Ueda, Y . Suzuki, W. Lee, and S. Imaoka

1065

Catalysis on Metals C- 1 1

C- 12

C- 13

Ru, Pt and Co Clusters in Zeolite Micropores; EX4FS/FTIWTPD Characterizationand Catalytic Behaviors in Methane Homologation M. Ichikawa, T . Tanaka, W. Pan, T. Ohtani, R. Ohnishi, and T. Shido

1075

The Relation between Pre-treatment of Promoted Copper Catalysts and their Activity in Hydrogenation Reactions D.S. Brands, E.K. Poels, and A. Bliek

1085

Cobalt Intermetallic Compoundsfor Selective Hydrogenation of Acetylene T . Komatsu, M. Fukui, and T. Yashima

1095

C- 14

Supported Pd-Cu Catalysts Preparedfrom Bimetallic Organo-metallicComplexes: Relation Between Surface CompositionMeasured by Ion Scattering and Reactivity A.J. Renouprez, K. Lebas, G Bergeret, J.L. Rousset, and P. Delichkre 1105

C- 15

Synergetic Effect of Pd and Ag Dispersed on MgO in the Reduction of NO by H2 at Room Temperature S. Naito and Y. Tanaka

1115

Heterogeneous Palladium Catalystsfor the Oxidation of Propylene to Propylene Glycol Acetates: Effect of Platinum and Rhodium as Promoters E.V. Gusevskaya, V.A. Likholobov, A.V. Karandin, A.I. Boronin, and E.M. Moroz

1125

Probing the Limits of Strucutre Insensitivity: Size-dependent Catalytic Activity of Al,O,-supported Iridium Clusters and Particlesfor Toluene Hydrogenation F.-S. Xiao, W.A. Weber, 0.Alexeev, and B.C. Gates

1135

The Kinetic Isotope Effectfor Alkane Dehydrocyclization B. Shi and B.H. Davis

1145

High-Thiotolerant Pt-Ge/Al,O, Naphtha Reforming Catalysts by in-situ Alloying T.F. Garetto, A. Borgna, and C.A. Apesteguia

1155

Metal-Support Interactions in Supported Platinum Catalysts: Zeolites and Amorphous Supports B.L. Mojet, M.J. Kappers, J.T. Miller, and D.C. Koningsberger

1165

The Role of Gaseous and Surface Isocyanates in the Reaction of Mixtures of H,, NO and CO over Supported Platinum Catalysts R. Diimpelmann, N.W. Cant, and A.D. Cowan

1175

C- 16

C-I7

C- 1 8

C- 19

C-20

C-2 1

xv

c-22

C-23

Reactions of Isobutane and Isobutylene over Silica- and L-Zeolite-Supported Pt/Sn and Pt/Sn/K Catalysts R.D. Cortright, E. Bergene, P. Levin, M. Natal-Santiago, and J.A. Dumesic

1185

Effect of Tungsten on Supported Platinum Catalysts J.L. Contreras and G.A. Fuentes

1195

General Papers C-24

C-25

C-26

On the Role of Microstructure of VanadiumPhosphorus Oxidesfor Propane Oxidation to Acrylic Acid N . Gribot-Perrin, J.-C. Volta, A. Burrows, C. Kiely, and M. Gubelmann-Bonneau

1205

Hydroconversion of Heavy Crude Oils Using Soluble Metallic Compounds in the Presence of Hydrogen or Methane A. Morales, A. Salazar, C. Ovalles, and E. Filgueiras

1215

Hydrodenitrogenation of Indole over NiMo Sulfide Catalysts L. Zhang and U.S. Ozkan

1223

Application of Theoretical Methods in Catalysis C-27

C-28

C-29

C-30

Quantum-chemicalStudy of the Nonclassical Carbonium Ion-like Transition States in Isobutane Cracking on Zeolites V.B. Kazansky, M.V. Frash, and R.A. van Santen

1233

Chemisw of Su1f.r Oxides on TransitionMetal Surfaces: A Bond Order Conservation - Morse Potential Modeling Perspective H . Sellers and E. Shustorovich

1243

Chemisorption and Decomposition of C, and C, Hydrocarbons on a Pd(l1 I ) Surface: A Periodic Density Functional Study J.-F. Paul and P. Sautet

1253

Density Functional TheoT Studies of Zeolite Acidity and Reactivity J.B. Nicholas

1263

Catalyst Synthesis

c-3 1 Control of Bulk and Surface Compositionof Doped Sm,Sn,O- Pyrochlore.

C-32

Relation between Formation of 0-Ba-CI Grajiings and C,-Selectivity in the Oxidative Coupling of Methane A,-C. Roger, C. Petit, S. Libs, and A. Kiennemann

1273

The Direct Room-Temperature Synthesis of Ce0,-based Solid Solutions: A Novel Route to Catalysts with a High oxygen Storage/Transport Cupacity F. Zamar, A. Trovarelli, C. de Leitenburg, and G. Dolcetti

1283

XVi

c-33

c-34

c-3s

C-36

c-37 C-3 8

c-39 C-40

Studies on Supported Metal Oxide-OxideSupport Interactions (An Incorporation Model) Y . Chen, L. Dong, Y.S. Jin, B. Xu, and W. Ji

1293

Zeolite-ConfinedMn Complexes of Cyclic Amines: New Selective Catalysts for Hydrocarbon Oxidation D.E. De Vos, J. Meinershagen, and T. Bein

1303

Platinum Indium Bimetallic in Silicalite: Preparation, Characterizationand Use in the VinylcyclohexeneTransformation P. MCriaudeau, A. Thangaraj, and C. Naccache

1313

Promotion of y-Alumina Dissolution by Metal Ions During Impregnation. Thermal Stability of the Formed Coprecipitates J.-B. d’Espinose de la Caillerie and 0. Clause

1321

Supported Catalysts Based on Carbon Fibrils M.S. Hoogenraad, M.F. Onwezen, A.J. van Dillen, and J.W. Geus

1331

Molecular Sieve Ti-UTD-I:A Novel Oxidation Catalyst K.J. Balkus, Jr., A. Khanmamedova, A.G. Gabrielov, and S.I. Zones

1341

Metal Composite Membranes: Synthesis, Characterization and Reaction Studies K.L. Yeung, R. Aravind, J. Szegner, and A. Varma

1349

Influence of the Process Parameters on the Extrusion of Ceramic Catalysts D. Ballardini, L. Sighicelli, C. Orsenigo, L. Visconti, E. Tronconi, P. Forzatti, A. Bahamonde, E. Atanes, J.P. Gomez Martin, and F. Bregani

1359

Synthesis Gas Conversion

c-4 1

C-42

c-43 c-44

c-45

Study on Catalytic Synthesis of Methanolfrom Syngas via Methylformate in Heterogeneous “One-Pot” Reactions H. Zhang, H . Li, G. Lin, Y. Liu, and K.R.Tsai

1369

The Effect of Zinc Oxide in Raney Copper Catalysts on Methanol Synthesis, Water Gas Sh$, and Methanol Steam Reforming Reaction D. Wang, L. Ma, C.J. Jiang, D.L. Trimm, M.S. Wainwright, and D.H. Kim

1379

Model Studies of Methanol Synthesis on Copper Catalysts J. Nakamura, I . Nakamura, T. Uchijima, T. Watanabe, and T. Fujitani

1389

Key Reactionfor Formation of Isobutene over ZrO, and Isoprene over CeO, in CO Hydrogenation K. Maruya, M . Hara, J. Kondo, K.Domen, and T. Onishi

1401

Promotion Effects in Methanol Synthesis over MgO-SupportedFe-Ir Catalysts Preparedfrom Mixed-Metal Clusters S . Marengo, R. Psaro, C. Dossi, S. Calmotti, and R.Della Pergola

1411

xvii C-46

Nanoscale Atfrition During Activation of Precipitated Iron Fischer-Tropsch Catalysts: Implicationsfor Catalyst Design A.K. Datye, M.D. Shroff, Y. Jin, R.P. Brooks, J.A. Wilder, M.S. Harrington, A.G. Sault, and N.B. Jackson

1421

Author Index

143 1

Subject Index

1435

This Page Intentionally Left Blank

xix

PREFACE

It seems highly appropriate that the Eleventh International Congress on Catalysis be held in Baltimore, USA, less than 200 km from the birthplace of these quadrennial events that began in Philadelphia in 1956. Planning for this 40th Anniversary Meeting has been coordinated by Gary L. Hailer, with the support of the Organizing Committee comprised of John N. Armor, Alexis T. Bell, W. Curtis Conner, Jr., Dady B. Dadyburjor, W. Nicholas Delgass, Sergio Fuentes, Richard D. Gonzalez, W. Keith Hall, Joe W. Hightower, Enrique Iglesia, Leo E. Manzer, James Maselli, Daniel E. Resasco, Kathleen C. Taylor, M. Albert Vannice, and Bohdan Wojciechowski. The PROCEEDINGS contain 145 papers - 7 plenary lectures and 138 submitted papers selected for oral presentation. The plenary lectures include five overviews of vital research areas by highly respected researchers and two overviews of advances in the science and technology of catalysis made during the last 40 years. The first group explores the forces that drive innovation in catalysis, constrained geometry in metallocene olefin polymerization, characterization and design of oxide surfaces, photocatalysis, and factors required in the molecular design of catalysts. Two others are presented by researchers who attended the first ICC meeting 40 years ago and who have been substantive contributors to science and engineering developments that have occurred since then. The 138 submitted papers were selected in the following manner. From a total of 521 submitted two-page abstracts, 156 were identified by peer review and evaluation of the Program Committee to be expanded into 10-page (maximum) camera-ready manuscripts. Submitted manuscripts were then peer-reviewed by at least two experts in the field according to standards comparable to those used for archival journals. Diversity in country origin was also considered, and an attempt was made to minimize multiple publications for individual research groups. Consequently, the 138 papers included herein should be considered as peer-reviewed publications that represent the worldwide state-of-the-art in catalysis research. These PROCEEDINGS of the International Congress on Catalysis differ from those published previously in two important ways. First, the papers were published PRIOR to the meeting for distribution to all delegates who attended the meeting in Baltimore. Second, none of the discussion at the meeting is included. With publication costs skyrocketing, we have elected to abandon the tradition of including the discussion, realizing that in so doing a valuable part of the meeting will be lost forever to posterity. However, it does have the advantage of smaller size (only two volumes, < 1,600 pages) which should make the books more attractive to libraries and other repositories of research literature.

XX

Finally, I would like to thank my co-editor, Nick Delgass, and Gary Heller for assistance in processing the paper revisions and to specially acknowledge Alex Bell and Enrique Iglesia, chair and co-chair of the Program Committee, for their excellent and timely management of the difficult but crucial task of paper selection. I am also grateful to Drs. Huub Manten of Elsevier Science for the fantastic cooperation the company provided in getting these two volumes printed. Most important, I wish to thank all the authors of the 145 papers appearing in these PROCEEDINGS for their diligence in faithfully meeting the exceedingly short deadlines that were necessary to get the material into print prior to the meeting. Joe W. Hightower, Editor

Houston, TX (USA)

J.W. Hightower, W.N. Delgass, E. Iglesia and A.T. Bell (Eds.) 11th International Congress on Catalysis - 40th Anniversary

Studies in Surface Science and Catalysis, Vol. 101 9 1996 Elsevier Science B.V. All rights reserved.

DRIVING FORCES FOR I N N O V A T I O N IN APPLIED C A T A L Y S I S lan E. Maxwell Koninklijke/ShelI-Laboratorium, Amsterdam (Shell Research B.V.), P.O. Box 38000, 1030 BN, Amsterdam, Netherlands 1. INTRODUCTION

Recent governmental sponsored studies in both the US and Europe have recognized the vital role of catalytic technologies for sustainable economic growth in the future. For example, it has been estimated that in developed countries catalysis contributes directly and indirectly through processes and products to some 20-30% of GDP (Gross National Product). Furthermore, catalytic environmental technologies such as automobile exhaust catalysts and the selective catalytic reduction (SCR) DeNOx systems in power plants have already significantly contributed to the reduction of environmentally harmful emissions into the lower atmosphere. In addition, these studies have identified catalysis as not only being pervasive but also offering significant scope for further innovative development of new and improved technologies for environmentally acceptable processes and products in the future. The spectrum of process industries which are directly impacted by catalysis include for example, oil refining, natural gas conversion, petrochemicals, fine chemicals and pharmaceuticals. Environmental catalytic technologies also play an important role in emission control systems for power generation, fossil fuel driven transportation, oil refining and chemical industries. Catalytic technologies typically embrace a wide range of disciplines such as heterogeneous and homogeneous catalysis, materials science, process technology, reactor engineering, separation technology, surface science, computational chemistry and analytical chemistry (Figure 1). Innovation in this field is therefore very often achieved by lateral thinking across these different disciplines. This presentation will attempt to develop this theme further by means of examples from recent commercial successes and from this platform provide some guidelines for multi-disciplinary approaches at the academic and industrial interface to further enhanced innovation in catalytic technologies in the future. 2. OIL REFINING AND NATURAL GAS CONVERSION

The discipline of materials science, for example, has a major impact on innovation in catalysis. Developments in the field of porous solids have led to some

tal

Figure 1. Disciplines of Prime Importance to Catalytic Technologies new catalytic processes in the refining area based on novel shape selective microporous materials. Two such new processes which have recently been commercialized based on these types of materials include the selective isomerization of nbutene to iso-butene [1] (MTBE precursor) and iso-dewaxing of lubeoils [2]. Interestingly, both groups of industrial researchers (Shell and Chevron groups, respectively) involved in these developments combined the disciplines of computational chemistry, materials science andheterogeneous catalysis to gain an in-depth understanding of the relationships between the detailed topology of the micro-porous materials and the shape selective catalytic performance. In the case of n-butene isomerization it was demonstrated (Figure 2) that the ideal micro-pore topology led to retardation of the C8 dimer intermediate and that the catalyst based on the ferrierite structure was close to optimal in this respect [ 1 ]. For selective isodewaxing a one-dimensional pore structure which constrained the skeletal isomerization transition state and thereby minimized multiple branching such as the SAPO-1 1 structure was found to meet these criteria. Clearly, these are ideal systems in which to apply computational chemistry where the reactant and product molecules are relatively simple and the micro-porous structures are ordered and known in detail. Another recent new application of a microporous materials in oil refining is the use of zeolite beta as a solid acid system for paraffin alkylation [3]. This zeolite based catalyst, which is operated in a slurry phase reactor, also contains small amounts of Pt or Pd to facilitate catalyst regeneration. Although promising, this novel solid acid catalyst system, has not as yet been applied commercially.

80 O

E

I

70I

\\

FER

60

(3

13) t-

5O

TON

40

MFI

.>_ 30 .1-, m (~

cr

20F 10

~ 1

:

-

2

3

MOR

= 4

5

Position TMP vs smallest zeolite ring [Angstrom] Figure 2. Comparison of Calculated Diffusional barriers for 2,4,4-trimethyl-3pentane (TMP) in Various Zeolites and Molecular Sieves A non-acidic isomerization catalyst system has unexpectedly emerged from recent studies by French workers [4] in the area of Mo-oxycarbides. Although at an early stage of development, these new materials exhibit high selectivities for the isomerization of paraffins such as n-heptane. An alternative non-carbenium ion mechanistic route to achieve isomerization of higher alkanes could potentially overcome some of the limitations of conventional solid acid based catalyst systems. Novel combinations of heterogeneous catalysis, reactor technology and separation technologies have also led to major innovations. Examples include catalytic distillation which is now widely applied for the manufacture of MTBE with other potential applications under development [5]. Another example of multidisciplinary synergy in this context which was recently commercialised is the socalled Synsat process [6] developed jointly by the Criterion and Lummus companies for enhanced deep hydrogenation and desulphurization of diesel fuels. This process employs a multiple catalyst bed system in a single reactor shell with intermediate by-product gas removal and optional counter-current gas/liquid flow in the bottom catalyst bed (Figure 3). Government legislation related to aromatics and sulphur contents of diesel fuels has become more stringent and global in recent years such that the development of improved catalytic hydrotreating processes is most timely. Catalytic membranes, which combine the disciplines of heterogeneous catalysis, separation technology, materials science and reactor engineering, which have for some time held considerable promise now appear to be gradually emerging as viable technologies. Promising potential applications include propane to aromatics [7] and catalytic oxidation of methane to synthesis gas using air as the oxidant [8]. In the former example, Japanese workers [7] applied a Pd-alloy membrane reactor (PMR)

Fresh feed ...... =i Make-up H2 ~

CatalystA

Recycled ~

CatalystB

liquid

CatalystC

~ ~

Reactor tc p / 1

Vaporto vapor/liquid separation/ recycle

Make-up H 2 - ~ J

ppH2

Reactor bottom

Dieselproduct Figure 3. The Synsat Process for Deep Hydrotreating of Diesel Fuels 85 8

9 o >

-

7s

J

65

m'~

55 ....... 20

I

40

PMR,~,,'4 ,,'" s// sss

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

\~" I

60

I

80

1O0

Conversion of Propane[%] Cat. Ga-H-ZSM5

773K

Figure 4. Comparison of Propane Aromatization Performances of a Palladium Membrane Reactor (PMR) and a Conventional Reactor (CR) using a Ga-H-ZSM-5 Catalyst to shift the equilibrium for the dehydroaromatization of propane. This PMR system resulted in a significant improvement in the selectivity towards the desired mixed aromatic products (Figure 4). For the partial oxidation of methane a catalytic membrane system under development by researchers at the Argonne National laboratory [8] effectively

separates oxygen from air which is then passed through the membrane in an anionic form to react with methane in the presence of catalyst. High conversion and selectivity levels to synthesis gas have been claimed, although space velocities have not yet been published. Such new catalytic processes can potentially reduce the costs of synthesis gas production and therefore positively impact the overall economics of natural gas conversion technology. New developments in the field of ceramic foam monoliths could also potentially provide new catalytic process technology for the conversion of methane into synthesis gas. For example, workers at Minnesota University [9] have achieved high synthesis gas yields at both high temperatures and space velocities using rhodium supported on a ceramic foam. 3. CHEMICALS

New materials are also finding application in the area of catalysis related to the Chemicals industry. For example, microporous [10] materials which have titanium incorporated into the framework structure (e.g. so-called TS-1) show selective oxidation behaviour with aqueous hydrogen peroxide as oxidizing agent (Figure 5). Two processes based on these new catalytic materials have now been developed and commercialized by ENI. These include the selective oxidation of phenol to catechol and hydroquinone and the ammoxidation of cyclohexanone to ecaprolactam. It was soon recognized that the TS-1 system has limitations particularly due to the small pore system which imposed restrictions on the molecular size of the OH

OH

ArH

,

ArOH

OH R ArOH

R,

R ,~:::0 + R'

R

H OH

HON,~

Figure 5. Range of Selective Oxidation Reactions Catalyze by the TS-1 Zeolite System Using Aqueous H202 as Oxidizing Agent

reactant molecules. More recently therefore titanium has been incorporated into larger pore zeolites [11] (e.g. beta and ZSM-48) and even mesoporous structures such as MCM-41 [12]. These larger pore materials also enable more bulky molecules such organic hydroperoxides to be used as oxidising agents. This field is still growing rapidly and would appear to hold promise for the development of new and improved heterogeneous catalyst systems for selective oxidation reactions of value to the chemicals industry. Base catalysis is another area which has received a recent stimulus from developments in materials science and microporous solids in particular. The Merk company, for example, has developed a basic catalyst by supporting clusters of cesium oxide in a zeolite matrix [13] . This catalyst system has been developed to manufacture 4-methylthiazole from acetone and methylamine. Heteropolyacids are also beginning to emerge from academic laboratories and find commercial applications. Showa Denko, for example, claim to have a process [14] for the direct oxidation of ethylene to acetic acid employing a bifunctional Pt/heteropolyacid catalyst system. The potential synergies between the disciplines of homogeneous and heterogeneous catalysis have also long been recognized but progress in the past has generally been frustrated by intangible technical problems. Particularly challenging is the goal of immobilizing homogeneous catalyst systems onto solid supports without incurring catalyst loss by leaching under reaction conditions. A particularly elegant approach to this problem involves the immobilization of a metal complex in a thin film of polar solvent (e.g. water) which is adsorbed on a high surface area hydrophilic support (e.g. silica). Such a system has been successfully applied in the laboratory [15] to immobilize a homogeneous water soluble chiral hydrogenation catalyst based on ruthenium (Figure 6). Using this catalyst a high degree of enantioselectivity was achieved for an important hydrogenation step in the synthesis of (S)-naproxen (an anti-inflammatory drug). 4. EMISSION CONTROL

Monolithic structures, often based on ceramic materials, are increasingly being applied in catalysis. The initial major thrust of monoliths was in the area of automobile exhaust catalyst systems where they are now applied exclusively. The demand for improved performance of these emission control systems, particularly under high temperature conditions is driving new developments such as ceramic foam technologies. Catalytic combustion, particularly for application in gas turbines, is another emerging field of technology where the developments in monolithic structures will be of growing importance. The Osaka Gas and Kobe Steel companies [16] have jointly developed a catalytic monolith which operates up to 1300 ~ and has been tested in a 160 kW gas turbine. New ceramic materials based on Mnsubstituted hexa-aluminates provide the high temperature stability required of catalytic monoliths for these demanding applications.

Ar2 CI P~ / P/Ru.~ Ar2

CI

" • •

Ru-BINAP(SO3Na), J = m - NaO~SC,~'-

-

H2 MeO

~ MeO

CO2H (S) - naproxen

Figure 6. Immobilization of Chiral Ruthenium Hydrogenation Catalyst in a Thin Hydrophilic Film on a Porous Glass Support Examples of multi-disciplinary innovation can also be found in the field of environmental catalysis such as a newly developed catalyst system for exhaust emission control in lean burn automobiles. Japanese workers [ 1 7] have successfully merged the disciplines of catalysis, adsorption and process control to develop a socalled NOx-Storage-Reduction (NSR) lean burn emission control system. This NSR catalyst employs barium oxide as an adsorbent which stores NOx as a nitrate under lean burn conditions. The adsorbent is regenerated in a very short fuel rich cycle during which the released NOx is reduced to nitrogen over a conventional three-way catalyst. A process control system ensures for the correct cycle times and minimizes the effect on motor performance. 5. FUTURE CHALLENGES

The above examples should serve to reinforce the multi-disciplinarity of catalytic technologies. However, to further exploit the significant potential of catalysis for innovation and renewal across a broad range of industries multi-discplinary approaches to problem solving will be vital to success. This ingredient for success is, in general, recognized within industrial research laboratories where multidisciplinary project teams are commonly deployed. However, this approach is traditionally less common in academic laboratories which generally tend to be more narrowly focussed in terms of disciplinary skills. Another element of concern at this interface is the perceived gap between academic basic research and industrial applied research. The recent trend towards shorter term goals within industrial research laboratories has further exacerbated

()

.~

oo

Sectors

Enabling Technologies

Multi-sector

Emerging Technologies

Multi-sector

Figure 7. Programme Model for Proposed UK National Institute of Applied Catalysis this situation whereby discontinuities are perceived to exist between basic and applied catalysis. Both these factors, which are likely retarding innovation and the potential synergies between industrial and academic laboratories in the field of catalysis, have been termed the "innovation gap". In Europe, particularly in the UK and the Netherlands, this mismatch has been recognized and some government supported initiatives are currently in progress. In the Netherlands, for example, a type of "virtual" organization called NIOK has been established to foster multi-disciplinary inter-university linkages and to strengthen the relationships with industry. More recently the UK has also launched a similar initiative with the intention of forming an organization based on both "virtual" and "hard core" components involving both academia and industry. This proposed new UK organization is termed NIAK (National Institute of Applied Catalysis). This NIAK organizational model is directly aimed at closing what is perceived to be a substantial "innovation gap" between industry and academia in the UK. A programme model has been recently developed for NIAC (Figure 7) which contains three separate components defined as emerging, enabling and sectors. The emerging and enabling components are envisaged to contain elements of common interest to all the industrial members whereas the sectorial programmes will be much more specifically oriented towards the individual needs of each industrial sector. Thus, in order to fully realize the potential of catalytic technologies not only will this require technical innovation in multi-disciplinary teams but also the appropriate organizational structures which maximize the synergies between academic and industrial research. The countries which recognize this potential and provide the

appropriate stimuli will likely have leading positions in catalytic technologies in the future. ACKNOWLEDGEMENTS

The author is most grateful for the valuable discussions with many of his Shell research colleagues in the field of catalytic technologies. In particular, as relate to this paper, discussions with Dr. G. Boxhoorn, Dr. K. de Jong, Ir. J. Naber and Dr. W. Stork are gratefully acknowledged. REFERENCES

1. H.H. Mooiweer, K.P. de Jong, B. Kraushaar-Czarnetzki, W.H.J. Stork and B.C.H. Krutzen, "Zeolites nd Related Microporous Materials: State of The Art 1994", Studies in Surface Science and Catalysis, Elsevier, 1994, Eds. J. Weitkamp et al, p2327 2. S.J. Miller, Microporous Mater., 1994, 2, 239-449 3. CM.A.M. Mesters, D. Peferoen, J.P. Gilson, C. de Groot, P.T.M. van Brugge and K.P. de Jong, submitted for publication to Appied Catalysis 4. M. Ledoux, C. Pham-Huu, H. Dunlop and J. Guille, "Proceedings lOth International Congress on Catalysis, eds. G.L. Solymosi and P. Tetenyi, Akademiai Kiado, Budapest, 1993, B, 955-967 5. V.J. D'Amico, P.H.O. Dixon & B.A. Strain, Paper AM-89-44, NPRA meeting, 1989, San Francisco. 6. A.J. Suchanek, E. L. Granniss, Presentation, AM-95-40, NPRA, Annual Meeting, March 19-21, 1995, San Francisco 7. Shokubai (Catalyst), 36 (4), 1994, 246 8. Argonne Research Laboratory Report, 1995 to be published. 9. D.A. Hickman and L.D. Schmidt, Science, 1993, 259, p343 10. B. Notari, Stud. Sur. Sci. Catal., 1988, 37, 413-425 11. C.B. Khouw, C.B. Dartt, J.A. Labinger, M.E. Davis, J. Catal., 1994, 149, p195-205 12. J.M. Thomas, Nature, Ti-MCM-41 13. F.P. Gortsema, B. Beshty, J.J. Friedman, D. Matsumoto, J.J. Sharkey, G. Wildman, T.J. Blacklock, S.H. Pan, Presentation at the 14th Conference on Catalysis of Organic Reactions, Albuquerque, April 27-28, 1992 14. M. Otake, Chemtech 1995, September, p36-41 15. K. Wan and M.E. Davis, Nature, 370, (1994), 449 16. H. Sadamori, T. Tanioka and T. Matsuhisa, "Proceedings of the International Workshop on Catalytic Combustion", Ed., H. Arai, Catalysis Society of Japan, (1994), p 158 17. S. Matsumoto, Pre-prints 2nd Japn-EC Joint Workshop on the Frontiers of Catalytic Science and Technology, 1995, vol 1, p39-42

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J.W. Hightower, W.N. Delgass, E. Iglesia and A.T. Bell (Eds.) 11th International Congress on Catalysis - 40th Anniversary

Studies in Surface Science and Catalysis, Vol. 101 1996 Elsevier Science B.V.

11

C o n s t r a i n e d G e o m e t r y and O t h e r Single Site Metallocene Polyolefin Catalysts: A Revolution In Olefin P o l y m e r i z a t i o n James C. Stevens Polyolefins and Elastomers Research and Development Laboratories, The Dow Chemical Company, 2301 Brazosport Boulevard, Freeport, TX 77541 1. ABSTRACT The polyolefins industry is at a crossroads. A new generation of single-site catalyst technology promises to revolutionize this multi-billion pound per year industry. Single-site catalysts have recently moved from laboratory curiosities to commercial success. Newly developed single-site c a t a l y s t s allow u n p r e c e d e n t e d control of polymer molecular a r c h i t e c t u r e , which yields products having improved properties. The consistency and control of polymer structure is allowing new discoveries to be made in f u n d a m e n t a l polymer research. This paper will touch on all aspects of single-site catalyst technology, including m e t a l l o c e n e m e t a l complexes, a c t i v a t i n g c o c a t a l y s t s (e.g., alumoxanes), cationic catalysts, as well as single site polymers, focusing on recent developments at Dow Plastics. 2. INTRODUCTION The metal catalyzed production of polyolefins such as high density polyethylene (HDPE), linear low d e n s i t y polyethylene (LLDPE) a n d polypropylene (PP) has grown into an enormous industry. Heterogeneous transition metal catalysts are used for the vast majority of PE and all of the PP production. These catalysts fall generally within two broad classes. Most commercial PP is isotactic and is produced with a catalyst based on a combination of titanium chloride and alkylaluminum chlorides. 1 HDPE and LLDPE are produced with either a titanium catalyst or one based on chromium supported on silica. 2 Most commercial t i t a n i u m - b a s e d PE catalysts are supported on MgC12. One of the most exciting developments in the polyolefins industry in recent years has centered on the development of commercial homogeneous single-site catalysts. These single-site catalysts produce olefin polymers with properties that are different when compared with traditional thermoplastic polyolefins. Homogeneous single-site catalysts based on bis-cyclopentadienyl derivatives of titanium have been known since the 1950's, although the catalytic activity of these early catalysts was too low for commercial practicality. 3 The key discovery by K a m i n s k y and Sinn t h a t m e t h y l a l u m i n o x a n e (MAO, [MeA10]n ) i n conjunction with Cp2TiMe2 and Cp2ZrC12 afforded extremely active catalysts for

12 PE and atactic PP lead to the recent explosion of interest in single-site catalysts.4, 5 The most valuable feature of single-site catalysts is the ability to logically control the structure of the polymer from the design of the catalyst. The most commonly used families of single-site catalysts are based on metal complexes shown in Figure 1. The polymer that is produced with these catalysts is strongly influenced by the catalyst structure. Catalysts with structure 1, having C2v symmetry produce atactic PP. By adding a bridging group between the cyclopentadiene ligands, catalysts having chiral C2 symmetry such as 3 produce isotactic pp.6,7 Linking the cyclopentadienes together to give a catalyst having Cs symmetry as in 4 produces syndiotactic PP. These relationships have led to the development of catalysts which produce isotactic-b-atactic PP by rotation of a substituted indenyl catalyst through C2v and C2 symmetry.8

R

2

~MX2 ~'~R

TiX3

HDPE, LLDPE Atactic PP

Syndiotactic PS

Isotactic PP

1

2

3

MX 2

i

Syndiotactic PP 4

Isotactic-b-atactic PP 5

F i g u r e 1. Single-site catalysts for olefin polymerizations. The bis-cyclopentadienyl-based, or metallocene, single-site catalysts are generally activated with MAO in relatively large molar amounts. The catalytic activity increases with increasing A1 : M ratio. Typically, at least 500 - 1000 molar equivalents of aluminum are required for acceptable activity. The high levels of MAO are a problem commercially, due to the relatively high cost of MAO. In addition, very high levels of MAO leave a large amount of aluminumcontaining "ash" in the polymer which can affect the product properties. The

13 catalytic activity is also a function of the transition metal. In general, the order of activity is Zr > Hf > Ti. Polymer molecular weight is also a function of the transition metal, generally following the order Hf > Ti > Zr. The ability to control the polymer from the design of the catalyst, coupled with high catalytic efficiency has led to an explosion of commercial and academic interest in these catalysts. Exxon started up a 30 million lb/yr ethylene copolymer demonstration plant in 1991 using a bis-cyclopentadienyl zirconium catalyst of structure 1. The Dow Chemical Company (Dow) began operating a 125 million lb/yr ethylene/1-octene copolymer plant in 1993 and has since expanded production capacity to 375 million lb/yr. This paper will focus on the structure / property relationships of the catalysts used by Dow to produce single-site ethylene a-olefin copolymers.

3. C o m m ~

Polyethylene

Commercial polyethylene falls within three general classes, as shown in Figure 2. LDPE is a highly branched dendritic polymer containing a range of short and long-chain b r a n c h e s , which r e s u l t from v a r i o u s r a d i c a l recombination processes in a high temperature (250 - 300 ~ high pressure process of up to about 45,000 psi. The numerous long-chain branches impart high melt strength and excellent processability to the polymer. In contrast, linear HDPE and LLDPE are produced using coordination catalysts and are characterized by a linear backbone containing no long-chain branching. LLDPE is primarily produced as a copolymer of ethylene with C4-C8 a-olefins up to about 15 weight percent. As a result, the linear molecules impart good toughness and strength properties, but are relatively more difficult to process than LDPE. Processability can be improved by broadening the molecular weight distribution, which tends to increase the number of low molecular weight molecules. The increased processability is generally achieved at the expense of strength and other physical properties due to the reduced number of high molecular weight molecules. Narrow molecular weight linear polyethylenes are relatively difficult to process. Focusing on LLDPE, traditional chromium or titanium-based Ziegler/Natta catalysts produce a product with a broad distribution of individual polymer molecules, each of which contributes to the overall properties of the resin. It is believed that the active catalysts contain active sites of various oxidation states and coordination environments, each of which exhibits different rates of propagation, termination, and comonomer reactivity. 9 The catalyst sites which incorporate comonomer have higher rates of chain t e r m i n a t i o n , and consequently the polymer molecules which contain more comonomer are lower molecular weight. A significant fraction of the polymer contains little if any comonomer and is generally of high molecular weight. The "mixture" of polymer molecules which results represents a limitation of conventional heterogeneous polyolefin catalysts in that the ability to control the individual component polymer molecules is rather restricted.

14

radicals H2C---CH 2 high temperature

I~PE

Z/N Catalysts H2C---CH 2

n ttDPE

H2C---CH 2

R Z/N Catalysts

+

LLDPE F i g u r e 2. Classes of commercial polyethylene. Metallocene catalyzed LLDPE is characterized by polymer molecules t h a t are the result of a single active catalytic site. As a result, all of the polymer molecules can in theory be made with statistically the same comonomer distribution. The molecular weight distribution is quite narrow, in the range of about 2.0 Mw/Mn. As a result of the ability to tailor the individual polymer chains, material scientists now have the ability to produce targeted polymer species and control the properties of the resin to a high degree. Great advances in such areas as product strength, clarity, toughness, and melting behavior can be commercially realized using metallocene catalysts. Single site polymers having a uniform comonomer and molecular weight distribution produced with low efficiency v a n a d i u m catalysts have been commercially available from Mitsui (Tafmer resins) since the 1980's. However, these polymers are relatively expensive and the ability to tailor the product is limited.

15 4. Constrained Geometry Catalysts Dow has developed a new family of ethylene-based polyolefins using constrained geometry catalyst technology. The catalyst and process technology has been commercialized under the tradename INSITE TM. INSITE Technology utilizes a family of new constrained geometry catalysts t h a t allows the production of unique polyolefin polymers in a relatively low pressure solution process. An important feature of the solution process is the need to operate the polymerization reaction above the melting point of the polymer in order to keep the product in solution. Solution LLDPE processes generally run at very high ethylene conversion and with a short reactor residence time of only a few minutes. Unfortunately, metallocene catalysts such as structures 1-5 produce low molecular weight products under such conditions, due to relatively facile ~hydride elimination. Constrained geometry catalysts, on the other hand, allow for the production of high molecular weight ethylene copolymers in a high temperature solution process. The key catalyst features are shown in Structure 6. The catalysts are monocyclopentadienyl Group 4 complexes with a covalently attached amide donor ligand. The amide ligand stabilizes the metal electronically, while the short bridging group (B) has the effect of sterically opening up one side of the complex, producing a stable but highly open and reactive active site upon activation with a variety of cocatalysts.

Rn N.-MX2 ]

R M = Ti, Zr, H f B = SIR2, C2H4, etc. R = alkyl, aryl, etc. X = halide, CH3 S t r u c t u r e 6. General structure of Constrained Geometry Catalysts.

In general, the open nature of the catalytic site in the constrained geometry catalysts does not allow for much steric control of the polymerization reaction, and homopoly a-olefins are generally atactic, although it has been reported that a small degree of tacticity can be introduced in polypropylene using such catalysts by selection of the substituents and conducting the polymerization at r e l a t i v e l y low t e m p e r a t u r e s . 1~ The degree of tacticity obtained under commercially useful conditions is so low that the catalysts can be considered to be atactic. The s t e r i c a l l y u n e n c u m b e r e d c a t a l y s t active site allows the copolymerization of a wide variety of olefins with ethylene. Conventional heterogeneous Ziegler/Natta catalysts as well as most metallocene catalysts are much more reactive to ethylene than higher olefins. With constrained geometry catalysts, a-olefins such as propylene, butene, hexene, and octene are readily incorporated in large amounts. The kinetic reactivity ratio, rl, is approximately

16 4 for the copolymerization of ethylene with 1-octene, which is approximately 2 orders of m a g n i t u d e more reactive towards octene t h a n some MgC12-supported heterogeneous catalysts. In addition, non-traditional olefins such as s t y r e n e can be incorporated in high levels. Styrene / ethylene copolymers containing significant a m o u n t s of styrene and having a high molecular weight have not been available in the past, as conventional polyolefin catalysts will e i t h e r not copolymerize ethylene with styrene to any appreciable extent, or the molecular weight is too low to be u s e f u l .

5. C o ~ a t a l y s t s and Polymerization Behavior C o n s t r a i n e d geometry complexes and metallocenes in general r e q u i r e the addition of a cocatalyst in order to become catalytically active. When activated with a large excess of MAO, catalyst efficiencies between 150,000 and 750,000 g of polymer per g r a m of metal are obtained, depending on the reactor t e m p e r a t u r e , specific catalyst, MAO level and other process variables. In general, these efficiencies a r e lower t h a n can be o b t a i n e d w i t h b i s - c y c l o p e n t a d i e n y l m e t a l l o c e n e s a n d MAO. In solution p o l y m e r i z a t i o n s , however, h i g h Mw polymers are obtained, even at t e m p e r a t u r e s as high as 160 ~ with 1-octene as comonomer. The Mw decreases with i n c r e a s i n g t e m p e r a t u r e , as shown in Figure 3, as a result of increasing ~-hydride elimination to give u n s a t u r a t e d chain ends. The Mw is sufficiently high u n d e r solution conditions t h a t H2 can be used as a Mw control, giving two i n d e p e n d e n t controls over m o l e c u l a r weight. 200000 , .

150000-

D

100000 -

50000I

I

I

I

0

I

I

0 t""

Reactor T, ~ F i g u r e 3. Mw data for ethylene 1-octene copolymerization. Catalyst = [(C5Me4)SiMe2N(t-Bu)TiC12 / MAO, 450 psi ethylene, 10 minute reaction time. 11

17 The open nature of the catalytic site allows for the incorporation of extremely high levels of comonomer, producing elastomers with over 20 weight % 1-octene comonomer. Increasing levels of comonomer depress the density, leading to ultra-low density elastomers. High Mw elastomeric ethylene/octene resins with densities between 0.87 and 0.85 g/mL can be obtained with high efficiencies. Figure 4 shows the relationship between density and weight % 1octene, determined using 13C NMR for such elastomers. Prior to the advent of metallocene catalysis, such extremely low density copolymers were not commercially accessible at low cost.

0.88

0.87

[]

[]

0.86

0.85 30

40

50

60

Weight % Octene

Figure 4. Ethylene-co-l-octene density as a function of 1-octene content for elastomeric copolymers. Comonomer incorporation, molecular weight, and catalytic efficiency are sensitive to the nature of the group bridging the Cp ring and the substituent on the amide ligand. Shorter bridges constrain the cyclopentadienyl ligand and amide group to adopt a particularly open and reactive catalytic environment. 12 Unlike the bis-Cp metallocenes, the titanium constrained geometry catalysts generally show the highest activity, comonomer incorporation, and Mw. While the constrained geometry catalysts exhibit many unique properties, the catalytic efficiency using MAO cocatalyst is relatively low for commercial applications, on the order of 104 - 10 5 g polymer per g of Ti. In contrast, a variety of cationic constrained geometry catalysts can be prepared which show extremely high activity, exceeding 10 7 grams of polymer per gram of transition metal. Cationic catalysts can be prepared with ammonium salts (Figure 5a) 13, oxidation of a corresponding Ti(III) complex (Figure 5b) 14, or abstraction of a hydrocarbyl group using B(C6F5)3 (Figure 5c). 15

18

[R3NH] [B(C6F5)4] A)

e2SiNN~T1Me2

NaN ~

-

CH 4

t_B/

[Cp2Fe][B(C6F5)4] ...... ~

~

t.Bu]

.

~ 1 ,CH Me2Si\ iTi'" \2

Me2SiXNITi.

[B(CGF5)4] G Me

/

B)

t-Bu~ M e 2 ~ N ~

\

C)

~ ~ " Me2Si\NIT1Me 2 t'Bu]

/

Cp2Fe Me2Si( ~~.~,,,C.H2 [B(C6F5)47

N-'"[

B(C6F5)3 S. ~ / ~ ~ ~ Me2 1\NIT1 ~ t-Bu/

| [CH3B(C6F5)3]

Me

Figure 5. Formation of Cationic constrained geometry catalysts: The use of B(C6F5)3 as a cocatalyst is particularly useful for solution polymerization, as this cocatalyst is soluble in the hydrocarbon polymerization solvent. The polymers produced in a continuous solution polymerization using these constrained geometry catalysts possess the expected properties of narrow molecular weight distribution and uniform comonomer distribution across the entire molecular weight range. In general, a narrow molecular weight and comonomer distribution would be expected to improve physical properties at the expense of processability. The polyolefins produced using INSITE Technology in a continuous solution polymerization process at high temperatures and high ethylene conversion give a polymer with high shear sensitivity, low melt fracture, high melt strength, and easy processability. The unusual properties of these INSITE polyolefins is the result of small but significant levels of longchain branching in an otherwise linear molecule. 16 These long-chain branches are postulated to result from the reincorporation of vinyl-terminated polymer molecules according to Scheme 1. The conditions of high reactor temperature

19 which leads to high vinyl termination, high comonomer reactivity, and high conversion with the concomitant high polymer concentration and low ethylene and comonomer concentrations in the continuous solution process produces the conditions favorable to long-chain branch formation.

Ti--CH2CH2--polymer A

Ti

+

Polymer--CH-CH 2 +

TimH

+

Polymer--CH--CH 2

monomers

PolymermCH--Polymer I

Polymer S c h e m e 1. Mechanism for formation of long-chain branching: Ti = active constrained geometry catalyst. 6. Conclusions Metallocene catalysts allow for the targeted control of polymer molecular structure to a degree which has not been possible previously. Rational structure-property relationships of these homogeneous catalysts are allowing new polymers to be produced which are finding large commercial markets. The high catalytic productivity possible with metallocene catalysts enable the production of polymers at competitive prices, even though the catalysts and cocatalysts are complex and relatively expensive. For ethylene/a-olefin copolymers, constrained geometry catalysts allow the production of a unique family of olefinic polymers. The proper selection of the metal, bridging group, and other substituents allows the control of product properties in a high temperature solution process. With the proper selection of catalyst variables, products ranging from high molecular weight elastomers to high density polyethylene can be produced.

References

1. J. Boor, Jr., Ziegler-Natta Catalysts and Polymerizations, Academic Press, 1979, New York, p. 108-129. 2. Ibid, p. 8.

20 3. D. S. Breslow and N. R. Newberg, J. Am. Chem. Soc., 1959, 81, 81. 4. H. Sinn,and W. Kaminsky, Adv. Organomet. Chem. 1980, 18, 99. 5. W. Kaminsky, M. Miri, H. Sinn, R. Woldt, Makromol. Chem., Rapid Commun. 1983, 4, 417. 6. o

8. o

J.A. Ewen, J. Am. Chem. Soc. 1984, 106, 6355. W. Kaminsky, K. Kiilper, H. H. Brintzinger, F. R. W. P. Wild, Angew. Chem., Int. Ed. Engl. 1985, 24, 507. G.W. Coates and R.M. Waymouth, Science 1995, 267, 217-218. J. Boor, Jr., Ziegler-Natta Catalysts and Polymerizations, Academic Press, 1979, New York, p. 262-269.

10. J. A. Canich, U.S. Patent 5,026,798 (1991). 11. J. C. Stevens, et al., European patent application 416,815 (1991). 12. J. C. Stevens, Stud. Surf. Sci. and Catal. 1994, 89, 277-284. 13. J. C. Stevens and D. R. Neithamer, US patents 5,064,802 (1991); 5,132,380 (1992). 14. R. E. LaPointe, et al., US patent 5,189,192 (1993). 15. R. E. LaPointe, et al., European patent application 520,732 (1991). 16. S.Y. Lai, et al., U. S. patent 5,272,236 (1993); 5,278,272 (1994).

j.w. Hightower, W.N. Delgass, E. Iglesia and A.T. Bell (Eds.) 11th International Congress on Catalysis - 40th Anniversary

Studies in Surface Science and Catalysis, Vol. 101 9 1996 Elsevier Science B.V. All rights reserved.

21

C h a r a c t e r i z a t i o n and C h e m i c a l Design of O x i d e Surfaces Yasuhiro Iwasawa Department of Chemistry, Graduate School of Science, The University of Tokyo, Hongo, Bunkyo-ku, Tokyo 113, Japan

This paper attempts to review recent and representative work dealing with single crystals, flat surfaces, and even designed surfaces relevant to oxide catalysis. These surfaces can provide novel information on the key issues in catalytic research such as the structure and composition of active sites, ensembles and phases, behavior of adsorbed active species, electronic property participating in catalysis, etc. which are well characterized by recent in-situ spectroscopy and also by traditional spectroscopy. The paper is also devoted to selected topics in the field describing the detail. Organized assembly of the knowledge of characterizations integrated over both molecular catalytic chemistry of powder catalysts and catalytic surface science of model surfaces would allow us to move toward the ultimate goal of rational catalyst design.

1. I N T R O D U C T I O N Metal oxides find application in a variety of technologies where surface science is critical to success, including catalysis, gas sensors, photoelectrolysis, electronic ceramics, semiconductor devices, pigments, cosmetics, etc. Metal oxides are tractable materials not only for spectroscopic techniques such as FT-IR, Raman, X-ray absorption fine structure(XAFS), etc., but also for techniques that might be disrupted by charging effects, including XPS, scanning tunneling microscopy(STM), etc., while special devices for sample preparation and careful interpretation of the spectra are essential. More critical is the tendency of many oxide surfaces to undergo thermal fracture, reconstruction, and particularly faceting[ 1]. These phenomena arise mainly from the need for charge balancing and minimization of surface polarity and energy, which may be relevant to the properties of oxide catalysts pretreated at different temperatures and ambient conditions. Understanding and controlling oxide surfaces are the key issues for the development of industrial oxide catalysts, but oxide surfaces are in general heterogeneous and complicated, and hence have been little studied so as to put them on a scientific basis by traditional approaches. While studies of the structure of surfaces have focused on metals and semiconductors over the past thirty years, the application of surface science techniques to metal oxides has blossomed only within the last decade[ 1-3]. An important future goal of catalytic surface science is to monitor the structure of surfaces and adsorbates at the molecular level in situ under catalytic reaction conditions, to model the more complex technical catalysts, and to undertake the design and tuning of new catalyst surfaces.

22 2. ACID-BASE AND REDOX P R O P E R T I E S OF MODEL S U R F A C E S

Acid-base reactivity is an important property of oxide catalysts, and its control is of interest in surface chemistry as well as being of importance in industrial applications. The exposed cations and anions on oxide surfaces have long been described as acid-base pairs. The polar planes of ZnO showed dissociative adsorption and subsequent decomposition of methanol and formic acid related with their surface acid-base properties[3]. Further examples related to the topic of acidbase properties have been accumulated to date[ 1,4-6]. In contrast to the extensive studies of heterogeneous acidic oxides, less effort has been given to the study of heterogeneous basic oxides. The first study of heterogeneous basic catalysts, in which sodium metal dispersed on alumina acted as an effective catalyst for double bond migration of alkenes, was reported by Pines and Ipatieff[7]. Now, a number of materials have been reported to act as heterogeneous basic catalysts; alkaline earth oxides, alkali metal oxides, rare earth oxides, ZaO2, ZnO and TiO2, alkali ion-exchanged zeolites, alkali metal ions on oxides, hydrotalcite, chrysotile, sepiolite, KF supported on alumina, etc[8,9]. A superbasic catalyst(~,-Al203-NaOH-Na) is prepared by addition of NaOH to alumina followed by further addition of Na. The resulting catalyst which possesses basic sites stronger than H_=37 and a distorted 13-Na~O2 phase at the catalyst surface as characterized by solid NMR and XPS, is industrially employed, with nearly 100% yield in the commercial plant for synthesis of 5ethylidene-2-norbomene (additive to ethene-propene copolymer rubber) from 5-vinyl-2norbornene that is obtained from dicyclopentadiene[10]. Considering the tendency of Na to donate electrons, it seems natural that Na dispersed on alumina acts as a heterogeneous basic catalyst. It is, however, found that the basic property of the oxide surface is not naturally proportional to the quantity of Na deposited on the surface[ 11 ]. Atom-resolved STM has visualized the mechanism producing structural sensitivity in the reaction of CO2, a reaction probe at basic sites, with the Na-deposited "I~O2(110) surface[ 12]. The amount of adsorption to form carbonates varies with Na coverage exhibiting an S-shaped dependence. An STM image of the 0.1 ML-Na deposited TiO2(l 10) surface depicts a dispersed geometry due to repulsive forces between the Na atoms. The surface is converted to a nearly complete c(4x2) overlayer via p(4x2) order locally formed on the surface by increasing Na coverage. The Na atoms in the c(4x2) surface are reactive to CO2 as shown in the STM topography in which the c(4x2) order disappeared and the chains for CO32 were locally ordered in a p(3x2) symmetry which appeared along the [001] direction. In contrast, CO2 does not adsorb on a surface with randomly dispersed Na atoms. The genesis of strongly basic sites is thus suggested not to be linearly correlated with Na quantity, but to be correlated with the ordered structure or at least suitable ensembles of Na. Recently, it has been demonstrated that coordination vacancies on the surface metal cations are relevant to the unique redox reactivity of oxide surfaces[2]. Oxidation of formaldehyde and methyl formate to adsorbed formate intermediates on ZnO(0001) and reductive C-C coupling of aliphatic and aromatic aldehydes and cyclic ketones on "HO2(001) surfaces reduced by Ar § bombardment are observed in temperature-programmed desorption(TPD). The thermally reduced "HO2(110) surface which is a less heavily damaged surface than that obtained by bombardment and contains Ti cations in the +3 and +4 states, still shows activity for the reductive coupling of formaldehyde to form ethene[ 13]. Interestingly, the catalytic cyclotrimerization of alkynes on TiO2(100) is also traced in UHV conditions, where cation coordination and oxidation states appear to be closely linked to activity and selectivity. The nonpolar Cu20(111) surface shows a

23 maximum selectivity for complete oxidative dehydrogenation of CH3OD to CO, while the polar Cu+-terminated (100) surface shows a maximum selectivity for partial dehydrogenation to CH20 [14]. Methanol on ZrO2(100) decomposes near 630 K to produce CO and CH4, whereas on the (110) surface the primary methoxide decomposition pathway is oxidation to produce CH20. This difference in reactivity can be related to the local atomic structure of each surface. Chemical and catalytic trends in oxide surfaces where metal atoms are isolated by oxide ligands resemble those observed in homogeneous metal oxo-complexes except for surface phenomena such as surface restructuring and transformation, diffusion of atoms and adsorbed molecules, etc[2]. 3. C H E M I C A L T U N I N G OF A C T I V E S I T E S The key properties of oxide surfaces are the coordination environment, oxidation state and acidic or redox properties of surface cations, and the basicity of surface anions. The longer term challenge to oxide surface science is to address important issues in selectivity in catalytic oxidation and acid-base reactions, in particular the principle of "tuning of metal reactivity by oxide ligands". There are a number of clear-cut examples in catalysis by mixed oxides, including the selective oxidation of propane to acrylonitrile with V-Sb oxide[15-17], Bi-V-Mo oxide [18,19], Mol.0V0.4Te0.2Nb0.104.65 oxide[20], and the selective oxidation of butane to maleic anhydride with VPO catalysts[21 ], where selectivity toward the desired products relies on the limited availability of oxygen at active ensembles of the metal oxide component and also on the properties and structure of defects, besides the nature of surface oxygen atoms. Notwithstanding the interesting industrial outlook, relatively few data exist either on the characterization of such catalysts or on the reaction kinetics and mechanisms. In the development of new catalysts, new chemical concepts regarding composition or structure are conceived. The requirements and design of quantitative ensemble sizes represent important but as yet unaddressed challenges to the field. Although the efforts on the design of excellent catalysts have been acutely difficult challenges, recently molecular-level catalyst preparation has become realistic on the basis o'f modem physical techniques and accumulated knowledge of oxide surfaces[22-25]. A Series of chemical designs of Nb structures on SiO2 is introduced as an example of a onecomponent tuning catalyst. Niobium has been considered to be a poor catalyst, but recently has attracted much attention as a key element for industrially important processes such as amrr,oxidation of propane[20], oxidative dehydrogenation of propane[26], etc. A guideline for the Nb-structure design is seen in a theoretical concept. We consider and extend this theoretical prediction to the case of ethanol oxidation, via a Nb-OC2H5 intermediate, on NbS+-oxide sites supported on a SiO2 surface[27]. In ~ C H elimination on dS-metal ethoxide complexes, the orbital interactions have to take place in such a manner that the electron donation from a(CH) to g*(MO) and the back-donation from g(MO) to c*(CH) are required to form the MH g and CO n bonds and break the CH a and MO c bonds. Also in case of d~ complexes, the presence of a weak M--H agostic interaction is predicted by the theoretical calculation, but the Ti-~CH angle is unfavorable for overlap of the occupied Ti d-orbital and the CH antibonding orbital. Furthermore, there is formally no d-electron available for the promotion of the CH bond scission. The predicted key factors are electron density of the d state and overlap of the two orbitals. The former boundary is satisfied by attaching Nb do ions to the SiO2 surface through Nb-O-Si bonding. The electronic structure of a distorted tetrahedral dioxo-Nb monomer on SiO2 calculated by the DV-Xct cluster method shows that the component of the Nb 4d orbitals is hybridized with the higher occupied O 2p levels, enabling the ~ C H breaking. The support electronically modifies the metal oxide

24

species through chemical or ionic bonds and induces the structural change of the surface metal oxides needed for catalysis. It predicts new catalysis involving I],-elimination of the CH bond by coordinatively unsaturated tetrahedral Nb monomers chemically attached to the SiO2 surface. The latter boundary of the orbital overlap seems to be less rigid for the tetrahedral Nb monomer structure. The SiO2-attached Nb monomer catalyst with a four-coordinate structure was prepared by the use of Nb013-C3H5)4 as precursor and characterized by extended x-ray absorption fine structure(EXAFS), FT-IR, Raman, ESR, and XPS. The monomer catalyst(l) exhibits high activity and selectivity for the dehydrogenation of ethanol to form acetaldehyde and H2 as shown in Table 1[27]. The activity is much higher than that of a usual impregnation Nb catalyst and the selectivity is as high as 95-100%, whereas the impregnation catalyst is less active and unselective. The dehydrogenation reaction proceeds via the Nb-ethoxide intermediate, but the intermediate is very stable and is not decomposed until 600 K. It is dehydrated to ethene and water above 600 K. On the other hand, when the Nb-ethoxide intermediate is exposed to ethanol, the dehydrogenation of ethanol proceeds at much lower temperatures such as 423 K, with ca. 100% selectivity. Thus the switchover of the reaction path from dehydration(~,-CH bond break) to dehydrogenation(l~-CH bond break) by the second ethanol molecule adsorbed on the Nb-ethoxide is observed. Even if an adsorbed species at the surface is too stable in vacuum or before catalysis, the catalytic reaction is able to proceed through the same species by activation of the intermediate by the reactant(reactant-promoted mechanism)[28]. Thus, it may be critical for the dehydrogenation on Nb sites to create a vacant site with an appropriate conformation for the transition state on which electron donation-induced activation of ~hydrogen of C2H50 group is favorable. , 0.307nm, ;~

, '
CH3CHO + H2PC) or a sequence of one-electron reactions, (e.g., C2HsOH + PC hv > CH3C'HOH + HPC,

C2HsOH +HPC

CHsC'HOH + PC hv , CH3CHO + HPC, CH3C'HOH + HPC

hv ; CH3C'HOH + H2PC, hv> CH3CHO + H2PC) [10]. .....

In contrast, step II is clearly a single two-electron reaction which proceeds with a rather low activation energy E a = 25 kJ/mol, but a high-negative activation entropy, ASa ~ - 200 J/K-mol. Note, that identical values were found for the rate constant kli of reaction II in four independent experiments, namely, by measuring: (1) - the rate of H2 evolution (with GC-method) in the overall photocatalytic reaction of Scheme 1, and (2), (3) and (4) - the rates of H 2 evolution (also using GC), disappearance of H2PC (using UV spectroscopy) and formation of HPC (using EPR) in the absence of illumination from/in solutions of PC that had

37 been prereduced to H2PC on Zn amalgam [10]. The value obtained for k n ~ 3.10 -4 sec 1 is rather small due to the high negative value for ASa. However, step II can be notably accelerated under the action of visible (including red) light. This light is known to excite the bands of intervalence charge transfer in reduced forms of HPAs, which makes the electron transfer from the HzPC frame to its internal protons more facile.

Photocatalysis and photogenerated catalysis with metal carbonyl, phosphine and other ligandso Excellent reviews of this field are given in [11-12]. Mechanistically, two types of photoinduced catalysis with the participation of transition metal complexes can be distinguished, namely, classical photocatalysis and photogenerated catalysis. The difference between these two types of catalysis are illustrated in Scheme 2. Scheme 2a presents a classical photocatalytic pathway for the water-gas-shift reaction in aqueous THF solutions (see [12] and refs. therein). For this pathway photon directly participates as a reagent in the catalytic cycle of olefine hydrogenation. Scheme 2b (see [11 ] and refs. therein) presents an example of photogenerated catalysis. Here photons do not participate in the catalytic cycle itself, but provide preparation of catalyst active form H2Fe(CO)3(olefine) from Fe(CO)5 precursor via photochemical dissociation of two CO ligands from the metal atom. CO CO2 _ hv ~,HFe(CO)4 ~ ~ HFe(CO)3 HFe(CO)4

--'72__

HFe(CO)4 + OH-

OH"

/

~1

"- 2Fe(CO)5 + Fe(CO)4 ~ . ~ hv

Fe(CO)5 ~

CO

H2Fe2(CO)7

2-~ Fe2(CO)7 2CO

H2

\~

Fe(CO4) ~ /

[I--Fe(CO)4 ~Qhv, H~ v~ CO H I II-Fe(CO)3

I Scheme 2b[

H H2 ~1-- Fe(CO)3

C3H8~

H I Fe(CO)3

~

\X/

38 Besides dissociation of ligands, photoexcitation of transition metal complexes can facilitate: (1) - oxidative addition to metal atoms of C-C, C-H, H-H, C-Hal, H-Si, C-O and CP moieties; (2) - reductive elimination reactions, forming C-C, C-H, H-H, C-Hal, Hal-Hal and H-Hal moieties; (3) - various rearrangements of atoms and chemical bonds in the coordination sphere of metal atoms, such as migratory insertion to C=C bonds, carbonyl and carbenes, c~- and [3-elimination, o~- and [3-cleavage of C-C bonds, coupling of various moieties and bonds, isomerizations, etc. (see [11, 12] and refs. therein). Such a broad range of classical elementary reactions of homogeneous catalysis with metal complexes, that can be facilitated by photons, make illumination of reaction solution a very useful instrument for substantial increase of the possibilities of homogeneous metal complex catalysis in organic synthesis. Particular examples of light-assisted syntheses will be given in section 3. Redox reactions with metal porphyrins (MPs) as photocatalysts. A spectacular example here is the reaction that couples upon illumination with the stmlight, methanol oxidation to formaldehyde with the formation of hydrogen peroxide in benzene-methanol mixture (90:10) hv, PC CH3OH + 02 > H2CO + H202 (2) in the presence of oxoalkoxoporphyrinatomolybdenum(V) as the photocatalyst [13]. Due to a large amount of energy provided by photons, some redox photocatalytic reaction can easily proceed at very low temperature. A characteristic example is the reaction between C C I 4 and triethylamine that,proceeds smoothly in the presence of porphyrinatozinc(II) and magnesium(II) photocatalysts even at 77 K in vitrified ethanol solutions (Scheme 3) [14]. Photoinduced electron transfer from Et3N to MP* and subsequent photoinduced electron transfer from M P to CC14 proceed at this temperature via electron tunneling, and the reaction products Et3N+ and CCI 3, Cl are separated in space [14].

Et 3 +

MP-

hv, CC14

[ Scheme3 ]

Et3N, h v "

\ MP

CC14

~- CC13 + C1-

2.2. Photocatalysis in organized molecular assemblies A prominent example here are photocatalytic systems based on lipid vesicles for watersplitting into H2 and 02. Design of such systems is based on both functional and structural mimicing of natural photosynthesis. Typically, photocatalytic cleavage of water in organized molecular assemblies (Scheme 4) is designed as a sequence of three key steps: charge separation, that is, formation of sufficiently strong oxidant (D § and reductant (A-) from intermediate electron acceptor A and electron donor D under the action of solar light in the presence of a photocatalyst PC, and subsequent catalytic reactions of oxygen evolution from water by oxidant D § and of hydrogen evolution by reductant A (see [6, 7, 15, 16] and refs. therein). The last two steps, to be

39

~ehv

Ill

I

AG~

eo

O

PC tpc+

T oxygen evolution

Q y charge

H20/H2 kJ + 89 = 237 mol

Q

Scheme 41

~r separation

hgdro0en evolution

accomplished, do not necessarily require light. The overall process is described by the equation 2 H 2 0 hv, PC, D ' A --, 2H 2 + 0 2 cat H2 ' cat 02

(3)

The major problem in accomplishing water splitting via the pathway of Scheme 4 is how to suppress the back recombination reaction D + + A ---> D + A, which is a simple exothermic bimolecular process and therefore typically proceeds much more rapidly than complex catalytic reactions of H2 and O2 evolution. An attractive way to overcome this problem is to use microheterogeneous photocatalytic systems based on lipid vesicles, i.e. microscopic spherical particles formed by closed lipid or surfactant bilayer membranes (Fig. 1) across which it is possible to perform vectorial photocatalytic electron transfer (PET). This leads to generation of energy-rich one-electron reductant A- and oxidant D+, separated by the membrane and, thus, unable to recombine. As a result of such PET reactions, the energy of photons is converted to the chemical energy of spatially separated one electron reductant and oxidant. C15H31 C15H31 O=C

C=O

I ~

t o

k -- H2C--- CH t O

i -

Clt2

i

O=P-OI

O

I CH2 I CH2 L N+

CH3 t CH3 CH3

Fig. 1. Structure of a lipid vesicle. As an example, let us consider a system, which contains a photocatalyst Ru(bipy)~§ in the inner volume of the vesicle, an electron carrier, e.g. cetylviologen (CI6V2+), in the membrane, and

40 an electron acceptor, Fe(CN)63- in the outer volume of the vesicle. Illumination into the PC absorption band induces electron transfer from Ru(bipy) 32. across the membrane to Fe(CN) 36 9The quanttma yield of the transmembrane electron transfer, % is related to the quanama yield of Ru(bipy) 33++ CI6V§ formation, % = 15%, as follows: q~ = q ) o k t / ( k t + kr) , w h e r e k t is the rate constant of electron transfer across the membrane by C16V§ radical and k, is the rate constant of recombination between Ru(bipy) 33. and Cl6V§ (see Fig. 2). The recombination and between Ru(bipy) 33, viologen radicals nearby the inner surface of the vesicle is fotmd to be strongly inhibited as compared to homogeneous solution, presumably due to a fast extraction of hydrophobic viologen radicals into the depth Ru(bipy)2+Gk;2+ ,P=,Po kt+k r of the membrane. Two mechanisms of transmembrane electron Fig.2. Mechanism of PET across a lipid membrane in the transfer were elucidated: (i) Ru(bipy) 32. _CV2+_Fe(CN) 63- system. via the translocation of viologen radical across membrane and (ii) via reaction of electron exchange between the radical and oxidized viologen located in the different monolayers of a bilayer membrane (see [16] and refs. therein). Note, that the dication C~6V2+ can not penetrate across the central core region of the membrane, presumably because of its unsufficient hydrophobicity. For various carriers of viologen type the rate constants k t and kr were measured, and the efficiencies b = k t / (k t + kr) of the transmembrane charge separation reaction were determined to be up to several percent. The rate constant k t was proved to be dependent on the substituent in viologen, the phase state (gel or liquid crystal) and the electrical polarization of vesicle membrane. Another promising way of transmembrane PET includes intramolecular electron transfer along bridge molecule D-PC-A which spans the bilayer and contains PC, D and A fragments linked by covalent bonds [17]. The membrane-separated reductant and oxidant formed upon PET can be used for accomplishment of various catalytic redox reactions which provide conversion of the chemical energy of a (D+...A-) pair into the chemical energy of a pair of more stable species such, e.g., as H2 and 02 molecules. This stored energy can be released when necessary in the form of high potential heat or electricity via combustion of H2 + 1/2 02 mixture in a furnace or fuel cell. It proves possible to anchor catalysts of H2 evolution to the outer and inner surface of the vesicle membrane. These catalysts are finely dispersed (10-20 A in diameter) metal Pt or Pd particles formed via reduction of appropriate salts in vesicle suspension (see [ 15, 16] and refs. therein). Among the viologen-type electron carriers a promising one is p-bis (1,2,5-triphenyl4-pyridil)benzene which possesses reduction potential low enough for water reduction at neutral pH. Recently, using this mediator we succeeded in H2 evolution conjugated with PET ,

41 across membrane in the "sacrificial" system which is schematically represented in Fig. 3 (V is an electron carrier, EDTA - ethylenediaminetetraacetic acid).

EDTAox

hv , Ru(bpy>~+C~ Ru(bpy)

"v"

EDTA~

Ru(bpy)3+

(Vv_lK::+

Fig.3. Mechanism of PET across a lipid membrane, coupled with H 2 evolution in EDTA-Ru(bipy) 32 _V2+_pd / lipid vesicle system. Photocatalyzed H2 evolution inside vesicle cavity over tiny anchored Pt particles also has been reported (see [15] and refs. therein). The membrane-bound catalyst for water oxidation to 02 can be prepared via oxidation of Mn(II) and Co(II) salts to Mn(IV) and Co(Ill) hydroxides, respectively, in the presence of lipid vesicles. Using these catalysts and photogenerated Ru(bipy)33. complex as an oxidant, it is possible to oxidize water to 02 in vesicle systems. One of such systems for 02 evolution is schematically represented in Fig. 4. Thus, vesicles serve as chemical 1 / 2 0 2 + 2H + hv | / 2 S 2 O a2" microreactors with transparent and electron-conducting walls. Molecular Ru(bpy)+ + engineering of such microreactors nowadays is one of the most fascinating areas of basic research in supramolecular photocatalysis. However, one should not SO~ underestimate the difficulties in OH 2 Ru(bpy)3+ designing vesicular catalytic reactors. The major problem comes from the Cat - (OH),Co.'" O+..z,; n = 100 fragility of the walls. Indeed, the so far available surfactant materials for the Fig.4. Photocatalytic $20 82--Ru(bipy) 32. walls are still not stable enough and - (OH)x-Co iu O(3n-x)n / lipid vesicle decompose rather easily upon n system for 02 evolution. illumination in the presence of strong

42 oxidants. Moreover, nonpolar small molecules, such, e.g., as 02 produced on cat o2 (see the notation in Scheme 4) attached to the outer surface of the membrane wall (Fig. 4), easily migrate through this wall. Contact with 02 destroys the system for H2 evolution on cat H2 attached to the inner wall of the microreactor. For this reason, the closed photocatalytic cycle of water splitting into H 2 and 02 has not so far been accomplished with vesicular systems, despite the fact that separately photocatalytic transmembrane electron transfer and H2 or 02 evolution, have been accomplished [ 15, 16]. Better isolation between the oxidative and reducfive parts of the photochemical system, perhaps, may be provided by solid layered structures, where two different interlayer compartments containing the catalysts for evolution of H 2 and 02, can be separated by a rigid layer of framework atoms [18]. It might also occur possible to avoid the inhibiting influence of 02 on the molecular apparatus for H 2 production, by means of evolving dihydrogen and dioxygen from water consecutively rather than simultaneously. In this case one may benefit from the technique [7], which provides the separation in time of the two steps: (i) photogeneration of a strong oxidant, which in this case must be conjugated with the simultaneous process of dihydrogen evolution, and (ii) subsequent oxidation of water with this oxidant. From what we know today about PET in biological and synthetic membrane or layered systems, we may expect that the non-biological apparatus providing photogeneration of spatially separated one-electron reductant and oxidant is likely to be developed in a rather universal way and may be expected to accomplish in the future not only water cleavage, but also various other redox reactions, such e.g., as photochemical synthesis of ammonia via the reaction N2 + 3H20 hv_~ 2NH3 + 3/202, photochemical synthesis of various organic compounds from CO2 and H20 (e.g., CO2 + 2H20 hv ~ CH3OH + 3/202 ), etc. Other examples of organized molecular assemblies of interest for photocatalysis are: (1) PC-A, PC-D or D-PC-A molecules where PC, A and D fragments are separated by rigid bridges; (2) host-guest complexes; (3) micelles and microemulsions; (4) surfactant monolayers or bilayers attached to solid surfaces, and (5) polyelectrolytes [19].

2.3. Photoeatalysis with semiconductors Semiconductor photocatalysts in a foma of colloids, powders, porous granules, thin films or bulk solids including single crystals (used in model studies) provide both liquid phase and gas phase transformations. Comprehensive reviews in this field can be found in monographs [4] (Chapters by N.S.Lewis and M.L.Rosenbluth; M.Gr/itzel; M.Schiavello and A.Sclafani; P.Pichat and J.-M.Herrmann; G.A.Somorjai; T.Sakata; H.Tributsch; M.A.Fox; H.A1-Ekabi and N.Serpone; D.F.Ollis, E.Pelizzetti and N.Serpone); [8] (Chapter by Yu.A.Gruzdkov, E.N.Savinov and V.N.Parmon) and [3]. The nature and general pathway of the photocatalytic action of semiconductors are nowadays well established (see Fig. 5), though numerous structural, thermodynamic and mechanistic peculiarities often make the detailed pathways for the same PC in different reactions or different PCs in the same reaction, also somewhat different [3, 4, 6, 8]. As seen from Fig. 5, upon absoption of photons with the energy hv > Eg, an electron | and hole Q centres are formed. They migrate to different sites on the PC surface, thus becoming spatially separated. Note, that what solid state physisists call surface electron and hole centers, in fact are some definite chemical species with strong reducing and oxidizing

43 properties, respectively. E.g., for TiO2 semiconductor surface electron and hole centers are suggested

to

be,

respectively,

{-Ti(III)-}surface and

{Ti(IV-O'- -Ti(IV)}surface

or

{Ti(IV)-O2-Ti(IV)} ()H sites [20]. Chemically pure semiconducor materials can absorb only those photons, the energy hv of which exceeds the band gap Eg. Therefore, Eg value determines the "red" boundary of the light that is used in photocatalytic action of these materials. By way of example, Table 1 presents the values of Eg and the corresponding values of boundary wave length 9~o= hc/Eg (where e is the velocity of light) for some semiconductor and dielectric oxides [2]. However, a semiconductor PC can be sensitized to light with ~ > ~o by Fig.5.The nature and general pathway of the chemical modifications of its surface photocatalytic action of a semiconductor layer or adsorption of certain catalyst particle. Eg is the band gap, E e shows molecules on its surface, provided that the direction of change of the energy for such treatment creates additional full electrons or empty electron levels in the band gap of the semiconductor material. Yet another approach to sensitizing PCs to a broader light spectrum is to use composite materials with a heterojunction (Fig. 6) between a narrow band gap and wide band gap semiconductors. A particular Table 1. photocatalyst of Fig. 6 smoothly provides the reaction of H2S Band gaps Eo and "red" boundaries Lo of the optical absorption f~3rtypical semiconductor and dielectric decomposition oxides [2] H2 S hv,PC ~ H 2 + S Oxide E z/eV* Lo/nm * To reduce Aad s the bottom of NiO 0.93 1340 the PC conduction band must be Cr203 1.4 890 located above the electron level of CuO 1.7 735 Aads, while to oxidize Dad s, the top CdO 2.1 595 of the valence band must be located below the level of Daas. If Fe203 2.2 570 a particular semiconductor PC TiO2 3.0 420 does not match these conditions, ZnO 3.2 390 one can shift the position of the MgO 7.2 178 conduction band up and that of SiO 2 8.6 145 the valence band down by just A1203 9.0 138 decreasing the size of the PC particle to ca. 10 2 ~ or smaller *The value refers to the pure bulk material and does not [21]. take into account the particle size and impurities.

44 The efficiency of semiconductor PCs in some reactions (such as dehydrogenation of organics, splitting of H20 and HES, etc.) can be enhanced by depositing tiny islands of additional catalysts, which facilitate certain reactions stages that may not require illumination. For example, islands of Pt metal are deposited on the surface of the composite photocatalyst in Fig.6 with the aim to facilitate the step of H2 formation. Fig.6. Photocatalytic cleavage of H2S over a The list of reactions provided by platinized composite sulfide semicondispersed semiconductors as photocatalysts ductor particle with a heterojunction. is very broad [22]. It includes: (1) splitting of water into the H2 and 02; (2) dehydrogenation of alcohols into H 2 and aldehydes or ketones; (3) oxidation of inorganic substrates (water into H202, NO2" into NO[, C N into CNO); (4) partial oxidation of organic substrates with 02 (benzene into phenol, toluene into methylphenols, - ~(CH2)4 ~ [--'( CH2)4 i .~

R

Ph Ph Ph O. Ph Ph Ph ph)= ~ a h ) - - O + ph)z---x + ph~CHCHO " ph)CH2 ~ p h ) - - O , etc.); (5) complete oxidation (mineralization) of hydrocarbons, oxigenated and chlorinated hydrocarbons, amines, nitrocompounds; (6) reduction of CO2 with H20 into CH3OH and C2HsOH; (7) nitrogen fixation: N 2 + 3H20 -~ 2NH 3 + 3/202, (8) syntheses of amines, e.g. NH 3 + 3CHaOH -}(CH3)3N + 3H20; (9) synthesis of aminoacids, e.g. NH 3 + PhCH2COCOOH --} PhCH2CHNH2COOH; (10) oligomerization and polymerization, e.g.

(11 ) structural isomerization of hydrocarbons, e.g.

(12) hydrogenation and cracking of alkenes and alkines with H20 or alkohols, e.g. C3H6 + H20 --~ CH4, CIH6, C3H8; (13) deposition of metals on semiconductors, e.g., 4Au 3§ + 6H20 --~ 4Au ~ + 302 + 12H § For some of these reactions detailed mechanistic studies were carried out. As an example, in Scheme 5 the pathway suggested for 1,1-diphenylethylene oxidation with 02 into benzophenon [23] is shown.

45

TiO2 hv _ _ @ + @ ;

Ph

@+02

~O2

02 (or 0~)

Ph

I Scheme 51 r Ph

ph>--- + (~

Ph

O~O"

TiO2 Ph P

~

+

"x

_@ I -" O--O

P

,A

O--O

I

! i_

Adsorption/desorption [24-25], as well as mass transfer phenomena [26] were also proved to be important for photocatalysis. An important advatage of semiconductor PCs compared to molecular ones, is a far greater stability of the former PCs. This is particular true for TiO2 photocatalyst. However, a notable disadvantage of TiO2 and most of other stable semiconductor PCs is that they have rather large band gaps Eg, and thus, are sensitive to only UV light. 3. PHOTOCATALYSIS IN ORGANIC SYNTHESIS The lists of reactions provided by molecular, supramolecular and semiconductor PCs

(vide supra) suggest certain perspectives for their practical applications for fine organic synthesis. Note, that sometimes PCs act selectively producing a sole product in rather exotic reactions; in other cases, a whole set of exotic products can be produced from the same reagents (vide supra). As additional examples note formation of glycols from alkohols ZnS (CH3OH > HOCH2-CH2OH and ZnS or (CH3)3OH > (CHs)zC(OH)CHzCHzC(OH)(CH3)2) and RuO2/TiO2 ZnS diamines from amines ((C2Hs)3N > CH3CH HCCH 3 [27]), as well synthesis of 1~(C2H5)2 J(C2Hs)2 aminoacids from hydroxocarboxylic acids (Scheme 6) or ketocarboxylic acids in ammoniawater solutions on CdS, ZnS and TiO2, or CdS and ZnTSPP (TSPP is the tetrasulfonatophenylporphyrinato ligand), respectively, with remarkably high quantum yields upto 35%. The aminoacids Ala, Gly, Phe, Leu and Tyr are produced efficiently in this way [27].

RCH(OH)COOH

Scheme 6 ]

hv,PC / ~ " ' - " ~

2H" ~ ~ . ~ . ~

RCOCOOH NH3

RCH(NH2)COOH

RC(=NH)COOH+ H20

46 Scheme 7 illustrates the transformations of a bicycloheptadiene in the presence of transition metal complexes as photocatalysts or photogenerated catalysts [ 11 ]:

,••,•

v

~hv,

Ni(CO)4/ hv /

Fe(CO)5 hv

.f.

Cr(CO)~= \ \(Ph3 P)2Ni(CO)2

+

Thus, photocatalysis and photogenerated catalysis indeed open up rather reach opportunities in fine organic synthesis, including some new reactions and nontraditional pathways for some known reactions. More efforts should be made in engineering of appropriate photocatalytic reactors for such synthesis. 4. PHOTOCATALYSIS IN ABATEMENT OF ENVIRONMENTAL PROBLEMS Photocatalysts demonstrate a remarkable ability to provide at room temperature and atmospheric pressure, complete oxidation (mineralization) of both organic (hydrocarbons and their derivatives containing O ' N, S, Cl and other heteroatoms) and inorganic (S 2-, SO 23

'

NO ~, CN) pollutants, remaining stable enough under the reaction conditions. The latter statement refers first of all to TiQ. All this make photocatalysis attractive for future commercial scale treatment of gaseous and aqueous waste streams, as well as hydrocarbon spills [3]. Intensive work in this field of applied photocatalysis is underway in various countries. Pilot plants have been tested with both electrically fed lamps and solar light (nonconcentrated or concentrated with mirrors). Their schemes and photographs can be found in monograph [3]. Economic feasibility studies suggest that even at the present state of the art photocatalytic technology indeed can be competitive with the traditional carbon adsorption or incineration technologies in treatment of contaminated soil vapor extraction vents and small scale VOCcontaining vents [28]. Rapid progress in basic and applied research in photocatalysis suggests

47 that in the near enough future photocatalytic technologies can be substantially improved and can become competitive with the traditional technologies also in other applications for environmental control. 5. PHOTOCATALYSIS AND THERMAL CATALYSIS INDUCED BY SOLAR RADIATION IN UTILIZATION OF SOLAR ENERGY At present, thermal catalysis induced by solar radiation is more ready for potential practical use in the energy production industry of the-future, than photocatalysis [7,9,29,30]. Fig.7 illustrates the scheme of the pilot plant for thermocatalytic solar-to-chemical energy

.i. N

/

i ••

.---

\

/~Storage vessel~ ~ for m e t h a n e ) / t 95%). Reductive alkylation of NEA with different aldehydes or ketones provides easy access to a variety of related modifiers [47]. The enantioselection occurring with the modifiers derived from NEA could be rationalized with the same strategy of molecular modelling as demonstrated for the Ptcinchona system. Not so long ago, the general opinion was that high enantioselectivity can only be achieved with natural, structurally unique, complex modifiers as the cinchona alkaloids. Our results obtained with simple chiral aminoalcohols and amines demonstrate the contrary. With enantiomeric excesses exceeding 80%, commercially available naphthylethylamine is the most effective chiral modifier for low-pressure hydrogenation of ethyl pyruvate reported to

59 date. Thus we can be confident that further research will lead to other effective modifiers extending the scope of enantioselective heterogeneous hydrogenation. 3.2 Chemoselective modification Liquid-phase hydrogenation has become a very established method in the production of fine chemicals. In the past decades most of the classical methods using chemical reducing agents have been substituted by catalytic hydrogenation methods using either non-modified or modified transition metal catalysts [48]. In contrast, progress in the catalytic liquid-phase oxidation with molecular oxygen [49,50] has been comparatively slow since its early discovery by D6bereiner [51]. The main reason for this is the often difficult control of selectivity and catalyst deactivation [50]. The latter is reflected by the extremely high catalyst/ reactant ratios applied in these oxidations, which render them frequently uneconomic. A considerable effort seems to be necessary to make this principally cheap and environmentally friendly method competitive to other oxidation methods applied in organic synthesis. Recently we have applied the concept of modifying the metal surface with adsorbed auxiliaries in the platinum-catalyzed oxidation of L-sorbose (scheme 2) to 2-keto-L-gulonic acid (2-KLG), intermediates in vitamin C synthesis [52]. This direct oxidation could offer a catalytic alternative to the presently used route for the oxidation of the C-I hydroxyl group of Lsorbose to the corresponding carboxylic acid, which involves three steps to achieve good selectivity due to the necessary protection (and subsequent hydrolytic deprotection) of the other four reactive hydroxyl groups. O H II I

OHH I I

HOH2C'-'C --(3 ""-"C--"C-CH2OH I

!

OH H

L-sorbose

I

OH

02

r suppo.

O H II I

OH I

H I

HO2C-'C -'C --G --C-CH2OH I

I

OH H

I

OH

2-KLG

Scheme 2 The direct oxidation of L-sorbose in aqueous phase with molecular oxygen over supported Pt and Pd catalysts is rather inefficient due to low selectivity paired with significant catalyst deactivation [53]. Modifying platinum metal catalysts by blocking some of the active surface metal atoms by deposition of inactive foreign metals, leading to a reduction of the ensemble size, has been successfully applied in designing catalysts for alcohol oxidation [54,55]. However, this strategy was not successful for L-sorbose oxidation. Upon modification with Bi or Pb, the initial rate increases, but the overall performance of the catalysts is unsatisfactory due to oxidation and subsequent dissolution of the promotors (modifiers) under reaction conditions [56]. An alternative approach to increase the oxidation rate is the use of alkaline solutions, because bases enhance the reactivity of L-sorbose and weaken the adsorption strength of 2KLG. Unfortunately, the rate enhancement at higher pH is accompanied by a drop in selectivity due to the poor stability of 2-KLG in alkaline solutions. To circumvent this problem, we have modified the platinum catalysts by adsorbed tertiary amines and carried out the oxidation in neutral aqueous solution [57]. This allowed to enhance the rate without increasing the pH of the bulk liquid, which leads to detrimental product decomposition. Small quantities of amines (molar ratio of amine : sorbose = I: 1700, and amine : Pts = 0.1) are sufficient for modification. Using amines of pKa ~ I0 for modification, resulted in a considerable rate enhancement (up to a factor of 4.6) with only a moderate loss of selectivity to 2-KLG. The rate enhancement caused by the adsorbed amines is mainly determined by their basicity (pKa). In contrast, the selectivity of the oxidation was found to depend strongly on the structure of the amine.

60 As concerns the nature of base catalysis in the L-sorbose ---- 2-KLG transformation, it is most likely that the oxidation of the aldehyde intermediate is accelerated via a rapid hydration to the corresponding geminal glycol [58]. The formation of a carboxyl group from the glycol intermediate is an oxidative dehydrogenation step, which is usually much faster on platinum metals than the direct oxygen insertion to the carbonyl group of the aldehyde [50,51 ]. In order to improve the selectivity, we have systematically changed the structure of the amine modifier and rationalized the interaction complex between modifier and L-sorbose using molecular modelling. The aim was to find an amine, strongly adsorbing on platinum and forming an interaction complex with L-sorbose, in which the oxidation of the C-1 hydroxyl group is sterically favored. This is the case with hexamethylentetramine (HMTA), which upon adsorption on Pt still possesses three accessible nucleophilic nitrogen atoms, enabling interaction with an OH group of sorbose, as illustrated in Fig. 6. Electrochemical model studies [59] revealed that hexamethylentetramine is adsorbed on Pt and not oxidized under reaction conditions. The co-adsorption of L-sorbose and hexamethylentetramine on Pt results in a tilted position of the reactant, in which only C-1 is exposed to oxidative dehydrogenation. This geometric constraints result in a strong improvement of the selectivity to 2-KLG upon modification of a Pt/C catalyst (Fig. 7).

Figure 6. Side view of energetically most favorable complex formed between L-sorbose and hexamethylentetramine (HMTA) auxiliary. Complex is stabilized by N-H-O hydrogen bond interaction.

Figure 7. Effect of HMTA auxiliary. Selectivity to 2-KLG as a function of Lsorbose conversion over 5 wt% Pt/C. Unmodified catalyst (O); catalyst modified with HMTA (O).

4. CONLUSIONS Using the design of titania-silica mixed oxides for epoxidation of bulky olefins, the combined use of the solution-sol-gel method and supercritical drying has been shown to be a potent and versatile tool for the structural and chemical tailoring of mixed oxides. The large variety of controllable sol-gel parameters together with the use of appropriate drying methods provide exceptional control of chemical and structural properties of as-derived mixed oxides. In the epoxidation of bulky olefins, amorphous titania-silica aerogels possess a considerable advantage compared to presently known Ti-substituted zeolites, due to the larger range in which the pore size can be tuned. The excellent epoxidation activity of the aerogels demonstrates that crystallinity is not a requirement for highly active Si-O-Ti connectivities. A powerful concept for improving the chemo- and stereoselectivity of metal catalysts is their modification by co-adsorbed auxiliaries. This concept should be particularly useful for

61 liquid-phase reactions performed at relatively low temperature, as applied in fine chemical catalysis. The auxiliaries, used in very small quantities (auxiliary : reactant ratio < 1:1000), can either be simply added to the reaction solution or brought onto the metal surface in a special pretreatment step. They have to possess a suitable anchoring group and the functions necessary for the desired interaction(s) with the co-adsorbed reactant(s). An interesting conceptual advantage of this method, compared to the well-known immobilization of auxiliaries (promotors), is the possibility of self-organization of the interaction between the mobile auxiliary and the reactant(s). With immobilized auxiliaries the geometric requirements for favorable interaction are difficult to control. The concept with co-adsorbed auxiliaries has been successfully applied for designing catalysts for the enantioselective hydrogenation of a-ketoesters and the platinum-catalyzed oxidation of L-sorbose to 2-keto-Lgulonic acid, and may open routes to several other applications, where chemo- and stereocontrol are demanding. Essential for further progress is a better understanding of the interaction of co-adsorbed surface species, particularly in the liquid phase, where this concept seems to be most promising. REFERENCES

.

3.

.

6. 7. 8. .

10. 11. 12. 13. 14. 15. 16 17. 18. 19. 20. 21. 22. 23. 24.

J.H. Block, A.M. Bradshaw, P.C. Gravelle, J. Haber, R.S. Hansen, M.W. Roberts, N. Sheppard and K. Tamaru, Pure Appl. Chem. 62 (1990) 2297; G.A. Somorjai, Surf. Interface Anal. 19 (1992) 493. G. Ertl, Surf. Sci., 299-300 (1994) 74; G.A. Somorjai, ibid., 849. E.W. Abel, F.G.A. Stone, G. Wilkinson, (eds.), Comprehensive Organometallic Chemistry II: a Review of the Literature 1982-1994, Pergamon, 1995. J.B. Moffat (ed.), Theoretical Aspects of Heterogeneous Catalysis, Van Nordstrand Reinhold, New York, 1990. J.M. Thomas, Angew. Chem., 12 (1988) 1735. R. Roy, Int. Ceram. Monogr. 1 (1994) 737. L. L. Hegedus and C.J. Pereira, Chem. Eng. Sci., 45 (1990) 2027. J. M. Thomas and K.I. Zamaraev (eds.), Perspectives in Catalysis, Blackwell Sci. Publ., Oxford, 1992. D.E. De Vos, F. Thibault-Starzyk, P.P. Knops-Gerrits, R.F. Parton and P.A. Jacobs, Macromol. Symp., 80 (1994) 157. M.E. Davis (ed.), Large Pore Molecular Sieves, Catal. Today, 19 (1994). M. Misono, Stud. Surf. Sci. Catal., 75 (1993) 69. L.T. Tejuca and J.L.G. Fierro (eds.), Properties and Applications of Perovskite-Type Oxides, Marcel Dekker, New York, 1993. A.T. Bell, Chem. Eng. Sci. 45 (1990) 2013.; V. Ponec and G.C. Bond (eds.), Catalysis by Metals and Alloys, Stud. Surf. Sci. Catal, Vol. 95, 1995. Y. Iwasawa, Catal. Today, 18 (1993) 21; J. Haber, Appl. Catal. A., 113 (1994) 199. D.A. Ward and E.I. Ko, Ind. Eng. Chem. Res., 34 (1995) 421; M. Schneider and A. Baiker, Catal. Rev. Sci. Eng., 37 (1995) 515. D. Dutoit, M. Schneider and A. Baiker, J. Catal. 153 (1995) 165. R. Hutter, T. Mallat and A. Baiker, J. Catal. 153 (1995) 177. R. Hutter, T. Mallat and A. Baiker, J. Catal. 157 (1995) 665. Brit. Patent No. 1'249' 079 (1971). R.A. Sheldon, J. Mol. Catal., 7 (1980) 107. R.A. Sheldon, in Aspects of Homogeneous Catalysis, R. Ugo (ed.), Vol 4, Reidel, Dordrecht, Holland, 1981, p. 3. U.S. Patent No. 4'410'501 (1983). B. Notari, Catal. Today, 18 (1993) 163. B. Notari, Stud. Surf. Sci. Catal., 37 (1988) 413.

62 M.A. Camblor, A. Corma, A. Martinez, and J. Prrez-Pariente, J. Chem. Soc. Chem. Commun., 589 (1992). 26. A. Corma, M.T. Navarro, and J. Prrez-Pariente, J. Chem. Soc. Chem. Commun., 147 (1994). 27. A. Corma, M.A. Camblor, P. Esteve, A. Martinez, and J. Prrez-Pariente, J. Catal., 145 (1994) 151. 28. J. Livage, M. Henry and C. Sanchez, Prog. Solid State Chem., 18 (1988) 259. 29. M. Aizawa, Y. Nosaka and N. Fujii, J. Non-Cryst. Solids, 128 (1991) 77. 30. B.E. Yoldas, J. Non-Cryst. Solids, 38 (1980) 81. 31. C.C. Lin and J.D. Basil, Mater. Res. Soc. Symp. Proc., 73 (1986) 585. 32. D. Dutoit, M. Schneider, R. Hutter and A. Baiker, J. Catal., in press. 33. J.P. Candy, B. Didillon, E.L. Smith, T.B. Shay and J.M. Basset, J. Mol. Catal., 86 (1994) 179. 34. H.U. Blaser, Tetrahedron Asymmetry, 2,843 (1991); R. Noyori, Asymmetric Catalysis in Organic Synthesis, Wiley, New York, 1994.. 35. Y. Nakamura, Bull. Chem. Soc. Jpn., 16 (1941) 367; Y. Izumi, Adv. Catal., 32 (1983) 215; T. Osawa, T. Harada and A. Tai, J. Mol. Catal., 87 (1994) 333. 36. Orito, S. Imai, S. Niwa and G-H. Nguyen, J. Synth. Org. Chem. Jpn., 37, 173 (1979); Y. Orito, S. Imai and S. Niwa, J. Chem. Soc. Jpn., (1979) 1118. 37. K.E. Simons, P.A. Meheux, S.P. Griffiths, L.M. Sutherland, P. Johnston, P.B. Wells, A.F. Carley, M.K. Rajumon, M.W. Roberts and A. Ibbotson, Recl. Tray. Chim. PaysBas, 113 (1994) 465. 38. A. Baiker and A. B laser, in Handbook of Catalysis, G. Ertl, H. Knrzinger and J. Weitkamp (eds.), Springer Verlag, Berlin, in press. 39. H.U. Blaser, H.,P. Jalett, D.M. Monti, A. Baiker, and J.T. Wehrli, Stud. Surf. Sci. Catal., 67 ( 1991) 147. O. Schwalm, J. Weber, J. Margitfalvi, and A. Balker, J. Mol. Struct., 297 (1993) 285. 40 41. O. Schwalm, B. Minder, J. Weber, and A. Balker, Catal. Lett., 23 (1994) 268. 42. G. Webb and P.B. Wells, Catal. Today, 12 (1992) 319. 43. B. Minder, T. Mallat, A. Baiker, G. Wang, T. Heinz A. Pfaltz, J. Catal.,154 (1995). 371. K.E. Simons, G. Wang, T. Heinz, A. Pfaltz, A. Balker, Tetrahedron: Asymmetry, 6 44 (1995) 505. B. Minder, M. Schiirch, T. Mallat, A. Balker, Catal. Lett. 31 (1995) 143. 45 46. T. Heinz, G. Wang A. Pfaltz, B. Minder, M. Sch~rch, T. Mallat and A. Baiker, J. Chem. Soc. Chem. Commun., 1995, 1421. 47. B. Minder, M. Schtirch, T. Mallat, A. Balker, G. Wang, T. Heinz, A. Pfaltz, J. Catal., in press. 48. M. Freifelder, Practical Catalytic Hydrogenation, Wiley-Intersci., New York, 1971. 49. H. van Bekkum, in Carbohydrates as Organic Raw Materials, F.W. Lichtenthaler (ed.), VCH, Weinheim, 1990, p. 267. 50. T. Mallat and A. Baiker, Catal. Today, 19 (1994) 247. 51. J.W. Drbereiner, Schweigers J. Chem. Phys., 38 (1823) 321. 52. T. Reichstein and A. Gruessner, Helv. Chim. Acta, 17 (1934) 311. 53. K. Heyns, Ann. Chem., 558 (1947) 177. 54. J. Ludec, Get. Patent No. 2'612'844 (1976); H. Fiege and K. Wedemeyer, Angew. Chem., 93 (1981) 812.52. 55. T. Mallat, Z. Bodnar, A. Baiker, O. Greis, H. Striibig and A. Relier, J. Catal., 142 (1993) 237, and references l- 14 therein. 56. C. Brrnnimann, Z. Bodnar, P. Hug, T. Mallat, and A. Baiker, J. Catal., 150 (1994) 199. 57. C. Brrnnimann, T. Mallat and A. Baiker, J. Chem. Soc., Chem. Commun., 1377 (1995). 58. Y. Ogata and A. Kawasaki, in" The Chemistry of the Carbonyl Group, J. Zabicky (ed.), Vol. 2, Intersci., London, 1970, p. 3. 59. C. Brrnnimann, Z. Bodnar, A. Aeschiman, T. Mallat and A. Balker, J. Catal., in press. 25.

J.W. Hightower, W.N. Delgass, E. Iglesia and A.T. Bell (Eds.) 11th International Congress on Catalysis - 40th Anniversary

Studies in Surface Science and Catalysis, Vol. 101 1996 Elsevier Science B.V.

63

A R e t r o s p e c t i v e V i e w of A d v a n c e s in H e t e r o g e n e o u s C a t a l y s i s : 1956-1996, Science Robert L. Burwell, Jr. Catalysis Laboratory, Northwestern University, Evanston, IL 60201, USA Most of the large number of acronyms added to BET between ICC I in 1956 and ICC 11 in 1996 represent new techniques for analyzing the products of reaction, for characterizing catalysts and for determining the structure of chemisorbed species. Use of these techniques has substantially altered the practice of catalysis. Only the elderly will remember the great improvement in the quality of life which resulted from the fulfillment of Emmett's prophecy, that gas chromatography appeared "destined to provide a rapid, accurate method for the analysis of complicated mixtures of products in a relatively simple, straightforward manner" (ICC 1, paper 65). To GC can be added scanning FTIR, NMR, EPR, microwave spectroscopy, and various MS. Adequate description of many catalysts will require a large number of bits of data since they are usually rather complicated materials rather than simple chemicals. Attempts at this were just beginning by ICC 1, but now, one expects authors to give specific surface areas and some details of the porosity of their catalysts. Automation of the former tedious point by point measurement of the N 2 adsorption isotherm has greatly facilitated this. Use of chemisorption to measure specific site numbers has become much more common, for example, CO for Cr203, and H 2 to measure H/M (usually equated to Ms/M where M is an atom of a supported Group VIII metal). If the metal particles are roughly spherical, the H/M can be converted into an average particle size, but modern high resolution electron microscopy and wide angle x-ray scattering (WAXS) provide particle size distributions rather than mere averages. Further, in favorable cases, the first can provide atomic level resolution and lattice details and the second can diagnose the degree of strain in metallic particles larger than about 2 nm. However, most of these techniques require significant funding. For example, for best results, WAXS needs cyclotron radiation. The author thought himself very fortunate when his department got him a Type K Potentiometer and a Podbielniak distilling column in 1941-42. Among other new techniques for study of catalyst structure, EXAFS diagnoses the atomic environment about catalytic sites of noncrystalline materials and the x-ray absorption line just before the EXAFS wiggles has helped to identify the oxidation number of catalytic sites as have EPS and NMR in appropriate cases. Provision of information on catalyst structure probably represents the major advance in catalysis since ICC 1, but

64 even so, exact details of the surface sites are stilllacking and particles smaller than 1 run can be elusive. Chemisorption on nonmetallic catalysts should provide the n u m b e r of catalytic sites and for comparative purposes a single M s can be taken as a catalytic site on metals. This permits the calculation of turnover frequencies which was a n e w concept in post ICC 1 and which permitted intercomparison of catalyst activities. For the firsttime then, one has been able, for example, quantitatively to discuss support effects in Rh/support catalysts. Except for support effects, structure sensitivity has usually appeared in one of two aspects, variation of rate with surface crystal face or with particle size. In ICC 1 Gwathmey reported in one of the first experiments with single crystal faces that different faces machined from Ni single crystal spheres catalyzed the hydrogenation of ethylene at different rates (ICC 1 paper 5). Many similar results have followed, Farnsworth (ICC 1, paper 15) studied 1 cm 2 nickel and platinum sheets employing UHV techniques. The much augmented rate of the hydrogenation of ethylene restdting from argon ion bombardment was drastically reduced by annealing, but that of H 2 + D 2 ,~ 2HD was unchanged. This may have been the first specific report of structure insensitivity. Despite many subsequent papers general agreement as to the origin of this counter intuitive phenomenon has not yet developed. In studying the structure of chemisorbed species, infrared absorption spectroscopy has been outstanding partictdarly as facilitated by scanning FTIR. Eischen's paper (ICC 1, 67) represented what was probably the first in situ examination of chemisorbed species during a catalytic reaction (IR of CO + O 2 on Ni-NiO). The other vibrational spectroscopies, laser Raman and magic angle spinning NMR, have also been useful. Despite its low resolution, high resolution EELS has been useful in UHV work for assessment of surface cleanliness and for the identification of adsorbed species. In ICC 1 there were only a few references to diffusional limitations, but they may have been present in a number of papers. Despite improved attention, problems may still exist particularly in systems involving transport from the gas to the liquid phase. Absent a demonstration that the rate of a hydrogenation was proportional to the amount of c a t ~ y s t one may suspect that C(H2)(liq.) was not in equilibrium with P(H2)(gas). Acidic, high area silica-alumina had received substantial attention in ICC 1, (52-58). Perhaps the most dramatic change in the subsequent cat~ytic literature was the debut of zeolites. Why acid catalyzed reactions are so much faster on zeolites than on silica-alumina has been extensively discussed but probably not conclusively. One should be able to know the exact structures of catalytic sites in zeolites, but initial hopes that this would do wonders for mechanistic understanding have not been fully realized. Super acids and carbonium ions came into heterogeneous catalysis from homogeneous chemistry and in special cases reaction via carbonium ions seems to occur,

CnH2n+2 + [H +] -~ CnH2n+3+ --) CnH2n+l + + H 2 Carbonium Carbenium

65 The carbenium ion so formed then reacts in the ICC 1 manner except perhaps for not abstracting a hydride ion from another alkane. Although initial views that zeolites in general were super acids have come into question, definite super acids have been found such as calcined H2SO 4 oZr(OH) 4 which catalyze the isomerization of alkanes at low T. Noble metal/acidic zeolite catalysts constitute "dual functional catalysts" whose basic chemistry resembles that of M/SiO 2 ~ A1203. They have been extensively investigated with particular emphasis upon determination of the exact size and location of the metal cluster. Although fusion of isolated noble metal atom clusters of the size of those which could exist in zeolite pores into large cryst~s would be very exothermic, conditions for preparing some metal/zeolite catalysts with clusters cont~ning ca. 6 atoms located in the pores have been developed. How such clusters differ electronic~ly (the quantum size effect) and catalytically from dusters of 200 atoms is still under investigation. Zeolites have led to a new phenomenon in heterogeneous catalysis, shape selectivity. It has two aspects: (a) formation of an otherwise possible product is blocked because it cannot fit into the pores, and (b) formation of the product is blocked not by (a) but because the transition state in the bimolecular process leading to it cannot fit into the pores. For example, (a) is involved in zeolite catalyzed reactions which favor a para-disubstituted benzene over the ortho and meso. The low rate of deactivation observed in some reactions of hydrocarbons on some zeolites has been ascribed to (b) inhibition of bimolecular steps forming coke. Full application of structural findings requires a complete determination of mechanism which in most heterogeneous cat~ytic reactions involves a closed cycle of adsorption-desorption steps sandwiching a sequence of elementary steps among chemisorbed intermediates. The problem is more difficult than in homogeneous catalysis in part because two or more surface sites of somewhat different properties are likely to be involved. Although the zeitgeist of the period favored structural studies rather than mechanism, advances, albeit only partial, were made in the mechanism of particular reactions. The general state of knowledge about elementary steps is relatively primitive, the mechanism of the hydrogenation of ethylene still lacks a consensus and, in particular, the role of slowly reacting surface species in hydrocarbon reactions is still unclear. Two relatively new areas which have flourished during the interval since ICC 1 and have strongly influenced ideas as to possible structures of chemisorbed complexes and possible elementary steps are organometallic chemistry-homogeneous catalysis and surface science (in which there were two papers in ICC 1, papers 5 and 46). Surface science has provided structures of various chemisorbed species on the surfaces of single crystal planes of metals and some reactions of such species. Applicability of these data was restricted by their UHV origin. However, improvement resulted from the development of pressure cells in which catalytic reactions could be run at 1 atm or above and followed by examination of the surface in UHV. The great advantage here over conventional catalytic work was the provision of detailed information about the state of the surface. However, even so reaction could not be followed in situ as

66 is possible with infrared or EPR. Further, species found on the surface in UHV work come without labels identifying them as reaction intermediates. Errors in proposing particular species as reaction intermediates have inevitably occurred, but, of course this is equally the case for species found by IR or EPR and it does not detract from the importance work in this area. ICC 1, paper 61, surveyed reactions homogeneously catalyzed by metal carbonyls and presented one of the earliest discussions of the application of the mechanisms of such reactions to heterogeneous catalysis, for example, of the structure of CO as a ligand to that as a chemisorbed species. Further, a new class of heterogeneous catalysts has appeared, heterogenized homogeneous catalysts. The chemistry of the new catalysts was usually similar to that of their homogeneous predecessors, but some were unparalleled among organometaUic catalysts. Some actinide complexes on alumina which appear to consist of adsorbed L2ThH+ (L is pentamethylcyclopentadienyl), rapidly catalyze isotopic exchange between D 2 and H 2 at-195~ (only Pt having a comparable rate), and between D 2 and ethane at 90~ and also hydrogenations of simpler olefins at -45~ and of benzene at 90~ Since the O.N. of Th is effectively held at +4, oxidative addition and reductive elimination steps cannot be involved in hydrogenation and exchange and evidence is strong for the following intermediates,

D.~...D

..""D

"Th " ~ H g+ A

~-

.... ----

D:::

...-H -....

"'"'""H

"'"'TM'""'" O

H:'"" "...

..;"R

""Th"" C

A (either a transition state or a free energy minimum) reacts to form B. In hydrogenation, R in C is an alkyl group formed by reaction of H with adsorbed olefin and the mechanism is not Horiuti-Polanyi-like. Rather reaction via both A and C are among the few relatively clear examples of Rideal-Eley processes. Charging problems have limited the UHV work largely to metals and absence of organometallic complex analogs to metals has limited applications of work in this area largely to nonmetals. Two zeolites with chemistry not analogous to acidic SiO 2 ~ have received particular attention. The low alumina ZSM-5, free or metal-loaded, catalyzes CH3OH -~ gasoline (on HZSM-5), propane -~ aromatics (on ZSM-5 loaded with gallium), and other reactions too numerous to detail. A material of the same crystal structure but containing about 1% of structural Ti and devoid of A1 (TS-1) has unique oxidative capacities for those hydrocarbons which can enter its relatively narrow pores. At about 50~ and with aqueous H20 2 as the oxidant, it catalyses the epoxidation of alkenes and the hydroxylation of alkanes and arenes. Pt/L-zeolite in which all acidic [H +] including that resulting from reduction of Pt 2+ is replaced by K + catalyzes conversion of hexane to benzene with high selectivity. A key feature is the very low rate of

57 deactivation. Shape selective favoring of the initial transition state is not necessary since a non-microporous PtIMgO-A120 3 is equally effective. For economic reasons the oxidative dimerization of CH 4 at about 750~ was extensively investigated. Li+/MgO was the most studied catalyst, but a wide variety of oxides are effective, none lamentably with enough selectivity to be interesting commercially. The reaction which probably involves oxidation at the surface with liberation of CH 3 to the gas phase where it dimerizes is, like the Rideal-Eley step, an exception to the earlier statement about mechanism. Strong metal support reaction SMSI was a popular subject for study. The type catalyst, Pt/TiO 2 reduced at 300~ behaves much like Pt/SiO 2, but reduction at 500~ largely eliminates its capacity for the chemisorption of H 2 and hydrogenation while inducing activity for the hydrogenation of CO. Ti suboxide formed by reduction encapsulates the Pt particles. In recent years research of possible utility in the production of fine chemicals has increased substantially and in part consequent to government policy. This work has been too variegated to summarize briefly. A flurry of work in the hydrogenation of CO also originated in government policy. It led to the elaboration of our understanding of these reactions, but it is not clear that it led to major developments. There are myriads of possible oxidations for which AG is negative but which are unknown non-catalytically except in some cases, enzymatically. This situation has stimulated a great deal of research in which one success, TS-I was mentioned above. Oxidative dehydrogenation has received extensive attention. In particular, many compounds of Mo and V have been studied because of their practical utility. The Mars and van Krevelen mechanism (1954) has become dominant in interpreting oxidations on altervalent oxides, i.e. the proximate oxidant is a metal ion in a higher oxidation state. Oxygen serves to return the reduced metal ion to its original state. The reaction, CH 4 + 1/202 --~ CO + 2H 2, has been found to occur with surprisingly high selectivity on Rh or Pt monoliths at 1000~ and very high space velocities. Many papers have employed the concept of spillover, most commonly of hydrogen dissociatively adsorbed on supported X onto the support. In type cases like reduction of MoO 3 by H 2 promoted by Pt on the MoO 3 or D 2 exchange with surface OH in Pt/A1203, the occurrence of spillover is clear but perhaps not the details of migration of H over the support. However, not all papers clearly establish the catalytic relevance of spillover even where it occurs. Further several papers have reported that the rate of a hydrogenation reaction on Pt/A120 3 or the like is increased by mixing some support into a batch of catalyst. Whence, H spills over onto the support to effect hydrogenation and the authors all quote an earlier paper with such a result. But the referees failed to point out to the authors that subsequently J. Catal. 24, 482 (1972) demonstrated that the augmented rate resulted from scavenging of poisons by the added support. It is difficult now-a-days to be conversant with all previous catalytic literature, but the level of ignorance appears to have increased substantially (and mea culpa). Too many papers report no assay of poisons in the feeds.

68 Stimulated by the growing importance of enantiomeric purity in drugs and insecticides, enantioselective hydrogenation has received augmented attention, primarily with homogeneous catalysts. However, the optically active alkaloid, cinchonidine, adsorbed on Pt/SiO 2 catalyses the liquid phase hydrogenation of pyruvate, C H 3 C O C O O M e , to methyl lactate, CH3CHOHCOOMe, of substantial optical activity. Presumably, adsorption of cinchonidine on Pt generates an optically active catalyst, but why should the rate be 30X that of the same catalyst without cinchonidine? Consequent to the work of m a n y and employing such techniques as structure variation, isotopic tracers, and stereochemistry, a large number of different adsorbed hydrocarbon fragments have been identified as key intermediates in various reactions of hydrocarbons. Correlation of these species with similar polynuclear organometallic species has been of interest. However, the author feels that mechanistic understanding has lagged behind some other aspects of catalysis. Finally, theory involving collective electrons and the like, characteristic of ICC 1, appears less frequently in the catalytic literature and theories involving local sites appear more oRen. The existence of coordination complex analogs m a y have played some part in this development. More recent theoretical developments relate primarily to chemisorbed species. Catalytic polymerization is surveyed in Heinemann's paper on advances in catalytic technology. Drastic sampling and subjective judgment was required to condense 40 years of research in heterogeneous catalysis into a few thousand words. There might be rather littleoverlap between this paper and that with the same title written by another author. References have been omitted to avoid presenting a sampling of names with the implication that they alone were responsible for the advances.

J.W. Hightower,W.N. Delgass, E. Iglesia and A.T. Bell (Eds.) 11th International Congress on Catalysis - 40th Anniversary

Studies in Surface Science and Catalysis, Vol. 101 9 1996 Elsevier Science B.V. All rights reserved.

69

A R e t r o s p e c t i v e V i e w of A d v a n c e s in H e t e r o g e n e o u s Catalysis: 1956-1996 T e c h n o l o g y Heinz Heinemann Materials Sciences Division, Lawrence Berkeley National Laboratory, University of California, Berkeley, CA 94720*

The first International Congress on Catalysis in 1956 occurred a t a time at which catalytic technology and industrial applications of catalysis were rapidly increasing at an almost exponential rate which accelerated through the next 20 years. One would have been hard put to envision the explosion of applied catalytic work, which, to some extent, was stimulated by the worldwide demand for goods following the recovery from the Second World War. It is difficult to do justice to the many developments in the few pages allotted to the author and this article will be limited to what the author considers the most outstanding new catalysts and catalytic processes. One objective of this paper is to show that, contrary to some statements that the rate of catalytic developments has slowed in recent years, there has been a rapid growth of technology, though the areas to which it is being applied have shifted. A major factor in the growth of new technologies, one that is not specific for catalytic developments, but applies equally to all technologies, as well as to theory, is the d e v e l o p m e n t and wide availability of more and more sophisticated computers. In the last few decades, computers have changed means of observation, reduced calculation times and simplified operations, both on the bench and on the industrial scale. The proceedings of the First ICC are divided into four major subjects and the same subjects could be applied to the current meeting. They were: "Chemistry and Physics of Solid Catalysts", "Homogeneous Catalysis and Related Effects", "Surface Chemistry and Its Relation to Catalysis" and "Techniques and Technology of Catalysis". About 20 of the 80 papers presented at the meeting were in the last category, a percentage that has not increased substantially at subsequent Congresses. 12 papers were concerned with homogeneous catalysis, none of which fitted into the category of technology. Homogeneous catalysis in technical applications has grown even faster than heterogeneous catalysis, though starting from a much lower base. While there are many homogeneously catalysed processes in use, each of them is limited to a relatively small volume of products. Present address: LBNLWashington, D.C. Project Office, 1250 Maryland Ave., S.W., Suite 500, Washington, D.C. 20024.

70 In view ot~ the limited space allowed, this paper will be essentially restricted to heterogeneous catalysis. The major initial driving force in the expansion of catalytic processing was the worldwide demand for energy and the availability of relatively cheap petroleum. This led to the development of major new processes in petroleum refining and in the petrochemical industry, as well as to inventions which revolutionized existing technology (Table 1). Table 1 Major new catalytic technology developments between 1956 and 1996 ,

Catalysts

...

Year

Process

1957

Polymerization

Ziegler-Natta catalysts

1962

Steam reforming

NiK2AI203

1964

Catalytic cracking with faujasite zeolites

x and y zeolites

1967

Bimetallic reforming

Pt-Re; Pt-Ir

1968

Selectoforming (shape selectivity)

Erionite

1972

Low pressure CH3OH

Cu-Zn-Al20 3

1974

Acetic acid v/a carbonylation

RhI

1976

Auto emission control for HC and CO

Pt-AI203

1980

Gasoline from methane

ZSM-5 zeolite

1982

NOx control

Pt-Rh for autos V205-TiO2 stack gas

1988

Selective oxidation

TiSiO2

1988

Chiral catalysis

Zeolites, or SO2 cinchonidine on supported Pt

1991

Polymerization

MetaUocenes

In the 1960s and subsequent years, catalytic processing was extended to new areas, particularly to the abatement of environmental concerns (Table 1). Automobile emission control catalysis became a major factor of catalytic processing and catalysts for this purpose today constitute a large portion of catalyst manufacture. Other environmental concerns resulted in stack gas catalytic conversion, primarily to remove NOx and SO2. It was also necessary to improve the selectivity of established processes to minimize toxic emissions and to increase energy efficiency. During the 1980s, a trend toward fine chemicals production and particularly toward pharmaceuticals became noticeable, which resulted in numerous new processes and reactions involving catalysis. There has thus been a

71 change from the very large volume of relatively few low-priced products to a large number of small volume, high-priced products. The importance of industrial catalysis is evidenced by the fact that over 18% of the U.S. GNP is achieved by catalytic processing. Catalyst manufacture has large markets of its own, as shown in Table 2. Table 2 Catalyst markets ($ billion)

Western Europe U.S.A. Worldwide

1989

2000 (projected)

1.3 1.9 5.0

1.9 2.4 6.5

Among the many new process inventions of the last 40 years, three stand out as having ushered in several new technologies. The first is the discovery of the catalytic properties of zeolites, the second is the use of precious metals in emission control and the third is selective polymerization. While catalytic cracking was a subject of major interest at the time of the first Congress in 1956, it was then generally assumed that only amorphous solids, such as clays or SIO2-A1203, possessed the necessary properties for acid-catalyzed reactions and that crystalline materials were of little, if any, interest. This myth was debunked by a paper which was presented at the second ICC in Paris in 1960. In the early 1960s, workers at Mobil demonstrated not only the greater activity but also the greater selectivity of cracking catalysts containing faujasite-type zeolites. Zeolite-containing catalysts dominate catalytic cracking today and their use has resulted in savings of petroleum of ~ 400 million barrels/year for the same amount of gasoline produced. At $17/bbl, this amounts to ~ $7 billion/year. In the mid-1960s, the shape selective properties of small and medium pore size zeolites were discovered at Mobil. The first commercial application was the "Selectoforming" process which selectively cracked straight-chain, low-octane number paraffins in gasoline. This was followed by new or improved petrochemical processes, such as the easier isomerization of xylenes at higher para selectivity, the disproportionation of toluene to benzene and xylene, and many others. Introduction of metals into the pores of zeolites at specific crystal sites made other processes possible. Today there are dozens of applications of zeolite catalyzed processing. While the discovery of the catalytic properties of zeolites was driven by the desire to improve industrial processing, the development of emission control catalysts was necessitated by governmental fiat. The first requirement was for 90+% removal of CO and of hydrocarbons, a goal which could not be met by oxidation with base metal oxides. To achieve the required specifications during automobile operations, it was necessary to develop supported platinum catalysts. Originally the support was alumina in pellet form. Later platinum on cordierite was used in honeycomb form, containing 200-400 square channels per square inch.

72 The original objective was the control of hydrocarbon and CO emissions. The danger of NOx emissions and the need to limit them became increasingly important in the 1980s. This was not possible with Pt on cordierite catalysts (usually containing rare earths, e.g., lanthana) but required the addition of rhodium (in quantities of approximately half of that of Pt) and a very careful control of the air/fuel ratio at 14.7 + 0.1 volumes. As shown in Fig. 1, there is a very narrow window in which both CO and NO removal are above 90%. An exhaust oxygen sensor and electronically controlled oxygen feed maintain this ratio in modern automobiles. The development of automobile exhaust catalysts resulted in these catalysts becoming the largest volume of catalysts produced, replacing catalytic cracking catalysts. Numerous other examples of environmental control by catalysis will be mentioned later. It must also be stated that in the attempt to minimize toxic by-products, not only their destruction but also avoidance of their production, has become important. Catalysis has and will continue to play a major role in improving the selectivity of chemical processes, thus eliminating or reducing undesirable by-products. Another field in which catalysis has played a big role in consumer products is polymerization. Ziegler-Natta polymerization of ethylene or propylene was invented just before the 1956 Congress, but became large-scale commercial in the late 1950s. The aluminum alkyl-titanium chloride catalysts originally employed in the high-pressure synthesis of polyethylene, polypropylene and other plastics were modified a n d / o r replaced by heterogeneous catalysts, such as chromia on silica, and a new low-pressure synthesis was developed. Books can and have been written on catalytic polymerization of various monomers (e.g., styrene, in addition to C2H4 and C3H6), on stereospecific polymers and control of the chain length of polymers. 100

o~

=s O

rt... r O

O

0 14.0

.j

14.5

'i' ...... 15.0

15.5

Simulated A/F Ratio

Figure 1. Principle of 3-way exhaust emission control. Catalyst efficiencies measured in the laboratory with a steady feed stream composition at various simulated air/fuel ratios. Catalyst: 0.042 wt % Pt/0.018 wt % Rh/alumina.

73 In 1991, metallocene-catalyzed polymerization became commercial technology, particularly for polyethylene and polypropylene. These two polyolefins account for 35% of all thermoplastics and elastomers. MetaUocene can also readily polymerize bulky monomers, such as styrene, to make novel polymers with physical properties competitive with nylon, polycarbonates and polyesters. Metallocenes are single site catalysts comprising a metal atom sandwiched between parallel planar cyclopentadenyl groups. The most common metals used for olefin polymerization are zirconium, titanium and hafmium. These systems have extremely high activity (as high as 40,000 kg polyethylene/g, metal/hour), which more than compensates for the relatively high catalyst cost and makes the system competitive with Ziegler-Natta catalysts. There are several features that distinguish metaUocene catalysts from other systems. They can polymerize vinyl monomers regardless of the monomers molecular weight or stress hindrance; they produce very uniform polymers of narrow molecular weight distribution; they can polymerize r with very high stereoregularity to give isotactic or syndistactic polymers. The production of some oxygen-containing monomers (e.g., various alcohols or aldehydes) later to be polymerized to fibers and plastics has been simplified and new chemistry has been opened up by the discovery in the early 1990s of the amazing properties of titanium silicates in crystalline structures as oxidation catalysts, which have been named TS-1, etc. Highly selective oxidations can be carried out using hydrogen peroxide (in the presence of oxygen) as oxidant to produce epoxides, among others (Fig. 2). These oxidations can be accomplished Reactants

~

Products

+ H202

-OH

OH

OH

+ H20 OH

+ H202

R~C= C

* H202

+ H20 O R~C/_.-~ C

OH

+ H20

O

C---~C~c._.C

+ H202

/\ C~C~c._. C

R-CH2-OH

+ H202

R-CHO

R~

+ H202

R/CH-OH

R~

R / C -- 0

+ H20

+ H20

+ H20

Figure 2. H202 oxidations catalyzed by TS-1.

74 with dilute aqueous H202 (e.g., 40 wt % H202) with no loss in selectivity which, in almost all cases, is higher than 80%. The synthesis of propylene oxide from propylene with aqueous H202 has a selectivity of 98%. This area is expected to grow rapidly in the next few years. While the titanium silicates have structural similarity to certain zeolites, they are not acidic zeolites and catalyze entirely different reactions. Another area of oxidation that has shown major growth is dehydrogenation and oxydehydrogenation, particularly in the conversion of butane and butenes for MTBE production. Several processes have been developed using metal oxide or precious metal catalysts but, in general, they are improvements of the pre-1956 Houdry butane dehydrogenation and the Texas Butadiene process. Reactor design and process engineering have greatly contributed to the improvements. The oxidation of butanes to compounds, such as maleic anhydride, in the presence of vanadium or molybdenum oxides has gained importance since 1956, again mostly due to design improvements. A new catalytic route to acetaldehyde from ethylene was introduced as the Wacker process. In a partially homogeneous reaction, the ethylene is oxidized with PdCI2 in the presence of water. Reduced Pd metal is oxidized back to PdCI2 with cupric chloride and the resulting cuprous chloride is oxidized with HCI and oxygen. Several other important commercial processes need to be mentioned. They are (not necessarily in the order of importance): the low pressure methanol process, using a copper-containing catalyst which was introduced in 1972; the production of acetic acid from methanol over RhI catalysts, which has cornered the market; the methanol-to-gasoline processes (MTG) over ZSM-5 zeolite, which opened a new route to gasoline from syngas; and ammoxidation of propene over mixed-oxide catalysts. In 1962, catalytic steam reforming for the production of synthesis gas a n d / o r hydrogen over nickel potassium alumina catalysts was commercialized. At times, it is difficult to distinguish between revolutionary new breakthroughs and evolutionary improvements of technology resulting in a replacement of previously used catalysts or processes. An example of this is the introduction in 1967 by Chevron of bimetallic reforming catalysts. Reforming with Pt-alumina catalysts was an important development at the time of the first ICC and many variants of catalyst composition and process operating conditions were introduced in the 1950s and early 1960s. The new bimetallic catalysts, mostly containing rhenium, in addition to platinum and alumina, greatly increased the stability and thus life of the catalysts and also permitted lower pressure operation, increasing the yield of reformate. An additional step was taken by Exxon with the introduction of Pt-Ir catalysts. In general, bimetallic reforming catalysts have largely replaced the Pt-AI203 catalysts. The new reforming catalysts were much more sulfur sensitive than previous ones and that led to hydrodesulfurization catalysts which would lower the sulfur level of reforming feeds to the part per billion level. Hydrocracking, though well-known in 1956, was undergoing major improvements in the 1960s and 1970s by using zeolite (faujasite)-based acidic components, along with tungsten or molybdenum oxides or sulfides as

75 hydrogenation components. Dependence on heavier petroleum has formed greater dependence on hydrocracking processes. Hydrodesulfurization has become vastly more important in the last 40 years. The great sensitivity of metal catalysts to sulfur has necessitated reducing sulfur levels in the feed to reforming processes to the parts per billion level, and the presence of aromatic sulfur and nitrogen compounds in heavy oils requires severe hydrotreating. The catalysts used are upgraded versions of oxides or sulfides of molybdenum, tungsten and cobalt. In the area of pollution control, the removal of NOx from stationary sources effluents, such as power plant stack gases, has been accomplished by use of titaniavanadia catalysts, which promote the reduction of NOx with NH 3 to produce nitrogen and water. During the last 20 years, emphasis on catalytic steps in the synthesis of pharmaceuticals and of agrochemicals has sharply increased. Major progress in novel catalytic reactions has, to some extent, shifted from hydrocarbon conversions and fuels to environmental control and to organic (and, to a lesser extent, inorganic) chemical synthesis. Several fine chemicals are now produced by catalytic reactions (often using zeolites) with greater selectivity and therefore better meet environmentally acceptable standards. Important among biochemicals is "chirality", the "handedness" or optical rotation of the molecule, which determines its bioactivity. One of the chiral isomers can be very effective, while the other may be ineffective or even harmful. A typical example is the case of thalidomide, which is a sedative for pregnant women in its R-form, but is a potent teratogen in its Sform, causing children to be born with deformities (Fig. 3). The U.S. market for chiral compounds is expected to grow from $0.5 billion in 1990 to $3 billion in 2000. Any short review is bound to be incomplete and the selection of successful processes and catalysts expresses only the opinion of the author. There is, however, no doubt about the magnitude of important developments of the last 40 years, with at least ten breakthrough accomplishments, and there is every indication that the trend of the last 40 years will continue. H

H 0 . ~ 0

0

0

0

(R) Sedative

0

(S) Teratogen

Figure 3. Physiological effects of chirality: thalidomide.

This Page Intentionally Left Blank

J.W. Hightower, W.N. Delgass, E. Iglesia and A.T. Bell (Eds.) 11th International Congress on Catalysis - 40th Anniversary

77

Studies in Surface Science and Catalysis, Vol. 101 9 1996 Elsevier Science B.V. All rights reserved.

An example of novel o x y n i t r i d e s or "A1PONs"

basic

catalysts:

the

aluminophosphate

A. Massinon a, E. Gu~guen a, R. Conanec b, R. Marchand b, Y. Laurent b and P. Grange a a Unit~ de Catalyse et Chimie des Mat~riaux Divis~s, Universit~ Catholique de Louvain, Place Croix du Sud 2/17, 1348 Louvain-la-Neuve, Belgium* b L a b o r a t o i r e de Chimie des Mat~riaux, URA 1496 CNRS "Verres C~ramiques', Universit~ de Rennes I, 35042 Rennes Cedex, France

et

New a l u m i n o p h o s p h a t e oxynitrides solid basic catalysts have been synthesised by activation under ammonia of an A1PO4 precursor. When the nitrogen content increases, XPS points out two types of nitrogen phosphorus bonding. The conversions in Knoevenagel condensation are related to the surface nitrogen content. Platinum supported on aluminophosphate oxynitride is an active catalyst for isobutane dehydrogenation. 1. INTRODUCTION The development of solid basic catalysts is nowadays a subject of increasing interest. Modified oxides [1], zeolites [1], hydrotalcites [2] and alkaline-substituted sepiolites [3] have shown interesting activities as solid basic catalysts. Phosphates, in particular aluminophosphates, are known to be good acid catalysts. A new family of aluminophosphate oxynitrides, called "A1PONs", has been prepared by nitridation of reactive aluminophosphate powders [4]. The presence of nitrogen in the three-dimensional network of the oxynitrides results in the creation of basic sites on their surface [5]. This paper reports surface characterisation of a series of aluminophosphate oxynitrides with variable nitrogen contents and their catalytic evaluations in the Knoevenagel condensation and isobutane dehydrogenation. 2. EXPER/MlgNTAL 2.1. Synthesis of oxide, oxynitrides a n d impregnation of Pt and Sn To prepare a high surface area amorphous phosphate precursor A1PO4, the citrate method was used [6]. To reach an Al/P ratio fixed at 1, 0.667 mole of Al(NO3)3.9H20 (Merck) and 0.667 mole of (NH4)H2PO4 (Merck) were dissolved * We acknowledge the financial support of the "R6gion Wallonne", Belgium, for this COST pro~am.

78 in distilled w a t e r u n d e r stirring. Both solutions were mixed t o g e t h e r at room t e m p e r a t u r e . After 1 hour stirring, an excess of citric acid (Merck) was added. The resulting aqueous solution was f u r t h e r stirred overnight. W a t e r was t h e n evaporated under reduced pressure and the obtained white gel was dried for 10 hours at 378 K in a vacuum oven (50 mbar). The gel swelled and looked like a "meringue". The sample was then calcined for 16 hours at 823 K. N i t r i d a t i o n of the oxide precursor was performed u n d e r p u r e a m m o n i a flow. Different aluminophosphate oxynitrides "A1PONs" with variable nitrogen contents were obtained by modifying the time and/or the t e m p e r a t u r e of nitridation (Table 1). A 1.5 wt.% Pt/A1PON catalyst was prepared by i m p r e g n a t i o n of p l a t i n u m on the A1PON support. The Pt/A1PON s a m p l e was o b t a i n e d by i n c i p i e n t wetness i m p r e g n a t i o n with a methanolic solution of H 2 P t C 1 6 . 6 H 2 0 (Merck). The excess of methanol was evaporated u n d e r an argon flow a n d the s a m p l e was then dried at 383 K overnight. The sample was decomposed u n d e r N2 flow w i t h a t e m p e r a t u r e r a m p of 2.5 K.min "1 up to 773 K and held at t h a t t e m p e r a t u r e for two hours, after which the s a m p l e was r e d u c e d in p u r e hydrogen at the same t e m p e r a t u r e during two hours. The addition of tin was accomplished by impregnation with a methanolic solution of Sn(CH3)4 (Merck) on the Pt/A1PON sample, using the same drying and reduction procedures as for Pt/A1PON. 2.2. P h y s i c o - c h e m i c a l c h a r a c t e r i s a f i o n The specific surface areas of the samples were m e a s u r e d by the single point BET method (p/p0=0.3). The total a m o u n t of n i t r o g e n (bulk n i t r o g e n of n i t r i d e - t y p e a n d h y d r o g e n a t e d N H x (x=l to 4) surface species) was d e t e r m i n e d by Grekov titration [7]. The principle of this chemical analysis of nitrogen is based on the reaction of the nitride ions N 3" with a strong base and the f o r m a t i o n of a m m o n i a t h a t is then titrated. In the traditional Kjeldahl method the alkaline a t t a c k occurs in solution. But in some cases, with r e f r a c t o r y nitrides, the products are not totally attacked under these conditions. This has been solved by using the same principle but by h e a t i n g the product at high t e m p e r a t u r e (673 K) with melted potassium hydroxide. The ammonia was dissolved in w a t e r and titrated with sulphuric acid. In the aim to estimate the a m o u n t of surface N H x species (p)S P - H, P- N(H, NH3, PO" . .NH4 + ) adsorbed at the surface the classical method of Kjeldahl was used. X-ray Photoelectron Spectroscopy analysis of the samples was performed with a Surface Science I n s t r u m e n t s spectrometer (SSI 100) with a resolution (FWHM Au 4f7/2) of 1.0 eV. The X-ray beam was a m o n o c h r o m a t i s e d A1Ka r a d i a t i o n (1486.6 eV). A charge n e u t r a l i s e r (flood gun) was a d j u s t e d at an energy of 6 eV. As the C l s spectra of these compounds were very complex, the binding energies were referenced to the binding energy of O l s , considered experimentally to be at 531.8 eV [8].

79 2~3. Catalytic evaluations

2~.1. Knoevenagel condensation F o u r mmoles of malononitrile and b e n z a l d e h y d e were introduced in a batch stirred t a n k reactor at 323 K with toluene as solvent (30 ml). Then 0.05 g of a l u m i n o p h o s p h a t e oxynitride was added. Samples were analysed by gas c h r o m a t o g r a p h y ( I n t e r s m a t Delsi DI200) using a capillary column (CPSi18CB25 m). Care was t a k e n to avoid mass or h e a t transfer limitations. Before the reaction no specific catalyst p r e t r e a t m e n t was done. 2~.2. Isobutane dehydrogenation Catalytic evaluations of isobutane dehydrogenation were carried out in a c o n v e n t i o n a l c o n t i n u o u s flow m i c r o r e a c t o r o p e r a t i n g at a p p r o x i m a t e l y a t m o s p h e r i c p r e s s u r e and using i s o b u t a n e in helium (Air Liquide 0.95% isobutane N25 in 99.05% helium N50) and hydrogen in helium (Air Liquide 1% H2 N50 in 99% He N50). The molar ratio between isobutane and hydrogen was m a i n t a i n e d in each case at 1/6. The reaction products were analysed using an on-line P a c k a r d (model 428) gas c h r o m a t o g r a p h equipped w i t h a flame ionisation detector, with helium as c a r r i e r gas. A 60m x 0.32ram (RSL 160 Alltech) column was used for the s e p a r a t i o n of the various compounds. The space velocity is given as the ratio between the weight flow of isobutane per hour and the weight of the catalyst. Total isobutane conversion is defined as the p e r c e n t a g e of i s o b u t a n e t r a n s f o r m e d into all products. The selectivity to isobutene is defined as the amount of isobutane converted into isobutene divided by the total isobutane conversion. Initial conversion is obtained by extrapolating the curve conversion versus time-on-stream to time zero. 3. R E S U L T S AND DISCUSSION

3.1. Characteristics of samples The characteristics of the studied oxynitrides are reported in Table 1. Table 1 Composition, nitridation t e m p e r a t u r e and time, surface area and nitrogen content of the "A1PONs" (~omposition 'Nitridation Nitridation Surface area Total N Surface N t e m p e r a t u r e (K) time (h) (m2.g-1) (wt.%) (wt.%) A1PO3.64N0.24 1073 3 275 2.8 2.7 A1PO3.55N0.30 1073 8 275 3.6 1.3 A1PO3.10N0.60 1073 40 235 7.2 1.0 A1PO2.67N0.89 1073 65 230 11 1.1 A1PO1.96N1.35 1073 120 215 17.5 2.4 A1PO1.71N1.53 1073 200 195 20 2.6

3.2. C h a n g e of s t r u c t u r e d u e to t h e n i t r i d a t i o n With XPS a good fitting of the N l s spectra is found with one peak of 85% Gaussian and 15% Lorentzian character.

80 It allows to observe a shift on the binding energies values. Indeed a break is evidenced between the behaviour of the samples with nitrogen content lower than 7% and the samples with higher nitrogen content (Figure 1). In the literature, Marchand et al [9] have shown that two N ls peaks exist in the phosphate glasses: a peak at 397.8 eV corresponding to P=N-P bonds and P

a peak at 399.3 eV corresponding to p>N- P bonds. The same trend is observed in our catalysts that suggests the presence of these two kinds of nitrogen coordination in the "A1PON". On the other hand the N l s binding e n e r g i e s of the a l u m i n o p h o s p h a t e O 398.8 oxynitrides "A1PONs" agree with the Z 398.6 N l s binding energies of bulk phosphate o oxynitride ceramic m a t e r i a l PON [10], >., 398.4 ea~ situated at 317.8 and 399.3 eV, r a t h e r 398.2 t h a n the N l s b i n d i n g e n e r g i e s of "A1ON" family [11,12], found between 398.0 396.3 and 396.6 eV. M o r e o v e r a = 397.8 progressive decreasing t r e n d of the I I I binding energies of P2p (from 134.1 to 397.6 0 5 10 15 20 133.5 eV) is observed when the atomic n i t r o g e n percent age increases. This Total nitrogen content (% bulk) behaviour can be explained by the better Figure 1. Variation of N l s binding nucleophilic c h a r a c t e r of n i t r o g e n c o m p a r e d to oxygen, r e d u c i n g the energy. positive charge a r o u n d p h o s p h o r u s atoms. Both observations indicate that, during nitridation, nitrogen seems to substitute more easily the oxygen atoms present in an environment of phosphorus r at her than in the environment of aluminium. >

399.0

cD

.,..

3dl. K n o e v e n a g e l condensation To evaluate properties of basic catalysts, the Knoevenagel condensation over aluminophosphate oxynitrides was investigated [13]. In this reaction usually catalysed by amines, the solid catalysts function by abstraction of a proton from an acid methylene group, which is followed by nucleophilic attack on the carbonyl by the r e s ul t ant carbanion, re-protonation of oxygen and e l i m i n a t i o n of water. The c o n d e n s a t i o n bet w een b e n z a l d e h y d e and malononitrile is presented below. To check if this liquid-phase reaction is not controlled by diffusion, the reaction is repeated with different weights of catalysts. A linear correlation is found between the initial rate of reaction and the weight of catalyst, indicating that the rate is not controlled by external or internal diffusion.

81

C +

\

NC

N

H

Base i

"

2

i

1

X

N

CN

H

/N +

H20

The i n t r i n s i c conversion (%.m "2) of m a l o n o n i t r i l e and b e n z a l d e h y d e v e r s u s time is shown in Figure 2. Commercial MgO (40 m2.g "1) used in the s a m e conditions as the "A1PONs" (i.e. w i t h o u t p r e t r e a t m e n t ) gives a low conversion [14]. W i t h o u t p r e t r e a t m e n t MgO is not an e x t r e m e l y basic compound but these results show t h a t "A1PONs" are more active t h a n MgO at the chosen conditions and such a c h a r a c t e r could be useful for i n d u s t r i a l applications.

10 O 9~

8

~

6

A

A

8

A

d

6 ~

0

~4 9~

,10

&

2:

"A

m

2 ~

2 D

0

0

100

200

O |

300

Time (rain) Figure 2. Conversion of malononitrile vs time in Knoevenagel condensation (O 2.8%N, 9 3.5%N, ~ 7.2%N, m l l % N , A 17.5%N, A 20%N).

10

0 ~ 20

Nitrogen content (%) Figure 3. Intrinsic conversion at 300 rain and surface N content vs total nitrogen content.

Figure 2 shows t h a t the conversion does not depend directly on the total n i t r o g e n content of the c a t a l y s t as it was seen before for the "A1PONs" synthesised by the sol-gel method [5]. W h e n the N H x surface groups (Table 1) are only considered, a good correlation with the malononitrile condensation is obtained (Figure 3). A c t u a l l y N H x surface species are not the only active species in the Knoevenagel reaction. Indeed the catalysts with 2.8, 17.5 and 20 wt.% bulk nitrogen have the same NHx surface species content but their activities are

82 very different. The 2.8 wt.% bulk nitrogen sample can be a s s u m e d w i t h o u t nitride bulk nitrogen because its total nitrogen content and its surface nitrogen content (2.7 wt.%) are equivalent. On the other h a n d the 17.5 and 20 wt.% bulk nitrogen content samples contain the highest quantity of nitride bulk nitrogen. The higher the nitride nitrogen content, the higher the activity. The s t r e n g t h of the basic sites is not yet evaluated. Corma et al. [3,15] d e m o n s t r a t e d t h a t using methylenic compounds with different acidities, a basicity range can be determined for their samples. However the condensation is carried out without solvent. Using toluene as solvent the reactions between b e n z a l d e h y d e a n d m a l o n o n i t r i l e ( C N - C H 2 - C N - pKa = 11.2 [ 1 6 ] ) o r e t h y l c y a n o a c e t a t e ( C 2 H 5 C O 2 C H 2 C N - pKa < 9 [16]) are studied on the "A1PONs". N e v e r t h e l e s s the conclusions are not so evident in toluene. In toluene the condensation of malononitrile is easier t h a n the condensation of ethylcyanoacetate whereas the opposite behaviour was expected. Indeed if the first step (abstraction of the acid proton) is limiting, the abstraction of acidic p r o t o n of e t h y l c y a n o a c e t a t e should be e a s i e r t h a n the a b s t r a c t i o n of malononitrile one due to the lower pKa. To check if this behaviour does not depend on our catalysts, both condensations are done using liquid bases: pyrrolidine (pKa = 11.3 [17]) and n-nonylamine (pKa = 10.6 [17]). The results found with both liquid bases present the same trend but experimental data do not lead to conspicuous results because a steric h i n d r a n c e with the ester function or the better charge stabilisation of ethylcyanoacetate can occur and the second step could become limiting ( c o n d e n s a t i o n b e t w e e n the two reactants). So both facts lead to a lower conversion with e t h y l c y a n o a c e t a t e because the conversion does not depend on the pKa of the reactifs. On the other h a n d , s o l v e n t effects could play a role d u r i n g the r e a c t i o n . F u r t h e r experiments are needed to solve this doubt. 3.4~ I s o b u t a n e d e h y ~ e n a t i o n The oxynitride "A1PON" used as support was synthesised as the others. It is well represented by the A1PO3.16N0.56 formula, its nitrogen content being 6.7 wt.%. This sample has a surface area of 310 m2.g -1. 3.4.1 Influence of t h e noble metals s u p p o r t e d o n a l u m i n o p h o s p h a t e oxynitride T h r e e catalysts with different compositions were used in this study. Catalyst A corresponds to the oxynitride precursor. Catalysts B and C contain 1.5 wt.% of platinum, however catalyst C contains 0.91 wt.% tin as well corresponding to an atomic ratio Pt/Sn of 1. For q u a n t i t a t i v e comparisons between catalysts, care was t a k e n to ensure t h a t the kinetic d a t a were not influenced by mass or heat transfer. Figure 4 shows the evolution of the initial conversion versus t e m p e r a t u r e at a space velocity of 0.03 h -1. The equilibrium conversion of i s o b u t a n e to i s o b u t e n e is 100% in our conditions. An increase of the conversion with t e m p e r a t u r e up to 773-823 K is observed. When metals were added, we also noted a large increase in isobutane dehydrogenation. Table 2 gives initial isobutane conversions, isobutene selectivities and yields of the reaction at 823 K for the three tested samples.

83

100 A

80

9

o

.0,,q

~

60

= o

40 20

I00 ~ ' - ~

&

me

80 ~

9

dD

60 40

T

A

I=

0.~ - $ ~ ~--o--o--o--c 373 473 573 673 773

o

20"

o

O

0

873

40

60 Time (h)

Temperature (K) Figure 4. Isobutane initial conversion vs temperature (C) A1PON, I Pt/A1PON, A Pt-Sn/A1PON

20

Figure 5. Isobutane conversion and isobutene selectivity vs time for Pt-Sn/A1PON (O conversion, 9 selectivity).

Table 2 Initial isobutane conversions, selectivities and ~elds of the reaction at T=823 K catalyst A B C Isobutane conversion (%) 2.7 70.5 84.5 Isobutene selectivity (%) 50.5 44.8 38.4 Yield of the reaction (%) 13.6 31.6 32.4 At first, we show that adding metals to the oxynitride affects significantly the conversion while keeping isobutene selectivity at a high level. The byproducts of the reaction are due to isomerisation, hydrogenolysis and coking reactions. The addition of tin to the Pt/A1PON catalyst increases the conversion. Platinum-tin alloy formation, as already suggested in literature [18], or improvement of the metal dispersion cannot yet be proposed. Figure 5 shows the evolution of conversion and selectivity versus time for catalyst C, which presents the best yield. This catalyst shows a significant deactivation probably due to the coke deposit. 3.4.2 I n f l u e n c e o f t h e s p a c e v e l o c i t y WHSV

Two space velocities, i.e. 0.03 and 0.3 h -1, have been used in the evaluation of catalytic activities of catalysts B and C at 823 K. Figure 6 shows a decrease in activity of the catalyst B when space velocity increases. The accessible sites are saturated at the lowest space velocity. This explains thus the lower conversion levels at a higher space velocity. However, for catalyst C, the evolution of the conversion, which is also depicted in Figure 6, is almost identical for both space velocities. This result could be explained by a better dispersion of the platinum due to the presence of tin.

84 lOO

~

so

~

60

~

40

~

9o

20

o

0

9I n ~

10

20

9

30

40 50 Time (h)

F i g u r e 6. I n f l u e n c e of WHSV on conversion at 823 K (O Pt-Sn/A1PON at 0.03 h -1, 9 Pt-Sn]A1PON at 0.3 h -1, Q P t / A 1 P O N at 0.03 h -1, i Pt/A1PON at 0.3 h "1)

In the presence of tin, the n u m b e r of active p l a t i n u m sites s e e m s to be superior compared to c a t a l y s t B a n d t h u s an increase in the space velocity by a factor of t e n does not seem to s a t u r a t e all t h e sites. T h e s e r e s u l t s show the importance of the role played by tin since p l a t i n u m loading was the same in both cases. It is r e a s o n a b l e to think t h a t in the case of catalyst B, due to the m e t h o d of d e p o s i t i o n , some aggregates of p l a t i n u m are formed on the surface of the c a t a l y s t . In t h e presence of tin, a part of the aggregates could d i s a p p e a r and some P t / S n alloy p a r t i c l e s , b e t t e r d i s p e r s e d at t h e surface, could be formed.

3.4~q. C o m p a r i s o n w i t h MgO F i g u r e 7 compares in function of time the activities of c a t a l y s t B a n d platinum s u p p o r t e d on magnesium oxide MgO, a well-known basic i n d u s t r i a l catalyst. 4O 30 20

10~I~ 0

0

oo oooooooooooo .......

1=====o~p

10

o ~

9

20 30 Time (h)

The MgO support has been synthesised by a c o p r e c i p i t a t i o n r e a c t i o n a n d presents a surface a r e a of 300 m2.g -1 t h a t is almost identical to the surface a r e a of A1PON (310 m2.g-1). P t / M g O h a s b e e n p r e p a r e d by t h e s a m e procedure as for Pt/A1PON. We could see t h a t the activity of Pt/MgO is very low compared to the activity of the Pt/A1PON under the same conditions.

F i g u r e 7. C o m p a r i s o n b e t w e e n O Pt/A1PON and 9 Pt/MgO. 4. C O N C L U S I O N A novel basic support and catalyst have been p r e p a r e d by activation of a l u m i n i u m p h o s p h a t e with ammonia. Fine control of time and t e m p e r a t u r e allows to adjust the O/N ratio of these oxynitride solids and thus to t u n e the acid-base p r o p e r t i e s . The a l u m i n o p h o s p h a t e o x y n i t r i d e s a r e active in K n o e v e n a g e l condensation, but a basicity r a n g e can not yet d e t e r m i n e d . Supporting Pt or Pt/Sn on A1PONs allows to prepare catalysts t h a t are highly active and selective in dehydrogenation reactions.

85 R~~CF~ o

o

3. 4. o

o

o

o

o

10. 11. 12. 13. 14. 15. 16. 17. 18.

H. Pines and W.M. Stalick, Base catalyzed reactions of hydrocarbons and related compounds, Academic Press N.Y. 1977. F. Cavani, F. Trifiro and A. Vaccari, Catal. Today, 2 (1991) 11. A. Corma and R.M. Martin-Aranda, J. Catal., 130 (1991) 130. R. Conanec, R. Marchand and Y. Laurent, High Temp. Chem. Process, 1 (1992) 157. P. Grange, Ph. Bastians, R. Conanec, R. Marchand, Y. Laurent, L.M. Gandia, M. Montes, J. Fernandez and J.A. Odriozola, Preparation of Catalysts VI, Elsevier, Amsterdam, p.381, 1994. Ph. Courty, H. Ajot, Ch. Marcilly and B. Delmon, Powder Technology, 7 (1973) 21. F.F. Grekov, J.Guyader, R. Marchand and J. Lang, Rev. Chim. Min., 15 (1978) 341. A. Massinon, J.A. Odriozola, Ph. Bastians, R. Conanec, R. Marchand, Y. Laurent and P. Grange, Appl. Catal., in press. R. Marchand, D. Agliz, L. Boukbir and A.Qu~merais, J. Non-Cryst. Solids, 103 (1988) 35. R. Marchand, Y. Laurent and A. Qu~merais, Rivista della Staz. Sper. Vetro, 5 (1990) 101. J.J. Benitez, M.A. Centeno, J.A. Odriozola, B. Viot, P. Verdier and Y. Laurent, J. Mater. Chem., submitted. H.M. Liao, R.N.S. Sodhy and T.W. Coyle, J. Vac. Sci. Technol. A, 11 (1993) 2681. B.P. Mundy and M.G. Ellerd, Name Reactions and Reagents in Organic Synthesis, Wiley-Interscience, New York, 1988. P. Grange, Ph. Bastians, R. Conanec, R. Marchand and Y. Laurent, Appl. Catal. A: 114 (1994) L191. A. Corma, V. Forn~s, R.M. Martin-Aranda and F. Rey, J. Catal.,134 (1992) 58. R.G. Pearson and R.L. Dillon, J. Am. Chem. Soc., 75 (1952) 2439 Handbook of Chemistry and Physics, Editors R.C. Weast and M.J. Astle, CRC Press, Florida, 61st Edition 1980-1981, D-161 R.D. Cortright and J.A. Dumesic, J. Catal., 148 (1994) 771.

This Page Intentionally Left Blank

J.W. Hightower, W.N. Delgass, E. Iglesia and A.T. Bell (Eds.) l lth International Congress on Catalysis - 40th Anniversary

Studies in Surface Science and Catalysis, Vol. 101 9 1996 Elsevier Science B.V. All rights reserved.

87

R e a c t i o n K i n e t i c s of t h e H y d r o d e n i t r o g e n a t i o n of D e c a h y d r o q u i n o l i n e o v e r NiMo(P)/A1203 C a t a l y s t s M. Jian and R. Prms Laboratory for Technical Chemistry, Swiss Federal Institute of Technology (ETH), CH-8092 Zurich, Switzerland

ABSTRACT The reaction mechanism and the kinetics of hydrodemtrogenation (HDN) of decahydroqumoline (DHQ) over NiMo(P)/A1203 catalysts were studied in the presence and absence of H2S. Cis- and trans-propylcyclohexylamine were identified as the most important reaction intermediates; their reactivity was found to vary under different reaction conditions. The kinetic constants in the HDN network of DHQ were calculated assuming a Langmuir-Hmshelwood mechanism. It was found that, in the phosphorus-containing catalysts, the adsorption of DHQ was enhanced and the rate constant of the first C-N bond cleavage decreased. As a consequence, the overall HDN reaction rate of DHQ decreased when phosphorus was added to a NiMo/A120,~ catalyst. The effect of H2S was the opposite of that of phosphorus: it increased the rate constants and decreased the adsorption constants. The inhibiting effect of DHQ on its own conversion was different from that on the hydrogenation of cyclohexene, proving that different catalytic sites are involved in these reactions.

1. I N T R O D U C T I O N Hydrodenitrogenation(HDN) is an important process in petroleum refining. It removes nitrogen from oil distillates, so that less NOx pollutes the air when oil is burned and poisoning of the subsequent refining catalysts is reduced when the oil is processed further. Although HDN has been studied intensively and different reaction mechanisms, catalytic active sites, and functions of the catalytic components have been proposed, there are still many questions to be answered in order to better understand the reaction and the catalyst (1-4). Three kinds of reactions are revolved in the HDN process (5): C-N bond cleavage, hydrogenation of the aromatic ring, and hydrogenation of the nitrogencontaining aromatic heterocyclic ring. C-N bond cleavage is usually the final and

88 most important reaction step m the HDN process, since the nitrogen atom can only be removed after the C-N bond is broken. However, in most kinetic and mechanistic studies, lumped Langmuix-Hinshelwood kinetics were used, and it was assumed that all HDN reaction steps (hydrogenation of aromatics and olefins, elimination of ammonia, etc.) take place on the same catalytic site (1). As a consequence, the functions of the catalytic components of the catalyst in the ldnetic network could not be completely unravelled; m a n y assumptions had to be made to simulate the ldnetic network, and the resulting kinetic constants differed from one author to the other (6, 7). In the widely used industrial NiMo-P/A120,~ HDN catalyst, the promotional effect of phosphorus has been demonstrated by its industrial performance. In model HDN compound studies, however, phosphorus promoted the HDN of quinolme (8), but decreased the HDN of decahydroqumoline and piperidme (9). A negative effect of phosphorus has also been observed on hydrogenation (10). The question, therefore, arises as to the role of phosphorus in an HDN catalyst. In the present study, the HDN of decahydroquinoline (DHQ) was studied over NiMo(P)/A120.~ catalysts in the presence and absence of H2S. The reaction took place at 593 K and 3.0 MPa, thus allowing us to observe the most important reaction intermediate, propylcyclohexylamme, and to calculate the kinetic constants from the experimental results. Rate and adsorption constants for the different reaction steps were determined by separate and by combined HDN studies of DHQ and cyclohexene.

2. E X P E R I M E N T A L The NiMo/A120:, catalysts (3 wt% Ni, 8 wt% Mo and 0 or 2 wt% P) were prepared by the incipient wetness impregnation method as described in ref. (10). The HDN reactions were carried out in a continuous-flow microreactor. A sample of 0.1 g catalyst diluted with 9.5 g SiC was first sulphided in situ with a mixture of 10% H,~S and H2 at 643 K and 1.5 MPa for 4 h. After sulphidation the pressure was increased to 3.0 MPa, and a solution of the reactant in n-octane was fed to the reactor by means of a high pressure pump. The initial reactant concentration (Ao) was adjusted by changing the reactant concentration. Dimethyldisulphide was added to the solution to generate H2S m the reaction stream (P(H2S)=6.53 kPa). The influence of the NH.~ product was determined by co-feeding pentylamine as a source of NH3 (P(NH3)=l.59 kPa). Reaction products were analyzed by on-line gas chromatography with a Shimadzu GC-14A gas chromatograph equipped with a 50 m CP Sil-5 fused silica capillary column and a flame ionization detector. Reaction intermediates were identified by GC-MS. Samples were taken after 50 h on stream when the activity of the catalyst was stable, with n-nonane and n-dodecane as internal standards. Space time was defined as ~ = e-Vc,d v ~ , where ~ is the void fraction of the

89

catalyst bed (e was assumed to be 0.4), Vat is the catalyst volume, and v$~ is the volume flow rate of the gas phase reactant.

3. R E S U L T S AND DISCUSSION The HDN reaction network of qumolme is shown in Fig. 1 (7, 8). A critical reaction step in this network is the breaking of the C-N bond through the consecutive THQ I ~ O P A ~ H C or D H Q ~ H C reaction (HC= hydrocarbons). Since equilibria can easily be established between Q, THQ5, THQ 1, and DHQ, all three main types of reactions in the HDN process (C-N bond cleavage, aromatics hydroge-nation, and aromatic heterocycle hydrogenation) are usually involved m the HDN of quinoline-type compounds. It was found that the rate constants of some reaction steps (such as T H Q 5 ~ D H Q and DHQ~PCHE) were of the same magnitude, so that a simple rate-limiting step treatment is not always possible (7, 11). Disagreement exists as to whether the HDN reaction proceeds mainly via THQ 1-->OPA~HC or via DHQ--->HC (1, 12-14). In almost all k~netic studies, several reactions were grouped together; thus the calculated ldnetic constants did not give a clear indication of reaction mechanism and rate-limiting steps and even less so of the catalytic sites and functions of the catalytic components. One of the most important reaction intermediates PCHA, has been observed only rarely (6). Q

THQ-1

OPA

It

It

It

THQ-5

DHQ

PCHA

PB

i

s,, o~

PCHE

,,-...., 1l PCH

Figure 1. HDN reaction network of quinolme-type compounds. Q=quinolme, THQ5=5,6,7,8-tetrahydroquinoline, DHQ=decahydroqumoline, THQ l=l,2,3,4tetrahydroqumi~ne OPA=ortho-propylaniline, PCHA=2-propylcyclohexylamine, PCHE=propylcyclohexene, PCH=propylcyclohexane, PB=propylbenzene. 3.1. T h e H D N r e a c t i o n n e t w o r k o f d e c a h y d r o q u i n o l i n e DHQ is an important intermediate in the HDN network of quinolme; only C-N bond cleavage is needed to remove the nitrogen atom from DHQ. However, since dehydrogenation of DHQ proceeds qmte fast, THQ5, THQ1, and Q are usually present as well, thus making the study of the HDN of DHQ difficult.

90 Table 1 Product compositions (%) in the HDN of DHQ at 593 K and 3.0 MPa (~= 0.20 sec) catalyst

H,_,S PCH PCHE

PB

PCHA DHQ T H Q 5 0 P A

Q

THQ1 others

NiMo

yes

1.81

2.20

0.02

2.89

89.1

2.85

0.03

0.15

0.26

0.72

NiMoP

yes

1.33

2.09

0.02

2.34

89.0

3.38

0.04

0.10

0.34

1.47

NiMo

no

0.57

0.78

0.07

0.16

92.6

4.43

0.04

0.18

0.35

0.86

NiMoP

no

0.81

1.21

0.13

0.16

91.7

4.38

0.08

0.25

0.42

0.82

In the present study our reaction system and sensitive analytic technique allowed us to perform the HDN reaction of DHQ under such reaction conditions t h a t only small amounts of Q, THQ-1, and OPA were formed (Table 1). This indicates that dehydrogenation of the carbocyclic ring of DHQ was slow and could be neglected. Therefore, the reaction network can be simplffied as in Fig. 2. THQ-5

DHQ

PCHA

KA

KB

PCHE

PCH

KC

Figure 2. Simplified HDN reaction network of decahydroqumoline The mass balance calculation showed that the performance of the reactor was reliable, and there were no diffusion limitations. The operating conditions (low temperature and high space velocity) allowed us to identify both c- and t-PCHA and to confirm t h a t they are the reaction intermediates between c- and t-DHQ and PCHE (Fig. 2) (gas chromatography gave three peaks which, according to the mass spectrum, all belong to PCHE). The mass spectrum of t-PCHA consists of a strong peak at m/e=141 (molecular ion) and small peaks at m / e = l l l and 56, while the mass spectrum of c-PCHA consists of a middle-sized peak at m/e=141, a strong peak at m/e=56, and small peaks at m/e=l 11, 98, 70, and 43. A further analysis of the steric isomers of PCHA (Table 2) shows that the reaction of c-DHQ to c-PCHA was faster than t h a t of t-DHQ to t-PCHA, that the isomerization of c- and t-DHQ was fast, and the isomeration of c- and t-PCHA was somewhat slower. The presence of PCHE in the HDN product (Table 1) indicates that at least part of the HDN reaction of PCHA proceeds through elimination of ammonia r a t h e r than by direct hydrogenolysis. Nevertheless, the direct product of the elimination reaction (allylcyclohexylamine) was not observed. This must be due to its strong adsorption and fast hydrogenation to PCHA.

91

Table 2 Ratios of trans-/cis-isomers at different space times over a NiMo/A1,_,O:3catalyst * space time (ms)

0

38

57

112

203

340

512

t-/c-PCHA

/

2.3

3.0

3.5

3.9

4.5

4.7

0.3

5.3

6.2

6.5

6.7

6.8

6.8

t-/c-DHQ

* 593 K, 3.0 MPa, H,_,S/H2 - 3.0x10 3 mol/mol. The concentration of PB m the reaction products (Table 1) is too high to be accounted for by dehydrogenation of PCHE, especially m the absence of H2S, since at thermodynamic equilibrium the PCH/PB ratio should be greater than 50 (15) under our experimental conditions. Furthermore, no toluene was observed m the simultaneous reaction of methylcyclohexene and DHQ under the same reaction conditions. Therefore, there must be another reaction path to account for the formation of PB m the HDN products of DHQ. This second reaction path can only be the reaction of D H Q ~ T H Q I ~ O P A ~ HC. It has been demonstrated that a relatively high concentration of PB is present in the HDN of OPA due to the direct hydrogenolysis of the C(spe)-N bond of OPA (16). 3.2. I n f l u e n c e o f s p a c e t i m e a n d i n i t i a l c o n c e n t r a t i o n The conversion of DHQ to C-N bond cleavage products as a function of space time (Fig. 3A) demonstrates t h a t PCHA is a p r i m a r y HDN product and that PCH is a secondary HDN product. ) 1 PCH J P C H E

A

0.20

0.15 PCHA x"

o- r ,x~ (..-

~" 0.05 O.,F~ 0.0

._,

r 0.2

. •

, 0.4

space time, s

~ PB 016

o . o o .,.

0.0

0:1

0:2 0:3 0:4 space time, s

0:5

0.6

Figure 3. Effect of space time on the HDN of DHQ at 593 K, 3.0 MPa and P(H,,S) = 6.53 kPa. A: product composition, B" first-order kinetic fitting. The ~ "" t plot (XDHQ -- PCHA 4- PCHE + PCH + PB) could be fitted with a straight line (Fig. 3B), but, under our reaction conditions (small x), this does not distinguish between first and zero order. Another way to distinguish

92 between first and zero order reactions is to measure the effect of the initial concentration on the rate. At low conversion (XDHQ < 0.1), the concentration of DHQ in the reaction stream can be regarded as constant, that is A,~,K~ = Ao. Therefore, assuming t h a t a Langmuir-Hinshelwood mechanism applies, t h a t all the reactants, intermediates, and products adsorb on one and the same site, and t h a t the adsorption of hydrocarbons, NH3, and HeS can be neglected, we obtain dA dt

k,KAA I+ZK~I

k,KAA I+K AAo

where A stands for the concentration of DHQ (Ao is the imtial concentration), k, is the rate constant, and KA is the adsorption constant of DHQ (cf. Fig. 2). The slope of the - I n ( 1 - X D H Q ) "~ t plot is equal to k~KA/(I+K.~Ao). By varying Ao, k, and KA can be calculated from the slope" against Ao plot. The plots are shown in Fig. 4, and the resulting kinetic constants are given in Table 3. The adsorption constants KA show that, in the present concentration range (Ao > 2 kPa), the order of the reaction is between zero and one. 25

25

2O

20, v =-

15

NiMoPIAI203

15

o v +

+

"" 0

|

,

,

,

2

4

6

8

Ao, kPa

5

10

10

Ao, kPa

Figure 4. Graphs used in the determination of the kinetic constants of the HDN of DHQ at 593 K and 3.0 MPa. I1: with H2S,o: no H2S. As we have already seen from the HDN product yields (cf. Table 1 for the compositions after 0.2 s), phosphorus has a negative effect on the HDN of DHQ in the presence of H2S but a positive effect in the absence of H2S. The kinetic constants (Table 3) show that introducing phosphorus to a NiMo/A120:3 catalyst increases its adsorptivity toward DHQ but decreases the rate constant of the first C-N bond cleavage in the presence of H2S. The effective rate constant k,K.~/(I+K~) of the first C-N bond cleavage (rate limiting step) in the presence of H,_,S is lower for the NiMoP/A1,_,O~ catalyst, which accounts for the negative effect of phosphorus. Likewise, the effective rate constant is higher for the NiMoP/A120~ catalyst in the absence of HeS.

93 Table 3. Kinetic constants of the HDN of DHQ at 593 K, 3.0 MPa catalyst

HeS

kl

KA

k2KB

k.3Kc

NiMo/A1203

yes

2.6

0.4

25

23

NiMoP/A120:3

yes

1.8

0.5

26

34

NiMo/A1903

no

0.4

1.4

360

/

NiMoP/A1203

no

0.6

2.0

540

130

The reaction p a t h D H Q ~ T H Q I ~ O P A ~ H C must be taken into account to explain the promotional effect of phosphorus in the absence of H2S. A strong promotional effect of phosphorus has been observed for the HDN of OPA over NiMo/A1,.,O~ catalysts, which could be explained by the larger adsorption constant of OPA on the P-containing catalyst. The HDN activity of OPA was even higher in the absence of H~S (16). If we assume that all the PB produced in the HDN of DHQ under our present reaction conditions was formed through the reaction of OPA, and that the PCH/PB ratio is the same for the reaction path through OPA in the HDN of DHQ and in the HDN of pure OPA, then the HC/PB ratios from the HDN of OPA can be applied to the HDN of DHQ under the same reaction conditions; the relative contributions of the two reaction pathways can be estimated. The results show that about 40% of the HDN reaction of DHQ takes place through the reaction path D H Q ~ T H Q I ~ O P A ~ H C in the absence of H,.,S but less than 10% in the presence of H,_,S. The very low concentration of PCHA (compared with that in the presence of H,_,S) also indicates that the rate limiting reaction steps might have changed in the absence of H2S.

3.3. A d s o r p t i o n c o n s t a n t s from h y d r o g e n a t i o n of CHE The inhibiting effect of DHQ and its NH~ product was studied on the final step in the network of Fig. 2, the alkene hydrogenation. To avoid confusion with the PCHE olefm formed from DHQ, cyclohexene (CHE) was used as the reactant, and pentylamine (PA) was used as the source of NH~. When the hydrogenation of CHE is performed in the presence of NH~, we have dC dt

kcKcC I+KcC+KNN

kcKcC I+KNN

(C stands for CHE and N for NH~)

Assuming that the adsorption of NH.~ and the hydrogenation of CHE take place on the same catalytic site, and that the adsorption of NH3 is much stronger than that of CHE, we obtain, with N - No = constant,

94

-In(1 - x c ) =

kcKc 9t l+KNN o

where kc is the rate constant and Kc the adsorption constant of cyclohexene. kcKc can be calculated from the first order hydrogenation of CHE alone, while KN can be calculated from the slope of the -ln(1-XCHE)~t plot of the hydrogenation of CHE in the presence of NHa.

.

NiMo/AI203

NiMoPIAI203

/

/

6

6

/

/ /

/

,V/

~//

!

v

2 0

0.0

"~" 2

0:2

024 space time, s

"

!

o 0.o

026

,/

-

,

-

-

i

O.2 0.4 space time, s

9

-

i

0.6

Figure 5. Hydrogenation of CHE at 593 K, 3.0 MPa. 9in the presence of H,_,S, ...... "in the absence of H,_,S; i - C H E alone, I : CHE+PA ( P A - 1.59 kPa) At small space times, the C5 hydrocarbons could not account for all the PA converted, although no PA was found in the reactor outlet; a similar observation was reported by La Vopa and Satterfield (17). But at higher space time the C5 hydrocarbons do account for more t h a n 90% of the PA converted. Thus only the data at very high space time were used to calculate KN (Fig. 5). The adsorption constants of DHQ were obtained in the same way. The m a s s balance of DHQ and CHE was always good. The resulting adsorption constants are given in Table 4. Table 4. Adsorption constants (kPa-~) obtained from the hydrogenation of CHE with H2S

no H2S

KDHQ

Kr,m

NiMo/A120.~

2.7

1.0

7.0

3.5

NiMoP/A1,_,Oa

4.1

1.7

11.0

4.8

catalyst

KDHQ

Kr,m

Table 4 shows t h a t the P-containing NLMoP/A1203 catalyst favours the adsorption of NH~ and of DHQ. The adsorption of NH~ is strongly enhanced m the absence of H2S for both catalysts. The effect of phosphorus, which is applied

95 in the form of phosphate, is certainly not restricted to an increase or decrease in the number of sites, since then only the rate constants, and not the adsorption constants, should have changed. Apparently, phosphorus changes the chemical properties of the catalytic sites. A comparison of the results in Tables 3 and 4 shows t h a t the adsorption constants differ substantially when different reactions are inhibited by the same molecule. The adsorption constants KDt~ are much smaller when they are determined from the HDN of DHQ itself than from the inhibition effect of DHQ on CHE. This confirms that different catalytic sites are needed for these chemically different reaction steps m the HDN process. This is also supported by the observation t h a t H2S promotes the C-N bond cleavage reaction of DHQ and piperidine, while it inhibits the hydrogenation of alkenes (10, 12). Therefore, one and the same catalytic site can not be responsible for these different reaction steps in the HDN kinetic studies. It also means that care should be exercised with the choice of the molecule and reaction when determining the adsorption constants from the inhibiting properties of compounds on certain reactions (17), because different catalytic sites might be involved for different reactions (HDN, HDS, hydrogenation, etc.). The presence of different catalytic sites for the first C-N bond breaking of DHQ and the hydrogenation of CHE is confirmed by the -ln(1-x~liQ) versus -In(1-x(nt~:) plot. In the simultaneous reactions of A and B, in which A and B are both adsorbed on the same catalytic site and follow a Langmuir-Hmshelwood mechanism, we have dA kAKAA dt = 1+ ZK,I

-

and

dB kBK~B - d--t-= 1+ ZK~I '

and thus

kAKA ln(l

In(l- x A) = kBK~

x B)

A plot of ln(1-xA) ~ In(1-xs) should result in a straight line if the assumption is true. Figure 6 shows the result of simultaneous reactions of DHQ and CHE. The curvature of the ln(1-XDHQ) ~ ln(1-XCHE) plot confirms that the adsorption sites for CHE and DHQ over the NiMo(P)/AI,,O~ catalysts are not the same.

1.0 0.8 W

-,-0.6 ,x0

Figure 6. Simultaneous reaction of DHQ and CHE at 623 K, 3.0 MPa, and H2S/I-I2 = 3.0x10 3 (tool/tool) over the NiMoP/Al~O3 catalyst.

~. 0.4 0.2 0.0 . . . . . . . . . . 0.0

0.1

0.2 0.3 -In(1 "•

0.4

0.5

95 4. C O N C L U S I O N S The present results show that the separate steps in an HDN reaction network can not be lumped together into one kinetic equation. The intermediate reactions may take place on different catalytic sites which differ m their ability to bind reactants, intermediates, and products. Phosphorus was found to modify the rate constants as well as the adsorption constants of the HDN reaction steps, indicating th at it changes both the number and nature of the active sites of NiMo/AI203 catalysts. The crucial reaction intermediate PCHA in the HDN network of quinolmetype compounds has been clearly observed. Formation of cis-PCHA was faster than that of trans-PCHA, but isomerization was relatively rapid. The presence of H2S m the reaction stream favours the cleavage of the first C-N bond in DHQ, but slows down the C-N bond cleavage in PCHA. The presence of H2S decreases the adsorption constants of DHQ and NH.~. It is concluded that ~-40% of the HDN reaction of DHQ takes place through the reaction path of D H Q ~ T H Q I ~ O P A ~ H C at 593 K and 3.0 MPa in the absence of H2S, while less than 10% takes place m the presence of H,_,S.

REFERENCES

1. M. J. Girgis and B. C. Gates, Ind. Eng. Chem. Res., 30 (1991) 2021. 2. H. Schulz, M. Schon and N. M. Rahman, Stud_ Surf. Sci. Catal., 27 (1986) 201. 3. R. Phns, V. H. J. de Beer and G./~ Somorjai, Catal. Rev. Sci. Eng., 31 (1989) 1. 4. T. C. Ho, Catal. Rev. Sci. Eng., 30 (1988) 117. 5. G. Perot, Catal. Today, 10 (1991) 447. 6. S. S. Shah, K. N. Mathur, J. R. Katzer, H. Kwart and A. B. Stiles, Prepr. Am. Chem. Soc., Div. Pet. Chem., 22 (1977) 919. 7. C. N. Satteriield and S. H. Yang, Ind. Eng. Chem. Proc. Des. Dev. 23 (1984) 11. 8. S. Eijsbouts, J. N. M. van Gestel, J. A. R. van Veen, V. H. J. de Beer and R. Prms, J. Catal., 131 (1991) 412. 9. M. Jian and R. Prms, Catal. Letters, 35 (1995) 193. 10. M. Jian, J. L. Rico Cerda and R. Prins, Bull. Soc. Chim. Belg., 104 (1995) 225. 11. J. L. Rico Cerda and R. Prms, Bull. Soc. China. Belg., 100 (1991) 815. 12. C. N. Satterfield and S. Gulteldn, Ind. Eng. Chem. Proc. Des. Dev. 20 (1981) 62. 13. S. H. Yang and C. N. Satterfield, Ind. Eng. Chem. Proc. Des. Dev. 23 (1984) 20. 14. K. S. Lee, H. Abe, J. A. Reimer and A. T. Bell, J. Catal., 139 (1993) 34. 15. J. F. Cocchetto and C. N. Satterfield, Ind. Eng. Chem. Proc. Des. Dev., 20 (1981) 49. 16. M. Jian and R. Prms, Catal. Today, in press. 17. V. La Vopa and C. N. Satterfield, J. Catal., 110 (1988) 375.

J.W. Hightower, W.N. Delgass, E. Iglesia and A.T. Bell (Eds.) 11th International Congress on Catalysis - 40th Anniversary

Studies in Surface Science and Catalysis, Vol. 101 9 1996 Elsevier Science B.V. All rights reserved.

97

Effect of SpiUover Hydrogen on Amorphous Hydrocracking Catalysts A.M. Stumbo*, P. Grange and B. Delmon Universit6 Catholique de Louvain - Unit6 de Catalyse et Chimie des Mat6riaux Divis6s, Place Croix du Sud, 2/17 - 1348 Louvain-la-Neuve, Belgium

Abstract The cracking of diphenylmethane on mixtures of sulfided CoMo/SiO 2 and amorphous silica-alumina particles was studied. The products were benzene and toluene. The addition of CoMo/SiO 2 to silica-alumina strongly increases the cracking rate and the OH-OD exchange, and diminishes the amount of coke formed. This is interpreted by a spillover of dissociated H 2 (or D 2) onto the silica-alumina, with spillover hydrogen forming Br6nsted sites and reacting with coke precursors, and spillover deuterium exchanging with the hydroxyls.

1. I N T R O D U C T I O N This work is a contribution to the understanding of the effect of spillover hydrogen in a type of catalyst of considerable industrial importance, namely that composed of transition metal sulfides and amorphous acidic solids. This is typically the case of sulfided CoMo supported on silica-alumina used for mild hydrocracking. Hydrocracking catalysts possess two functions, they consist of a hydrogenationdehydrogenation component (noble metals or transition metal sulfides) and an acidic support (silica-aluminas or zeolites). In the mechanism generally accepted to explain the hydrocracking reaction (the so-called "ideal hydrocracking"), an alkane is initially dehydrogenated to an olefin on the metal phase and then adsorbed on an acidic site, where it is converted to an alkylcarbenium ion. The latter, often after a rearrangement to a more stable form, is cracked via a 13-scission mechanism, forming a lighter olefin and an ion that are hydrogenated to the corresponding paraffins [1 ]. Spillover is a phenomenon that involves the formation of an active species on one phase and the migration of this species onto another phase which is not able to form this species itself. It has been demonstrated that spillover hydrogen (Hsp) can create protonic acidic sites on many kinds of solids, enhancing their activity for a variety of acid catalyzed reactions. This is the case of benzene cracking on silica activated with Pt/A120 3 [2], butane isomerization over Pt/SOn2--ZrO2 [3], cracking of n-hexane [4] and n-heptane [5] on erionite and toluene disproportionation on Fe-HY [6]. Nevertheless, many studies concerning hydrocracking catalysts do not allude to a possible influence of spillover, because they are not * Acknowledges the financial support from the CNPq (Conselho Nacional de Desenvolvimento Cientffico e Tecnol6gico - Brazil) and the Federal Science Policy Office of Belgium (IPA Program).

98 conceived in a way that allows the detection nor the investigation of that phenomenon, since they employ model molecules that need the simultaneous presence of both functions mentioned above to react. In order to investigate the effect of Hsp on amorphous silica-aluminas, we selected diphenylmethane (DPM) as model molecule. The cracking by the "ideal hydrocracking" mechanism is not possible, because DPM does not offer the possibility of a 13-scission. Since the mechanism by which this molecule reacts does not seem to clearly fit into a precise category, we shall refer to it throughout this work using the general term "cracking". Hattori et al. [7], studying DPM cracking over a large variety of catalysts, found three possible routes for this reaction: (i) via carbocations intermediates, formed on acidic sites, producing mainly benzene and toluene and also diphenylethane as a polymerization by-product; (ii)through hydrogenation of both aromatic rings, followed by the cleavage of the C-C bond, producing cyclohexane and methylcyclohexane, the latter undergoing further cracking to produce methane; (iii) via simple cleavage of C-C bond, forming benzene and toluene. The products of the reaction on silica-aluminas, Mo/SiO 2 and Ni-Mo/A1203 catalysts matched those of route (i). Route (ii) is typical of reduced Fe and Ni supported on silica, alumina, titania or zirconia and route (iii) of the same catalysts in the sulfided form. These observations were confirmed by Shimada et al. [8,9]. When catalysts like those used in this work are employed, the previous hydrogenation of the aromatic ring is not necessary to the cracking of DPM and carbocation formation on acidic sites is the main reaction pathway. This reaction does not correspond to a bifunctional mechanism. DPM is thus a suitable tool to study specifically a possible influence of spillover hydrogen on the acidity of silica-aluminas. Our method was to physically isolate the two functions of the catalyst, i.e., the acidic function and the functions necessary for hydrogen activation (hydrogenation-dehydrogenation and, possibly, production of spillover hydrogen). For that, we used mechanical mixtures of amorphous silica-aluminas of different compositions with a CoMo/SiO 2 catalyst. Silica, a support with only a very weak acidity, has been chosen to avoid that the acidity of the support of the donor contribute significantly to total acidity and thus cause interferences with the reactions of the acidic phase. XPS measurements were used to check whether the phases remained unchanged after mixing and sulfiding. The effect of Hsp will be revealed by the comparison between the activities of the pure phases and those measured for mechanical mixtures of both components in different proportions. Since spillover phenomena have been most directly sensed through the use of IR in OHOD exchange [10] (in addition, in the case of reactions of solids, to phase modification), we used this technique to correlate with the catalytic results. One of the expected results of the action of Hsp is the enhancement of the number of Br6nsted sites. FTIR analysis of adsorbed pyridine was then used to determine the relative amounts of the various kinds of acidic sites present. Isotopic exchange (OH-OD) experiments, followed by FTIR measurements, were used to obtain direct evidence of the spillover phenomena. This technique has already been successfully used for this purpose in other systems like Pt mixed or supported on silica, alumina or zeolites [10]. Conner et al. [11] and Roland et al. [12], employed FFIR to follow the deuterium spillover in systems where the source and the acceptor of Hsp were physically distinct phases, separated by a distance of several millimeters. In both cases, a gradient of deuterium concentration as a function of the distance to the source was observed and the zone where deuterium was detected extended with time. If spillover phenomena had not been involved, a gradientless exchange should have been observed.

99 2. EXPERIMENTAL 2.1. Preparation of the mechanical mixtures Three commercial silica-aluminas were used as acidic phases: we call them SA6 (6.5 wt.% A1203 , 500 mE.g "l, average pore diameter: 80 .~), SA12 (12 wt.% A1203, 500 mE.g-1, 80 ~,) and SA60 (60wt.% A1203, 500mE.g -1, 66 .~,). The precursor of the Hsp generator, a CoMo/SiO 2 catalyst (14 wt.% MoO 3 and 3% CoO, 220m2.g ~ 115 .~), was prepared by successive impregnation. Silica (Kali-Chimie AF-125, 270m2.g -1, 115 .~) was first impregnated with an aqueous solution of cobalt acetate (Merck, ultra pure) and subsequently with an aqueous solution of ammonium heptamolybdate (Merck, ultra pure). After each impregnation step the sample was dried overnight at 393 K and calcined at 673 K for 2 hours, under a stream of air (Air Liquide, S). Mechanical mixtures were prepared according to the following procedure, in order to insure intimate mixing and good mutual contact. The pure phases were grounded and sieved to obtain particles of sizes under 40 ~tm. The powders, mixed in the desired proportions, were suspended in n-pentane (15 ml/g solid), placed in an ultrasonic bath for 5 rain and then vigorously mechanically stirred (Ultra Turrax T-50, 3000 rpm) for 10 rain. The n-pentane was evaporated at room temperature, under a flux of Ar and continuous magnetic stimng. After drying at 393 K overnight, the powder was pressed (10 ton.cm-2), grounded and sieved to obtain particles between 0.315 and 0.5 mm. The pure phases were submitted to the same treatment. The samples will be identified by their relative weight content of silica-alumina, named R m, defined as:

Rm _ -

wt. % SiO 2 - A1203 xl00 wt. % CoMo / SiO 2 + wt. % SiO 2 - A1203

(1)

All samples were sulfided in situ prior to catalytic tests and characterizations. A flow (100 ml.min -l) of argon (Air Liquide, N46) was first established, the temperature raised to 423 K, at 10 K.min -1, and maintained at this value for 30 min. The gas was then changed to a mixture of 15% (vol.) HES (Air Liquide, N28) in H 2 (Air Liquide, N30), at the same flow rate. The temperature was raised to 673 K, at l0 K.min -1, and kept at that level for 2 hours. 2.2. X-Ray Photoelectron Spectroscopy (XPS) measurements A Surface Science Instruments SSX-100 spectrometer (model 206), equipped with an aluminum anode whose radiation was monochromatized (AIKct, 1486.6 eV) and focalized, was used. The positive charge developed at the surface of the samples was compensated with a charge neutralizer adjusted at an energy of 8 eV. The sulfided samples were pressed and transferred to the spectrometer protected by a meniscus of iso-octane, to avoid oxidation by atmospheric oxygen. The lines corresponding to C ls, OEs, SiEs, SEp, A12p, MO3d, SEs and Co2p3/2 were analyzed. Their binding energies were determined taking the position of the C is line, corresponding to carbon in a C-C or C-H environment, as 284.8 eV. The intensities were estimated by calculating the area of each peak. Apparent atomic concentrations, that take into account the amount and the dispersion of the elements on the surface, were calculated using the sensibility factors determined experimentally by Weng et al. [13] for our spectrometer. The results were compared to the

100 theoretical values calculated considering that a simple "dilution" of the pure phase containing the metals takes place, according to the following expression: oorotic ,

=

(2)

2.3. Catalytic activity tests The cracking of diphenylmethane (DPM) was carried out in a continuous-flow tubular reactor. The liquid feed contained 29.5 wt.% of DPM (Fluka, >99%), 70% of n-dodecane (Aldrich, >99%; solvent) and 0.5% of benzothiophene (Aldrich, 95%; source of H2S, to keep the catalyst sulfided during the reaction). The temperature was 673 K and the total pressure 50 bar. The liquid feed flow rate was 16.5 ml.h -1 and the H 2 flow rate 24 1.h-1 (STP). The catalytic bed consisted of 1.0 g of catalyst diluted with enough carborundum (Prolabo, 0.34 mm) to reach a final volume of 4 cm 3. The effluent of the reactor was condensed at high pressure. Liquid samples were taken at regular intervals and analyzed by gas chromatography, using an Intersmat IGC 120 FL, equipped with a flame ionization detector and a capillary column (Alltech CP-Sil-8CB). The results of the catalytic tests were expressed as DPM total conversion. These experimental results were compared to theoretical values (Ct), calculated considering, as an approximation, a zero-order reaction and the absence of interactions of any kind between both phases, according to the following expression:

ct - Rm 100 XCRm=l + E1- ~Rm ] x CRm=0

(3)

where CRm=I and CRm=0 are, respectively, the experimental conversions corresponding to pure silica-alumina and pure CoMo/SiO 2. After 24 h of reaction, the catalytic bed was retrieved and sieved to separate the catalyst from the diluent. The used catalyst particles were placed in a Soxhlet apparatus, washed with n-hexane for 8 hours and then dried overnight at 393 K. Their carbon content was determined by automatic titration of the CO 2 formed by burning the washed sample, in a Str6hlein Coulomat 702 apparatus. 2.4. FTIR of adsorbed pyridine The samples were ground and pressed (2 ton.cm -2, for 15 s) in the form of 13 mm diameter wafers, weighing between 3 to 5 mg. They were placed in a specially designed cell, that allowed the heating of the sample under vacuum or controlled atmosphere. IR spectra could be taken through N aCI windows. The samples were submitted to the sulfidation procedure described above, followed by 2 h of heating at 673 K, under vacuum (about 2x10 -3 Pa). After cooling under vacuum, pyridine was adsorbed at room temperature for 30 minutes. The samples were then outgassed in three steps of 1 h: the first one at room temperature and the others at 423 K and 523 K. Spectra were taken before pyridine adsorption and after each outgassing step, with a FTIR spectrometer Bruker IFS-88 (spectral resolution set at 1 cm -1). Each spectrum represented the average of at least 50 scans.

101 The amount of Brrnsted sites was evaluated by measuring the surface of the characteristic band at 1540 cm -1. Corrections have been made to take into account the differences in weight and surface of the wafers. After each test, the wafer was weighted and its cross section was measured with a planimeter. The results were corrected to represent those of a "standard wafer" (A c) of 5 mg and 25 units of area, according to the following expression: 25 5 A c = A e x ....... x Sw mw

(4)

where A e is the experimental integral absorbance of the band considered, S w is the cross section of the wafer (in arbitrary units) and m w is the weight of the wafer (mg).

2.5. Isotopic exchange (H-D) experiments Wafers were prepared, sulfided and evacuated (2 h at 673 K) as described above. The temperature was then set at 423 K and 80 kPa of purified deuterium (Air Liquide, N28) was admitted into the cell. The purification procedure consisted of passing the gas through a moisture filter (Chrompack Gas Clean 7971), an oxygen filter (Chrompack Gas Clean 7970) and a liquid nitrogen trap. Several spectra were taken at regular intervals, using the analysis conditions mentioned above. Before each measurement, the samples were cooled to room temperature. Preliminary experiments had shown that no exchange took place at that temperature. The amount of deuterium exchanged was measured by the total area of the OD bands situated between 2800 and 2100 cm-1. These results were also corrected according to Equation 4.

3. RESULTS 3.1. XPS measurements No new peaks were observed in the mechanical mixtures. The binding energies of all elements were the same in the pure phases and in the mixtures [14]. Figure 1 shows the apparent atomic percentages of molybdenum and cobalt, as given by XPS, on the surface of the sulfided pure phases and mechanical mixtures. In both cases, the experimental results are close to the theoretical values calculated according to Equation 2. ,,

O

,i,

.....

,

O

0.5

"i.... I.. i

0

50 Rm

" "iii

lOO

0-..l...

0

I

0

50

""m

100

Rm

Figure 1. Molybdenum (left) and cobalt (fight) apparent contents (atomic %), determined by XPS, in SA6 (m), SA12 ( . ) and SA60 (o) series, compared to the theoretical values calculated by Equation 2 (dashed lines).

102

3.2. Catalytic activity tests The main products of diphenylmethane (DPM) cracking were benzene and toluene. Very small amounts of polymerized by-products have been found (< 0.5%), but no cyclohexane or partially hydrogenated compounds like cyclohexylphenylmethane were detected. Figure 2 shows the conversions obtained with the three series studied, as a function of the mechanical mixtures composition, one hour after the beginning of the reaction and at the steady-state. Each series presents a maximum of activity, but at a different composition. SA6 series has a maximum between R m values of 50 and 75, whereas SA12 series has a maximum around R m = 50, and SA60 series near R m = 75. The dashed lines on the figures represent the sum of the individual contributions of the pure phases, calculated according to Equation 3. A very important synergetic effect is observed in all series, i.e., the activity of the mixtures is considerably higher than the calculated values (increase by 200% to 750%). 100 Pure silica-aluminas are strongly deactivated, losing about 80% of their activity before reaching the steady-state. The loss in 50 pure CoMo/SiO 2 catalyst is much less pronounced (about 15%). Mechanical mixtures represent an intermediate case; they 0 lose between 35% and 50% of their activity. Table 1 shows the experimental carbon 100 contents of the used samples, compared to the theoretical values obtained by adding the O "~ 50 k.. contributions of the individual pure phases, ~D taking into account their proportions in the O mixtures. All mechanical mixtures present r,.) 0 experimental carbon contents considerably lower than these calculated values. 100

3.3. FTIR of adsorbed pyridine 50

0

50

100

Rm

Figure 2. DPM conversion as a function of the mechanical mixtures composition, after 1 h (..) of reaction and at the steadystate (e), compared to the theoretical values calculated by Equation 3 (dashed lines).

Figure 3 shows the amount of Brrnsted sites, as measured by the surface of the characteristic IR peak at 1540cm -1 after outgassing at 523 K, as a function of the composition of the mechanical mixtures. The dashed lines represent the addition of the contribution of the pure phases, calculated as in Equation3. An enhancement of the amount of Br/Snsted sites on the mixtures, when compared to the theoretical values, is observed. This effect is not very clear in SA6 series, but it is more evident in SA12 and SA60 series. The reproducibility of the experiments has been checked; the variation between different wafers of the same sample was always inferior to 10%.

103

Table 1 Carbon content (wt.%) after catalytic test SA12 Series SA6 Series Theor. Experim. Rm Experim. Theor. 0.45 0 0.45 25 11.3 7.3 50 75 16.4 26.3 22.1 100 34.9 -

70

0.6

SA60 Series Experim. Theor. 0.45 2.3 6.7 4.5 8.4

12.9 19.0

25.2

-

3.4. I s o t o p i c e x c h a n g e e x p e r i m e n t s

All the samples showed two kinds of OD bands: a relatively sharp peak 35 0.3 around 2760 cm ~ and one broad band at lower frequencies, whose limits were always within approximately 2700 and 00 2100 cm-1. The former is assigned to isolated OD species and the latter to OD 60 0.8:5 species in interaction [ 15]. "=" 1SA12 II ] The pure silica-aluminas exchanged O the lowest amounts of deuterium among = 30 0.4 all the samples. Their OD bands were not very intense and increased very :0 9 O0 slowly. The pure CoMo/SiO 2 catalyst, where a potential source of spillover 130 0.6 hydrogen is present, exchanged a higher amount of deuterium than the pure silica-aluminas. The three series of 65 0.3 mechanical mixtures had bands located exactly at the same positions as those of the pure phases. However, their 0-0 intensities were considerably higher. 0 50 100 After 20 hours of exchange, pure silicaRm aluminas and CoMo/SiO 2 exchanged virtually no more deuterium. In the case of the mechanical mixtures, the surface Figure 3. OD bands surface (m) and Br6nsted of the OD bands was still growing after sites band surface (o) as a function of sample that period, but at a much lower rate. In composition, compared to the theoretical values spite of this slow growth, the values calculated as in Equation 3 (dashed lines). obtained at that point were chosen to compare quantitatively the amounts of deuterium exchanged. The reproducibility of the experiments has been checked; the variation between different wafers of the same sample was always inferior to 5%. Figure 3 shows the total surface of the OD bands, after 20 h of exchange, plotted against the composition of the mechanical mixtures. The dashed lines represent the sum of the individual contributions of

104 the isolated phases, calculated as in Equation 3. Each series presents a maximum of exchange, located approximately at the same compositions as those of DPM cracking activity. An important synergetic effect is observed. In all cases, the experimental results are considerably higher (increase by 200% to 1000%) than the theoretical values.

4. DISCUSSION XPS results show that no new peaks are formed or significant energy shifts occur, which indicates that the identity of the two kinds of particles contained in the catalyst are preserved. A significant sintering or redispersion of Co or Mo during the preparation and/or sulfidation of the mechanical mixtures can also be excluded, since the surface compositions correspond to the amounts expected due to the simple "dilution" of the CoMo/SiO 2 phase with the pure silica-alumina. These mechanical mixtures are therefore suitable to study the possible effect of spillover hydrogen. Any change observed in the catalytic behavior or physico-chemical properties of the surface when comparing the mixtures with the individual components must therefore be attributed to some interaction between the unmodified phases. The products obtained from DPM cracking in the present work agree with the results from the literature, mentioned in the Introduction, which indicate that the reaction proceeds via carbocation formation on acidic sites. This implies that the decomposition of DPM does not need the successive intervention of two catalytic sites, like in the "ideal hydrocracking" mechanism. Only acidic sites are sufficient to carry out the reaction. The improved activity of the mixtures when compared to the pure phases must therefore be explained differently. The ability of the unsupported or supported transition metal sulfides to adsorb and dissociate molecular hydrogen from the gas phase is well known [16-19]. It is therefore natural to suppose that the mobile hydrogen species thus formed (Ho, formed by the homolytic scission of H 2, or H + and H-, formed by heterolytic scission) could migrate (spillover) from the sulfides to the surface of the silica-alumina, where they could create new active sites and also contribute directly to the cracking of DPM. The first step of that reaction would be similar to that of a dealkylation [20]. Initially, DPM is adsorbed on a BrtJnsted acidic site and protonated on an aromatic carbon to give the corresponding carbonium ion. The latter can crack to give benzene in the gas phase plus a benzyl carbonium ion adsorbed on the surface. This cation can then react with two hydrogen atoms or with a hydride and desorb, giving toluene in the gas phase (R § + H- ---) RH or R § + 2H ---) RH + H§ A proton regenerates the Brtinsted site. By-products with higher molecular weight can be produced by transalkylation reactions [20]. This mechanism could explain the synergy observed between the two phases. This interpretation of the experimental data is supported by the differences observed in the deactivation patterns and carbon contents after test, since one notorious effect of Hsp is the capacity to diminish the deactivation caused by coke deposition on the active sites [21,22]. This is supposed to be due to a reaction with the coke precursors, very likely a hydrogenolysis. In pure silica-aluminas, where no source of spillover is present, no special protection against deactivation should be observed. Indeed, the silica-aluminas lose most of their activity (about 80%) before reaching the steady-state and present the highest carbon contents after catalytic test. On the other hand, in the case of the mechanical mixtures, where spillover hydrogen is continuously produced by the CoMo/SiO 2 phase and can migrate to the silica-alumina surface, the predicted protection effect is noticed. The relative losses of activity are much lower

105

(between 35% and 50%). Carbon contents are also lower than the simple addition in adequate proportions of the individual contributions of the components of the mixtures. In pure CoMo/SiO 2, the source of Hsp is in the immediate vicinity of the few weak active sites that SiO 2 contains. The distance that activated hydrogen should migrate to reach the coke precursors is very small; in consequence, the protection offered by this mobile species can be most effective. We observe accordingly that this sample presents the lowest relative loss of activity and the lowest carbon content. The results obtained by FTIR measurements of adsorbed pyridine also support the explanation proposed to the observed synergy effect. They indicate an enhancement of the amount of Brtinsted acidic sites of the mechanical mixtures, when compared to the sum of the individual contributions of the pure phases. In the SA6 series the effect cannot be shown unequivocally, but it was clearly detected in SA12 and SA60 series. The action of spillover hydrogen is the most likely explanation to the creation of these acidic sites. In this context, it is striking that the compositions where the amount of these sites was maximum correspond to those where the catalytic activity is the highest (R m = 50 in SA12 and R m = 75 in SA60). The final proof is brought by the isotopic exchange experiments. The interpretation of deuterium exchange results needs some care, since species like HDO or D20 may be responsible for a non-spillover exchange mechanism [23]. The authors proposing this mechanism also indicate that these species could be formed by the reaction between oxygen traces and deuterium, in the presence of metals. One possibility, in the present work, is therefore that HDO and/or D20, that could be present in the deuterium as impurities, would be responsible for the observed exchange. This is not likely, because the gas had been extremely carefully purified. Besides, no IR bands corresponding to these hypothetical oxygenated species (1445 and 1218 cm -1 for HDO and D20, respectively [15]) have been detected on the pure phases or mechanical mixtures. The explanation based on the role of HDO or D20 should be rejected on the ground of experimental evidences. If we accepted this reasoning, we should have observed a significant exchange in the pure silica-aluminas, where there was no Hsp source. This was not the case. Practically no deuterium was detected on their surfaces. We must therefore conclude that there is no or only a negligible direct exchange through the gas phase. Consequently, we must conclude that our results represent a direct evidence for the existence of spillover phenomena in our mixtures. D 2 is adsorbed and dissociated on the metal sulfides. The species created in this way can then migrate (spillover) to the surface of the silica (primary spillover) and of the silica-alumina (secondary spillover), making the exchange with the hydroxyls of the latter possible. The shape of the curves in Figure 3, representing the amount of deuterium exchanged after 20 h, is similar to those of the catalytic activity at the steady-state in Figure 2. This strongly suggests that the magnitude of the synergy in catalytic activity is related to the generation of spillover hydrogen. The maxima found in both cases (catalytic tests and isotopic exchange) correspond to a situation where there is an ideal balance between the activity of the donor (supported metal sulfides) to produce Hsp and the capability of the acceptor (silica-alumina) to exchange with Hsp. In the range of R m values below that optimum composition, the growth of the amount of deuterium exchanged (or of catalytic activity) reflects the increase in the number of active sites that can be generated, due to the increasing silica-alumina content. When this content is superior to that optimum value, the donor is not able any more to produce enough Hsp to fully activate the acidic phase. This explains the decrease in the number of deuteroxyls (and in catalytic activity) observed beyond that ideal composition.

106 5. CONCLUSIONS A very important synergy effect in the cracking of diphenylmethane is observed when particles of a sulfided CoMo/SiO 2 catalyst are mixed with particles of an amorphous silicaalumina. This can be attributed to the action of spillover hydrogen, which can create new active sites and partially protect them from deactivation by coke deposition. Isotopic exchange experiments, showing that the presence of CoMo/SiO 2 brings about the exchange of deuterium with the hydroxyls of the silica-alumina, constitute a direct evidence of this phenomenon.

REFERENCES

1. J.A. Martens, P.A. Jacobs and J. Weitkamp, Appl. Catal., 20 (1986) 283. 2. M. Lacroix, G.M. Pajonk and S.J. Teichner, Bull. Soc. Chim. Fr., 7-8 (1981) 265. 3. H. Hattori, in "New Aspects of Spillover Effect in Catalysis" (T.Inui et al., eds.), p. 69, Elsevier, Amsterdam, 1993. 4. F. Roessner, U. Roland and T. Braunschweig, J. Chem. Soc. Far. Trans., 91 (1995) 1539. 5. I. Nakamura, K. Aimoto and K. Fujimoto, AIChE Symp. Ser., 85 (1989) 15. 6. I. Nakamura, R. Iwamoto and A. I-ino, in "New Aspects of Spillover Effect in Catalysis", (T.Inui et al., eds.), p. 77, Elsevier, Amsterdam, 1993. 7. H. Hattori, K. Yamashita, T. Tanabe and K. Tanabe, Proc. 9th Int. Congr. Catal., p. 27. 8. H. Shimada, M. Kurita, T. Sato, Y. Yoshimura, T. Hirata, T. Konakahara, K. Sato and A. Nishijima, Chem. Lett., (1984) 1861. 9. H. Shimada, T. Sato, Y. Yoshimura, J.Hiraishi and A.Nishijima, J. Catal., 110 (1988) 275. 10. W.C. Conner, Jr., G.M. Pajonk and S.J. Teichner, Adv. Catal., 34 (1986) 1. 11. W.C. Conner, J.F. Cevallos-Candau, N. Shah and V. Haensel, in "Spillover of Adsorbed Species" (G.M. Pajonk et al., eds.), p. 31, Elsevier, Amsterdam, 1983. 12. U. Roland, R. Salzer and S. Stelle, in "Zeolites and Related Microporous Materials: State of the Art 1994" (J. Weitkamp et al. ,eds.), p. 1231, Elsevier, Amsterdam, 1994. 13. L.T. Weng, G. Vereecke, M.J. Genet, P. Bertrand and W.E.E. Stone, Surf. Interf. Anal., 20 (1993) 179. 14. A.M. Stumbo, P. Grange and B. Delmon, to be published. 15. E. Baumgarten and E. Denecke, J. Catal., 95 (1985) 296. 16. M. Karroua, H. Matralis, P. Grange and B. Delmon, J. Catal., 139 (1993) 371. 17. S. Giraldo, P. Grange and B.Delmon, in "New Aspects of Spillover Effect in Catalysis" (T. Inui et al., eds.), p. 345, Elsevier, Amsterdam, 1993. 18. X. Chu and L.D. Schmidt, J. Catal., 144 (1993) 77. 19. N.M. Rodriguez and R.T.K. Baker, J. Catal. 140 (1993) 287. 20. P.A. Jacobs, in "Carboniogenic Activity of Zeolites", Elsevier, Amsterdam, 1977, p. 113-120. 21. G.M. Pajonk, in "2nd Conference on Spillover" (K.-H. Steinberg, ed.), p. 1, Leipzig, 1989. 22.J. Kapicka, N.I. Jaeger and G. Schulz-Ekloff, Appl. Catal., 84 (1992) 47. 23. D. Bianchi, D. Maret, G.M. Pajonk and S.J. Teichner, in "Spillover of Adsorbed Species" (G.M. Pajonk et al., eds.), p. 45, Elsevier, Amsterdam, 1983.

J.W. Hightower, W.N. Delgass, E. Iglesia and A.T. Bell (Eds.) l l t h I n t e r n a t i o n a l C o n g r e s s on Catalysis - 40th Anniversar 3,

Studies in Surface Science and Catalysis, Vol. 101 9 1996 Elsevier Science B.V. All rights reserved.

107

H y d r o d e s u l f u r i z a t i o n of B e n z o t h i o p h e n e Catalyzed by M o l y b d e n u m Sulfide Cluster E n c a p s u l a t e d into Zeolites M. Taniguchi, a S. Yasuda, ' Y. Ishii, b T. Murata, b M. Hidai b and T. Tatsumi ~ aEngineering Research Institute, Faculty of Engineering, The University of Tokyo, Yayoi, Tokyo 113, Japan bDepartment of Chemistry and Biotechnology, Faculty of Engineering, The University of Tokyo, Hongo, Tokyo 113, Japan

A cationic molybdenum sulfide cluster [Mo354(H20)9] 4§ with incomplete cubane-type structure and a cationic nickel-molybdenum mixed sulfide cluster [Mo3NiS4CI(H20)9]3§ with complete cubane-type structure were introduced into zeolites NaY, HUSY and KL by ion exchange. Stoichiometry of the ion exchange was well established by elemental analyses. The UV-visible spectra and EXAFS analysis data exhibited that the structure of the molybdenum cluster remained virtually intact after ion exchange. MoNi/NaY catalyst prepared using the molybdenum-nickel sulfide cluster was found to be active and selective for benzothiophene hydrodesulfurization.

1. INTRODUCTION The great practical significance of the processes of hydrodesulfurization (HDS) of petroleum feedstocks has developed a keen research interest in the structure and synthesis method of the active component of HDS catalysts. All HDS processes have been performed on catalysts of the same type, based on molybdenum (or tungsten) sulfide supported on high surface area 7-aluminas with cobalt or nickel added as a promoter. In recent years much attention has been focused on the deep HDS of light oil, since the reduction in sulfur concentration in fight oil is effective in reducing particulate and NOx in the exhaust gas of Diesel engines. The main problem in the deep HDS of fight oil is the conversion of alkyldibenzothiophenes, e.g., 4,6-dimethyldibenzothiophene, which is one of the most difficult compounds to desulfurize among dibenzothiophene derivatives [1]. One approach to this problem is to rearrange and/or to eliminate alkyl groups by using catalysts with enhanced acidity. Zeolite-supported molybdenum with nickel or cobalt as a promoter is promising in this regard. Molybdenum/zeolite catalysts prepared by impregnating zeolites with ammonium heptamolybdate solution generally give rise to poor dispersion of molybdenum [2]. In contrast, ion exchange would be an ideal method for loading active metal species onto supports. Few cationic forms are available as simple salts of molybdenum of high oxidation

108

state, however. Furthermore, most of them exist only in strongly acidic solutions where many zeolites are unstable and where exchangeable cations must compete with protons. Therefore few studies have succeeded in introducing cationic molybdenum compounds into zeolites by ion exchange [3-5]. Molybdenum/Y-zeolites were prepared by aqueous ion exchange with Mo2(ethylenediamine)4Ch [3] and MoO2CI: [4]. Lunsford et al. also reported solid-state exchange of HY and ultrastable form of HY (HUSY) with MoC15 [5]. A molybdenum sulfide cluster, [Mo3S4(H20)9] 4§ (1) with incomplete cubane-type structure was recently prepared [6-8]. It is stable in water and air and easily forms mixed metal clusters with cubane-type Mo3MS4 cores (M = Fe, Co, Ni, Pd, etc.) [9]. Mixed metal clusters with a Mo3PdS4 core have been derived from 1, showing intriguing reactivities at the Pd site toward alkenes, CO, isonitriles, and alkynes [10,11]. A molybdenum-nickel cluster with a Mo3NiS4 core has also been found to uptake CO stoichiometrically to give a new cluster where one CO molecule combines with the nickel atom [12]. Thus these clusters are of considerable interest in connection with potential application to a variety of catalytic reactions. It occurred to us that the molybdenum and mixed metal clusters could be incorporated into zeolites by ion exchange and promote unique catalytic reactions. Preparation of molybdenum/zeolites using the molybdenum sulfide cluster cation as a precursor by ion exchange is considered to have several advantages. First, molybdenum species could be loaded with high dispersion. Secondly, molybdenum is loaded on zeolites as sulfide, not as oxide; presulfiding is unnecessary before HDS reactions. To our best knowledge, there is no such report that succeeded in introducing molybdenum species into zeolites in its sulfide state and applying the resulting catalyst to HDS reactions. Thirdly, by loading molybdenum-nickel cluster with a Mo3NiS, core into zeolites, molybdenum and nickel could be introduced in the homogeneously well-mixed state. This paper describes the successful incorporation of molybdenum and molybdenum-nickel clusters into zeolites with 12-membered ring by aqueous ion exchange and application of the resulting materials to HDS reaction of benzothiophene. Stoichiometry of the ion exchange was examined by elemental analysis. UV-visible spectroscopy and EXAFS measurements were carried out to investigate the structure of molybdenum species loaded on zeolites.

2. EXPERIMENTAL

2.1. Catalyst preparation The chloride salt of cluster 1 was synthesized by the reported method [8]. As supports, NaY (Nikka Seiko, SK-40; Si/AI = 2.3), HUSY (Tosoh, HSZ-330HUA; Si/A1 = 3.1) and KL (Tosoh, TSZ-500KOA; Si/AI = 3.1) zeolites with 12-membered ring were used. To a suspension of zeolite (4.58 g) in water (91.6 g) was added dropwise a 0.01 M aqueous solution of the chloride salt of 1 (86.9 ml) with vigorous stirring at 313 K. In the case of NaY the color of zeolite changed from white to brown and the solution was colorless. For HUSY and KL, brown-colored species remained in the solution after ion exchange. The metal cluster-containing zeolite was separated by filtration and washed with distilled water, followed by drying at 313 K under reduced pressure. Resulting materials are referred to as Mo/NaY, Mo/HUSY and Mo/KL, respectively. The chloride salt of cluster [Mo3NiS4CI(H20)9]3+ (2) was also synthesized according to the reported method [9] and

109 similarly treated with NaY to afford MoNi/NaY. MoNi/A1203 was also prepared by impregnating A1203 (JRC-ALO-4) with the solution of chloride salt of 2. 2.2. Characterization Molybdenum and sodium concentrations in the filtrate were determined by ICP on a Nippon Jarrell-Ash ICAP-575 spectrophotometer in order to estimate molybdenum loadings after ion exchange and to confirm the stoichiometry of ion exchange. Sulfur contents were determined by the LECO method. Chlorine contents of the catalysts were determined by ion chromatography. Powder X-ray diffraction patterns of the catalysts were collected on a Rigaku RINT 2400 X-ray diffractometer. Crystallinity of the zeolites were determined by the average of the relative intensities of selected seven intense peaks, using starting material NaY, HUSY and KL as the reference. UV-visible spectra were recorded on a Hitachi U-4000 spectrophotometer. Mo K-edge X-ray absorption spectra were collected on a R.igaku R-EXAFS 2100S (30 kV, 280 mA) instrument using a Ge(400) crystal monochromator at room temperature in air atmosphere. Four samples were applied to the EXAFS measurement: (a) 5% physical mixture of cluster 1 and NaY, (b) Mo/NaY, (c) sample (b) treated in flowing helium at 373 K for 0.5 h and (d) sample (b) treated in flowing helium at 573 K for 0.5 h. The sample (a) was set on a sample holder with 0.3 mm thickness, and the samples (b)-(d) were pressed into self-supporting wafers. The method of EXAFS data analysis utilized REX, an EXAFS data analysis software provided by Rigaku. The sample (a) was used as a standard for Mo-S, Mo-O and Mo-Mo interactions. Crystallographic data (N and R of Mo-S, Mo-O and Mo-Mo, where N is the coordination number and R the radial distance from the absorber to the backscatter atom) of the cluster 1 [13] were used in order to determine the EXAFS parameters o and AE0 (where o is the Debye-Waller factor and AEo the inner potential correction of the edge position) which were used for the samples (b)-(d). Once a close fit was obtained for the parameters N and R with k3-weighting, the fit was optimized by allowing a few of the parameters to be fitted while the rest of the parameters were held fixed. 2.3. HDS reaction The catalytic activities of the zeolite- and alumina-supported Mo and Mo-Ni catalysts for benzothiophene HDS were measured in a flow system incorporating a microreactor made of stainless steel tube containing 0.2 g of catalyst. The catalysts were pretreated in flowing helium by ramping the temperature from room temperature to 573 K at a rate of 3.3 K/min and then in flowing hydrogen for 0.5 h at 573 K. When presulfided, the catalysts were treated in flowing 5% H2S/H2 mixed gas for 2 h at 573 K following hydrogen pretreatment. Benzothiophene in decane (0.5% or 5% as sulfur) was supplied to the reactor at a rate of 4.5 ml/h. The reaction conditions were 573 K, 3.0 MPa, H2/benzothiophene = 379 or 3790 (mol/mol) and W/F = 38.2 or 382 g-cat.h/mol-benzothiophene.

110 3. RESULTS AND DISCUSSION

3.1. Chemical composition The results of elemental analysis of the molybdenum and molybdneum-nickel/zeolite catalysts are shown in Table 1. After ion exchange of NaY with the solution of 1, almost all molybdenum was found to be transferred from the solution to NaY; the resulting Mo/NaY contained 5.0 wt% of molybdenum, which corresponds to the presence of 0.41 Mo3S4 cores per a supercage of the FAU type structure. The size of cluster 1 was estimated to be 0.72 nm, smaller than the diameter of the aperture (0.74 nm) [14] of supercages of the FAU type structure. The C1/Mo ratio of Mo/NaY was 0.11, suggesting that most of the cluster 1 acted as a tetravalent cation in the ion exchange. Fled sodium ion per loaded Mo cluster in the filtrate was 4.6 times (ideally 4 times) as much as the loaded cluster 1. The slight excess of lost sodium over the supported cluster may be due to the acidity of the cluster salt. The S/Mo ratio was maintained after ion exchange. In the preparation of Mo/HUSY, the cluster 1 amounting to 2.5 wt% (as molybdenum metal) of HUSY was added to the suspension of HUSY; 92% of the molybdenum was loaded onto HUSY. The CI/Mo ratio of Mo/HUSY was found to be 0.34, suggesting that in ion exchange the cluster 1 acted as a trivalent cation on the average. These findings indicate that the protons in HUSY are less exchangeable by the cluster cation than the Na cations in NaY. In the preparation of Mo/KL, the addition of 1 was stopped when pH of the solution was lowered to about 3. The resulting Mo/KL contained only 2.1 wt% (74% of added molybdenum clusters) of molybdenum. Chlorine was absent, which indicates that the cluster 1 acted as a tetravalent cation. The LTL structure is characterized by a monodimensional system of channels, whose diameter (0.70 nm) [14] is close to the size of the cluster 1. It is conceivable that, once the cluster 1 was incorporated into an LTL main channel and present at a site near the external surface, the sites deep in the channel are no more accessible to another cluster. The cluster 2 was also introduced into NaY by ion exchange. After the ion exchange, 99% of molybdenum and nickel were loaded on NaY; the Ni/Mo ratio did not change. The CI/Mo ratio suggests that excess chlorine was present on the catalyst.

Table 1 Elemental analysis of Mo and Mo-Ni / zeolite catalysts Catalyst

Si/A1 ratio

Loaded % loaded % loaded Na CI/Mo Mo (wt%) Mo a Ni" released b ratio

Mo / NaY

2.3

5.0

Mo / HUSY

3.1

2.3

Mo / KL

3.1

2.1

MoNi / NaY

2.3

5.0

99

100

S/Mo ratio

-

4.6

0.11

92

-

-

0.34

n.d.c

74

-

-

0.00

n.d.C

4.2

0.96

n.d.c

99

a Fraction of Mo or Ni loaded based on Mo or Ni added to the solution. b Mol Na fled into the solution per mol cluster loaded. c Not determined.

1.3

111

Table 2 Crystallinity of Mo and Mo-Ni / zeolite catalysts Catalyst

Si/A1 ratio

pH after ion exchange

Crystallinity (%)

Mo / NaY Mo / HUSY Mo / KL

2.3 3.1 3.1

4.6 4.2 3.2

75 85 93

MoNi / NaY

2.3

4.7

72

3.2. Crystallinity of zeolites It is reported that introduction of molybdenum cations into zeolites often gives rise to destruction of the zeolite skeleton because of their acidity [5,15]. The crystaUinity of the ion exchanged zeolites are shown in Table 2. Although the solutions during ion exchange were slightly acidic (pH = 3.2 - 4.7), the X-ray powder diffraction patterns were only slightly attenuated, indicating the zeolites kept their crystallinity without significant destruction. The decrease in crystallinity was not related to the pH after ion exchange but to the Si/AI ratio of the zeolites.

(a)

O

< (

I

200

300

I

1

1

I

400 500 600 700 Wavelength (nm)

1

800

900

Figure 1. UV-visible spectra of (a) physical mixture of NaY and the chloride salt of 1, (b) Mo/NaY, (c) Mo/HUSY and (d) Mo/KL.

112 3.3. Structure of the clusters UV-visible spectra of the cluster 1 and molybdenum/zeolite catalysts are shown in Figure 1. The cluster 1 showed bands at 300, 390 and ca. 650 nm. Similar bands were observed for the spectrum of each molybdenum/zeolite catalyst, suggesting that the structure of cluster 1 was practically unchanged after ion exchange. In order to obtain more structural information about the molybdenum species in Mo/NaY, EXAFS measurements of the cluster 1 and Mo/NaY were carried out. The Fourier transforms of the EXAFS data are shown in Figure 2. Structural parameters (Table 3) showed no change of the Mo-O, Mo-S and Mo-Mo distances, suggesting that there is no significant structural difference between the cluster 1 and the molybdenum compound in the Mo/NaY. From these EXAFS parameters and the UV-visible spectra, it is considered the structure of cluster 1 remained virtually intact after ion exchange.

14

14 (a)

12

(b)

12 .~10

~10

...a

'~

8

6 a~ 4

u2 4

0

0.1

0.2

0.3 0.4 R (nm)

0.5

0

0.6

0.1

0.2

0.3 0.4 R (nm)

0.5

0.6

14

14

(c)

12

(d)

12

~10

.~10

'2

'~ e~

8

8

6

6

a~ 4

4

0

0 0

0.1

0.2

0.3 R (nm)

0.4

0.5

0.6

0

0. I

0.2

0.3

0.4

0.5

0.6

R (nm)

Figure 2. Fourier transforms of EXAFS data of (a) the chloride salt of 1, (b) Mo/NaY, (c) sample (b) treated in flowing He at 373 K and (d) sample (b) treated in flowing He at 573 K.

113

Table 3 Structural parameters from EXAFS of the Mo/NaY catalysts Mo-S Sample Cluster I

Treatment -

N~

Mo-Mo R (nm) b

(3.0) ~ (0.230) ~

Na

Mo-O R (nm) b

(2.0) c (0.274) c

Na

R (nm) b

(3.0) ~ (0.218) ~

3.5

0.232

2.0

0.276

3.9

0.217

Mo / NaY

He, 373 K

3.1

0.232

1.7

0.277

3.3

0.218

Mo / NaY

He, 573 K

1.2

0.238

0.69

0.276

1.3

0.214

Mo / NaY

-

"Coordination number. b Radial distance from the absorber to the backscatter atom. ~ Taken from ref. 13 and treated as fixed parameters.

After treatment at 373 K in helium flow, the coordination numbers of Mo-O, Mo-S and Mo-Mo were not largely different from those in the cluster 1 and the distances of Mo-O, Mo-S and Mo-Mo were hardly changed. After treatment at 573 K in He flow, however, the color of NaY changed from brown to black and the curve fitting results of the EXAFS data exhibited lower coordination numbers of all interactions than those of the cluster 1. The decrease in the coordination number seems to be not due to the decrease in sulfur amount in the catalyst but to the disordering of each interaction since the S/Mo ratio hardly decreased after thermal treatment. These results show that the structure of the cluster 1 loaded on NaY was maintained at 373 K, but lost at 573 K.

3.4. HDS reactions The results of hydrodesulfurization of benzothiophene (0.5 wt% as sulfur) in decane are shown in Table 4. Over Mo/NaY and Mo/HUSY, selectivity for ethylbenzene was low and main products were dihydrobenzothiophene and alkylbenzothiophenes. The latter seems to be produced by the reaction of unreacted benzothiophene with alkenes generated from decane cracking on acid sites of the zeolite. For Mo/I-/USY, decane cracking and following alkylation of benzothiophene took place appreciably. Mo/KL showed high HDS activity compared to Mo/NaY in spite of its lower molybdenum loading. This is consistent with our supposition that molybdenum clusters were loaded on near external surface of KL, as described above. Since the KL zeolite has no strong acidity, decane cracking and following alkylation of benzothiophene proceeded only to a small extent. Involvement of Ni resulted in great enhancement of HDS activity; MoNi/NaY proved to be a highly active and selective catalyst for HDS of benzothiophene to ethylbenzene, and other products were scarcely observed. No change in the activity was observed during 4 hours. Table 5 shows HDS product distributions over several catalysts prepared by using the molybdenum-nickel cluster 2. Sulfur content in decane was adjusted to 5.0 wt% in these experiments. MoNi/NaY was found to be more active than MoNi/Al203. It is to be noted that during the high temperature pretreatment the original cluster structure would have been changed. However, the high activity of the MoNi/NaY catalyst for benzothiophene HDS is probably due to the formation of active sites derived from this particular mixed metal cluster,

114 Table 4 Effect of nickel and acidity of zeolites on the benzothiophene hydrodesulfurization ~ Selectivity (%) Catalyst

Conversion (%)

EB b

DHB"P

Others

Mo/NaY

53

8.3

76

16

Mo / HUSY Mo/KL

39 75

6.1 38

66 52

28 9.4

MoNi / NaY

97

88

1.2

5.7

a Reaction conditions: 573 K, 3.0 MPa, W/F = 382 g-cat.h/mol-benzothiophene, S content 0.5 wt%. b Ethylbenzene. Dihydrobenzothiophene.

Table 5 Catalytic performance of MoNi catalysts for benzothiophene hydrodesulfurization ~ Selectivity (%) Catalyst

Conversion (%)

EB b

DHBT:

Others

MoNi / NaY MoNi / NaY (presulfided)

36 19

51 23

39 55

11 4.7

MoNi / A1203 MoNi / A1203 (presulfided)

23 30

42 39

52 45

5.1 16

a Reaction conditions: 573 K, 3.0 MPa, W/F = 38.2 g-cat.h/mol-benzothiophene, S content 5.0 wt%. b Ethylbenzene. c Dihydrobenzothiophene.

together with high dispersion. The pore structure of the zeolite support may also play a role in the stabilization of active species reminiscent of the cluster structure. Presulfiding effect was quite different between two supports; although MoNi/AI203 was activated by presulfiding, MoNi/NaY was slightly deactivated by presulfiding, suggesting that the active species in MoNi/NaY is different from that in MoNi/A1203. The S/Mo ratio of the cluster 1 is 1.3. It is suspected that the increase in the HDS activity of MoNi/AI203 by sulfiding resulted from the increase in the sulfur content, in agreement with the generally accepted belief that the active component of A1203-supported conventional HDS catalysts is MoS2 [16], where the S/Mo ratio is 2. Study on applicability of MoNi/zeolite catalysts to the HDS reactions of other sulfur compounds is ongoing in our laboratory.

115 4. CONCLUSIONS A cationic molybdenum sulfide cluster 1 with incomplete cubane-type structure and a cationic molybdenum-nickel sulfide cluster 2 with complete cubane-type structure were employed as new precursors for Mo/zeolite and MoNi/zeolite catalysts. Mo/NaY, Mo/HUSY, Mo/KL and MoNi/NaY were prepared by aqueous ion exchange with the cluster 1 and 2. Stoichiometry of the ion exchange was well established by elemental analysis. Though pH of the solution during ion exchange was lowered, crystallinities of the zeolites were preserved. The UV-visible spectra and EXAFS analysis data exhibited that the structure of the cluster 1 remained virtually intact after ion exchange. In the HDS of benzothiophene Mo/HUSY and Mo/KL showed higher activity than Mo/NaY. The use of molybdenum-nickel mixed cluster resulted in a great enhancement of HDS activity. By presulfiding, MoNi/AlzO3 was activated, whereas MoNi/NaY was not. These results suggest that the active species in MoNi/NaY is different from that in MoNi/AI203.

ACKNOWLEDGMENT We thank Professor K. Asakura, Faculty of Science, The University of Tokyo, for his support with the EXAFS analyses.

REFERENCES

1. A. Ishihara and T. Kabe, Ind. Eng. Chem. Res., 32 (1993) 753. 2. e.g., M. Laniecki and W. Zmierczak, Zeolites, 11 (1991) 18. 3. M.B. Ward, K. Mizuno and J.H. Lunsford, J. Mol. Catal., 27 (1984) 1. 4. M. Huang and R. F. Howe, J. Catal., 108 (1987) 283. 5. P.E. Dai and J.H. Lunsford, J. Catal., 64 (1980) 173. 6. F.A. Cotton, Z. Dori, R. Llusar and W. Schwotzer, J. Am. Chem. Soc., 107 (1985) 6734. 7. M. Martinez, B.L. Ooi and A.G. Sykes, J. Am. Chem. Soc., 109 (1987) 4615. 8. T. Shibahara, M. Yamasaki, G. Sakane, K. Minami, T. Yabuki and A. Ichimura, Inorg. Chem., 31 (1992) 640. 9. P.W. Dimmock, G.J. Lamprecht and A.G. Sykes, J. Chem. Soc., Dalton Trans., (1991) 955 and references therein. 10. T. Murata, H. Gao, Y. Mizobe, F. Nakano, S. Motomura, T. Tanase, S. Yano and M. Hidai, J. Am. Chem. Soc., 114 (1992) 8287. 11. T. Murata, Y. Mizobe, H. Gao, Y. Ishii, T. Wakabayashi, F. Nakano, T. Tanase, S. Yano, M. Hidai, I. Echizen, H. Nanikawa and S. Motomura, J. Am. Chem. Soc., 116 (1994) 3389. 12. T. Shibahara, S. Mochida and G. Sakane, Chem. Lett., (1993) 89. 13. H. Akashi, T. Shibahara and H. Kuroya, Polyhedron, 9 (1990) 1671. 14. W.M. Meier and D.H. Olson, Atlas of Zeolite Structure Types (Third edition), Butterworth-Heinemann, Stoneham, 1992.

116 15. e.g., R. Cid, F.J.G. Llambias, J.L.G. Fierro, A.L. Aguro and J. Villasenor, J. Catal., 89 (1984) 478. 16. A.N. Startsev, Catal. Rev.-Sci. Eng., 37 (1995) 353.

J.W. Hightower, W.N. Delgass, E. Iglesia and A.T. Bell (Eds.) 11th International Congress on Catalysis - 40th Anniversary

Studies in Surface Science and Catalysis, Vol. 101 9 1996 Elsevier Science B.V. All rights reserved.

117

Role o f a d s o r b e d h y d r o g e n species on r u t h e n i u m and m o l y b d e n u m sulfides. C h a r a c t e r i z a t i o n b y inelastic n e u t r o n scattering, t h e r m o a n a l y s i s m e t h o d s and m o d e l reactions. M. Lacroix a, H. Jobic a, C. Dumonteil a, P. Afanasiev a, M. Breysse a and S. Kasztelanb. aInstitut de Recherches sur la Catalyse, 2 avenue Albert Einstein, 69626 Villeurbanne cedex France bInstitut Frangais de P6trole, l&4 avenue de Bois Pr6au, BP 311, 92506 Rueil-Malmaison, France.

Abstract

The interaction of hydrogen over unsupported MoS 2 and RuS 2 has been investigated as a function of the sulfur to metal ratio. On these solids the presence of sulfur deficient sites is required to generate an activity and to allow hydrogen chemisorption. The nature of the adsorbed species differs depending on the catalyst under investigation. On RuS2, two types of hydrogen were evidenced by thermoflash desorption and inelastic neutron scattering: one was assigned to hydrogen adsorbed on surface sulfur anions while the other one is retained on coordinatively unsaturated ruthenium cations. By contrast, only SH groups were detected on MoS2. ESR measurements have shown that a fraction of chemisorbed hydrogen induces a modification of the concentration of paramagnetic Mo(V) and Mo(III) species. Thus both solids behave differently towards an hydrogen atmosphere. RuS2 has a pseudometallic comportment whereas for MoS 2 redox or acid base properties are involved.

1. INTRODUCTION Sulfide catalysts find extensive use as hydrotreating catalysts. Hydrotreating is a general term that includes hydrogenation and hydrogenolysis reactions i.e. hydrodesulfurisation, and hydrodenitrogenation. As these reactions require hydrogen, the interaction of hydrogen with the solid is important in order to understand their catalytic properties. Current industrial catalysts are supported systems containing basically a group VI sulfide like MoS 2 or WS 2 promoted by cobalt or nickel sulfides. The properties of these multicomponent solids are complex because of the observed synergy between groups VI and VIII sulfides and of the interaction with the support which could affect the hydrogen reactivity. From a basic point of view it appears sensible to simplify the industrial catalysts by reducing them into their component parts.

118 Polycrystalline molybdenum sulfide has been frequently used as model catalyst. This solid crystallizes in the hexagonal system in which the molybdenum ions are located at the center of a trigonal prismatic unit formed by six surrounding S2" anions. Assembling these moieties leads to a lamellar structure with the basal planes presenting a continuous layer of sulfur. These [hk0] planes were found poorly reactive towards hydrogen while the edges of the crystal may easily undergo partial desulfurisation generating sulfur vacancies which are seen as the basis of the catalytic and adsorbing properties [1, 2]. The interaction of hydrogen with this anisotropic surface has been the subject of many investigations. The review of Moyes and the paper of Komatsu and Hall [3, 4] have shown that the amount of adsorbed hydrogen changes drastically from one study to another. For instance, using the same preparation method (decomposition (NH4)2MoS4) the values of H/Mo vary between 0.012 and 0.37. This important variation is probably due to various catalysts pretreatment prior to adsorption measurements leading to different number of vacancies. This hypothesis was supported by Jalowiecki et al [5, 6] who reported that a fully sulfur saturated slab was inactive and did not chemisorb hydrogen while both the activity and the hydrogen adsorption capacity increased and went through a maximum as far as sulfur was removed from the catalyst upon reduction. This showed that unsaturation is a necessary prerequisite for giving rise to a reactive surface. Moreover, the observed results suggested that both properties required an optimal concentration of vacancies and sulFar species. Therefore, it has been envisaged that hydrogen is heterolytically adsorbed on a Mo-S pair leading to the formation of Mo-H and SH groups [5-7]. Quantum chemical calculations have also suggested that heterolytic adsorption is the favored route for hydrogen adsorption on MoS2 [8]. From this theoretical study it was also found that sorbed hydrogen is able to move from the edge sites to the basal plane and the resulting vacant site could be replenished by further heterolytic adsorption. However, characterization of the adsorbed species by inelastic neutron scattering (INS) and 1H NMR have detected the presence of SH groups while the signal ascribed to Mo-H was never observed even for a hydrogen saturated slab [4, 9, 10]. Recently, the interaction of hydrogen with a model RuS2 catalyst was studied using thermoanalysis methods, IH NMR and INS [ 1113]. The systematic applications of these techniques for the characterization of the adsorbed species present on different reduced states proved the existence of both Ru-H and SH entities whose relative population drastically depended on the sulfur to metal ratio. Coupling these results with catalytic activity measurements, it was shown that the H2-D2 exchange and the 1butene hydrogenation rates were directly related to the concentration of Ru-H species suggesting that a hydride type hydrogen is the active species involved in hydrogenation. Taking into account these experimental data, the following questions could be addressed: i) Is the presence of hydride type hydrogen essential for explaining the activity of a given sulfide catalyst .9, ii) Why was the Mo-H bond never detected in MoS2? In order to get insight into these fundamental questions, a similar methodology as previously utilized for RuS2 was undertaken and results obtained on both systems were systematically compared.

2. EXPERIMENTAL Unsupported ruthenium sulfide was prepared by precipitation at room temperature from an aqueous solution of RuCI3 by pure H2S and by further sulfidation under an H2S flow at

119 673K for 2h. The solid was then cooled to room temperature under the same sulfiding atmosphere, flushed with an oxygen free inert gas and stored in sealed bottles. Physicochemical characterizations confirmed that the obtained solid has the expected pyrite structure and elemental analysis indicated a S/Ru ratio equal to 2.25. Microcrystalline molybdenum sulfide was prepared by thermal decomposition of 0NIHa)2MoS4 (ATTM). The later was synthesized using the method described by Dieman et al [14]. UV-Vis spectra of aqueous solutions of ATTM have shown that this precursor does not contain any absorption bands corresponding to oxothiomolybdates. MoSx was obtained by heating the ATTM crystals at 673K for 4h in flowing mixtures of 15% HzS in H2. The x value determined gravimetrically on a microbalance (Sartorius 4433) by oxidizing the sulfided sample in 02 at 773K to form MoO 3 was found to be close to 2.27. The desulfurization of the catalysts was carried out in a dynamic microreactor at 1 atmosphere hydrogen pressure and at different temperatures. The setup included a specific UV-photodetector equipped with a 10.2 eV light source which allowed the detection of about 1 ppm of H2S. The catalyst was first flushed with nitrogen at room temperature and then contacted with an hydrogen flow (50 cm3/min). The reactor temperature was then linearly increased to the desired value and left at this temperature for 2h. The amount of H2S released from the solid was quantified after calibrating the detector with a known concentration of H2S diluted in argon. The degree of reduction was defined as the ratio between the amount of H2S eliminated during the reduction process and the total sulfur content. The catalytic properties of the reduced samples have been determined at 273K or 323K using the H2-D2 exchange reaction. This reaction which involved hydrogen activation was chosen because it proceeds at temperatures lower than those required for solid reduction. In a typical run, the solid was reduced at a given temperature and then cooled down to the reaction temperature. The reactor was flushed with nitrogen and then isolated to allow the stabilization of the reactant flows. The catalyst was then submitted to an equimolar mixture of H 2 and D 2 diluted in argon. The H E and D E partial pressures were 76 torr. The variation of the H2-D2 composition was analyzed by means of a mass spectrometer (FISONS Instruments) equipped with a quadrupole analyser working in a Faraday mode. A silica capillary tube heated at 453K continuously bled off a small fraction of the gas phase into the spectrometer. Conversions were calculated either with respect to the decrease of the D2 or of the H 2 signal. For both solids the conversion was kept lower than 20% by adjusting the contact time. It has been verified that under these experimental conditions no H2S, HDS nor D2S was released during a catalytic run indicating that sulfur to metal ratios were stable during the reaction course. The amount of hydrogen retained by the solids after a given pretreatment was determined by thermodesorption by heating abruptly the solid from room temperature up to 573K. Preliminary experiments had shown that this temperature is high enough to desorb the hydrogen present on both solids. Species leaving the catalyst surface were detected by a gas chromatograph equipped with a TCD detector. It was checked that on both solids no H2S was removed during these desorption experiments whatever the degree of reduction of the solids and several adsorption-desorption cycles could be done without changing the adsorption properties of the reduced catalysts. Using this technique two different sets of experiments were performed. The amount of hydrogen retained by the solids during the reduction step was first determined, then the desorbed solids were again submitted to a hydrogen flow at room

120 temperature in order to quantify the hydrogen retained in similar experimental conditions than those used to determine their catalytic properties. The neutron spectra were obtained at the ISIS spallation neutron source, at the Rutherford Appleton Laboratory, U.K., using the spectrometer TFXA. It is a time-of-flight instrument with an inverse geometry and a time-focusing analyser; it gives reasonable counting rates and good energy transfer resolution (AE/E ~, 2%) over a wide range of energy transfers. The precision on the frequencies is _+ 10 cm "1. All the spectra were recorded at 20 K, the evacuation and the loading of the samples being performed out of the cryostat. ESR characterization was performed in situ in order to avoid any contact of the pretreated solids with air. Spectra, recorded as the first derivative of the absorption, were obtained at room temperature or 77K using a Varian E9 spectrometer working in the X band. The g values were measured relative to a DPPH reference (g = 2.0036). The sample tubes were filled with the solid to a height greater than the depth of the resonant cavity and the number of paramagnetic species was calculated by double integration of the recorded spectra normalized to that of Varian Strong Pitch sample (g = 2.0028, 3.1015 spins, cm-l).

3. RESULTS Figure 1 reports the evolution of the degree of reduction of both solids as a function of the temperature of reduction. RuS 2 can be completely reduced at 773K while only 8 % of the sulfur content of M o S E is released at this temperature. For RuS E previous physicochemical characterizations have evidenced that the pyrite phase is stable up to 623K [13]. At higher temperatures, the solid sinters with the concomitant formation of a metallic phase. For MoSE, only a slight decrease in surface area (from 55 to 40 mE/g) has been observed above 573K. 1.0

=

.=o

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RuS2

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273 373 473 573 673 773 Reduction t e x t u r e

(K)

Figure 1. Evolution of the degree of reduction Figure 2. HE-D2 activity ofMoS 2 (at 323K) as a fianction of the reduction temperature, and of RuS 2 (at 273K) versus the reduction temperature.

r

121 The changes in the catalytic properties as a function of the temperature of reduction are reported in Figure 2. For both solids, it is clear that as far as the stability of the solids is preserved an increase of the degree of reduction brings about an increase of the H2-D2 conversion. These results confirmed that the presence of coordinatively unsaturated sites is required for hydrogen activation. It should be underlined that in order to determine accurate activity measurements the reaction temperature was not the same for both solids. Taking into account the apparent activation energies (-- 12 kcal/mol) the activity of RuS2 is about two orders of magnitude higher than that of M o S 2. Figures 3 and 5 give examples of hydrogen TPD profiles observed on both catalysts reduced at 473K. The more intense patterns concern the desorption of the hydrogen fixed on the solids during the reduction step (curve a) while the weaker peaks (curve b) are related to the hydrogen retained after readsorption at room temperature. Hydrogen adsorption is thus an activated process. 50

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Figure3. Hydrogen TPD patterns observed on Figure 4. Amount of adsorbed H l and a reduced RuS2 catalyst, catalytic activity (at 273K) versus the degree of reduction of RuS2. On RuS 2 the TPD profiles clearly evidence the existence of at least two different adsorbed species denoted by H x and H 2. Previous work has shown that their relative amount vary with the reduction state of the solid [11-13]. The intensity of the high temperature peak H 2 continuously decreased with desulfurisation while the concentration of H l goes through a maximum (see Figure 4). Interestingly, the variation of both the H 1 concentration and the catalytic activity with the degree of reduction follow a similar trend suggesting that H ~ is the active species involved in this model reaction. By contrast, the smooth profile observed on MoS2 (Figure 5) suggests that only one adsorbed species exists on such a solid. As shown in Figure 6 the exchange activity reasonably correlates with the overall amount of hydrogen retained by the solid during the reduction or after a readsorption at room temperature.

122

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Figure 5. Hydrogen TPD patterns observed on Figure 6. Amount of hydrogen adsorbed and a reduced MoS2 catalyst, catalytic activity (at 323K) versus the degree of reduction of MoS2. The INS spectrum recorded at 25K of a reduced and hydrogen desorbed RuS 2 catalyst showed in Figure 7 (spectrum a) exhibits only a weak signal in the range 650-720 cm 1 which can be assigned to bending modes of residual SH. The cell was then taken out of the cryostat, and hydrogen was slowly introduced at room temperature until a pressure of 0.5 bar was reached in the cell. The sample was left overnight to equilibrate and the cell was placed back into the cryostat. In these conditions the amount of adsorbed hydrogen present in the cell was about 150 cm 3. Spectrum b shows that chemisorption gives rise to four peaks found at 540, 648, 719 and 821 cm ~. The intensity enlargement of the peaks at 648 and 719 cm 1 clearly demonstrates that the interaction of hydrogen with the catalyst surface increases the SH group density. The two other bands at 540 and 823 cm ~ cannot be ascribed to new SH groups because these frequencies are well out of the range of usually observed SH bendings. As these vibrations correspond fairly well with metal-hydrogen bending modes already observed on transition metal hydrido-carbonyl complexes, they were assigned to Ru-H species. Figure 8 reports the INS spectra observed on two different reduced states of MoS2. The first one, sample A, was heated under hydrogen flow up to 473 K. The second one, sample B, was reduced at a higher temperature: 673 K. For these experiments about 25g of MoS2 were utilized in order to have a similar or even higher amount of adsorbed hydrogen in the neutron beam with respect to RuS2. The INS spectrum of sample A is similar to the one previously reported by Moyes et al. for the as prepared catalyst [ 15]. The main features can be observed in Figure 8 spectrum a: the peak at 660 cm ~ is assigned to SH bending modes, overtones of these modes are found at 1320 and 2000 cm -~. However, after hydrogen adsorption, some differences can be noticed with respect to the spectra published by the same group [ 16]. In ref. 16, the intensity below 400 cm I was similar in the as prepared catalyst and in the spectrum obtained after adsorption of 1 bar of hydrogen, whereas the peak at 650 cm -~, the SH bending modes, became stronger after hydrogen adsorption. In the present work, the intensity of the SH bands does not increase as the hydrogen pressure is raised, but a peak near 120 cm -~, previously assigned to molecular hydrogen [16], is already observed at 0.5 bar (spectruna b)

123 whereas it appeared only at 10 bars in ref. 16. Furthermore, this peak is splitted into two components at 104 and 117 cm l , indicating a stronger interaction of molecular hydrogen with the catalyst. From the INS intensities, it is found that the amount of SH groups on MoS 2 and RuS2, after reduction at similar temperatures, = 500 K, is comparable on both sulfides. On MoS2, when the temperature of reduction is increased, the number of SH groups decreases (spectrmn c), in agreement with the TPD measurements. Unlike the previous results obtained with RuS2, no clear indication of the presence of Mo-H species can be evidenced by INS. 0

500

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Energy transfer E l-E2 (meV)

Energy transfer E l-E2 (meV) Figure 7. INS spectra observed on RuS2. a) degassed solid, b) degassed solid + 0.5 bar of H 2.

Figure 8. INS spectra observed on MoS 2. a) sample A, b) sample A + 0.5 bar of H 2, c) sample B + 0.5 bar of H 2. 10

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124 The evolution of the ESR spectra of various reduced and non degassed MoS 2 samples is presented in Figure 9. The ESR spectrum of the freshly decomposed ATTM exhibits an axial signal with gl - 2.04 - 2.06 and a g2 value at 1.98. This signal coincides fairly well with that previously reported for Mo(V) species in a sulfur environment [ 17]. Reduction of the solid up to 573K only affects the signal intensity without any noticeable modification of the g values. At higher reduction temperatures the shape of the ESR is significantly modified by the appearance of a new feature at g --- 2.00. As shown in Figure 10, the presence of this new signal also induces a discontinuity in the evolution of the number of paramagnetic species.

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Figure 11. Observed ESR signals after reduction, hydrogen desorption and air exposure. ESR spectra were also recorded after hydrogen desorption. The effect of such catalyst pretreatments on the ESR signal of a solid reduced at 573K and at 773K is illustrated in Figure 11. For solids reduced at TOY 573K, the elimination of the adsorbed hydrogen provokes an increase of the Mo(V) signal. However the resulting number of Mo(V) species remains lower than over the non reduced solid suggesting that both sulfur removal and the presence of adsorbed hydrogen contribute to a decrease of the concentration of Mo(V). As far as hydrogen chemisorption is concerned, these results could be interpreted assuming that a fraction of Mo(V) is transformed into Mo(IV) according to the following overall reaction : Mo(V) + 89H 2

r

.Mo(IV) + 89H §

(1)

As a matter of fact, Mo(IV) has a d 2 electronic configuration and therefore does not give rise to an ESR signal. An opposite effect was observed for solids reduced above 573K since hydrogen removal decreases the signal intensity mostly for the high field species. To account for these results reaction (1) could simply be rewritten by replacing Mo(V) and Mo(IV) respectively by Mo(IV) and Mo(III). While direct evidence of the presence of Mo(III) species was obtained in sulfided catalysts, the observed signal at g -- 2.00 is similar to the one

125 observed by Casewit and Rakowski DuBois on Mo(III) complexes coordinated with sulfur ligands [18]. Furthermore, air exposure of high temperature reduced samples yield to a noticeable change of signal intensity and peak shape while on samples reduced at Tr 573K neither the signal nor the peak shape are affected. These results are in agreement with the expected sensitivity of Mo(III) species towards mild oxidation.

4. DISCUSSION AND CONCLUSION The interaction of dihydrogen with the surface of two model transition metal sulfides have been investigated for various sulfur to metal ratio. Both solids exhibit some similarities as well as great differences. Whatever the solid, some superficial sulfur atoms must be removed to generate an activity and to allow hydrogen chemisorption. However the nature of the hydrogen species interacting with the reduced catalysts differs depending on the nature of the solid under investigation. On RuS 2 two hydrogen species can be detected by thermodesorption and their relative amount depend drastically on the S/Ru ratio. The concentration of the strongly adsorbed species reaches a maximum for a low degree of reduction (-- 0.15) and decreases down to an undetectable level when the surface of the catalyst becomes almost sulfur depleted (-- 0.4). By contrast both the activity and the amount of weakly bonded species reach a maximum for this surface composition. INS characterization has demonstrated that the interaction of hydrogen with a reduced solid increases the SH groups density and leads to the appearance of two Ru-H bending modes at 540 and 823 cm 1. It has been observed that the intensity of both the SH and the high frequency Ru-H species are not significantly affected upon evacuation while the 540 -1 cm band intensity decreases. Therefore this bending mode could be ascribed to the weakly bonded hydrogen species which is probably connected with the catalytic properties. Accordingly, two different adsorption mechanisms may account for the observed data. At low solid reduction an heterolytic dissociation of dihydrogen on a Ru-S pair leads concomitantly to the formation of a Ru-H and SH groups while at high surface reduction an homolytic dissociation on some reduced Ru centers would lead to weakly bonded Ru-H species. Results obtained on MoS2 are quite different. By contrast with RuS2, TPD experiments showed that only one hydrogen species is present on MoS2 independently of its degree of reduction. INS spectra have evidenced the formation of SH groups as well as the presence of molecular hydrogen interacting with the catalyst surface. Unlike the previous results obtained with RuS2, no Mo-H species can be detected. ESR measurements have shown that MoS2+• obtained by decomposition of ammonium thiomolybdate possesses unpaired electrons having a (+V) oxidation state. At moderate temperatures, sulfur removal decreases the concentration of such paramagnetic defects as well as hydrogen adsorption does. The desorption of hydrogen partly restored the Mo(V) signal indicating that an electronic transfer occurs during an adsorption-desorption cycle. At high temperature, a similar adsorption process may also involve paramagnetic Mo(III) species. However the number of such paramagnetic sites never exceeds 0.5 % of the total molybdenum content present on the initial overstoichiometric MoS227. Consequently only a small fraction of the adsorbed hydrogen is activated in such a way and other forms of hydrogen have to be taken into account to assess the amount detected by TPD. For an initial S/Mo ratio equal to 2.27, about 10 % of the sulfur anions should be in a

126 (-I) oxidation state in order to maintain the electrostatic neutrality of the slabs. Therefore this sulfur excess could be either SH groups, bridged anions as well as $22 as suggested by Polz et al. Among these different species, the homolytic adsorption of hydrogen on $22- leading to two SH groups could occur without changing the oxidation state of molybdenum. The obtained results evidence the differences between RUSE and MoS2. The former solid has a pseudometallic behavior whereas for MoS2 redox or acid base properties are involved. This demonstrates the difficulties encountered when the interaction of the reactants and the catalyst are described at a molecular level.

5.REFERENCES

.

.

5.

10. 11. 12. 13.

14. 15. 16. 17. 18. 19.

R.R. Chianelli, Int. Rev. Phys. Chem., 2 (1982) 127. R.R. Chianelli, A.F. Ruppert, S.K. Behal, A. Wold and R. Kershaw, J. Catal., 92 (1985) 56. R.B. Moyes, in: Hydrogen effects in catalysis (Z. Paal and P.G. Menon, eds), Dekker, New York and Basel (1988), pp. 583-607. T. Komatsu and W.K. Hall, J. Phys. Chem., 95 (1991) 9966. L. Jalowiecki, A. Aboulaz, S. Kasztelan, J. Grimblot and J.P. Bonnelle, J. Catal., 120 (1989) 108. A. Wambeke, L. Jalowiecki, S. Kasztelan, J. Grimblot and J.P. Bonnelle, J. Catal., 109 (1988) 320. X.S. Li, Q. Xin, X.X. Guo, P. Grange and B. Delmon, J. Catal., 137 (1992) 185. A.B. Anderson, Z.Y. A1-Saigh and W. K. Hall, J. Phys. Chem., 92 (1988) 803. C. Sampson, J.M. Thomas, S. Vasudevan and C.J. Wright, Bull. Soc. Chim. Belg., 90 (1981) 1215. P. Sundberg, R.B. Moyes and J. Tomkinson, Bull. Soc. Chim. Belg., 100 (1991) 967. M. Lacroix, S. Yuan, M. Breysse, C. Dor6mieux-Morin and J. Fraissard, J. Catal., 138 (1992) 409. H. Jobic, G. Clugnet, M. Lacroix, S. Yuan, C. Mirodatos and M. Breysse, J. Amer. Chem. Soc., 115 (1993) 3654. M. Lacroix, C. Mirodatos, M. Breysse, T. D6camp and S. Yuan, Proc. 10th Int. Cong. Catal. (L. Guczi, F. Solymosi and P. T6t6nyi, eds.) Elsevier, Budapest, 1993, pp. 597609. E. Dieman and A. Mialler, Coord. Chem. Rev., 10 (1973) 79. P. Sundberg, R.B. Moyes and J. Tomkinson, Bull. Soc. Chim. Belg., 100 (1991) 967. P.N. Jones, E. Kn6zinger, W. Langel, R.B. Moyes and J. Tomkinson, Surface Sci., 207 (1988) 159. B.G. Silbemagel, T.A. Pecoraro and R.R. Chianelli, J. Catal., 78 (1982) 380. C.J. Casewit and M. Rakowski DuBois, Inorg. Chem., 25 (1986) 74. J. Polz, H. Zeilinger, B. Mi~ller and H. Kn6zinger, J. Catal., 120 (1989) 22.

J.W. Hightower, W.N. Delgass, E. Iglesia and A.T. Bell (Eds.) 11th International Congress on Catalysis - 40th Anniversary

Studies in Surface Science and Catalysis, Vol. 101 1996 Elsevier Science B.V.

127

Characterization of a Zeolite Membrane for Catalytic Membrane Reactor Application Anne Giroir-Fendler, J6r6me Peureux, Henri Mozzanega and Jean-Alain Dalmon Institut de Recherches sur la Catalyse, 2 av. Albert Einstein, 69626 Villeurbanne Cedex France

Abstract This paper describes the morphological and transport properties of a composite zeolite (silicalite) - alumina membrane. Some advantages obtained in combining the membrane with a conventional fixed-bed catalyst are also reported.

1. INTRODUCTION One of the most studied applications of Catalytic Membrane Reactors (CMRs) is the dehydrogenation of alkanes. For this reaction, in conventional reactors and under classical conditions, the conversion is controlled by thermodynamics and high temperatures are required leading to a rapid catalyst deactivation and expensive operative costs. In a CMR, the selective removal of hydrogen from the reaction zone through a permselective membrane will favour the conversion and then allow higher olefin yields when compared to conventional (nonmembrane) reactors [ 1-3 ]. Following this principle, when using a membrane, lower temperatures can be used leading to a longer catalyst lifetime and energy and cost saving. However such a membrane should be stable at high temperature (ca 500~ highly permselective (loss of reactant should be of course avoided) and permeable enough (the permeation rate should be in the same range as the reaction rate). Dense Pd-based membranes have been first used for CMRs applications [4]. They are indeed highly selective for H 2 permeation but are expensive, sensitive to ageing and poisoning and are strongly limited by their low permeabilities. Classical commercial ceramic porous materials, as those obtained via sol-gel processes, generally have adequate permeabilities but could present some drawbacks. They indeed have a limited thermal stability and are generally not permselective enough: their pores are in the mesoporous range and maximum separation factors correspond to Knudsen diffusion mechanisms. To ensure a better separation, molecular sieving will act much better. This size exclusion effect will require an ultramicroporous (i.e. pore size D < 0.7 nm) membrane. Such materials should be of course not only defect-free, but also present a very narrow pore size distribution. Indeed if it is not the case, the large (less separative and even non separative, if Poiseuille flow occurs) pores will play a major role in the transmembrane flux (Poiseuille and Knudsen fluxes vary as D 2 and D respectively). The presence of large pores will therefore cancel any sieving effect. A zeolite membrane, where the pores originate from the structure, presents only one type of (ultramicro)pore and therefore seems to be a good candidate for CMRs application. Moreover the structural origin of the pores should induce a much better thermal stability of the

128 membrane when compared to sol-gel systems where only the texture originates the pores and controls size evolutions. The preparation of defect-free zeolite membranes is the subject of a recent and intensive research for gas separation and CMRs applications [5]. Beside their use in equilibrium-restricted reactions, CMRs have been also proposed for very different applications [6], like selective oxidation and oxidative dehydrogenation of hydrocarbons; they may also act as active contactor in gas or gas-liquid reactions. This paper describes the morphological and transport properties of a composite zeolite (silicalite) - alumina membrane. Results in CMRs applications are also briefly given.

2. EXPERIMENTAL 2.1 Preparation of the zeolite membrane Commercial porous ceramic tubes (SCT/US Filter Membralox T1-70 [7]) were used in this study as support for the zeolite material. They are made (Figure 1) of three consecutive layers of macroporous a-A1203 with average pore sizes decreasing from the external to the internal layer. A thin toplayer made of mesoporous y-Al20 3 was also present in some samples. For gas permeability, gas separation and catalytic measurements the tubes were first sealed at both ends with an enamel layer before zeolite synthesis. Tubes with porous lengths up to 20 cm were used in this study.

Figure 1: Schematic of a cross-section of a commercial SCT tube used as support. Layers 1, 2 and 3 are made of ~-AI203 and have respective thicknesses of (Bm): 1500, 40, 20 and average pore sizes of (~tm): 12, 0.9, 0.2. Layer 4 (optional) is made of ~/-AI203 and has a thickness of 34 pm and average pore size of 4.5 nm.

The silicalite-alumina membrane was prepared after adding a solution containing the silicalite precursor (i.e. silica + template) to the above-mentioned porous tube (hereafter called support) and a specific hydrothermal treatment performed [8]; under the chosen conditions no material is formed in the absence of the porous support. The tube is then calcined at 673 K for removing the template. 2.2 Characterization of the zeolite membrane SEM an& SEM-EDX analyses have been used in order to observe how and where the new material forms on the alumina support. XRD and 29Si MASNMR studies have been performed for its identification. Porous characteristics of the composite material have been explored using N 2 adsorption-desorption experiments (Micromeritics ASAP 2000M). Transport properties have been studied before and after Si deposition using a rig similar to the one for catalytic testings (Figure 2). Pure gas permeabilities (H2, He, N2, normal and isobutane) were studied by measuring the flux passing though the membrane as a function of temperature and pressure for a constant transmembrane differential pressure (no sweep gas).

129

Figure 2. Rig for gas transport and catalytic measurements Gas separation performances (H2/n-butane , n-hexane/2-2 dimethylbutane) have been measured using a sweep gas (countercurrent mode) in order to increase the permeation driving force (no differemial pressure was used); permeate and retentate compositions (see Figure 2) were analysed using on line gas chromatography. Catalytic testings have been performed using the same rig and a conventional fixed-bed placed in the inner volume of the tubular membrane. The catalyst for isobutane dehydrogenation [9] was a Pt-based solid and sweep gas was used as indicated in Fig. 2. For propane oxidative dehydrogenation a V-Mg-O mixed oxide [10] was used and the membrane separates oxygen and propane (the hydrocarbon being introduced in the inner part of the reactor).

3. RESULTS

3.1 Morphology studies The robe has been first characterised just after the hydrothermal treatment step (i.e. before the calcination). The N 2 isotherm is typical of macroporous materials (Figure 3, curve 1) and the tube is gas-tight.

V(cm31g)

1 0

0:.2

0.)4

P/Po

016

j 0'.8

1

Figure 3. N 2 isotherms before (curve I) and after (curve 2) template removal

130 All the other characterization studies have been performed after the calcination step. XRD ext?eriments have shown that the material formed during the synthesis has the MFI structure. 29Si and 27A1 MASNMR spectra indicated that this phase has a Si/AI ratio varying between 550 and 30 as a function of the sample prepared and also that no extra framework AI is present. SEM micrographs (Figure 4) reveal the presence of small crystallites (Figure 4, C) deposited on the large ot-Al20 3 particles of the thicker layer of the support (layer n~ Figure 1) and of larger ones on the external surface of the ot-Al20 3 layer n~ (Figure 4, B). However, when using as starting material a ceramic tube with a ,/-AI203 toplayer (Figure 1), the formation of crystallites at the external surface is not observed.

Fi~mare4. SEM micrographs of the silicalite-alumina composite material. A: cross-section of the tube. B and C: magnifications of the inner surface of the tube and of the first otAI203 layer.

131 Crosswise SEM-EDX analyses show that the global Si/AI ratio (zeolite + support) varies in a very different way (Figure 5) according to the presence or absence of the ~/-AI203 toplayer in the support.

1 SilAI 0.8

0.4 I

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30 L (pm)

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I

50

Figure 5. Crosswise SEM-EDX analysis in the composite membranes (L, radial distance from the inner surface of the tube). Curve 1 (e) corresponds to support without the ~,AI203 toplayer, curve 2 (r-l) with the 7-A1203 toplayer. N 2 adsorption leads to a type I isotherm (Figure 3, curve 2), typical of microporous systems. The corresponding pore size distribution, as calculated using the Horvath-Kawazoe equation [ 11] is given in Figure 6. A sharp maximum, near 0.6 nm is observed. Other methods of isotherm analysis, such the DFT method [ 12], lead to very similar results in the microporous range, but also reveals the mesoporous domain and part of the macroporous one (pore diameters < 200 nm). Very few pores are present in the mesoporous range. Hg porosimetry experiments completed this characterization and have shown [ 13] that pores corresponding to the intermediate ot-Al203 layers (n ~ 2 and 3, see Figure 1) almost completely disappeared after the zeolite synthesis. 0.003 Porous

volume 0.002 (cm3/g) 0.001

0.5

1.0 Pore

1.5

2.0

2.5

diameter (rim)

Figure 6. Pore size distribution (m) and cumulative pore volumes (..-). Microporous domain, Horvath-Kawazoe equation.

132

3.2 Transport studies The zeolite-alumina tube is no more gas-tight after the thermal treatment. The presence or absence of the y-A1203 toplayer in the starting support does not influence the gas transport properties of the final zeolite-alumina tube. For similar temperature conditions the permeance for nitrogen varies with the applied pressure in a different way according to nature of the sample (i.e. the two starting supports or the zeolite-alumina composite, Figure 7).

1000-

Permeance

100 9

(lO'Tmole/s.Pa.m2 )

_

..2

.

_

10-

i

2

s

Internal pressure (1# Pa)

Figure 7. N2 permeances at 30~ in the tx-Al203 support alone (1, O), in the ot-Al203 support plus the ,/-AI203 toplayer (2, II) and in the zeolite-support composite (3, A). In the zeolite-alumina composite the behaviour of different gases (permeance of pure gases as a function of the temperature, Figure 8) behave very differently from those predicted by ideal Knudsen diffusion processes.

10- ~ c L " ' ~ ' c ~ . . . . . 5 3 . . " Permeance ~ (lO'7mole/s.Pa. 2 )" w _

o' I

i

100

i

i

I

I

200 300 Temperature ('C)

I

I

400

Figure 8. Permeances in the zeolite-support composite for H2 (1, r'l), N 2 (2, A), He (3, II), n-butane (4, O), isobutane (5, o). Intemal pressure 1.2 105 Pa. Gas separations also show non-Knudsen behaviors. In the case of the H2/n-butane mixture, the temperature has a drastic effect on the main permeating gas, at low temperature almost only butane, the heavier component, permeates (Figure 9).

133

When introducing a mixture of n-hexane and 2-2 dimethylbutane (45/55 molar ratio), almost only n-hexane permeates: the permeate contains up to 99.5% of the linear isomer (Figure 10).

Flux 100(10 mole/h)

100

200

300

400

Temperature (~

Figure 9. H 2 (D) / n-butane ( t ) separation with the composite zeolite-alumina membrane (fluxes in the permeate as a function of the temperature). A mixture of hydrogen, n-butane and nitrogen (12 : 14 : 74) was fed in the tube (Fig. 2) with a flow rate of 4.8 l/h. Sweep gas (N2), countercurrent mode, flow rate 4.3 l/h.

12-

Flux

(104mol~)

8-

$_

w

60

i ~

..w

100 Temperature (~

I--

160

l

200

Figure 10. N-hexane (rl) / 2-2 dimethylbutane (m) separation with the composite zeolite-alumina membrane (fluxes in the permeate as a function of the temperature). A mixture of n-hexane, 2-2 dimethylbutane and nitrogen (5 : 6 : 89) was fed in the tube (Fig. 2) with a flow rate of 2 l/h. Sweep gas (N2), countercurrent mode, flow rate 0.5 l/h.

3.3 Catalytic studies Most of the results have been already partly presented in [9] (isobutane dehydrogenation) and [10] (propane oxidative dehydrogenation). Let us recall that the membrane presented in this paper has been associated with a fixed bed catalyst placed within the tube. In the isobutane dehydrogenation the catalytic membrane reactor allows a conversion which is twice the one observed in a conventional reactor operating under similar feed, catalyst and temperature conditions (and for which the performance corresponds to the one calculated from thermodynamics) [9]. In the propane oxidative dehydrogenation, where the membrane separates the two reactants, a 20% increase in the yield was observed with respect to a conventional reactor working at isoconversion [ 10].

134 4. DISCUSSION

4.1 Morphological properties The starting alumina support (Figure 1) is either macroporous or, in the presence of the y-Al203 toplayer, mesoporous. These two materials are highly permeable (Figure 7). The fact that the tubes become completely impervious to any gas aiter the hydrothermal synthesis suggests that the material originated from the synthesis has plugged up the support, resulting in a defect-free gas-tight tube (no cracks or pin-holes). The presence of macropores (Figure 3) suggests however that the porosity (non-passing through pores) has not completely vanished. After the calcination step, experimental data (XRD, 29Si MASNMR) show that a zeolite with the silicalite structure has been formed. 29Si MASNMR indicates for the zeolite material a Si/A1 ratio depending on the sample prepared: it has been observed that both the natures of the silicon source and of the alumina supports may originate these fluctuations. SEM micrographs (Figure 4) show the deposition on the ct-Al203 grains of small crystaUites with the typical hexagonal shape of silicalite. The pore size distribution, as deduced from N 2 adsorption, presents a very narrow peak centred on 0.5 nm, also in good agreement with the pore diameter of silicalite-type zeolites. All these data confirm that a well-defined zeolite silicalite-type crystalline phase has been formed in the presence of the alumina porous tube (which seems indispensable for the zeolite synthesis, as no material is formed in its absence). SEM-EDX analyses (Figure 5) underline the difference of the final material whether the support has a y-Al203 toplayer or not. In its absence, the large zeolite crystals growing on the inner surface of the tube (Figure 4) result in high surface Si/Al ratios. In the porous domain corresponding to layers 2 and 3, the Si/Al ratio is more or less constant and could correspond to a complete filling up of the porous volume (layers 2 and 3 have similar porosities). When analysing the much thicker layer n~ where SEM micrographs show that the zeolite does not fill up the pores, the Si/Al ratio decreases. When using a support with the ~/-AI203 toplayer, the absence of zeolite crystals on the surface, as observed by SEM, is confirmed by the EDX analysis: the Si/Al ratio strongly decreases at the surface. This means that the separative zeolite layer has been here formed only in the bulk of the porous support and not mainly deposited on the outer surface of the support (some infiltration in the support being possible), as is generally the case in zeolite membranes described in the literature [ 14-19]. The presence of the zeolite only within the support could lead to some advantages as far as mechanical and thermal stabilities are concerned. Indeed the separative layer is here protected by the resistant t~-AI203 (thus avoiding damage when introducing for instance catalyst pellets in the tube). The thermal stability of the present zeolite membrane has been also checked: after calcination at 700~ the microporous character remains [ 13 ] and transport properties are not significantly modified after oxidative dehydrogenation of propane at 600~ [ 10] and also after several hours at 850~ under air/steam mixtures [20]. Let us also underline that the zeolite membrane described here is prepared in only one hydrothermal synthesis step. This is not always the case for other preparations reported in the literature for which several successive zeolite synthesis could be needed with the same support to suppress defects.

4.2 Gas transport properties Figure 7 shows that N 2 permeability strongly depends on the pore size. For the macroporous support (curve 1) Poiseuille flow occurs, leading to an increase of the permeance

135 with the pressure. For the mesoporous support (curve 2), Knudsen diffusion rules the transport and permeance becomes pressure independent. When using the microporous zeolite membrane (curve 3) the N 2 permeance decreases when the pressure increases: such a behaviour can be accounted for by activated diffusion mechanisms [21], which are typical of zeolite microporous systems. In such systems the diffusivity depends on the nature and on the concentration of the diffusing molecule which interacts with the surface of the pore. For gases with low activation energies of diffusion, a decrease of the permeability can be observed [22]. Figure 8 indicates that the permeance of n-butane goes through a maximum when increasing the temperature. This observation agrees with the above-mentioned mechanism. At low temperature the nC 4 concentration is high in the zeolite and permeance is low due to the low probability of finding a free site ensuring the mobility. An increase of temperature will then favour the transport. However, when the temperature is too high the coverage of the interacting species decreases and becomes the limiting factor, leading to a global decline of the transmembrane flux. Such phenomena have been also recently described in other zeolite membranes [23], [ 18]. The changes in H2/n-butane separation when increasing the temperature (Figure 9) can be explained on the same basis. At low temperature n-butane interacts strongly with the zeolite and owing to its size completely blocks the pores. Hydrogen has no more room to penetrate the pore and only n-butane diffuses. When increasing the temperature, the occupancy of the zeolite porous volume by n-butane progressively decreases and the hydrogen flux increases. However even at high temperature, the observed separation factor is less than the one expected from Knudsen diffusion. These results underline that the experimental separation factor may strongly differ from the one expected from pure gas permeabilities measurements. However this is not always the case, especially when the two components weakly interact with the surface. When using the membrane to separate a H2/isobutane mixture, the permeation of isobutane, due to its size, is restricted over the entire temperature range and the transmembrane fluxes of the two components of the mixture better follow the permeabilities of the pure gases. Separation factors are here much higher (factors up to 80 have been measured). Molecular sieving effect of the membrane has been evidenced using a mixture of two isomers (i.e. no Knudsen separation can be anticipated), n-hexane and 2-2 dimethylbutane (respective kinetic diameters 0.43 and 0.62 nm). Figure 10 shows the permeate contains almost only the linear species, due to the sieving effect of the zeolite membrane (pore size ca 0.55 nm). This last result also underlines that the present zeolite membrane is almost defect-flee.

4.3 Catalytic properties In isobutane dehydrogenation, when using the zeolite membrane associated with a fixed-bed Pt-based catalyst placed in the inner volume of the tube, the yield is twice the one observed in a conventional reactor [9]. This is especially due to the high separative performance of the membrane which selectively removes the produced hydrogen from the reactor. When using a less selective membrane, for instance the starting support with the 7AI203 toplayer (for which only Knudsen separation occurs), the performance of the CMR is limited by the loss of isobutane through the membrane [9]. The zeolite membrane has also been used in the oxidative dehydrogenation of propane. The goal is here to limit the undesirable complete oxidation owing to the controlled addition of oxygen diffusing through the membrane. Following this mode, it is possible to keep a low oxygen/hydrocarbon ratio all along the fixed bed and then limit complete oxidation. Using the

136 zeolite membrane a 20% increase in the propene yield has been obtained [ 10]. Here also, when using only the starting support as diffusion barrier, it is not possible to control the 0 2 addition and the performance is quite similar to the one observed in a conventional reactor [ 10].

CONCLUSIONS The synthesis of a well-crystallised silicate-type zeolite within the porous volume of an ot-Al20 3 tubular support leads to a defect-flee zeolite membrane showing high thermal and mechanical stabilities. In gaseous mixtures, the separative performance of the membrane greatly depends on the temperature and the nature of the components, especially when strong interactions take place. When occurring, molecular sieving leads to high separative performances. These properties, combined with its high thermal stability, suggest that the zeolite membrane is a very promising material for CMR applications.

Acknowledgements We are indebted to SCT - US Filter for providing alumina support tubes.

REFERENCES [1 ] N. Itoh, AIChE J., 33 (1987) 1576. [2] J.N. Armor, Appl. Catal., 49 (1989) 1. [3] A. Champagnie, T.T. Totsis, R.G. Minet, I.A. Webster, Chem. Eng. Sc., 45 (1990) 2423. [4] V. M. Gryaznov, Plat. Met. Rev., 30 (1986) 68. [5] J.L. Falconer, R.D. Noble and D.P. Sperry Eds, Membrane Separations Technology: Principles and Applications, S.A. Stern and R.D. Noble (Eds), Elsevier, 1994. [6] J-A. Dalmon, Handbook of Heterogeneous Catalysis, G. Ertl, H. Kn6zinger and J. Weitkamp (Eds), VCH, Chap. 9.3 (in press). [7] R. Sofia, Catal. Today, 25 3-4 (1995), 285. [8] J. Ramsay, A. Giroir-Fendler, A. Julbe and J-A. Dalmon, F. Pat. 94 05562 (1994). [9] D. Cazanave, J. Peureux, A. Giroir-Fendler, J. Sanchez, R. Loutaty and J-A. Dalmon, Catal. Today, 25 3-4 (1995), 309. [ 10] A. Pantazidis, J-A. Dalmon and C. Mirodatos, Catal. Today, 25 3-4 (1995), 403. [11] G. Horvarth and K. Kawwazoe, J. Chem. Eng. Japan, 16-6 (1983) 470. [ 12] J.P. Olivier, W.B. Conklin and M.v. Szombathely, Characterization of Porous Solids III, Studies Surf. Sc. Cat., Elsevier, 87 (1994) 81. [ 13] D. Uzio, J. Peureux, A. Giroir-Fendler, J-A. Dalmon and J.D.F. Ramsay, Characterization of Porous Solids III, Studies Surf. Sc. Cat., Elsevier, 87 (1994) 411. [14] H. Suzuki, US Pat 4 699 892 (1987). [15] A. Ishikawa, T.H. Chiang and F. Foda, J. Chem. Soc. Chem. Com., 1989, 12, 764-765. [ 16] M.D. Jia, B. Chen, R.D. Noble and J.L. Falconer, J. ofMembr. Sci., 90 (1994) 1. [17] Y.M. Ma and S. Xiang, US Pat. 5 258 339, 1993. [ 18] F. Kapteijn, W. Bakker, J. Moulijn, H. van Bekkum, Catal. Today, 25 (1995), 213. [ 19] P. M6riaudeau, A. Thangaraj and C. Naccache, Microporous Mat., 4 (1995) 213. [20] A. Giroir-Fendler and J-A. Dalmon (to be published). [21 ] J. Caro, H. Jobic, M. Bialow, J. Karger and B. Zibrowius, Adv. Cat., 39 (1993) 351. [22] J. Peureux, A. Giroir-Fendler, H. Jobic and J-A. Dalmon (to be published). [23] Z.A.E.P. Vroort, K. Keizer, H. Verweij and A.J. Burggraaf, Proc. ICIM4, 1994, 503.

J.W. Hightower, W.N. Delgass, E. Iglesia and A.T. Bell (Eds.) 11th International Congress on Catalysis - 40th Anniversary

Studies in Surface Science and Catalysis, Vol. 101 9 1996 Elsevier Science B.V. All rights reserved.

137

Catalytic Reduction of SO3 Stored in SOx Transfer CatalystsA Temperature-Programmed Reaction Study Gwan Kim a and Michael V. Juskelis b aGrace Davison Division and bResearch Division, W.R. Grace & Co.-Conn. 7500 Grace Drive, Columbia, MD 21044-4098 USA

Abstract A laboratory test method has been developed to assess the efficiency of SO x transfer catalysts for the reduction of SO 3 stored in the catalysts. The test method is based on the temperature-programmed reaction (TPR) with propane over a presulfated catalyst. The reaction products released during the test were determined by means of mass spectrometry (MS). Using this technique, we have appraised the oxides of four transition metals, V, Cr, Fe, and Ce, each impregnated on microspheres of a ternary oxide containing Mg, La, and A1. Based solely on the onset temperature for H2S release, the ranking in catalyst efficiency found was V > Ce > Fe > Cr. In good contrast to this, the ranking based on thermogravimetric analysis (TGA) test for the SO 2 oxidation / SO 3 trapping was Ce > Cr > V > Fe. Using the TPR/MS test method, we have also examined a sample of DESOX TM, the leading commercial SO x transfer catalyst consisting of (MgO)2A1203 and oxides of Ce and V. The onset temperature for H2S release was approximately 450~ when propane was the reactant. It was approximately 580~ when H2 or CH4 was used. These data imply that the rate of H2S release is limited by the supply of a c t i v e hydrogen. Thus, the data further suggest that the primary role of catalytic metal oxide for the reduction of SO3 stored in the catalyst is to facilitate C-H bond cleavage on the catalyst by activating hydrocarbons. 1. I N T R O D U C T I O N The three key steps that determine the performance of a SO x transfer catalyst which controls SO x (SO 2 and SO3) emission from a fluid catalytic cracking unit (FCCU) are (1) the oxidation of SO 2 to SO 3 under the FCCU regenerator condition, typically at 7 0 0 - 730~ (2) the trapping of SO 3 on the catalyst in the form of sulfates, and (3) the reduction of sulfates to release sulfur as H2S in the FCCU riser, typically at 520 - 530~ While the reactions involved in each of these steps are thermodynamically favorable [ 1, 2], steps 1 and 3 must be catalyzed in order to achieve a sufficiently high rate of SO x emission control. The performance of catalysts for steps 1 and 2 combined can be easily assessed by running a TGA test at 700~ [3]. However, no satisfactory test method has been developed for appraising the catalyst efficiency in the laboratory for step 3 under a condition similar to the FCCU riser environment. Some [3, 4] applied TGA to follow the weight loss resulting from the reduction of sulfates with pure or diluted H 2 at 650~ Their data are sufficient to show the reducibility of the sulfates with molecular H 2 under that particular condition, which is far different from a typical FCCU riser condition [5]. Others [6] attempted to simulate their test by including multiple gas oil cracking-stripping-regeneration cycles. Their results, however, lack reproducibility and correlation with catalyst performance in FCCUs. We have developed a laboratory test method whereby the efficiency of a catalyst for step 3 can be readily assessed.

138 To illustrate this test method, we present some of the results obtained from a study of the oxides of four transition metals, V, Cr, Fe, and Ce, which are known to be active for catalyzing step 1 [7].

2. EXPERIMENTAL 2 . 1 . Catalyst Preparation The four catalysts (Samples B-E, Table 1) were prepared by the usual incipient-wetness impregnation with either oxalate or acetate solutions on an oxide containing Mg, La, and A1, followed by 115~ drying and 2h heating in 538~ air. The chemical compositions of all samples were determined by using inductively coupled plasma (ICP) analysis. A master batch of the ternary oxide (Sample A consisting of 37.4% MgO, 18.9% La203, and 42.8% A1203 by weight, Table 1) was prepared by spray drying of a hydroxide slurry, followed by washing, drying and 2h air calcination at 718~ The slurry was prepared by a single-stage coprecipitation at pH 8.5 in a mix-pump reactor, feeding magnesium nitrate, sodium aluminate, and sodium hydroxide solutions simultaneously, followed by aging at pH 9.5. The details of such a preparation have already been disclosed [8]. 2 . 2 . Sulfation of Samples for TPR Test A 0.40g sample was sulfated by running 3h SO 2 oxidation at 704~ in a down-flow Vycor glass reactor with flowing N 2 (126 ml/min) containing 9.50 vol.% O 2 and 0.6000 vol.% SO 2. The sample was cooled in flowing N 2 prior to discharge.

2 . 3 . Propane-TPR/MS Test In the first set of comparison test runs, a 0.10g of each sulfated sample was examined by TPR in a soak (300~ (up to 850~ at 20~ mode, using propane at 14.2 ml/min as the reactant. In the second set of propane-TPR comparison test runs, another 0.10g of each sulfated sample was examined by a soak (300~ (30~ (up to 530~ mode. The reaction products such as SO 2 and H2S released during the course of a TPR run were determined by means of mass spectrometry (Hiden Analytical, HAL-2) in MID mode, monitoring mass numbers 48 for SO fragment from SO 2, and 34 for H2S. 2 . 4 . H 2 and CH4-TPR/MS Tests using H 2 o r C H 4 in Place of Propane In these soak (300~ (up to 900~ mode tests, a 0.10g of sulfated DESOX sample--the leading catalyst based on magnesium aluminate spinel with an excess MgO catalyzed with Ce and V oxides [9]Bwas allowed to react with either n 2 (5%)/A/" (50 ml/min), undiluted n 2 (70 ml/min) or CH 4 (14 ml/min).

2 . 5 . TGA Test for SO 2 Oxidation and SO 3 Trapping Using a TGA unit identical to one already described elsewhere [4], an 8 - 16 mg sample was subjected to flowing N 2 (225 ml/min) at 700~ until the weight stabilized. To this was added an O2/N2 gas mixture followed by an SO~fN2 gas mixture to provide a feed containing 0.73% O 2 and 0.27% SO 2. The weight gain over a period of 15 min exposure to SO 2 oxidation was recorded. 3 . R E S U L T S AND DISCUSSION

3 . 1 . Catalyst Properties The chemical composition from ICP analyses and nitrogen porosimetry data obtained by Quantachrome AUTOSORB-6 are summarized in Table 1 for all the five samples. Note that the catalytic metal oxide loadings in Samples B-E were adjusted so that the efficiency of each catalyst for step 3 can be directly compared on the same basis, per gram atom of metal. The

139 four catalyst samples (B-E) exhibit higher surface areas than Sample A, reflecting a creation of new surfaces resulting from post-impregnations. It is reasonable to expect, however, that the total SO 3 storage capacity as well as the number of SO 3 trapping sites on the surface remain nearly identical for all the five samples. It is also reasonable to assume that none of the four catalytic metal oxides contributes significantly to the SO 3 storage capacity. The powder X-ray diffraction (XRD) data reveal the presence of microcrystaUine MgO and La203 only for all samples. The size of the catalytic metal oxide crystallites is presumed to be quite small at such a low loading, well below the monolayer coverage. Table 1 Property of Catalysts Sample Metal Oxide Loaded (Wt.%) Source S.A. (N2), m2/g Median Pore Diameter, nm

A None 0

B V205 2.49 Oxalate 167 8.3

151 14.5

C D C r 2 0 3 Fe203 2.14 2.13 Acetate Oxalate 193 186 6.6 7.2

E CeO 2 4.22 Acetate 186 6.9

3.2. Propane-TPR/MS Study The propane-TPR/MS results obtained from the soak-ramp-soak mode tests are presented in Figure 1 for the five samples. Some of the results obtained from the soak-ramp-soak mode tests are shown in Figure 2, which emphasizes H:S release at relatively low temperatures

1.0 .

.

.

.

.

.

:

/

---C

I

_ ~()4h

= k-

~/550

"

"

114 I

/ i i(-'"-\V".\~ I .i.."1i

'} I

~'~""+~

600 700 TEMPERATURE ( ~ )

800

Figure 1. Results from propane-TPR/MS tests in soak-ramp mode of sulfated samples,

0

200

400 TIME ( sec. )

600

Figure 2. Results from propane-TPR/MS tests in soak-ramp-soak (530~ mode of sulfated samples.

below 530~ All the data taken from the propane-TPR/MS tests are summarized in Table 2. These data clearly establish V205 to be the best catalyst for step 3, exhibiting the lowest onset temperature for H2S release, approximately 460~ Well-formulated catalysts containing both V and Ce such as DESOX exhibit an even lower onset temperature for H2S release,

140 Table 2 Performance of Catalysts in Laboratory Tests by Propane-TPR/MS and TGA Sample Metal Oxide H2S Release in Propane-TPR Test Onset Temp. (~ Take-off @ or below 530~ Total H2S Released, I.t mol Weight % Gain in TGA Test

A None 600 No 118 3.0

V205 Cr203 Fe203

B

C

E CeO 2

460 Yes 235 5.3

500 No 229 11.3

580 No 212 9.8

D 520 Yes 171 4.5

approximately 450~ which is well below the typical FCCU riser temperature of 520-530~ The next best in the rank solely based on the H2S release onset temperature is CeO v Notice in Figure 2, however, that there is practically no take-off in H2S release over Sample E as long as the temperature remains at 530~ This represents a situation where SO 3 trapped on the surface or stored in the bulk [ 10] is released very slowly, thus creating a condition where the rate of SO x emission control is limited by the number of SO 3 trapping sites. Considerably trailing behind V205 but only slightly behind CeO z is Fe20 3, which shows an onset temperature for H2S release of approximately 520~ This means Fe203 is not expected to be able to adequately catalyze step 3 in the FCCU riser environment because of short contact time [5], even though the temperature at the very bottom of the riser exceeds 530~ In fact, the result of our pilot plant test of such a catalyst is in agreement with this assessment. Thus, it is quite clear that the onset temperatme for H2S release is more critical than the rate of take-off in determining the catalyst efficiency for step 3. Two samples, one (Sample C) with Cr203, the other (Sample A) without any catalytic metal oxide, showed no release at all below 530~ Judging from the pilot plant experience with Fe203-containing catalysts, these two are not expected to be able to function as SO x transfer catalysts. The total amount of H2S released per 0.10 g of sulfated sample during the soak-ramp (up to 800~ mode test, i.e., the integrated area up to 800~ under each plot in Figure 1, is included in Table 2. Because each sample was 3h sulfated and all three steps (1 - 3) are reflected in these data, it is not surprising to see that the results do not parallel with the ranking based on TGA data (also included in Table 2) which cover a 15-rain period of the steps 1 and 2 combined. Nevertheless, we found these data quite useful especially when TGA data for steps 1 and 2 were not immediately available.

3.3. H2-TPR/MS and CH4-TPR/MS Studies The result obtained from a H 2 (5%)/At (95%) - TPR/MS in a soak-ramp mode test is shown in Figure 3 for a sample of DESOX. The onset temperature found for H2S release in this case, approximately 580~ is substantially higher than 450~ the typical onset temperature found in the propane-TPR/MS test. The result was essentially identical in terms of the onset temperature for H2S release even when undiluted H 2 was used as the reactant. Unlike the propane-TPR/MS tests, where the reaction products are essentially H2S only with virtually negligible amounts of SO 2, H2-TPR/MS tests always showed both SO 2 and H2S. These data, notably the pattern of change in the rates of SO 2 and H2S released with temperature in Figure 3, clearly demonstrate, as expected, that the reduction of S § to S 2 in step 3 is a consecutive reaction. The result obtained from a CH4-TPR/MS test of the same DESOX sample is shown in Figure 4. As in H2-TPR, the H2S release is preceded by the evolution of a large amount of SO 2. The onset temperature (580~ of H2S evolution is also comparable to that observed in H2-TPR. It has been presumed that syngas formation [ 11 ] is associated with the generation of H2S in this case.

141

o5-

.,..,

so2 ..

>r

r

~.

"-

SO2 :-"-..

4-

,. .. :9i : :.

H2S i

o 3

9 o

.

~.1

r~

..

2-

r~


95

95

Pd and Pd-Au palladium-based catalysts and 1,2 ~3C-ethylene. w/KOAc > Pd w/KOAc). The total amount of vinyl acetate produced per pulse agreed with the VAM STY data obtained from the fixed bed reactor with the exception that the Pd-Au catalyst produced more vinyl acetate in the TAP reactor than the Pd-Au w/KOAc catalyst. This can be attributed to the Pd-Au catalyst's VAM production rate being limited by the desorption of VAM when operated at elevated pressures and with a constant flow of C2H 4 and 02. Several long term multipulse experiments were used to help to bridge the gap between the steadystate/high surface coverage fixed bed reactor work included in Table 1 and the transientflow surface coverage pump-probe TAP reactor experiments of Figure 3. The long term multipulse experiments were conducted with a mixture of ethylene, d4-acetic acid, and oxygen characteristic of VAM synthesis conditions. This mixture was pulsed into a 9 cm3/min stream of nitrogen at an intensity of 2.6 torr/pulse of gas. Each experiment consisted of 200 pulses spaced 1 second apart from one another. As can be seen from Figure 4, the Pd-Au w/KOAc catalyst produced the most vinyl after 200 pulses in direct agreement with the fixed bed reactor results. It is also seen that vinyl acetate production is limited at the early stages of this experiment when KOAc is absent from the catalyst mix. Again, this can be attributed to the slow desorption rate of VAM when KOAc is not present. Figure 4. d3-vinyl acetate (MW=89) response curves derived from a long term multipulse sequence with a premixed d3-acetic acid + ethylene + oxygen feed.

196 Figure 5 shows VAM evolution curves which depict the desorption rate of VAM as a function of acetic acid concentrations. These results were obtained when unlabelled ethylene and oxygen are pulsed into a carrier gas that contains various amounts of acetic acid. As the concentration of acetic acid in the feed stream is increased, the desorption of vinyl acetate is enhanced. Thus, the role of KOAc seems to be linked with its ability to keep acetic acid on the catalyst surface. This Figure 5. Vinyl acetate transient response curve (MW=86) for result complements nicely the obtained from a typical pump-probe experiment with the hypothesis of Tamura and palladium-based catalysts and unlabeled ethylene. Results were Yasui (1979), who theorized obtained with the Pd-Au w/KOAc catalyst. that KOAc forms double salts (e.g., KH(OAc)2) with the acetic acid on the surface of the catalyst. Moreover, these complexes would form a molten salt layer on the surface of the catalyst which would set-up an environment referred to as supported liquidphase catalysis. Carbon dioxide can form as a result of combustion of ethylene, acetic acid, or vinyl acetate. The 1,2 t3Cethylene experirnents provided evidence on the role of Au and KOAc and its influence over the formation of carbon dioxide in the process, t2CO2 would be a characteristic of acetate decomposition, while 13CO2 would be characteristic of ethylene decomposition. Figure 6 compares the 12CO2 response curves obtained from each of the catalysts used. The maximum in the ~CO 2 response curve is attributed to the combustion Figure 6. 12CO2transient response curves (MW--44) obtained of acetic acid, while the from a typical pump-probe experiment with the palladium-based extended tail of the Pd catalysts and 1,2 ~3C-ethylene. w/KOAc and Pd curves is attributed to a secondary combustion pathway. The secondary acetate decomposition pathway is most likely attributed to the conversion of carbon monoxide to carbon dioxide.

197 Figure 7 compares the transient responses for evolution for riCO obtained from a typical TAP pumpprobe experiment with the four catalyst samples. 12CO can be directly attributed to the combustion of acetic acid. The combustion of acetic acid (CH3COOH) is believed to evolve carbon dioxide via the carboxylic (COOH) fragment of the molecule. Then, the remaining carbon (CH 3 group) is converted to carbon monoxide at low oxygen coverages and onward to Figure 7. nCOtransient response curves (MW=12) obtained carbon dioxide at higher from a typical pump-probe experiment with the palladium-based oxygen coverages [Davis and catalysts and 1,2 ~3C-ethylene. Barteau, 1991; Aas and Bowker, 1993]. This is supported by the TAP reactor experiments since the maximum of the riCO response curve occurs between 5.5-6.0 seconds, while the CO 2 evolution occurs earlier between 4.55.0 seconds. It is also seen that both KOAc and Au enhance the formation of CO. This agrees with the earlier observations with the CO 2 response curves, and supports the hypothesis that the secondary CO 2 peak occurs from the conversion of CO to CO 2. The results from Figures 6 and 7 support the observation that acetic acid combustion is accelerated by the presence of Au and KOAc. The evolution of carbon dioxide is enhanced by both Au and KOAc, while the evolution of carbon monoxide is enhanced by the presence of KOAc and suppressed by the presence of Au. This shows that acetic acid combustion is more complete with a Pd-Au alloy versus Pd alone which is important since carbon monoxide can act as a temporary catalyst poison in the process. These results agree with Nakamura and Yasui's (1980) on acetic acid oxidation which showed an increase in acetic acid combustion when KOAc is added to a Pd catalyst. Figure 8 shows the t3CO2 (MW--45) response curve for the standard pump-probe experimental conditions. ~3CO2would be generated from either the combustion of 1,2 13C2H 4 or the product 1,2 ~3C-vinyl acetate. Since the curve has a maximum at 4.5 seconds which is before the evolution of vinyl acetate at 5.5-7.5 seconds, ethylene combustion seems to be the pathway that can be attributed to the formation of this peak. Moreover, on interpretation of Figure 7, it is apparent that both the Pd-Au w/KOAc and the Pd w/KOAc catalysts produced less 13CO 2 than their Pd-Au and Pd counterparts. Therefore, it can be concluded that KOAc impeded the combustion of ethylene.

198

Figure 8. ~3CO2 transient response curves (MW--45) obtained from a typical pump-probe experiment with the palladium-based catalysts and 1,2 ~3C-ethylene. Figure 9 shows the results from experiments with the Pd-Au w/KOAc catalyst and the dependence of acetic acid and ethylene combustion on surface oxygen coverage. Figure 9a shows the ~3CO2 evolution curve, while Figure 9b shows the ~2CO2 evolution curve. The experiments were conducted with the same size ~3C2H4 pulse intensity while the 02 pulse intensity was increased so that the C2H 4 to 0 2 molar ratios were varied between 2 to 1 and 4 to 1. It shows that increased oxygen concentration enhances both the combustion of acetic acid and ethylene. As can be seen from the figure, ethylene combustion is virtually eliminated at low oxygen coverages. This suggests that the main route to carbon dioxide formation in the synthesis of vinyl acetate is the combustion of acetic acid. This agrees with previous researchers [Davidson et al., 1984; Crawthome et al., 1994] who had observed similar behavior on other palladium-based catalysts.

Figure 9. ~3CO2(MW--45) and ~2CO2 (MW--44) response curves for a ethylene + oxygen pulse to a continuous acetic acid saturated nitrogen stream after it is passed over the Pd-Au w/KOAc catalyst with different ethylene to oxygen ratios (e.g., 4 to 1).

199

4.

CONCLUSIONS AND SUMMARY

The effect of the catalyst composition upon the catalyst activity, selectivity, and reaction pathways was examined using a conventional high pressure fixed bed reactor and a TA_P reactor. Particular emphasis was placed upon the effect of Au and KOAc on the acceleration or impedance of the pathways associated with vinyl acetate synthesis. A summary of the key findings is given below: (1) KOAc enhanced the combustion of acetic acid to carbon dioxide and carbon monoxide (2) KOAc suppressed the combustion of ethylene to carbon dioxide (3) KOAc and Au enhanced the desorption rate of vinyl acetate (4) Au and KOAc enhanced the formation of vinyl acetate (5) Au suppressed the formation of carbon monoxide from the combustion of acetic acid (6) Au enhanced the conversion of carbon monoxide to carbon dioxide On conclusion of this experimental work, key insight was learned as to the role of Au and KOAc in the synthesis of vinyl acetate. Additionally, this work helped to identify how Au and KOAc physically perturbed the chemistry. For example, Neurock (unpublished results) has shown by density functional theory calculations that the binding energy of CO and other adsorbates are higher on Pd-Au clusters than on Pd clusters alone. This supports the hypothesis that Au has an electronic stabilization effect upon Pd complexes in the system which increased the catalyst's ability to turnover these moieties to vinyl acetate. On the other hand, Crathorne et al. (1994) have shown that acetic acid exits in high concentrations near the surface of Pd-KOAc catalysts. It is believed that KOAc created a molten salt layer with the acetic acid on the surface of the catalyst which promoted high Pd-OAc surface coverages. This multi-layer coverage provided a protective coating on the surface of the catalyst which impeded ethylene combustion by isolating discrete palladium sites for preferential adsorption of ethylene and oxygen that leads to VAM formation rather than combustion. The formation of vinyl acetate is viewed to occur either on small clusters of palladium acetate dissolved in the "supported liquid phase" or on palladium acetate dense surfaces. The overall picture of the role of a Pd-Au-KOAc-SiO 2 catalyst in the synthesis of vinyl acetate is included as Figure 10.

Figure 10. Role of Pd-Au-KOAc-SiO 2catalyst in the manufacturing of vinyl acetate

200 5.

ACKNOWLEDGMENTS

The authors would like to acknowledge the contributions of several individuals for their insight and hard work in achieving the data included in this report: J. Scott McCracken (TAP reactor), Kevin S. Slusser (fixed bed reactor), and Tom Borecki (catalysis synthesis). The authors would also like to thank DuPont's vinyl acetate business and manufacturing teams for allowing this work to be published. 6. REFERENCES N. Aas and M. Bowker, J. Chem. Soc. Faraday Trans., 89 (1993), 1249. S.M. Augustine, and J.P. Blitz, J. Catal., 142 (1993) 312. W.J. Bartley, S. Jobson, G.G. Harkreader, M. Kitson, and M. Lemanski, U.S. Patent 5,274,181 (1993). T.C. Bissot, U.S. Patent 4,048,096 (1977). E.A. Crathorne, D. MacGowan, S.R. Morris, and A.P. Rawlinson, J. Catal., 149 (1994) 254. J.M. Davidson, P.C. Mitchell, N.S. Raghavan, Front. Chem. React. Eng. (Proc. - Int. Chem. React. Eng. Conf.), 1 (1984), 300. J.L. Davis and M.A. Barteau, Surf. Sci., 256 (1991) 50. H. Debellefontaine, and J. Besombes-Vailhe, J. Chim. Phys., 75 (1978) 801. Farbwerke Hoechst Artiengesellschaft, Great Britain Patent, 1,188,737 (1967). J.T. Gleaves, J.R. Ebner, and P.L. Mills, paper presented at USPC-2, St. Louis, Mo., Sept. 1995. T. Kunugi, K. Fujimoto, H. Arai, T. Kono, and A. Namatame, Kogyo Kagaku Zasshi, 71 (1968) 2007. I.I. Moiseev, M.N. Vargaftik, J.K. Syrkin, and Aka Doklady, Nauk SSR, 133 (1960), 377. M. Nakamura, Y. Fujiwara, and T. Yasui, U.S. Patent 4,087,622 (1978). S. Nakamura, and T. Yasui, J. Catal., 17 (1970) 366. S. Nakamura, and T. Yasui, J. Catal., 23 (1971) 315. S. Nakamura, and T. Yasui, J. Japan Pet. Inst., 23 (1980), 416. M. Neurock (unpublished results). B. Samanos, P.Boutry, and R. Montarnal, J. Catal., 23 (1971) 19. Sennewald, U.S. Patent 3,631,079 (1971). W. Schwerdtel, Chem. Ind., (1968), 1559. M. Tamura, and T. Yasui, Shokubai, 21 (1979), 54. M.N. Vargaftik, V.P. Zagorodnikov, and I.I. Moiseev, Kinet. Katal, 22 (1981), 743. P. Wirtz, K. Woemer, F. Wunder, K. Eichler, G. Roscher, and I.Nicolau, Canadian Patent Application 2,071,699 (1992).

J.W. Hightower, W.N. Delgass, E. Iglesia and A.T. Bell (Eds.) 1 lth International Congress on Catalysis - 40th Anniversary

Studies in Surface Science and Catalysis, Vol. 101 9 1996 Elsevier Science B.V. All rights reserved.

201

F o r m a t i o n of c i t r a c o n i c a n h y d r i d e b y v a p o r - p h a s e d e c a r b o x y c o n d e n s a t i o n of p y r u v i c acid M. Ai a and K. Ohdan b aDepartment of Applied Chemistry and Biotechnology, Niigata Institute of Technology, 1719 Fujihashi, Kashiwazaki 945-11, Japan bUbe Laboratory, UBE Indstrial Ltd., 1978 Kogushi, Ube 755, Japan

Citraconic anhydride (Methyl maleic anhydride) was found to be produced from pyruvic acid by an oxidative decarboxy-condensation. The best catalyst is iron phosphate with a P/Fe atomic ratio of 1.2. The presence of oxygen is required to promote the reaction. The main side-reaction is formation of acetic acid and CO2 by oxidative C-C bond fission. The best results are obtained at a temperature of 200~ The yield of citraconic anhydride reaches 71 mol% at a pyruvic acid conversion of 98%. 1. INTRODUCTION Pyruvic acid is the simplest homologue of the a-keto acid, whose established procedures for synthesis are the dehydrative decarboxylation of tartaric acid and the hydrolysis of acetyl cyanide. On the other hand, vapor-phase contact oxidation of alkyl lactates to corresponding alkyl pyruvates using V205- and MoO3-baseds mixed oxide catalysts has also been known [1-4]. Recently we found that pyruvic acid is obtained directly from a vapor-phase oxidativedehydrogenation of lactic acid over iron phosphate catalysts with a P/Fe atomic ratio of 1.2 at a temperature around 230~ [5]. CH3-CH(OH)-COOH + 0.502 ; CH3-CO-COOH + H20 The one-pass yield of pyruvic acid reached 50 mol%. The main side reactions are the formation of acetaldehyde and CO2 by oxydative C-C bond fission of lactic acid and that of acetic acid and CO2 by oxidative C-C bond fission of the produced pyruvic acid. CH3-CH(OH)-COOH + 0.502 ; CH3CHO + CO2 + H20 CH3-CO-COOH + 0.502 , CH3COOH + CO2 Over the V205- and MoO3-based mixed oxide catalysts, the main part of lactic acid is converted to form acetaldehyde and CO2.

202 In the reaction of lactic acid to form pyruvic acid over the iron phosphate catalysts, formation of a new c o m p o u n d was observed. As the extent of reaction increased, the amount of pyruvic acid increased to a m a x i m u m and then decreased, while that of the new compound increased steadily. It was therefore concluded that the new compound is formed from pyruvic acid in parallel with acetic acid and CO2. According to gas-mass analyses, the molecular weight was determined as 112. However, there are many c o m p o u n d s with molecular weigth of 112. After the NMR analyses and X-ray diffraction analyses for the single crystal, the new compound was determined to be citraconic anhydride, i.e., mono-methyl maleic anhydride. 2 CH3-CO-COOH + 0.502

; CH3-C-C--O \ + CO2 + 2 H 2 0

II .,o

HC-C=O Citraconic anhydride and citraconic acid are used as raw materials of various chemicals, resines, surface-active agents, and dyes. as is maleic anhydride. They have generally been produced from itaconic acid which is produced by fermentation of cane sugar and grape sugar. Therefore, the price is much higher than that of maleic anhydride. Indeed, by analogy with the formation of maleic anhydride from n-butene, there have been attempts to produce citraconic anhydride by vapor-phase contact oxidation of olefinic hydrocarbons with carbon number of more than five, such as isoprene, 2-methyl butene, toluene, and xylene, using V205-based mixed oxide catalysts. However, the yields were clearly lower than that oberved in the oxidation of n-butene. Citraconic anhydride formation from pyruvic acid by oxidative decarboxycondensation has not been k n o w n prior to these studies. Therefore, in this paper, we attempted to get more insight into the new reaction.

2. EXPERIMENTAL An iron phosphate catalyst with a P/Fe atomic ratio of 1.2 used in this study was prepared according to the procedures described in the previous studies [6-8]. On the other hand, a V-P oxide catalyst with a P/V atomic ratio of 1.06 and pumice supported 12-molybdophosphoric acid (H3PMo12040) and its cesium salt (Cs2HPMo12040) catalysts were the same as used in a previous study [9]. Pumice supported WO3-based mixed oxide catalysts were the same as used in a previous study [101. The contact oxidation of pyruvic acid was carried out with a continuous-flow system. The reactor was made of a stainless steel tube, 50 cm long and 1.8 cm i.d., mounted vertically and immersed in a lead bath. Air or a mixture of nitrogen and oxygen was fed in from the top of the reactor and an aqueous solution containing 100 g of pyruvic acid in 1000 ml was introduced into the preheating section of the reactor by means of a syringe pump. The feed rates of pyruvic acid, oxygen, nitrogen, and water vapor are either 10.5, 70, 280, and 480 or 10.5, 13.4,

203 350, and 480 m m o l / h , respectively, unless otherwise indicated The reaction t e m p e r a t u r e was in the rage of 180 to 270~ The extent of reaction was varied by changing the amount of catalyst used from 1 to 36 g, while fixing the feed rates. The effluent gas from the reactor was led successively into four chilled scrubbers to recover the water soluble compounds. The recovered solution (about 50 ml) and the effluent gas were analyzed by four GC's.

3. RESULTS 3. 1 Performance of various oxide catalysts Since no information has been reported on the reaction, it seems necessary to make a character sketch of the catalytic function. Therefore, various kinds of metal oxides were tested as the catalysts. The results are s u m m a r i z e d in Table 1. It is clear that the V-P oxide and Mo-P heteropoly compound catalysts are not effective for the production of citraconic anhydride. The acidic catalysts such as Table 1 Performance of oxide catalysts for production of citraconic anhydride

Catalyst (Atomic ratio) P/V

(1/1.06)

H3PMo12040 Cs2HPMo12040

Pyruvic acid conversion

Citraconic anhyd. yield (mol%)

(g)

(%)

3 5 10 20 20

23 72 16 77 31

5 5 0 4 6

38 54 77 89 52

0 4 6 0 0

P/Si AI/Si

(1/9) (15/85)

K/Si

(5/95)

20 5 10 20 1

P/Ni P/Fe

(1/1) (1.2/1)

5 10

30 86

2 40

W P/W Mo/W Ti/W Ti / W Sn/W K/W

(5/95) (5/95) (10/90) (70 / 30) (10/90) (10/90)

20 20 20 20 20 20 20

92 25 83 93 59 92 93

61 20 57 60 36 61 52

Reaction temperature = 230~ = 10.5, 350, and 480 mmol / h.

Feed rates of pyruvic acid, oxygen, air, and water

204 Si-P and Si-A1 are not effective, either. The basic catalyst such as Si-K is very active for decomposition of pyruvic acid, but citraconic a n h y d r i d e is not produced. The best results are obtained with W and W-based catalyst. The onepass yield of citraconic anhydride reaches about 60 mol% at a pyruvic acid conversion of about 92%. The combination of P to W decreases the activity markedly. The combination of Mo or K decreases the selectivity to citraconic anhydride. The combination of Ti or Sn is scarcely effective, w h e n the a m o u n t is less than 10 atomic %. When the a m o u n t of Sn or Ti is high, the selectivity falls. It should also be noted that the next best results are obtained with the iron phosphate catalyst u n d e r the reaction conditions used. 3.2. Performace of WO3-based oxide catalysts Since the best results were obtained with the W and W-based oxide catalysts, the reaction was studied in more detail using 20 g portions of these catalysts. The reaction was performed at 230~ with feed rates of pyruvic acid, air, and water = 10.5, 350, and 480 m m o l / h . The contact time defined as volume of catalyst (ml)/rate of gaseous feed (ml/s) was about 5.2 s. The main products were citraconic anhydride and CO2. The amount of acetic acid was very small. No other products were detected except for very small amounts of CO, acetone, and acetaldehyde. A relatively large discrepancy was observed between the amount of consumed pyruvic acid and that of the sum of produced citraconic a n h y d r i d e and acetic acid. This discrepancy was defined as "loss". The product distributions are shown in Figure 1 as a function of the time-onstream. At the begining of reaction, that is in the first 1 h on stream, the conversion of pyruvic acid reached 92% and the yields of citraconic anhydride and acetic acid were 61 and 5 mol%, respectively. The amount of loss was about 26 mol%. It should be noted that the catalytic activity falls markedly with an increase in the time-on-stream. Similar falls in catalytic activity were also observed in the cases of the other W-based mixed oxides, such as W-Sn, W-Ti, and W-Mo. It was also found that the deactivated catalysts can easily be regenerated by a heattreatment at 500~ in air for I h.

3. 3. Performance of iron phosphate catalyst The reaction was studied using the iron phosphate catalyst at 230~ with feed rates of pyruvic acid, air, and water = 10.5, 350, and 480 m m o l / h . The main products were citraconic anhydride, acetic acid, and CO2. W h e n the a m o u n t of catalyst used was 10g, that is, when the contact time is about 2.6 s, the conversion of pyruvic acid reached 95% and the yields of citraconic anhydride and acetic acid were 50 and 28 mol%, respectively; the loss was about 17 mol%. The selectivity to citraconic anhydride is clearly lower and that to acetic acid is higher than in the case of the W-based oxide catalysts. However, the catalytic activity was very stable. No clear change in the yield of citraconic anhydride was observed during the reaction for 10 h. The reaction was then performed using different amounts of catalyst from I to 20g. The yields of citraconic anhydride and acetic acid and the loss are plotted as a

205 100 Catalyst = 20g T = 230~ 80

Conversion

Citraconic Anhyd.

30

Citraconic

0

anhyd.

- T = 230oc

40

x,

60

50

/

\

~ 4O

20

1

AcOH /

~

>, Loss

20

T

0

0

AcOH

1

10! NA.~ A |

-

2

3

7.-!

4

Time-on-stream / h Figure 1. Performance of WO3 catalyst

I I 0 -! 40 60 80 100 Conversion of pyruvic acid / %

Figure 2. Performance of iron phosohate catalyst

function of the conversion of pyruvic acid in Figure 2. The slopes of lines from the origin indicate the selectivities to each product. The selectivities to citraconic anhydride, acetic acid, and loss are about 51, 29, and 20 mol%, respectively. It should also be noted that the selectivities remain unchanged with a large variation in the extent of reaction. This indicates that the citraconic a n h y d r i d e and acetic acid produced are stable enough under the reaction conditions used. 3. 4. Effects of oxygen concentration on the reaction o v e r i r o n phosphate catalyst The reaction was performed over the iron phosphate catalyst by changing the feed rate of oxygen from zero to 350 m m o l / h , while fixing the sum of feed rates of oxygen and nitrogen at 350 m m o l / h . The feed rate of pyruvic acid was fixed at 10.5 m m o l / h . The yields of citraconic anhydride obtained at a temperature of 230~ and a short contact time of 0.52 s (amount of catalyst used = 2 g) are plotted as a function of the feed rate of oxygen in Figure 3. The formation of citraconic a n h y d r i d e increases with an increase in the feed rate of oxygen up to about 70 m m o l / h (air). However, with a further increase in oxygen feed rate, the formation of citraconic anhydride levels off. It is clear that the presence of oxygen is required to form citraconic anhydride from pyruvic acid. Another series of tests were performed in a low oxygen concentration: the feed rates of pyruvic anhydride, oxygen, nitrogen, and water were 10.5, 13.4, 350, 480 m m o l / h , respectively. The extent of reaction was varied by changing the a m o u n t

206

~9,

"~ 9

20

50

02 feed (mmol/h) 13. O , 70. o 9

40

15

o~.

~

../ ~

O// . /

t"

//~itraconic Anhyd.

31-

_

6 I"

-

o

-~

t( ~

0 fj

I~

20 40 60 o.,,~ oxygen feed / m m o l / h

Figure 3. Effect of oxygen on the rate

o

Aco8 '

'

-

,

40 60 80 100 Conversion of pyruvic acid / % Figure 4. Effect of oxygen on the selectivity

of catalyst used. The yields of citraconic anhydride and acetic acid are compared with those obtained with a oxygen feed rate of 70 m m o l / h by using air as the oxygen source (Figure 4). It is clear that the selectivity to citraconic anhydride increases and that to acetic acid decreases with a decrease in the oxygen concentration. 3. 5. Effects of reaction temperature on the reaction over iron phosphate catalyst The reaction was performed over the iron phosphate catalyst by changing the reaction temperature from 180 to 270~ and the amount of catalyst used from 1 to 36 g, while fixing the feed rates of pyruvic anhydride, oxygen, nitrogen, and water at 10.5, 13.4, 350, 480 m m o l / h , respectively. The yields of citraconic anhydride obtained at different temperatures are plotted as a function of the conversion of pyruvic acid in Figure 5. And the yields of acetic acid at different temperatures are also plotted in Figure 6. The selectivity to citraconic anhydride decreases and that to acetic acid increases as the temperature is raised. The results indicate that the activation energy for the formation of citraconic a n h y d r i d e is much lower than that for the formation of acetic acid. The selectivity to acetic acid decreases steadily with a lowering of the temperature. However, the highest selectivity to citraconic anhydride is obtained at 200~ Possibly vaporization of pyruvic acid may become difficult at temperatures below 200~ The yield of citraconic anhydride reached 71 mol% and that of acetic acid was 7 mol% at the pyruvic acid conversion of 98%; the loss was about 20 mol%.

207

!

9o*c

70

50F

-

- 70

/& 40

-

"0 20 "~ 10

o

20

~-/ , ,

10

O[~

~ ~ ! I 20 40 60 80 100 Conversion of pyruvic acid / %

Figure 5. Effect of temperature on the yield of citraconic anhydride

20

0 40 60 80 100 Conversion of pyruvic acid / % Figure 6. Effect of temperature on the yield of acetic acid

4. DISCUSSION Since formation of citraconic anhydride from pyruvic acid is one of "acid to acid type" transformations, such as reactions from isobutyric acid to methacrylic acid and from lactic acid to pyruvic acid, the required catalysts must be acidic [11]. If the catalysts are basic, it may be impossible to obtained acidic products, because basic catalysts activate selectively acidic molecules and, as a result, they show a very high activity for the decomposition of acidic products [11]. As mentioned before, two reactions of pyruvic acid take place in parallel as follows: 2 CH3-CO-COOH + 0.502 2 CH3-CO-COOH + 02

-

~ citraconic anhydride + CO2 + 2 H 2 0 ~ 2 acetic acid + 2 CO2

WO3-based oxide catalysts show a high selectivity to form citraconic anhydride even at a relatively high temperature of 230~ This means that their activity for the oxidative C-C bond fission is completely suppressed. However, they lose quickly the catalytic activity. Possibly, they cannot generate enough of the oxygen species required for the reaction, because WO3 possesses redox function in a very poor extent unlike MOO3, V205, and iron phosphate. As seen in Figure 3, the presence of oxygen is required to promote the formation of citraconic anhydride.

208 Similarly, acidic oxides without redox function, such as Si-A1, Si-P, and Ni-P oxides, are not effective as catalysts for this reaction. MOO3- and V205- based oxides possess both acidic and redox functions. Therefore, they show a good performance as catalysts in many partial oxidations for producing especially acidic compounds. However, they possess double bond oxygen species, that is, M=O species, which promote the oxygen insertion reactions as well as the dehydrogenation. Therefore, in the reaction of lactic acid to pyruvic acid, they promote preferentially the C-C bond fission by oxygen insertion to form acetaldehyde and CO2 [5]. Similarly, in the reaction of pyruvic acid, they promote preferentially the formation of acetic acid and CO2 by the C-C bond fission rather than the condensation to form citraconic anhydride. On the other hand, iron phosphate possesses both acidic and redox functions, though the functions are considerd to be much lower than those of MoO3 and V205. However, it possesses no double bond oxygen species unlike MoO3 and V205. Therefore, its function to promote oxygen insertion processes is very weak. As a result, in the reaction of pyruvic acid, the formation of acetic acid and CO2 by the C-C bond fission is suppressed, to a certain extent, and a relatively good performance is obtained for production of citraconic anhydride. The results shown in Figure 4 indicate that the oxygen dependency of the acetic acid formation is higher than that of the citraconic anhydride formation. Therefore, use of a low oxygen concentration is beneficial to the selectivity to citraconic anhydride, though it is disadvantageous to the reaction rate. It should also be noted that the iron phosphate catalyst is inactive for consecutive decomposition of the produced citraconic anhydride and acetic acid. As may be seen in Figures 5 and 6, when iron phosphate is used as the catalyst, the selectivity is dependent largely on the reaction temperature. Therefore, the activation energy for the citraconic anhydride formation is considered to be much lower than that for the acetic acid formation. Indeed, the selectivity to acetic acid decreases steadily with a decrease in the temperature. However, the selectivity to citraconic anhydride shows a maximum at about 200~ Possibly, the vaporization of pyruvic acid may become difficult at temperatures below 200~ As for the reaction path from pyruvic acid to citraconic anhydride, it is considered that a condensation reaction first takes place by a reaction between an oxygen atom of carbonyl group and two hydrogn atoms of methyl group in another molecule, followed by oxidative decarboxylation to form citraconic acid. The produced citraconic acid is dehydrated under the reaction conditions used. The proposed reaction path is shown in Figure 7.

209

H3C,"

H3C

#O

o..C-C.oH

>

I,-I .,(3 H-C-ClJ C"OH FI 0

H3C\c- C~""O H + II C:_.OH

H,C-

"~0

zO

"C- C~.IOH

H"C-C..O il C"OH

-I- H20

0.5 02

-

>

0

COz

H3C,, C - C ~ 0 II /"~ /O +

H~C- C % 0

H20

Figure 7. Estimated reaction path to citraconic anhydride

REFERENCES

1. T. Yokoyama and K. Matsuoka, Jpn. Patent No. 56-19 854 (1981). 2. Y. Yamaguchi, Jpn. Patent No. 57-24 336 (1982). 3. S. Sugiyama, N. Shigemoto, N. Masaoka, S. Suetoh, H. Kawami, H. Kawami, K. Miyaura, K. Miyaura, and H. Hayashi, Bull. Chem. Soc. Jpn., 66 (1993) 1542. 4. H. Hayashi, N. Shigemoto, S. Sugiyama, N. Masaoka, and K. Saitoh, Catal. Lett., 67 (1994) 273. 5. M. Ai and K. Ohdan, Chem. Lett., (1995) 405. 6. M. Ai, E. Muneyama, A. Kunishige, and K. Ohdan, Bull. Chem. Soc. Jpn., 67 (1994) 551. 7. E. Muneyama, A. Kunishige, K. Ohdan, and M. Ai, J. Mol. Catal., 89 (1994) 371. 8. E. Muneyama, A. Kunishige, K. Ohdan, and M. Ai, Appl. Catal., 116 (1994) 165. 9. M. Ai, J. Catal., 89 (1984) 413. 10. M. Ai, in Proc. 10th International Congress on Catalysis, Budapest 1992, eds.. L. Guci, F. Solymosi, and P. T6t6nyi, Akademiai Kiadpo, Budapest, 1993, p. 1199. 11. M. Ai, in Proc. 7th International Congress on Catalysis, Tokyo 1980, eds., T. Seiyama and K. Tanabe, Elsevier, Amsterdam, 1981, p. 1060.

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J.W. Hightower,W.N. Delgass,E. Iglesia and A.T. Bell (Eds.) 11th International Congress on Catalysis - 40th Anniversary

Studies in Surface Science and Catalysis, Vol. I01 9 1996ElsevierScienceB.V. All rights reserved.

211

Heterogeneous enantioselective dehydration of butan-2-ol Saskia Feast, Donald Bethell, Philip C. Bulman Page, M.Rafiq H.Siddiqui, David J. Willock, Graham J. Hutchings, Frank King ~, and Colin H. Rochester b Leverhulme Centre for Innovative Catalysis, Robert Robinson Laboratories, Department of Chemistry Liverpool University, P.O. Box 147, Liverpool, L69 3BX, U.K. "ICI Katalco, Research and Technology Group, P.O.Box 1, Billingham, Cleveland, TS23 1LB, U.K. b Department of Chemistry, University of Dundee, Dundee, DD 1 4HN, U.K.

ABSTRACT: Zeolite Y modified with chiral sulfoxides has been found catalytically to dehydrate racemic butan-2-ol enantioselectively depending on the chiral modifier used. Zeolite Y modified with R-1,3-dithiane-1-oxide shows a higher selectivity towards conversion of S-butan-2-ol and the zeolite modified with S-2-phenyl-3-dithiane-l-oxide reacts preferentially with Rbutan-2-ol. Zeolite Y modified with dithiane oxide demonstrates a significantly higher catalytic activity when compared to the unmodified zeolite. Computational simulations are described and a model for the catalytic site is discussed.

Introduction:

The activity of molecules used as pharmaceuticals, agrochemicals and food additives can depend strongly upon stereochemistry. Accordingly the synthesis of pure enantiomers is becoming increasingly important in the development of new products. Although enantioselective catalysis in homogeneous systems is fairly well established ~2. The problems associated with these systems are the drawbacks normally associated with homogeneous systems i.e. product recovery and catalyst separation. Recently some progress has been made using heterogenised homogeneous catalysis by immobilising homogenous systems 3. Some success has also been reported by the modification of zeolites as catalyst for heterogeneous enantioselective

212 catalysis in the liquid phase *~. In this presentation we describe a method by which zeolite structures can be made chiral by the inclusion within them of a chiral species. Specifically we discuss new catalytic systems consisting of zeolite Y chirally modified by 1,3-dithiane-l-oxide with R=H and R=phenyl ( Fig. 1). With these modified zeolites we have observed the enantioselective dehydration of butan-2-ol in the gas phase and in particular, we demonstrate the use of the combination of experimental studies and computational simulation to address a complex catalytic process.

Experimental: MATERIALS AND REACTION SYSTEM: The method for the synthesis of zeolite Y and the incorporation of the chira] modifier R- 1,3-dithiane- 1oxide Fig. 1 has been reported 7. A ~~ similar method was used to modify s . y s . oSvS zeolite Y with S-phenyl- 1,3-dithianeR 1-oxide. Catalytic experiments were I II carried out in a conventional glass microreactor using on-line GC analysis. The activity of the catalyst (zeolite Y modified with 1,3-dithiane-l-oxide) and the comparison with unmodified zeolite Y, zeolite Y modified with 1,3-dithiane, silica/alumina modified with 1,3- dithiane oxide, and boron nitride modified with 1,3- dithiane-l-oxide was carried out using the following conditions. For a typical catalyic run, catalyst (0.3 g) was reacted with prevaporized butan-2-ol (3.3 x 10 ~ mol h "1) using nitrogen (3.7 x 10 .2 tool h 1) as a diluent. The enantioselectivity reactions were tested using the following conditions. The catalyst (zeolite Y modified with enantiomerically enriched dithiane oxides) 0.1 g, was reacted in a glass microreactor with prevaporized racemic butan-2-ol (7.35 x 10 ~ mol h 1) diluted in nitrogen (6.7 x 10 ~ mol hl). The products were analysed using an on-line GC with a 40m capillary T-cyclodextrin column with trifluoroacetyl stationary phase. The GC was temperature programmed from 25 - 70 ~ with a split ratio of 120:1 to achieve the separation of butan-2-ol enantiomers. TECHNIQUES: X- ray diffTaction (XRD) was carried out using a Hilton Brooks modified Philips 1050W diffractometer with a Cu Ka source (40 keV and 20 mA). Thermogravimetric analysis (TGA) was carried out using a Perkin Elmer TGA 7 analyser. All solid state MAS NMR spectral data were aquired on a Bruker MSL 400 MHz spectrometer. The 13C solution NMR spectra were recorded on a B ~ e r AMX 400 spectrometer and referenced to TMS. B IOSYM molecular modelling was carried out on a Silicon Graphics Indigo 2 work station, using Monte Carlo docking and energy minimisation Figure 1

§

/

213 progromme Discover to visualise the interaction of the sulfoxides within the zeolite fromework.

R E S U L T S AND D I S C U S S I O N : Characterization by X-ray diffraction demonstrated that the modifier did not affect the crystallinity of the zeolite, and studies using 1~C MAS NMR spectroscopy showed that the dithiane oxide had been molecularly adsorbed into the zeolite. The dithiane oxide was remarkably stable inside the acidic zeolite: it did not undergo acid catalysed elimination, hydrolysis or rearrangement. Heating the modified zeolite in flowing nitrogen at 180 ~ did not lead to any decomposition, and the dithiane oxide could be recovered intact in high yield by solvent extraction. Thermogravimetric analysis showed that the dithiane oxide modifier either decomposed or desobed between 500-600 ~ Of particular importance is that such t r e a t m e n t of enantiomerically enriched R-dithiane oxide did not result in racemization. Decomposition of compound I (R = H) adsorbed on the zeolite was only observed at temperatures well in excess of 400 ~ The initial experiments were carried out using zeolite Y modified with racemic 1,3-dithiane-l-oxide, and this was compared as a catalyst for the dehydration of racemic butan-2-ol with a control sample of zeolite Y prepared by an analogous method but without the modifier. The results (Table 1) demonstrate that the dithiane oxide modified catalyst was considerably more active than the control sample by several orders of magnitude, and t h a t this activity was maintained over several days of testing without significant loss of the enhanced activity. Under the reaction conditions studied, the control somple only become active at temperatures above 150 ~ and required a reaction temperature of 225 ~ to achieve 90% conversion, whereas the modified zeolite gave 90% conversion at 115 ~ For the modified zeolite, the distribution of butene isomers is close to that expected for equilibrium; 8 subsequent experiments revealed that the modified zeolite was also an active catalyst for butene isomerization. To confirm that the rate enhancement observed was due to an interaction between the modified zeolite and the substrate, a n u m b e r of control experiments were carried out (Table 1). Zeolite Y was modified using 1,3-dithiane II, and, under the same conditions used for the sulfoxide-modified zeolite, no rate enhancement was observed, indicating t h a t the sulfoxide oxygen atom is a vital feature of this catalyst system. Investigation of a non microporous silica/alumina catalyst t h a t had a Si/A] ratio identical to that of the zeolite Y used in this study indicated t h a t modification by the dithiane oxide acted as a poison. The microporous aluminosilicate framework is therefore also important. In addition, 1,3-dithiane-l-oxide supported on the

214 i n e r t m a t e r i a l boron nitride was found to be totally inactive for butan-2-ol dehydration, i n d i c a t i n g t h a t it is the combination of the d i t h i a n e oxide w i t h the zeolite w h i c h is essential.

Table 1 Reaction of racemic butan-2-ol over modified catalysts Catalyst

Y"

TempfC Conversion/% Selectivity/% but- 1-ene E-but-2-ene Z:but:2:ene

115 0 -

Y-SO b

Y-S ~

SiO/

AI,0~'

BN-SO f

AI~.O.~ ~

SiOJ

225 90

115 90

110 3.6

175 35.6

200 31.5

200 24.5

115/225 0

17.7 40.6

8.3 53.8

15.4 53.8

14.7 46.5

16.5 31.7

21.6 33.9

-

................... : ........... 4 1 . 6

...... 3 7 . 8

........... . 3 . 0 . 7 ......... 3 8 . 8

.......... 5 1 = 8 .............. 4 6 . . 5 ......................... =..............

"Zeolite Y ( u l t r a s t a b i l i s e d LZY 82, Union Carbide) b Zeolite Y modified w i t h (+/-)-l,3-dithiane-l-oxide, 1 molecule per s u p e r cage (7.6 wt%) ~Zeolite Y modified w i t h 1,3-dithiane, 1 molecule per supercage (7.2 wt%) d Non microporous silica alumina, SiO]A1203 =5.7 ' Non microporous silica alumina, SiO2/A1203 =5.7, modified w i t h 1,3-dithiane1-oxide (6.9 wt%). f Boron nitride modified with 1,3-dithiane-l-oxide (7.4 wt%). Boron nitride alone also showed no activity. In a f u r t h e r set of e x p e r i m e n t s racemic butan-2-ol w a s reacted w i t h zeolite Y modified w i t h chiral sulfoxides. In the first set of e x p e r i m e n t s racemic butan2-ol was r e a c t e d over zeolite Y modified by R-1,3-dithiane-l-oxide (83% ee) a n d the results s h o w n in Table 2 indicate t h a t S- butan-2-ol reacts preferentially. In a second set of experiments, the s~me reaction was i n v e s t i g a t e d for zeolite Y modified w i t h S-2-phenyl-l,3-dithiane-l-oxide (99% ee) 9 . The results, shown in Table 2, confirm t h a t the s~me effect is observed, however, in this case the R- butan-2-ol reacts in preference to the S- enantiomer. To u n d e r s t a n d the origin of this high catalytic activity and the enantioselectivity a series of computer simulation studies was carried out to d e t e r m i n e the m o s t favourable locations for 1,3-dithiane-l-oxide (R=H and

215 R=phenyl) w i t h i n th e zeolite framework. The Biosym docking p a c k a g e x~w a s used to g e n e r a t e a r a n d o m s et of ten energetically favourable configurations

Table 2 a Reactivity of r a c e m i c butan-2-ol over zeolite Y modified by homochiral d i t h i a n e oxides. Modifierb

Temp.

Convc.

Butan-2-ol conv./x 10~ tool h "1

~ % {R} R=H 110 0.5 {R} R=H 120 1.3 {R} R=H 150 9.9 {S} R=phenyl 110 4.2 {S} R=phenyl 120 7.5 note : {R} and {S} refer to the chiral centre at the

R S 0.002 0.035 0.006 0.094 0.018 0.707 0.276 0.015 0.415 0.081 sulfoxide sulfur.

Relative rate" 17.5 15.7 39.3 18.4 5.1

a Reaction conditions: 0.1 g of the zeolite Y modified catalyst, tested in a conventional glass microreactor with racemic butan-2-ol (7.35 x 10 .3 tool h-l), prevaporized in a nitrogen diluent (6.2 -6.7 x 10 .3 tool h-l). Products were analyzed u si n g on-line GC w i t h a 40m capillary y- cyclodextrin column with trifluoroacetyl s t a t i o n a r y phase, t e m p e r a t u r e p r o g r a m m e d from 25-70 ~ with a split ratio of 120:1. bZeolite Y modified w i t h R- 1,3-dithiane- 1-oxide/S-2-phenyl- 1,3-dithianeoxide ~ conversion of R- an d S-butan-2-ol d Relative ratio of p r e f e r e n t i a l enantiomeric reaction of R- and S-butan-2-ol for each of the d i t h i a n e oxide molecules in a host zeolite Y s t r u c t u r e w i t h a Si:A1 ratio of 1:1. E a c h s t r u c t u r e was energy- minimized a n d the lowest energy conformation t a k e n as the most likely d i t h i a n e position in the zeolite. Within this s t r u c t u r e , the f r a m e w o r k proton n e a r e s t to the sulfoxide oxygen was t r a n s f e r r e d to form the hydroxy sulfonium cation in the case of R=H and the R - h y d r o x y - l , 3 - d i t h i a n e for R=phenyl; the cation position w a s t h e n reoptimized (Figures 2 and 3). D u r i n g this process the atomic charges for the adsorbed molecules were t a k e n from a Mulliken a n a l y s i s of a DZP basis set H a r t r e e Fock calculation (using CADPAC Xlon the MOPAC ( P M 3 ) o p t i m i z e d molecular structure. Table 3 gives the cation c h a r g es for the sulfoxide group and C2 for each cation. In both cases the sulfur and t r a n s f e r r e d proton are positively charged but at sulfur the

216 addition of the phenyl group at the two position has reduced the charge by almost 0. le.

Fig. 2 Calculated low energy conformation of the protonated dithiane oxide cation (R=H) in zeolite Y (Si/Al = 1). The bottom view shows a view through the twelve ring containing the deprotonated framework oxygen, the top view is perpendicular to this. For clarity the zeolite framework is shown using a stick model and the adsorbed molecule is drawn in space filled form represented by the Van der Waals radii for the atoms being in the order S>O>C>H.

When comparing the relaxed structures of figures 2 and 3, it should be noted that the phenyl group has two effects: it greatly increases the size of the dithiane molecule, leading to greater steric hindrance, and it reduces the effective charge at the sulphoxide sulfur, lowering the electrostatic interaction between the sulfoxide group and the deprotonated framework oxygen. These effects are reflected in the relaxed geometries; in both cases the two closest molecule/framework contacts occur between the S and H atoms of the sulfoxide cation and the deprotonated framework oxygen, the S..O distances being 3.01/k in the phenyl case compared to 2.84A for R=H. For the 1,3-dithiane-1-oxide (R=H) case molecular dynamics simulations at the experimental temperature revealed that the R-hydroxysulfomum cation was considerably more stable than the more weakly adsorbed 1,3-dithiane molecule TM. We consider that the hydroxydithiane cation may act as a proton transfer agent and this may account for the enhanced reactivity of this system.

217

Fig. 3 Calculated low energy conformation of the protonated dithiane oxide cation (R=phenyl) in zeolite Y (Si/A1 = 1). The views and presentational details are as for figure 2.

Table 3 Cation charges for the sulfoxide group and C2 for each cation.

C2 (inter sulfur C) $3 (S of SOH) 0 (O of SOH) H (H of SOH)

R=H charge (le l)

R=phenyl charge (le [)

-0.504 +0.825 -0.611 +0.393

-0.513 +0.727 -0.624 +0.404

The structure shown in Figure 2 was used as the host for f u r t h e r docking calculations to introduce each of the enantiomers of butan-2-ol; the docked structures were energy minimized and the lowest energy structures selected as the most likely butan-2-ol binding positions; these are shown in Figure 4. These structures involve the framework and the R-hydroxysulfonium cation acting in concert to bind the butan-2-ol molecules, and calculations indicate that the S-enantiomer is bound more tightly t h a n the R form ( 16.4 k J mo1-1) and we consider that this may play some role in the observed enantioselective reactions with these catalyst systems.

218 F i g u r e 4 Calculated low energy conformations for enantiomers of butan-2-ol in dithiane oxide (R=H) loaded zeolite Y (Si/A1 = 1). Each of the two butan-2-ol enantiomers was docked into the zeolite Y/dithiane oxide cation system (see Figure 2), and the lowest energy structures are shown here. Binding energies for butan-2-ol: S form, -64.7 kJmol-1; R form,-48.3 kJmo1-1. Two views of each structure are shown, with most of the zeolite framework cut away for clarity; the bridging oxygen atom which donated the proton is highlighted as a small sphere. Throughout this work the zeolite potentials and atomic charges were taken from the consistent force field library provided by Biosym Technologies 13'14. The potentials for the guest molecules were also taken from this library but with charges assigned from an SCF calculation on the isolated molecule performed with a DZP basis set within the CADPAC package TM Taken together, these results provide the first example of a gas phase enantioselective reaction heterogeneously catalysed by a zeolite. Although the initial approach has been to design a catalyst that consumes chiral molecules, we have demonstrated that a catalyst can be designed that is capable of discriminating between enantiomers and preferentially transforming one of them. This important effect is achieved by enantioselective rate enhancement, i.e. both enantiomers react faster in the chiral environment than in the absence of the chiral modifier, but one reacts faster than the other. In view of the vast range of microporous materials and potential enantiomerically pure modifiers available, together with the recent advances in theoretical methods, we believe that this approach will provide the basis for a general advance in the design of ultraselective chiral catalysts. We are indebted to the SERC Catalysis and Interfaces Imtiative and ICI Katalco for financial support. Computational results were obtained using software programs from Biosym Technologies of San Diego. References 1 H.B. Kagan in Asymmetric Synthesis, ed. J. D. Morrison, Academic Press, Orlando vol. 5 (1985)

219 2 S. Akutagawa, Appl. Catal. A., 128 (1995) 171 3 K.T. Wan and M. E. Davis, Nature, 370 (1994) 449 4 W. Reschetilowski, U. Bohmer, and J. Wiehl, Stud. Surf. Sci. Catal.,84 (1994) 2021 5 R. Mahrwald, U. Lohse, I. Girnus, and J. Caro, Zeolites, 14 (1994) 486 6 A. Corma, M. Iglesias, C. del Pino, and F. Sanchez, Stud. Surf. Sci. Catal., 75C (1993) 2293 7 S. Feast, D. Bethell, P. C. B. Page, F. King, C. H. Rochester, M. R. H. Siddiqui, D. J. Willock, and G. J. Hutchings, J. Chem. Soc., Chem. Comm., (1995) in press 8 A.C. Butler and C. P. Nicolaides, Catal. Today, 18 (1993) 443 9 P.C.B. Page, M. T. Gareh, and R. A. Porter, Tetrahedron: Asymmetry, (1993) 2139. 10 Catalysis Users Guide, release 236, Biosym Technologies, San Diego (1994). 11 CADPAC5: The Cambridge Analytic Derivatives Package Issue 5, A suite of quantum che~mistry programs with contributions from IL Alberts, JS Andrews, SM Conwell, NC Handy, D Jayatilaka, PJ Knowles, R Kobayashi, N Koga, KE Laidig, PE Maslen, CW Murray, JE Rice, J Sanz, ED Simandiras, AJ Stone and M-D Su, (1992). 12 D. Willock, S. Feast, P. C. B. Page, D. BetheU, and G. J. Hutchings, Topics in Catalysis, (1995) in press. 13 J. R. Hill and J. Sauer, J. Phys. Chem., 98 (1994) 1238 14 J. R. Maple, T. S. Thatcher, U. Dinur, and A. T. Hagler, Chemical Design Automation News, 9 (1990) 10 15 R. D. Amos and J. E. Rice, CADPAC: The Cambridge Analytic Derivatives Package 5, issue 4.0, Cambridge, (1987)

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J.W. Hightower, W.N. Delgass, E. Iglesia and A.T. Bell (Eds.) 11th International Congress on Catalysis - 40th Anniversary

Studies in Surface Science and Catalysis, Vol. 101 9 1996 Elsevier Science B.V. All fights reserved.

221

Enantioselective hydrogenation catalysed by palladium T.J. Hall, P. Johnston, W.A.H. Vermeer, S.R. Watson and P.B. Wells School of Chemistry, University of Hull, Hull, HU6 7RX, United Kingdom The enantioselective hydrogenations of five unsaturated carboxylic acids (including tiglic and angelic acids) catalysed by 1% Pd/SiO2 and of methyl pyruvate catalysed by 4% Pd/Fe203 are described. Enantioselectivity was induced by adsorption of the alkaloids cinchonidine and cinchonine onto the catalysts. Reactions occurred at room temperature and 10 bar pressure. In all reactions, modification with cinchonidine gave S-product in excess, whereas modification with cinchonine gave R-product in excess. Best values of the enantiomeric excess were 27% in acid hydrogenation and 15% in ketoester hydrogenation. The unsaturated acids undergo selective enantioface adsorption adjacent to the modifier; the process has been modelled and the sense of the observed enantioselectivity interpreted. A D-tracer experiment has shown that the ketoester undergoes dissociative adsorption. The reaction kinetics and the deuterium distribution in the product are consistent with the major route to product being conversion to adsorbed-enol and subsequent C=C hydrogenation.

1. INTRODUCTION The enantioselective hydrogenation of the methyl and ethyl esters of pyruvic acid catalysed by cinchona-modified Pt and Ir has been the subject of detailed study and the mechanism of chiral induction is understood [ 1-6]. In these reactions, modification of the catalyst by the adsorption of cinchonidine gives R-lactate in excess in the product, whereas cinchonine gives S-lactate in excess. High values of the enantiomeric excess, above 80~ are achievable. In 1988, Blaser and co-workers reported reversed enantioselectivity over a Pd/C catalyst, i.e. adsorption of cinchonidine induced an enantiomeric excess of 4% in favour of the S-enantiomer [7]. This remarkable result is here confirmed using Pd/Fe203 as catalyst and it is shown that the Pdcatalysed reaction differs in every significant particular from the Pt-catalysed reaction. It became evident during this study that enolisation of the reactant was occurring and that the reaction was one of C=C hydrogenation. Accordingly, we examined for comparison the hydrogenation of several esters of unsaturated carboxylic acids (e.g. methyl tiglate, methyl angelate) but observed no enantioselectivity. However, following Nitta and co-workers' report of the enantioselective hydrogenation of E-a-phenylcinnamic acid catalysed by cinchonidine-modified Pd [8], the free acids were examined and enantioselectivity observed. This paper accordingly reports several examples of the enantioselective hydrogenation of the carbon-carbon double bond.

222 2. EXPERIMENTAL

2.1 Materials 1% Pd/SiO2 was prepared at ICI Katalco and 4% Pd/Fe203 at Johnson Matthey. Reactants used are shown in Table 1. Trifluorotiglic acid and its ester were synthesised by D. Hunter at Glasgow University. The other unsaturated carboxylic acids (Aldrich), methyl pyruvate (Fluka), but-3-en-2-one (Aldrich), cinchonidine (Aldrich), cinchonine (Hopkin and Williams), ethanol (AnalaR grade, BDH) and THF (non stabilised, Fisons) were used as received. Table I Reactants Ri/_

COOH

R2x ..--c ~/ R 3

O/~

H3C

acids (see below)

/OCH3

--c

%0

methyl pyruvate

H2C~-~-CH

H3c / but-3-en-2-one

R1

R2

R3

Name

Trivial description

H H Me Et Et

Me CF3 H H H

Me Me Me Me Et

E-2-methyl-but-2-enoic acid E-2-methyl-4,4,4-trifluoro-but-2-enoic acid Z-2-methyl-but-2-enoic acid Z-2-methyl-pent-2-enoic acid Z-2-ethyl-hex-2-enoic acid

Tiglic acid Trifluorotiglic acid Angelic acid 2MP2 acid 2EH2 acid

2.2 Procedure: hydrogenation of unsaturated carboxylic acids and their methyl esters 0.3 g samples of Pd/SiO2 were evacuated and reduced in 1 bar hydrogen at room temperature for 0.5 h. The catalyst was then wetted by injection into the reduction vessel of a 10 ml aliquot of alkaloid solution (0.05 g in 50 ml THF). The thoroughly wetted catalyst was transferred to the glass liner of a Baskerville steel autoclave and the remaining 40 ml alkaloid solution added. The liner was positioned in the autoclave, 5 mmol reactant added, and the autoclave sealed and purged with pure nitrogen. Immediately thereafter hydrogen was introduced to the desired pressure and stirring at 1200 rpm commenced. Hydrogen pressure was maintained at the set value by a computer-controlled admission system and a data system logged the progress of reaction. Reactions were stopped after 20 h reaction time.

2.3 Procedure: hydrogenation of methyl pyruvate 0.1 g samples of Pd/Fe203 were evacuated and reduced for 1 h in 1 bar hydrogen at the required temperature (normally 293 K). After brief evacuation of the hydrogen the catalyst was wetted by injection into the reduction vessel of a 10 ml aliquot of alkaloid solution (0.2 o in 40 ml ethanol). The thoroughly wetted catalyst was transferred into an open glass beaker, the remaining 30 ml alkaloid solution was added and the slurry was stirred in air for 1 h. The catalyst was then separated by centrifugation and decantation, transferred to a Fischer Porter glass high-pressure reactor, and 20 ml fresh solvent and 113 mmol pyruvate added. The

223 reactor was then sealed, purged, hydrogen introduced, and stirring commenced A computercontrolled hydrogen admission system maintained the set pressure and recorded the course of reaction as a function of time. When unmodified catalysts were examined the procedures were as described above except that the alkaloid was omitted.

2.4 Analysis When the desired hydrogen uptake had been achieved, the vessel was opened, catalyst separated by filtration, and the reaction solution analysed by chiral gas chromatography (column: Cydex B, 50 m, SGE Ltd). Analysis gave conversion and enantiomeric excess. Enantiomeric excess is defined a s [ R - S [/(R+S).

3. RESULTS

3.1 Hydrogenation of unsaturated acids and their methyl esters The hydrogenation of the esters methyl tiglate, methyl trifluorotiglate, and methyl angelate occurred slowly over 1% Pd/SiO2 at room temperature and l0 bar pressure. Modification of the catalyst by cinchonidine or cinchonine induced no enantioselectivity under any conditions Typical results are given in Table 2. Hydrogenation of the free acids over unmodified catalyst occurred slowly, proceeded to completion in 20 h and gave racemic product as expected Enantioselective hydrogenation occurred at a slower rate over alkaloid-modified catalyst, cinchonidine modification providing an excess of S-product and cinchonine an excess of R-product. Racemic and enantioselective hydrogenations of tiglic acid each exhibited an apparent activation energy of 17 kJ tool"1 (268 to 308 K). Enantiomeric excess was constant at 20 to 23% over the range 273 to 308 K but lower, 13%, at 268 K. Enantioselective hydrogenation of trifluorotiglic acid exhibited an activation energy of 23 kJ toolI (253 to 323 K) and a temperature-independent enantiomeric excess of 13 • 2%. No geometrical isomerisation of tiglic acid to angelic acid, or vice versa, accompanied hydrogenation Enantioselective hydrogenation Z-2-methyl-pent-2-enoic and Z-2-ethyl-hex-2-enoic acids occurred over alkaloid-modified PdYSiO2 as described in Table 2. Enantioselectivity was favoured by an increase in hydrogen pressure to 50 bar. The enantiomeric excess of 27% in Z2-methyl-pent-2-enoic acid hydrogenation was the highest value recorded in this study. Enantioselectivity in acid hydrogenation was not sensitive to the reduction temperature of the catalyst Modification by N-benzylcinchonidinium chloride substantially deactivated the catalyst and eliminated enantioselectivity. 3.2 Hydrogenation of Methyl Pyruvate The hydrogenation of methyl pyruvate proceeded over 4% Pd/Fe203 at 293 K and 10 bar when the catalyst was prepared by reduction at room temperature. Racemic product was obtained over unmodified catalyst; modification of the catalyst with a cinchona alkaloid reduced reaction rate and rendered the reaction enantioselective. S-lactate was formed in excess when the modifier was cinchonidine, and R-lactate when the modifier was cinchonine

224 Table 2 Hydrogenation of various unsaturated acids and esters catalysed by 1% Pd/SiO2 and 10 bar pressure and 293 K

Reactant

Modifier a

Initial Conversion rate /% /mmol h'lg "1

Methyl tiglate Ethyl trifluorotiglate Methyl angelate

CD CD CD

10 2 8

35 6 33

Tiglic acid Tiglic acid Tiglic acid Trifluorotiglic acid

none CD CN CD

18 7 5 14

100 65 60 95

0 22 (S) 20 (R) 15 (S)

Angelic acid Angelic acid Angelic acid

none CD CN

29 20 18

100 66 56

0 15 (S) 14 (R)

2MP2 2MP2 2MP2 2MP2

CD CD CN CN

18 (10 33 (50 15 (10 19 (50

61 65 38 60

20 (S) 27 (S) 19 (R) 22 (R)

none CD CN

29 18 15

100 62 51

0 12 (S) 10 (R)

acid acid acid acid

2EH2 acid 2EH2 acid 2EH2 acid

bar) bar) bar) bar)

ee /%

0 0 0

aCD = cinchonidine; CN = cinchonine (Table 3). Enantioselective reaction was of order 0.7 in hydrogen by the initial rate method (over the range 2 to 50 bar, 293 K, cinchonine modifier) and 0.2 in pyruvate (0.1 to 3.0 M, 293 K, 10 bar pressure, cinchonine modifier). Enantiomeric excess was independent of reactant concentrations within these ranges. Reactions exhibited self-poisoning so that complete conversion was not achieved within 20 h reaction time. As the quantity of cinchonine modifier added to the catalyst was increased from zero to 1 gram per gram so the initial reaction rate fell from 180 to 50 mmol h l g 1 and enantiomeric excess rose to 15% (293 K, 10 bar pressure).

225

Table 3 Hydrogenations of methyl pyruvate (A) and ofbut-3-en-2-one 03) over 4% Pd/Fe203 at 293 K Catalyst reduction T/K

Modifer a

293 293 293 293 293 293 673

None CD CN CN CN CN CN

A A A A A A A

10 10 2 10 20 50 10

180 36 57 102 147 193 v.slow

50 25

293

CD

B

10

>9,500

100 b

Reactant

Pressure /bar

Initial rate /mmol h l g -1

Cony. /%

Enantiomeric excess /% 0 5 (S) 12 (R) 13 (R) 13 (R) 14 (R) 0

4

~CD = cinchonidine; CN = cinchonine bProduct = butan-2-one, 100%

Table 4 Deuterium distributions in methyl pyruvate and methyl lactate (X - H or D)

x =

CX3COCOOCH3/% 0 1 2

-dx obs. 6 -dx calc. a 12

39 38

47 38

3

0

8 12

15 3

CX3CX(OX)COOCH3/% 1 2 3 4 13 16

31 31

32 31

8 16

2 3

"binomial distributions for complete exchange at each site. H = 50%, D = 50%

A D-tracer experiment was conducted under standard conditions (293 K, 10 bar pressure) using D2 in place of H2 and C2HsOD as solvent in place of unlabelled ethanol. Dincorporation in reactant and each enantiomer of the product was determined by chiral-gc/ms. Exchange of up to three H-atoms for D occurred in the reactant. The product enantiomers gave identical mass spectra indicating the same D-content and distribution; the product contained 0 to 5 D-atoms. Fragmentation in the mass spectrometer showed that no deuterium was located in the ester group. Thus, the retrieved reactant was CX3COCOOCH3 and the product was CX3CX(OX)COOCH3 (X = H or D). The deuterium distribution in the reactant and products was determined from the mass spectra by application of 13C, 180 and ion fragmentation corrections in the usual way. The isotopic distributions so obtained are shown in Table 4. This fragmentation correction is made on the (usual) assumption that the

226 probability of rupture of H-C, D-C, H-O, and D-O bonds in the mass spectrometer is the same, which is unlikely to be the case. For this reason, and the fact that the parent ion currents of the lightest ions are subject to a greater degree of fragmentation correction, the values quoted for the concentrations of pyruvate-do and lactate-do and -dl are less reliable than those of the more extensively exchanged species. The experimental D-distributions are in modest agreement with those calculated for the cases in which (i) pyruvate has undergone complete exchange at three positions with a pool of adsorbed 'hydrogen' of composition H=50%, D=50%, (ii) lactate has undergone complete exchange at five positions with the same pool of adsorbed 'hydrogen' (Table 4). Pd/Fe203 prepared by reduction at elevated temperature was less active and enantioselective than samples reduced at 293 K (Table 3, entry 7).

4. DISCUSSION

4.1 Hydrogenation of unsaturated carboxylic acids The prerequisite for enantioselective hydrogenation at a metal surface is that selective enantioface adsorption of the prochiral reactant should occur at metal atom sites in the neighbourhood of the adsorbed chiral modifier. The modifiers cinchonidine and cinchonine, in the free state, each exhibit three configurations of comparable minimum energy [1, 9]. When these alkaloids are adsorbed at a Pt surface, one of these minimum energy configurations provides the chiral environment for the selective enantioface adsorption of methyl pyruvate. In that reaction, the rate at the enantioselective sites is enhanced over that at sites elsewhere on the surface that catalyse racemic reaction, and values of enantiomeric excess above 80% are obtainable. The enhanced rate has been attributed to an effect of H-bonding between the OHgroup of the adsorbed half-hydrogenated state and the quinuclidine-N of the adsorbed modifier

[1]. In this Discussion the assumption is made that the adsorption of cinchonidine and cinchonine on Pd is similar to that on Pt [10]. In the hydrogenation of the unsaturated acids and esters over cinchona-modified Pd the key observation is the failure to achieve enantioselectivity in ester hydrogenation and the success encountered in the hydrogenation of the free acids (Table 2). Also, the enantioselective hydrogenation of E-o~-phenylcinnamic acid over cinchonidine-modified Pd has been reported [8, 11 ]. It appears either that the free acid is able to undergo selective enantioface adsorption in the vicinity of the modifier in a manner forbidden to the methyl ester for steric reasons, or that an acid-base interaction between reactant and alkaloid occurs as a precursor state to selective enantioface adsorption which, again, would be unavailable to the ester. Such an acidbase interaction is most easily envisaged as H-bonding between the acid-H atom of the reactant and the quinuclidine-N of the alkaloid. The latter situation has been modelled [12]. Calculations have been carried out in which tiglic acid has been docked with cinchonidine in such a way as to simulate such H-bonding. The docking was carried out for the specific condition in which (i) the alkaloid molecule was in its appropriate minimum energy state, and (ii) the quinoline moiety of the alkaloid and the carbon/oxygen skeleton of the tiglic acid were maintained in the same plane. [This relationship is the closest approximation achievable to that which might obtain if the complex of alkaloid and acid was adsorbed at a plane metal surface.

227 No allowance for the presence of the surface has been made in these calculations.] Two configurations meet these criteria; they are shown schematically in Figure 1. Figure I a shows tiglic acid adsorbed by the enantioface which, on hydrogenation would give S-2-methyl butanoic acid, whereas Figure l b represents the enantioface that would give R-2-methyl butanoic acid. The interconversion of these two states by rotation of the whole tiglic acid molecule about the =--N. . . . H- hydrogen bond was modelled, and showed that the energy of the state represented in Figure l a was lower than that in Figure l b by about 5 kcal mol l moreover, the state represented in Figure l a is located within a broad region of minimum energy. It can thus be inferred that, for the adsorption of tiglic acid adjacent to adsorbed cinchonidine, selective enantioface adsorption occurs such that, on hydrogenation, S-product formation is favoured, in agreement with experiment. H

Hj

0 ~ / C H 3 (S)-2-Methyl Butanoic Acid CH3

(a) H

""'H

H

(R)-2-Methyl Butanoic Acid CH3

(b)

Figure 1. Possible precursor states to the selective enantioface adsorption of tiglic acid

228 The carbon-carbon double bond that undergoes hydrogenation is remote from the modifier and no rate enhancement for the enantioselective process is to be expected. None was observed. Moreover, since the rate at the enantioselective sites is the same as that at other sites on the surface that experience no chiral environment and so give racemic product, the overall enantiomeric excess should be modest, as is the case. To obtain higher enantioselectivity it would be necessary selectively to poison the sites for racemic hydrogenation. The values of enantiomeric excess observed under comparable conditions vary on passing from one acid to another (Table 2) but show no significant trends. The mechanism proposed for the hydrogenation of tiglic acid is applicable to the other acid hydrogenations studied.

4.2 Hydrogenation of methyl pyruvate Although cinchona-modified Pd showed no enantioselectivity in the hydrogenation of the methyl esters of the unsaturated acids, the hydrogenation of methyl pyruvate occurred with a modest enantiomeric excess. The reduction of this ester over Pd d~fered from the corresponding reaction over Pt in every important particular. Enantiomeric excess was low (high over Pt) and in the reverse sense (e.g. cinchonidine modification provided an S-excess in the product over Pd but an Rexcess over Pt). Enantioselective reactions underwent self-poisoning over Pd (proceeded to completion over Pt), were of non-integral order (integral over Pt) and proceeded more slowly than reaction over unmodified catalyst (enhanced rate over Pt). Enantioselective reaction was solvent-specific over Pd (not over Pt) and was favoured by low catalyst reduction temperature (high reduction temperature for Pt). The exchange of H for D in the pyruvate methyl group, which occurred over Pd but not over Pt [13], holds the key to these profound differences in behaviour. This exchange indicates that adsorbed pyruvate (species A, Figure 2) underwent dissociative adsorption at the Pd surface (species B is one formulation of the product of dissociation) and interconversion of species A and B in the presence of a pool of adsorbed-H and -D brought about complete exchange in that methyl group. Occasional formation of the adsorbed enol, species C, is expected as an alternative product of hydrogen-atom addition to species B. Lactate will be formed by hydrogenation of species A or species C, or both, depending upon their relative reactivity. When the C=C and C=O functions were present in the same molecule, but-3-en-2-one, hydrogenation of the former function occurred exclusively and at a rate almost too fast to measure (Table 3). Thus removal of species C is expected to be kinetically fast and that of species A kinetically slow. Strong evidence that the lactate product was formed by enol hydrogenation is provided by the observed order of 0.7 in hydrogen. The rate determining step would necessarily be the H-atom addition to species B to give C; supposing the H-coverage to be described by the appropriate Langmuir equation for dissociative hydrogen adsorption, a fractional order in hydrogen in the region of one-half is expected. It is thus concluded that methyl pyruvate hydrogenation over Pd is a kinetically fast hydrogenation of adsorbed enol formed via the dissociative adsorption of the ot-ketoester. Pd catalysts were active and enantioselective only when reduced at low temperature, suggesting that dissociative adsorption of the reactant was dependent on the presence of surface Pd atoms in a positive oxidation state.

229

O•C

H3C/

HO R

+2H Slow

A

H ~C--

R

H3~

o\ (D)H,--)/C H2c

"/

R B r.

§

HO C

H2

~, ~"

+2H _ Fast -

HO H 9 C

R

H3C/

Figure 2. Mechanism for exchange in methyl pyruvate and for its hydrogenation to methyl lactate via an enol intermediate. R = COOCH3 Enolisation of ketones is favoured in alkanol solution, and the observed solvent specificity in this reaction may indicate that the formation of the enolic species C is favoured when ethanol is used as solvent. No rate enhancement of the enantioselective hydrogenation pathway is expected, in the manner adduced for the Pt-catalysed reaction, because the process is not one of simple H-atom addition across a carbon-oxygen double bond. The self-poisoning character of the reaction was most evident when the modifier was present, from which it may be inferred that adsorbed cinchonidine and cinchonine were convened to an ineffective form and poisoned the surface as reaction progressed. In view of the very high activity of this Pd catalyst for C=C hydrogenation, it may be that the quinoline system of the alkaloids was partially hydrogenated under reaction conditions. Such a proposal would be consistent with our observation that H/D exchange in cinchonidine (over a different Pd catalyst) was accompanied by ring hydrogenation [ 10]. Cinchona alkaloid derivatives having a partially saturated quinoline ring system are poor modifiers of Pt [ 14]. No discussion is offered, concerning the sense of the observed enantioselectivity in this reaction because of a need to be cautious. We have observed that the sense of the enantioselectivity can be reversed simply by a variation of the procedure used for catalyst modification [15], and this has been confirmed by others [16]. Thus it appears that the state of the Pd surface, as well as the nature of the species adsorbed upon it and their spatial relationship, contributes to chiral direction in this reaction.

230 ACKNOWLEDGMENTS

We thank EPSRC, ICI Katalco, Zeneca, and Johnson Matthey for financial support. Acid hydrogenation was conducted as part of the EPSRC/DTI LINK Programme in 'New Catalysts and Catalytic Processes': our colleagues were G. Webb, E. Colvin, E. Allan, D. Hunter and N. Young of the University of Glasgow, S.D. Jackson of ICI Katalco, W. Moss and G. Robinson of Zeneca Pharmaceuticals and S. Korn of Zeneca Fine Chemicals. Pyruvate ester hydrogenation was carried out as part of a programme funded by EPSRC and Johnson Matthey: our colleagues were I.L. Dodgson, A. Fulford, K.G. Griffin and B. Harrison from the Company. We also acknowledge stimulating discussions with A. Ibbotson.

REFERENCES

1.

2. 3. 4. 5. 6. 7. 8. 9. 10. 11. 12. 13. 14. 15. 16.

K.E. Simons, P.A. Meheux, S.P. Griffiths, I.M. Sutherland, P. Johnston, P.B. Wells, A.F. Carley, M.K. Rajumon, M.W. Roberts and A. Ibbotson, Recl. Trav. Chim. PaysBas, 113 (1994) 465. O. Schwalm, B. Minder, J. Weber and A. BaJker, Catal. Letts., 23 (1994) 271. K.E. Simons, A. Ibbotson, P. Johnston, H. Plum and P.B. Wells, J. Catal., 150 (1994) 321. H-U. Blaser, H.P. Jalett, D.M. Monti, A. Balker and J.T .Wehrli, Stud. Surf. Sci. Catal., 67 (1991) 147. H-U. Blaser, Tetrahedron: Asymmetry, 2 (1991 ) 843. G. Webb and P.B. Wells, Catal. Today, 12 (1992) 319. H-U. Blaser, H.P. Jalett, D.M. Monti, J.F. Reber and J.T. Wehrli, Stud. Surf. Sci. Catal., 41 (1988) 153. Y. Nitta, Y. Ueda and T. Imanaka, Chem. Lett., (1994) 1095. K.E. Simons, PhD thesis, University of Hull (1994). G. Bond and P.B. Wells, J. Catal., 150 (1994) 329 J.R.G. Perez, J. Malthete and J. Jacques, C.R. Acad. Sci. Paris Serie II, (1985) 169. S.R. Watson, PhD thesis, University of Hull (1995). I.M. Sutherland, A. Ibbotson, R.B. Moyes and P.B. Wells, J. Catal., 125 (1990), 77. H-U. Blaser, H.P. Jalett, D.M. Monti, A. Baiker and J.T. Wehrli, Stud. Surf. Sci. Catal., 67 (1991) 147. S.P. Gfiffiths, P. Johnston and P.B. Wells, unpublished work. P. Collier, J. Iggo and R. Whyman, personal communication.

J.W. Hightower, W.N. Delgass, E. Iglesia and A.T. Bell (Eds.) 11th International Congress on Catalysis - 40th Anniversary

Studies in Surface Science and Catalysis, Vol. 101 9 1996Elsevier Science B.V. All rights reserved.

231

S t e r e o c h e m i c a l S t u d i e s of the E n a n t i o - d i f f e r e n t i a t i n g H y d r o g e n a t i o n of Various Prochiral Ketones over Tartaric Acid-Modified Nickel Catalyst T. Sugimura, a T. Osawa, bt S. Nakagawa,ar T. Harada c and A. Tai a* aFaculty of Science, Himeji Institute of Technology; Kamigori, Ako-gun, Hyogo 67812 Japan bFaculty of General Education, Tottori University; Koyamacho-minami, Tottori 680 Japan CFaculty of Science and Technology, Ryukoku University, Seta, Otsu 520-21 Japan Abstract

The stereochemistry and enantiomeric excess (ee) of the products of the enantiodifferentiating hydrogenation of various prochiral ketones over (R,R)-tartaric acid modified Raney nickel catalyst (TA-MNi) were studied in detail. The results were consistent with our newly introduced concept of stereocontrol in that there were two types of interaction modes between adsorbed TA and the substrate on (R,R)-TA-MNi; a two point interaction to give a product of R configuration and an one point interaction to favor a product of S configuration. The relative contribution of these two modes determined the optical purity and stereochemistry of the product depending on the structure of the substrate. In the hydrogenation of a 3-oxoalkanoate where the contribution of two points interaction predominated, elimination of the minor contribution of the one point interaction by tuning the substrate structure enabled us to achieve 96% of ee, the highest value so far obtained using a heterogeneous catalyst.

1. INTRODUCTION Asymmetrically modified nickel catalyst (MNi) is one of the promising and wellinvestigated heterogeneous catalysts for the enantioface-differentiating hydrogenation of prochiral ketones. Many kinds of catalyst preparations and mechanisms of enantio-differentiation with this catalyst have been proposed by various research groups [1]. The best catalyst available at present is TA-NaBr-MRNi (tartaric acid-NaBr-modified Raney nickel catalyst) developed [2] and improved [3] by our group. Under well-optimized conditions, this catalyst gave 86% enantiomeric excess (ee) in the hydrogenation of methyl acetoacetate (MAA) to methyl t Present address: Facultyof Science,ToyamaUniversity,Gofuku,Toyama930 Japan ~:ReserchFellow fromToyoKasei KogyoCo. Ltd., Sone,Takasago676 Japan

232 3-hydroxybutanoate (MHB) and 80% of ee in the hydrogenation of 2-octanone to 2-octanol in the presence of an excess amount of pivalic acid [4]. The reaction site of TA-MNi could not be homogeneous judging from its existence on the solid surface. Through various investigations, it has been deduced that there are two types of sites on the catalyst; one is an enantio-differentiating site (E), where an optically active product is produced with the aid of an adsorbed chiral auxiliary (TA) and the other is a non-enantio-differentiating site (N), where the chiral auxiliary has no effect and the product becomes racemic. When the intrinsic enantiodifferentiating ability of the E site (factor i), which is related to the efficiency of the mutual interaction between TA and the substrate, is taken into account, the observed enantio-differentiating ability (e.d.a.) determined as the ee of the product is formulated according to eq. 1 (reaction site model) [5].

(1)

e.d.a. (%) = [i E / ( E + N)]xl00

The improvements in MNi so far achieved were mostly due to our efforts to eliminate the N site from MNi by changing preparation variables of the catalyst,. while the other important factor i has not been satisfactorily considered. In the present study, hydrogenation of various prochiral ketones with TA-MNi almost freed from the N site were carried out in order to gain insight into the mode of stereo control on MNi, which was expected to determine the stereochemistry of the reaction and to take part in the origin of the factor i.

2. RESULTS Two series of the reactions carried out in this study are shown in eqs 2 and 3.

H2 CH3 C(CH2)n COOR O

= (R,R)-TA-MNi

CH3 C*H(CH2)n COOR

(2)

OH

n = 0,1,2,3

RlCCH2COOR 2 O

H2 (R,R)-TA-MNi

R1 C.HCH2COOR 2

(3)

OH

R, R1,R2 : alkyl group of various chain lengths and branching The results of the series of reactions shown in eq. 2 are listed in Table I together with our early reported data on the h y d r o g e n a t i o n of 2-octanone (7) [4]. The hydrogenation on all substrates proceeded smoothly and gave the corresponding chiral secondary alcohol. In the case of 3_, 4_,5_, and 6_, some amounts of lactone were produced as by-product. From this study, quite interesting stereochemical behavior

233 T a b l e 1. Enantioface-Differentiating H y d r o g e n a t i o n of Various Keto Esters a n d 2Octanone reaction conditions .......p. r ~ configuration entry substrate temp.(~ addition of confi ee Pivalic acid am

ii

l i,,

i,

o .,o,~,

1

Jl

1

i

i

i,

i

100

no a

R

15

100

y esb . . . . . R .

14

I I

2

O

3

O

O

4

~O,K,.~

2

.,

.....

.......

5

.,O,r

6

0

0

...

no

R

85

ye s

R

.........72

100

no

R

38

100

yes . . . .

R

5

100

s .

..

100 ,,,,,

,,,,,,

.

100

no

4

120

yes

S

33

R = i-Bu

100

yes

S

49

10

80

yes

S

58

11

60

yes

S

61

12

40

yes

S

58

7

8

O

O

R.O.JJ,..A.~

9

61

13

s

R = Me

60

yes

14

6

R = n-Bu

60

yes

100

no

S

9c

120

yes

S

50 c

17

100

yes

S

62 c

18

80

yes

S

66 c

19

60

yes

S

74 c

20

40 . . . . . . .

yes

S

74 c

15 16

O ~

N

7

.....

a: A trace amount"0f acetic acid (2% w t / w t ) was added to the substrate. b: Pivalic acid (50% w t / w t ) was added to the substrate. c: The results reported in ref. 4.

59

.. . . . . . . .

, i,

234 T a b l e 2. Enantioface-Differentiating H y d r o g e n a t i o n of Various Alkyl 3oxoalkanoates entry

substrate

o

0

temp.

reaction

(~

time (h)

ee a

100

4

85

60

31

84

40

43

81

100

18

91

60

34

94

40

75

94

100

55

88

60

71

96

40

95

96

100

30

_C

60

67

_C

100

36

87

60

52

90

14

40

66

87

15

100

36

69(PA) b

16

60

43

71(PA) b

100

48

85

100

24

85

60

45

87

100

33

84

40

88

. .0 % e L

2

O

O

--oX..fl-./ ~0,.~~ O

O

..o.J~ O

10 11 12

O

O

13

N

O O

~

"-0~ O A c

17

O

18

O

19

-J-o

20

../

10

11

O

12

O 13

O,,

O ,,

21 )K'O'A""~ 14 60 a: All products were R excess except the product from 9. b: Pivalic acid (9 g) was added to a solution of the substrate in THF (10 ml). C: No reaction.

235 of the MNi was observed. In the reactions carried out in the absence of pivalic acid, the ee of the R configuration in excess was very low in I (n = 0), extremely increased in 2 (n = 1), decreased in 3 (n = 2), and then diminished in 4 (n = 3), whereas in 7 a slight excess of the S configuration was observed (Table I entries 1, 3, 5, 7, 15). When a l a r g e a m o u n t of pivalic acid was added in the reaction media, each substrate showed a characteristic change in stereochemistry and ee of the reaction. With the addition of pivalic acid, the ee in R excess was u n c h a n g e d in 1_, significantly decreased in 2 and 3_, while an appreciably high ee in S excess was observed in 4 as well as in _7 (Table 1, entries 2, 4, 6, 9, 17). The change in the alkyl group of the alkoxy side in 4 from/-butyl to methyl (5) or n-butyl (6) resulted in no essential change in ee (Table 1, entries 11, 13, 14). Our early study indicated that the ee in reaction 7 with pivalic acid s h o w e d significant temperature dependence [4]. The present study showed that reaction 4 with pivalic acid also resulted in a temperature dependence similar to that of Z (Table 1, entries 8 to 12). The results of a series of reactions shown in eq 3 are summarized in Table 2. All reactions proceeded to over 90% conversion and afforded almost a quantitative yield of 3 - h y d r o x y e s t e r with more than 85% ee except for the reaction of 1___00. The hydrogenation of 10 which carried a bulky t-butyl group next to the carbonyl group did not proceed (Table 2, entries 10 and 11). The elongation of the chain or branching of the alkyl group at the acyl side tended to increase the ee of reaction (Table 2, entries 1, 4, 7, and 12). Except for the reaction with 2 of which the ee showed no temperature dependence, the ee of the reaction with 8, 9_, and 1__!increased with a decrease in reaction temperature and reached to plateau at 60 ~ (Table 2, entries 1 to 9 and 12 to 14). The hydrogenation of 1__! in the presence of pivalic acid also decreased the ee same as in the case of _2 (Table 2, entries 15 and 16 and Table 1, entry 4 ). Introduction of a bulky alkoxy group in acetoacetate, 13 and ~ gave almost the same result as 2 (Table 2, entries 18 to 21).

3. D I S C U S S I O N S

In o r d e r to explain the stereochemistry of the enantioface-differentiating hydrogenation over TA-MNi, we have proposed a two point interaction (2P) mode for MAA (2) [6] and a one point interaction (1P) mode for 2-octanone (7) [7]. Figs 1 and 2 s h o w a detailed sketch of each model provided by (R,R)-TA-MRNi. These models can be simplified as shown in Figs 3 and 4, respectively. Both models have been p r o p o s e d based on the concept that the adsorption mode of TA should be the same in each case and one of the OH groups in TA located close to the Ni surface (site 1) interacts with C=O of the substrate to be hydrogenated at the immediate surface of Ni through hydrogen bonding. The 2P mode achieved when the substrate possesses a second functional group to associate with the second OH in TA (site 2) as is the case with M A A gave a product of R configuration. If the substrate has no additional functional group as is the case with 2-octanone, association of TA and the substrate becomes 1P mode. In this case, a weak steric repulsion takes place between the OH group of site 2 and the alkyl chain of the substrates and the orientation of the adsorbed substrate favors the formation of the S configuration to some extent. Pivalic

236 acid employed as an additive in the reaction system is expected to enhance the steric effect by forming a bulky repulsive fence at site 2 with the association and making the stereo-control ability of 1P effective [4]. The contribution of these two modes to the stereo-controlling ability of the E site is considered to be different. The contribution of the 2P mode directly relates to the formation of the R product, while the 1P mode contribution varies from the formation of an almost racemic product to the formation of a product with an appreciably high excess of S depending on type of substrate and the reaction conditions. Pivaric acid,

0

9 C H

site 2 site 2

o

MAA

(R,R)-TA

(R,R)-TA

site 1

H21L ~

~

H,

Nickel Surface

Nickel Surface

2-,

Figure 1. Two-point interaction Model (2P) of MAA

,r

/

Figure 2. One-point interaction Model (1 P) of 2-Octanone C

\ site 2 O ..... O ~

,r

site 2

.

..... o ' ~

'i o

"i' 0 iHi

Figure 3. Two-point model (2P)

OH

~L

i

|l lllll

H,,i

,ill

Figure 4. One-point model (1P)

Taking the above mentioned characteristics of the two modes into consideration, we introduced the concept of stereo-control in the enantio-differentiating hydrogenation of various functionalized prochiral ketones on TA-MNi based on the coexistence of 2P and 1P on the E site of the catalyst. That is the 1P function counteracts the 2P function when 1P and 2P coexist, and the relative contribution of the two modes determine the stereochemistry of the product produced in excess and also relates qualitatively to the i factor. From the study of a series of substrates, I to 4 and 7_, summarized in Table 1, the (R,R)-TA-MNi found to show quite interesting stereochemical behavior in connection with the concept of stereocontrol mentioned above. From the stereochemistry of the

237 hydrogenation product in the absence of pivalic acid (Table 1, entries, 1, 3, 5, 7, 15), it was expected that the contribution of 2P was in excess for K to 3, and the 2P contribution became m a x i m u m in 2_, while the contribution of 1P and 2P were compensated in 4_, and exclusive 1P participation in 7 resulted in a slight S excess. From the ee of each reaction product (Table 1, entries 1, 3, 5, 7, 15), participation of 2P and 1P was expected to change systematically depending on the degree of fit between the two carbonyl groups in the substrate and the two interaction sites in TA as shown in Figs 3, 5, 6, and 7, respectively. The results of the reactions in the presence of pivalic acid (Table 1, entries 4, 6, 9, 17) clearly showed that pivalic acid in the reaction media decreased the 2P participation by competing with the interaction at site 2 with the substrate and enhanced the 1P contribution and its efficiency by forming a steric fence in the proximity of site 2. It is noteworthy that, in the absence and presence of pivalic acid, the ee of the hydrogenation of 3 changes from an appreciable R excess to negligible R excess and that of the hydrogenation of 4 changes from zero to appreciable S excess, respectively. On these two substrates, the m o d e of stereocontrol was transient from 2P to 1P contribution depending on the presence or absence of pivalic acid. In our early study of the hydrogenation of 2-alkanone with pivalic acid, it was reported that the ee was significantly temperature-dependent and this behavior was characteristic of the 1P m o d e [4]. A similar t e m p e r a t u r e dependence of ee found in 4_ and 7 (Table 1, entries 8 to 12 and 16 to 20) is also additional support of the presumption that the hydrogenation _4 with pivalic acid proceeds mostly by the 1P mode. The decrease of ee with increased temperature in the reaction by the 1P mode was explained by a decrease in steric repulsion due to the decrease in the tightness of h y d r o g e n - b o n d i n g at site 1, a decrease in the associative ability of pivalic acid with site 2, and also an increase of the mobility of alkyl chain of the substrate with the increased temperature. , o"~

[

AN

site 2 ........... 0

9 .r site 2.. .... 0..~/0

/

site 2 ....."....

/

,

0

. . . .

i,o=

Figure 5. Possible 2P mode

Figure 6. Possible 2P mode

Figure 7. Possible 2P mode

for I

for 3

for 4

A series of studies with various 3-oxoalkanoates provided further information on the factor i. As listed in Table 2, all substrates give more than 85% ee in the R configuration in excess. These facts clarify that a strong contribution of 2P is common to all 3-oxoalkanoates. The other important finding is that elongation or branching of the alkyl moiety at the acyl side of a 3-oxoalkanoate tends to increase the ee of hydrogenation. An increment in ee from 2 to 9 is especially significant (Table 2, entries 2 and 8). Fig 8 and Fig 9 are schematic sketches of the S configuration favored the 1P expected in 2 and 9_, respectively. In 2_, no significant steric interaction is

238 expected between the methyl group at the acyl side and the OH of site 2, whereas in 9_, the isopropyl group at the acyl side and the OH of site 2 come into contact, making 1P contribution difficult. In this context, the participation of 1P became less in reaction of 9_ than in reaction of 2 and hence 9__gave higher ee than 2 due to the increase in the relative participation of 2P. The temperature dependence of ee found in reaction 9 suggested that reaction of 9 still involved a small portion of the 1P mode at high temperature. Results of other substrates carrying a long chain at the acyl side; 8_,11 (table 2, entries 4 to 6, 12 to 14,) are also rationalized in the same way as that of 9. Through the series of discussions described above, the various forms of stereochemical behavior of MNi could be explained rationally on the basis of our newly proposed concept. A"x. sitex2 Si~oH~C~-- CH 3

V site"i'O= o 0 Figure 8.1P mode for 2 /

Figure 9.1P mode for 9 /

The 96% ee is the highest so far achieved by the enantio-differentiating hydrogenation over an asymmetrically modified heterogeneous catalyst. Although it is difficult to separate factor i and E~ (E+N), the present results indicated that both of them are well optimized and become almost unity.

4. EXPERIMENTAL 4.1. Preparation of TA-NaBr-MRNi TA-NaBr-MRNi was prepared by the reported method [3]. RNi (W-1 type) was prepared from 1.9 g of Raney nickel alloy (Kawaken Fine Chemical Co., Ni/A1 = 42/58). To wash out the excess base and aluminum salts, a sufficient amount of deionized water was used with ultrasonic irradiation. The modifying solution was prepared by dissolving of (R,R)-tartaric acid (1 g) and NaBr (6 g to 10 g) in 100 ml of water and adjusting the pH to 3.2 with 1N NaOH aqueous solution. RNi was heated in the modifying solution at 100 ~ for 1 hour, washed with water (50 ml), methanol (50 ml, twice), and THF (10 ml). The TA-NaBr-MRNi obtained by this method was immediately used for the hydrogenation. 4.2. Preparation of the Substrates The following substrates were obtained from commercial sources, methyl pyruvate (1), methyl acetoacetate (2), methyl 4-oxopentanoate (3), and methyl 3-oxopentanoate (8). Alkyl 5-oxohexanoates (4, 5 and 6) were prepared by condensation of methyl acetoacetate and methyl acrylate followed by acidic hydrolysis, decarboxylation, and esterification [8]. Methyl 3-oxo-4-methylpentanoate

239 (9), methyl 3-oxononanoate (11), and methyl 3-oxo-lO-acetyloxydecanoate (12) were prepared from Meldrum's acid and 2-methylpropanoyl chloride, heptanoyl chloride, and 8-acetyloxyoctanoyl chloride, respectively by a reported method [9]. i-Propyl acetoacetate (13) and t-butyl acetoacetate (14) were prepared from diketene and the corresponding alcohols in the presence of triethylamine [10]. Each substrate was purified either by fractional distillation or column chromatography before use. 4.3. Hydrogenation

In an autoclave of 100 ml capacity, TA-NaBr-MRNi and a solution of 1.5 g to 10 g of the substrate in 10 ml of THF were placed. In some cases, 9 ml of pivalic acid was added, and in the other cases, 0.2 ml of acetic acid was added. Hydrogen was charged into this at ca. 107 Pa as the initial pressure. Reaction temperatures and times of the hydrogenation are listed in Tables I and 2 in the text. After cooling, the reaction mixture was filtered and purified either by distillation or column chromatography on silica gel depending on the nature of the product. Structural characterization of each purified product was carried out using NMR 0EOR GX 400 spectrometer) and IR (JASCO IR-88 spectrometer) spectra. 4.4. Determination of ee of the product

The enantiomeric excess (ee) of the hydrogenated products was determined either by polarimetry, GLC equipped with a chiral column or 1H-NMR with a chiral shift reagent. Methyl lactate and methyl 3-hydroxybutanoate, obtained from I and 2_., respectively, were analized polarimetry using a Perkin-Elmer 243B instnmlent. The reference values of [(Z]D(neat) were +8.4 ~ for (R)-methyl pyruvate and -22.95 ~ for methyl 3-hydroxybutanoate. Before GLC analysis,/-butyl 5-hydroxyhexanoate, methyl 5-hydroxyhexanoate, and n-butyl 5-hydroxyhexanoate, obtained from 4_, 5_, and 6, respectively, were converted to the pentanoyl esters, methyl 3-hydroxybutanoate was converted to the acetyl ester, and methyl 4-methyl-3hydroxybutanoate obtained from 9 was converted the ester of (+)-~-methyl-(x(trifluoromethyl)phenyl acetic acid (MTPA). The GLC analysis with a Shimazu GC 17A equipped with a column of CPChirasil DEX CB (25 m, 0.25 mm id, GL Science, Japan) was applied for the determination of the ee of the following compounds at the temperature stated in parentheses: i-butyl 5-pentanoyloxypentanoate (150 ~ methyl 5-pentanoyloxypentanoate (130 ~ n-butyl 5-pentanoyloxypentanoate (150 ~ and methyl 3acetyloxybutanoate (100 ~ the MTPA ester of 4-methyl-3-hydroxybutanoate (160 ~ obtained by the above mentioned derivation, and methyl 3-hydroxypentanoate (80 ~ methyl 3-hydroxynonanoate (140 ~ i-propyl 3-hydroxybutanoate (90 ~ and t-butyl 3-hydroxybutanoate (100 ~ obtained directly from ~, !1, 13, and respectively. In all cases, enantiomers showed completely separated peaks, and repeated analysis showed an error within 0.2%. The ee of methyl 10-acetyloxy-3hydroxydecanoate obtained from 12 was determined by 1H-NMR in the presence of (+)-Eu(hfc)3. In the cases of methyl pyruvate and methyl 3-hydroxybutanoate, the ee value was determined by polarimetry and GLC fell within the range of 0.2%.

240 ACKNOWLEDGEMENT This work was partially supported by a Grant-in-Aid for Scientific Research No. 06640697 from the Ministry of Education, Science and Culture, Japan. We also thank Toyo Kasei Kogyo Co. Ltd., who permitted us to use their facilities.

REFERENCES

1. A. Tai and T. Harada in "Tailored Metal Catalysts" (Ed. Y. Iwasawa), D. Reidel, Dordrecht, p. 265 (1986). W. M. H. Sachfler in "Catalysis in Organic Reactions" (Ed. L. Augustine), Chem. Ind., 22, 189 (1985). 2. T. Harada, M. Yamamoto, S. Onaka, M. Imaida, H. Ozaki, A. Tai, and Y. Izumi, Bull, Chem. Soc. Jpn., 54, 2323 (1981). 3. A. Tai, T. Kikukawa, T. Sugimura, Y. Inoue, S. Abe T. Osawa, and T. Harada, Bull. Chem. Soc. Jpn., 67, 2474 (1994). 4. T. Osawa, T. Harada, and A. Tai, J. Cat. 121, 7 (1990). 5. T. Harada, A. Tai, M. Yamamoto, H. Ozaki, and Y. Izumi, Proc. 7th Int. Congr. Catal. Tokyo,p 364 (1980). 6. A. Tai, T. Harada, Y. Hiraki, S. Murakami, Y. Izumi, Bull. Chem. Soc. Jpn., 56,1414 (1983). 7. T. Osawa, T. Harada, and A. Tai, J. Mol. Catal. 88, 333 (1994). 8. S. Okamoto, T. Harada, A. Tai, Bull. Chem. Soc. Jpn., 52, 2670 (1979). 9. Y. Oikawa, K. Sugano, O. Yonemitsu, J. Org. Chem., 43,2087 (1978). 10. T. Kato, T. Chita, Chem. Pharm. Bull., 23, 2263 (1975).

J.W. Hightower, W.N. Delgass, E. Iglesia and A.T. Bell (Eds.) l l t h International Congress on Catalysis - 40th A n n i v e r s a r y

Studies in Surface Science and Catalysis, Vol. 101 L996 Elsevier Science B.V. .

.

241

.

Enantio-differentiation over heterogeneous catalysts. The shielding effect model. J6zsefL. Margitfalvi, Mihhly Heged0s and Ern6 Tfirst Central Research Institute for Chemistry of the Hungarian Academy of Sciences 1025 Budapest, Pusztaszeri ut 59-67, Hungary This paper deals with the origin of enantio-differentiation over heterogeneous catalysts. A new model is proposed, in which the modifier provides a specific shielding effect. A prochiral molecule, due to the specific character of shielding, can adsorb onto the metal surface by its unshielded site resulting in enantio-differentiation. As emerges from computer modeling, quantum chemical and quantum mechanical calculations made on the catalytic system: cinchonidine - a--keto esters - Pt the shielding effect can be responsible both for the rate acceleration and induction of enantio-differentiation. This model is based on an earlier proposition, which suggest (i) the formation of a weak complex between the modifier and the substrate in the liquid phase and (ii) the hydrogenation of either the shielded or unshielded forms of a-keto ester over the Pt sites. The shielded form gives the optically active, while the unshielded one the racemic product. Further support for this model was obtained in kinetic experiments and kinetic modeling. I. INTRODUCTION There is a growing interest for enantioselective reactions and there is a real need to develop both homogeneous and heterogeneous catalysts with high enantioselectivity. There are only two heterogeneous catalytic reactions in which high optical yields (up to 95 %) were obtained: hydrogenation of 13-keto esters and cx-keto esters over Ni-tartrate and cinchona-Pt/Al203 catalysts, respectively [1-6]. Despite the extensive studies devoated to the above two hydrogenation reactions the exact nature of the enantio-differentiation (ED) is still not really known. The lack of knowledge of the nature of ED involved in heterogeneous catalytic hydrogenation reactions is the main reason that we are far away to design heterogeneous catalysts with the desired enantio-differentiation ability. In this work a new approach is desribed, which can help to understand ED over heterogeneous catalysts. We also hope that this approach can be used to find new modifiers for enantioselective heterogeneous catalytic reactions. The basis for this approach is the steric shielding known in organic chemistry [7,8]. A chiral template molecule can induce shielding effect (SE) in such a way that it preferentially interacts with one of the prochiral sites of the substrate. If a substrate is preferentially shielded its further reaction can take place only from its unshielded site resulting in ED.

2. EXPEREVIENTAL

2.1. Computer modeling Computer modeling was applied to investigate the ability of commonly used modifiers to create SE. In this respect the following calculations were carried out: (i) conformational

242 analysis of the modifier and substrate, (ii) molecular docking; i.e formation of the [modifier substrate] complex, (iii) adsorption of the [modifier-substrate] complex onto Pt surface. The corLformational analysis of the modifier and substrate was performed using the Hypercube: Hyperchem 3.1 programs with the MM+ forcefield. The minimum energy conformations were determined using the Polak-Ribiere (conjugate gradient) minimization algorithm, and the conformational analysis was carried out with rigid quinoline and quinuclidine parts. The molecular docking calculations were performed with BIOSYM programs. The optimization of the geometry was carried out using the Discover (cvff forcefield, VA09A minimization method) and the Ampac/Mopac (AM1 semiempirical method) modules of the InsightlI program package. These calculations were performed for the whole modifier-substrate system from different starting positions. The adsorption of the modifier-substrate complex onto Pt (111) surface was investigated using the Solids-Docking module of the InsightlI package. This module determines the conformations of the adsorbed molecules by a combined approach of high temperature molecular dynamic simulations with molecular mechanics minimization. All the calculated structures were visualized on a Silicon Graphics workstation.

2.2. Hydrogenation of methyl and ethyl pyruvate The hydrogenation of ehtyl pyurvate (EtPy) was carried out at 23 ~ in a SS autoclave equipped with an injection chamber for separate introduction of the modifier. Cinchonidine (CD) and Troger's base (TB) was used as modifiers. Different batches of EtPy, (Fluka) and Pt/Al203 catalysts (Engelhard E 4759, 5 %w Pt, Dpt = 25 %) were used. Experimental details incliding GC analysis can be found elsewhere [3,12]. The optical yield was calculated as e.e. = (~]-[S])/([R]+[S]). The e.e. values were corrected for the amount of racemic product formed in minor amount in the reactor prior to the injection of CD. 2.3. Kinetic modeling The material balance was calculated for EtPy, ethyl lactates (EtLa) and CD by solving the set of differential equation derived form the reaction scheme. Adam's method was used for the solution of the set of differential equations. The rate constants for the hydrogenation reactions are of pseudo first order. Their value depends on the intrinsic rate constant of the catalytic reaction, the hydrogen pressure, and the adsorption equilibrium constants of all components involved in the hydrogenation. It was assumed that the hydrogen pressure is constant during the kinetic run. Some of the rate constants (kr and kina) were approximated from independent kinetic measurements. The starting values of the differential equations were obtained from the first measured concentration of EtPy and EtLa. The initial ratio of the open and closed forms of the modifier ([CDclose,]o/[CDopen]o) was estimated. The estimation was based on NOESY NMR spectra.

3. THE PRINCIPLE OF CHEMICAL SHIELDING Recently it has been evidenced that a large aromatic substituent, such as naphtyl, can provide a intramolecular steric shielding for an c~-keto ester moiety [9] resulting in enantio-

243 differentiation in the hydrogenation of the 0t-keto group. No ED was observed if the naphtyl ring was substituted for a phenyl one. The ED was attributed to the SE induced by the large aromatic moiety. Similar results were observed in the enantioselective hydrogenation of ethyl pyruvate over Pt/AI203 catalyst in the presence of new types of modifiers [ 10]. In the presence of these new modifiers the ED was completely lost if the naphtyl or quinolyl ring was replaced by phenyl or pyridyl group. It should also be mentioned that in the hydrogenation of ot-keto esters over CD-Pt/AI203 catalysts the ED was partially or fully lost if the quinoline ring of the modifier was partially or fully hydrogenated [11 ]. As emerges from the these results large aromatic substituents might play an important role in the induction of ED. The similarities of above experimental results inspired us to investigate the role of SE in heterogeneous catalytic enantioselective hydrogenation reactions. In heterogeneous catalytic reaction the SE means that a given template molecule interacts with the prochiral substrate in the liquid phase in such a way that one of the prochiral sites is preferentially shielded. If the substrate is shielded then its adsorption onto the metal can take place with its unshielded site resulting in ED. An organic molecule can induce SE if it has (i) an asyrmnetry center (A), (ii) an appropriate bulky functional group 03) for weak interaction with the substrate, (iii) bulky but planar group (C) to induce the steric shielding. If the above requirements are fulfilled the modifier can form a week complex with the substrate. In the above complex the modifier should have an umbrella like conformation with high extent of 't~oncavity" as shown in Figure 1. The role of "concavity" in chemical shielding has been discussed earlier [7,8]. The shielded form of the [substrate - modifier] complex formed in the liquid phase can maintain its entity even after adsorption onto the metal surface. It should be added that based on reaction kinetic data the formation of [substrate - cinchonidine] complex in the liquid phase was suggested in our earlier and recent studies [4, 12]. There are also important requirements for the heterogeneous catalysts: (i) the catalyst should not hinder the formation of the [substrate - modifier] complex, (ii) the modifier should not adsorb irreversible onto the catalyst; (iii) the catalyst should be inactive in the transformation of the modifier into a new derivative, (iv) the catalyst should be resistant towards poisoning by modifier, substrate or product.

4. RESULTS AND DISCUSSION

4.1. Computer modeling The principles of the SE were applied for two enantioselective hydrogenation reactions: (i) hydrogenation of 13-keto esters over Ni-tartrate and (ii) hydrogenation of ct-keto esters over cinchona-PffAl203 catalysts. In this respect the tartaric acid - f3-keto ester system gave a negative result. Neither the substrate nor the modifier have bulky substituents required for SE. The first approach applied for [cinchonidine (CD) - ct-keto ester] complex was also unsuccessful. In the open conformation CD cannot provide the required steric shielding. In open form either the quinuclidine or the quinoline moiety of CD will interact with the substrate. It has already been demonstrated that the quinuclidine moiety has a crucial role both in the rate acceleration and the induction of ED [13].

244 In earlier kinetic and computer modeling [ 1, 2, 14] the open form of CD (CDopen) was used to illustrate the adsorbed [CD - ct-keto ester] complex. In this complex the quinuclidine nitrogen was involved in the interaction with the substrate directly or via a proton bridge. We have modelled the [CDopen - methyl pyruvate] complex. The result is shown in Figure 2. In this complex there is no steric hindrance to prevent the free rotation of the substrate around the quinuclidine nitrogen. Thus, in complex shown in Figure 2. there is no preferential stabilization of the substrate. In earlier computer modeling it was suggested that Pt is involved in the stabilization of the [CDopen--~-keto ester] complex, i.e. the Pt surface prevent the free rotation of the substrate, however the driving force for enantio-differentiation, i.e. for preferential adsorption of the substrate, was not discussed [ 14]. In our second approach the closed form of CD (CDclosed) was used for modeling (see Figure 1). It was found that CD in its closed conformation can provide the concave, umbrellalike form required for steric shielding. The calculated [CDclosed-methyl pyruvate] complex is shown in Figure 3. This complex (complex (R)), after subsequent hydrogenation over Pt should result in (R) -lactate ester. The conformational change of CD from open form to closed one requires the rotation of the quinuclidine ring around the C-(9)- (C4') axis, i.e. to change the torsion angle (C4')-(C9)-(C8-(N1) from 159.42 ~ to 52.32 ~ The energy map for CD was calculated by changing the torsion angles (C3')-(C4')-(C9)-(C8) and (C4')-(C9)-(C8)-(C7). The conformational analysis indicates that CD can exist at least in four different forms. The energy needed to change the conformation of CD from the open form (O1) (it is the crystallographic form) to the most stabile closed one (CI 1) is less than 5 kcal/mole. Figure 4. shows the [CDclosed-methyl pyruvate] complex, in which the substrate is rotated by 180~ The above complex upon hydrogenation will result in (S)-lactate. The major difference between complexes (R) and (S) is the mode of interaction between the loan pair of electrons of the quinuclidine nitrogen and the keto-carbonyl group. In complex (R) the "directionality" [15] of the nucleophilic attack by quinuclidine nitrogen towards the keto carbonyl group is very favourable for the interaction with the keto carbonyl group. The orbital steering theory states [16] that a proper 'reaction window" or "reaction cone" can result in perturbation of the reacting group. We suggest that this perturbation leads to a pronounced rate increase. Thus, in complex (R ) the favourable directionality promotes the perturbation of the keto carbonyl group resulting m the observed rate acceleration. Contrary to that in complex (S) the interacting groups are misaligned. Due to this misalignment no rate acceleration can be expected, i.e. the hydrogenation of (S) complex is not accelerated. Variety of ct-keto esters, such as methyl and ethyl pyruvate, methyl mandalate, dihydro-4,4 - dimethyl-2,3 furanedione were used to calculate the shielded form of [CDclosed- ct-keto ester] complexes leading to the formation of ( R ) or (S) product, respectively. The details of these results will be a subject of a subsequent paper [ 17]. As emerges from these calculations the favourable 'ttirectionality" is maintained in complexes (R), even for dihydro-4,4 - dimethyl2,3 furanedione. Monte-Carlo simulation method was used to investigate the interaction of the [CDclosedMePy] complexes with Pt (111) surface. The result shown in Figure 5 indicates that the shielded complex can maintain its entity even after adsorption. Further computer modeling indicated that there are other molecules with the ability to induce SE. In this respect Troger's bases are of particular interest. The calculated Troger's base-methyl pyruvate complex (R form) is shown in Figure.6.

245

Figure 3. The shielded form of [CDmethyl pyruvate] complex, (R) form

Figure 4. The shielded form of [CDmethyl pyruvate] complex, (S) form

4.2. Hydrogenation experiments Hydrogenation of ethyl pyruvate in the presence of cinchonidine. In our previous studies [3, 4,14] variety of experimental data were obtained, which could not be explained by existing models [1,2] proposed earlier. These results are as follows [3,4,12]: (i) the monotonic increase type behaviour of the optical yield - conversion dependencies, (ii) the complexity of the reaction kinetics, (iii) side reactions catalyzed by CD. It was also demonstrated that the enantio-differentiation can be induced if the modifier is injected into the reactor during racemic hydrogenation. Upon injection of CD into the reactor during racemic hydrogenation the rate acceleration was always instantaneous, while the optical yield vs. conversion dependencies showed a monotonic increase type behaviour as seen in Figure 7. In acetic acid the increase part of the above dependence is so fast that it hardly can be followed by our sampling technique. At low concentration of modifier the optical yield passes through a maximum. In this case the

246 decreasing part is due to the transformation of CD during the hydrogenation reaction. Parallel formation o f CD derivatives with saturated quinoline ring was evidenced by thin layer chromatography.

Figure 7. Typical optical yield - conversion dependencies. [CD]o, M: O - 6.8 x 10.6 9 - 3.40 x 10 5, 9 - 0.8 x 10 -4 + 5.0 M AcOH.

Figure 8. The shielded form o f [ C D methyl pyruvatesyn] complex, ( R ) form

We have compared, the rate acceleration effect induced either by the CD and different moieties originated from CD, i.e quinuclidine and quinoline. These experiments were carried out in ethanol. If the relative rate of racemic hydrogenation is equal to one the following relative rates has been measured: quinoline = 2, quinuclidine = 3, cinchonidine = 40. In the presence of quinoline a short induction period was needed to observe the rate acceleration. It is suggested that during this period quinoline was partly hydrogenated. Other tertiary nitrogen bases, such as triethylamine, triethylenediamine, etc. resulted also rate acceleration with relative rate = 2-4.

247 The above comparison indicates that the rate acceleration induced by CD is more pronounced than that of the other tertiary nitrogen bases. This fact also indicates that in CD a cooperative effect should exist between the quinuclidine nitrogen and the quinoline ring. The

cooperative effect is in force if the modifier is in a shielded form. One of the most interesting side reactions taking place during the enantioselective hydrogenation is the transesterification of the substrate or the reaction product. If the enantioselective hydrogenation of ethyl pyruvate was performed in methanol as a solvent the formation of methyl pyruvate and methyl lactate was observed. CD appeared to be an effective catalyst for the above transesterification reaction. The transesterification reaction can be attributed to the perturbation of the ester carbonyl group in the [CDclosed-substrate] complex. The possibility of this side reaction was predicted by earlier quantum-chemical calculations [ 18]. These results indicated that the reaction pocket in methyl pyruvate for the nucleophilic attack is situated between the two carbonyl groups, i.e. both carbonyl groups can be perturbed by a nucleophile provided both carbonyl groups have the fight "directionality". However, the fight "directionality" for both carbonyl groups can be obtained if they are in syn position. The conformational analysis of methyl pyruvate shows that it can have two conformers. In the second conformer the two carbonyls are in syn position. The anti-syn conformational change requires 3 kcal. The [CDclosed- methyl pyruvatesyn] complex ((R) form) was also calculated and shown in Figure 8. In the above complex the "directionality" of the lone pair of electrons of the quinuclidine nitrogen is advantageous for interactions with both the keto and the ester carbonyl groups. Table 1. Hydrogenation of ethyl puruvate in the presence of Troger's base No

Pressure, bar

Modifier .....

Acid added, M

Rate contant min "1

Conv. '~" %

Optical '~* yield,

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

50 50 50 50 10 50 50

TB TB TB TB CD* CD*

AcOH, 5.0 AcOH 0.5 TFAc, 0.5 AcOH, 5.0 AcOH, 5.0 AcOH, 5.0

0.0041 0.0075 0.0057 0.0050 0.0045 0.0560 0.1320

10.9 12.5

0.088

21.2

0.242

11.2 12.9 97.9 99.5

0.383 0.248 0.860 0.931

Solvent: Toluene, T = 10 ~ * measured at 23 ~

TFAc- trifluor acetic acid,

** Conversions and optical yields measured after 60 minutes.

H y d r o g e n a t i o n of ethyl puruvate in the presence of Trogers base. It has been demonstrated by computer modeling that other organic molecules, such as Troger's base (TB) can induce ED in the hydrogenation of ethyl pyruvate over Pt/AI203. The results obtained in the presence of TB are summarized in Table 1. TB results in ED only in the presence of acetic acid. This modifier, contrary to the cinchona alkaloids, did not result in rate acceleration and the enantio-differentiation was moderate.The main difference between CD and TB is that in the [TB - substrate] complex there is no interaction between the tertiary nitrogen and the keto

248 carbonyl group. Consequently, there is no pronounced rate acceleration and the optical yield is relatively low. The relatively low optical yield can be attributed to the small size of the shielding group. Further experiments, including synthetic works and kinetic measurements are in progress to optimize the optical yield induced by Troger's bases. 4.3. Kinetic Modeling The shielding effect model suggest that both the rate acceleration and the induction of ED is attributed to the formation of a shielded [substrate-modifier] complex in the liquid phase (reaction (1)). In this complex CD is in closed form. The above complex, referred as [X]cl upon hydrogenation will result in (R) lactate (reaction (2)). Reactions (1) and (2) strongly resemble the corresponding steps in enzyme catalyzed reactions. The formation of complexes [X] takes place in an equilibrium reaction. The corresponding unshielded [substrate-modifier] complex, [X]closed can also be hydrogenated, however it gives racemic product similar to the free substrate, see reactions (4) and (5), respectively. The rate of reaction (2) is much higher than that of reactions (4) and (5). The equilibrium reactions for the conformational changes of CD and for the transformation of [X]close d into [X]open are also included into the reaction scheme ( reactions (6) and (7), respectively). The hydrogenation of both forms of CD is also taken into account. This reaction, in which the quinoline ring of CD is hydrogenated is responsible for the decrease of the optical yield with conversion observed at very low concentration of CD. The deactivated form of CD formed in the above reaction is referred as

CDma. The simplified reaction scheme for the enantioselective hydrogenation of ot-keto esters over cinchona-Pt/Al20 3 catalyst can be written as follows: Substrate

+

[CD]cl



[X] cl

+

H2/Pt

--k* .... >

Substrate

+

[CD]o p



[X]op

+

H2/Pt

-k*r---> (R)-lactate + CDop

(4a)

[X]op

+

H2/Pt

--k'----> (S)-lactate + CDop

(4b)

Substrate

+

H2~t

-kR.... > (R)-lactate

(Sa)

Substrate

+

H2/Pt

-ks .... >

[X]c I

(1)

(R)-lactate + [CD]ci

(2)

[X]o p

(S)-lactate

(3)

kR=ks=kr

(Sb)

[CD]op



[CD]cl

(6)

[x]~l

.

[X]op

(7)

[CD]cl

+

H2/Pt

--kina...... >

[CD]ma

(8a)

[CD]o p

+

H2/Pt

--kina...... >

[CD]ina

(8b)

In this scheme, due to the rate acceleration effect, the enantioselective hydrogenation is much faster than the two racemic hydrogenation reactions (k* > kr, k* > k'r). Please note that the rate constants for the hydrogenation reactions of are pseudo fist order, which contains in a

249 certain form the intrinsic rate constant and the adsorption equilibrium constants of all components involved in the hydrogenation. More detailed reaction network will be needed, which will take into account all of the adsorbed forms of substrate, modifier, reaction product and by-products. Two kinetic experiments with different CD concentrations were used for kinetic modeling. In this simulation all of the rate constants not involved in the hydrogenation step were not altered. The calculated and simulated kinetic curves and optical yield-conversion dependencies are shown in Figure 9a and 9b. The results of kinetic modeling indicates that the whole kinetic curve and the optical yield - conversion dependencies can be well described by a kinetic model derived from the shielding effect model.

.0

-

'

1.0

'-"- '-

i

0.8"

0.8

~, 0.6

"~, 0.6

g 0.4

=.0.4

0.2

0.2 1

0.

'

0

20

40

60

80

Time [min]

100

120

.

l

.

l

,

1

0.0~

0.0

,

'

0.2

0.4

0.6

0.8

1.0

Conversion

Figure 9. Results of kinetic modeling, a. Conversion vs. time, b. Optical yield vs. conversioia dependencies. I I - [CD]o = 6.8 x 10-6 M, 9 - 3.4 x 10-5 M, [Substrate]o = 0. 98 M., [Rlactate]o = 0.01, [S-lactate]o = 0.01 M. [CDopen]o/[CDelosed]o = 9/2. The continuous curves were obtained by kinetic modeling.

CONCLUSIONS A new model is proposed to undestand the origin of ED over heterogeneous catalyts. The model was applied for the enantioselective hydrogenation of ct-keto esters. The shielding effect model is based on steric shielding provided by a large aromatic ring. The new model can explain all of the observations, which could not be interpreted by existing models [ 1,2]. It is the first working model, which can explain (i) the instantaneous rate acceleration, (ii) the monotonic increase character of the optical yield - conversion dependencies and (iii) the appearence of the maximum in the optical yield - conversion dependencies. The shielding effect model can explain the functional behaviour of different parts of the cinchonidine molecule. This model can be used to design new modifiers for the enantioselective hydrogenation of otketoester. The computer modeling with molecular mechanics calculations appeared as a powerful tool to give qualitative explanations for the modifer-susbtrate interactions taking

250 place in the liquid phase. The next step should be to give quantitative data for the above interactions. These studies are in progress in our laboratories.

ACKNOWLEDGMENT Financial support given by OTKA (Grant No: T1801 and T4340) is greatly acknowledged. REFERENCES

1. I.M. Sutherland, A. Ibbotson, R.B. Moyes and P.B.J. Wells, J. Catal., 125, (1990) 77. 2. M. Garland and H.U. Blaser, J.Am.Chem. Soc., 112, (1990) 7048. 3. J.L. Margitfalvi, B. Minder, E. T/das, L. Botz and A. Baiker, New Frontiers in Catalysis, (Proc. 10th Int. Cong. Catal. Budapest, July 1992), Guczi, L. et al. (eds), Elsevier, Amsterdam (1993) 2471. 4. J.L. Margitfalvi, Chem. Ind. (Marcel Dekker), 62 (Catal. Org. React., Scaros, M.G., Pmnier, M.L. (eds)), (1995) 189. 5. Y. Izumi, Adv. Catal., 32 (19830 215. 6. A.Tai and T.Harada, Taylored Catalysts, Y. Iwasawa (ed.), D.Reidel, Dordrecht, (1986) 265. 7. H. Bushman, H.D. Scharf, N. Hoffmann and P. Esser, Angew. Chem. Int. Ed. Engl., 30 (1991)477. 8. G. Helmchen and R. Schiere, Angew.Chem. Int. Ed. Eng. 22 (183) 237. 9. U.Maitra and P. Mathivanan, Tetrahedron: Asymmetry, 5 (1994) 1171. 10. K.E. Simons, G. Wang, T. Heinz, T. CJiger, T. Mallat, A. Pfaltz and A. Baiker, Tetrahedron: Asymmetry, 6 (1995) 505. 11. J.L. Margitfalvi, P. Marti, A. Baiker, L. Botz and O. Sticher, Catal. Lett., 6, (1990) 281. 12. O. Schwalm, B. Minder, B., J. Weber, and A. Baiker, Catal. Lett. 23 (1994) 245. 13. J.L. Margitfalvi J. and M. Hegedus, 8th International Congress on the Relation Between Homogeneous and Heterogeneous Catalysis, Balatonfured, Hungary, Sept. 1995. N~ oral presentation; accepted for publication in J. Mol.Cat. 14. H.U. Blaser and M. Muller, stud. Surf. Sci. Catal. 59 (1991) 73. 15. F. Manger, Tetrahedron, 39 (1983) 1013. 16. D.R. Storm and D.E. Koshland, Jr., J.Am.Chem.Soc. 94 (1972) 5805. 17. E. Tfirst and J.L. Margitfalvi, to be published. 18. O. Schwalm, J. Weber, J. Margitfalvi and A. Baiker, J. Mol. Structure, 297 (1993) 285.

J.W. Hightower, W.N. Delgass, E. Iglesia and A.T. Bell (Eds.) 11th International Congress on Catalysis - 40th Anniversary

251

Studies in Surface Science and Catalysis, Vol. 101 9 1996 Elsevier Science B.V. All rights reserved.

R a c e m i z a t i o n of ( l S ) - ( - ) - e x o - 2 , 4 - D i d e u t e r o a p o p i n e n e o v e r Pd" E v i d e n c e for an I n t r a m o l e c u l a r 1 , 3 - D e u t e r i u m Shift Gerard V. Smith a'b, Boris Rihter a, Agnes Zsigmond b'c, Ferenc Notheisz b'c, and Mih~ly Bart6k c aDepartment of Chemistry, bMolecular Science Program, Southern Illinois University, Carbondale IL 62901, United States of America CDepartment of Organic Chemistry, J6zsef Attila University, D6m t6r 8, 6720 Szeged, Hungary

1. INTRODUCTION The hydrogenation and isomerization of alkenes can usually be described by the classical Horiuti-Polfinyi mechanism. According to that mechanism, in a deuterium atmosphere, double bond migration incorporates deuterium into the allylic position.

Scheme 1: Horiuti-Poldnyi (classical) mechanism for double bond migration. /

, i

D2

.4 ,

/!

D

a

Early experiments [1 ] using (+)-apopinene and deuterium showed, however, that in the isomerized molecules the deuterium content was very low and the isomerization was much faster than deuterium incorporation into the allylic position. Therefore it seemed probable that isomerization takes place through an intramolecular hydrogen shift. A sigmatropic 1,3hydrogen shift was suggested, in which the allylic e n d o - H shifted (top shift) [2].

Scheme 2: 1,3-top side hydrogen shift

252 Although this type of reaction is symmetry forbidden in an unadsorbed molecule, theoretical calculations showed that in a molecule adsorbed on transition metals, such a shift is allowed [3-5]. Later, other theoretical calculations suggested another type of 1,3-hydrogen shift, one in which the allylic exo-hydrogen is abstracted by the surface from an adsorbed alkene (either 1,2-diadsorbed or 7r-complexed) and the resulting x-allyl species moves over the abstracted hydrogen in such a way that it adds to the former vinylic position and causes, in effect, a stepwise intramoleeular 1,3-hydrogen shift (bottom shift) [6].

Scheme 3: ~-allyl shift mechanism for double bond migration

To distinguish between these two hydrogen shift mechanisms, (1S)-(-)-exo-2,4dideuteroapopinene was constructed as a probe molecule. A top shift of the allylic endo-H will not affect the deuterium content of the molecule and no change should occur in the hydrogen content at any position, but a bottom shift of the allylic exo-D will decrease deuterium in the vinylic position (C2) and increase deuterium in the allylic position (C4).

Scheme 4: Comparison of topside and bottom side hydrogen shifts

H

either classical or ~-allylic shift

ki

D

~

ki

(+)-2d,4d-ap~

D

topside suprafacial sigmatropicshift

2. E X P E R I M E N T A L By a ten-step route from ct-pinene [7], (1S)-(-)-exo-2,4-dideutero-apopinene was synthesized thrice with different deuterium concentrations. The hydrogen contents (19~ 15.1%, and 6.1%) of the molecules at the C2 position were determined by 200 MHz proton

Scheme 5: Numbering in apopinenes 5

4

1

2

(-)-apopinene

253 N M R with the C3 proton used as an intemal standard. Determination of hydrogen content at C4 was not possible on the 200 MHz instrument due to overlapping peaks. Double bond migration (racemization) within (1S)-(-)-exo-2,4-dideutero-apopinene was studied on Pd-black, 0.46% Pd/SiO2, 1.17% Pd/SiO2, 0.4% Pd/AI203, 1.0% Pd/A1203, and PdsoSi20 metallic glass catalysts in both deuterium and hydrogen. Hydrogenations were run on neat apopinene (except where noted) in the liquid phase at one atmosphere of hydrogen or deuterium and were stopped at different percentages to furnish several different percentages of racemization. In the recovered alkenes from these reactions, the hydrogen contents at C2 were determined. The apparatus and the methods for studying the hydrogenaton and for measuring percent dispersion (%D) of the catalysts by hydrogen chemisorption have been described earlier [8].

3. R E S U L T S AND DISCUSSION Essentially, the same results are found for reactions in deuterium and hydrogen (Table 1). As double bond migration (racemization) proceeds, the hydrogen content at C2 increases.

Table 1. The racemization of (1S)-(-)-exo-2,4-dideuteroapopineneon Pd catalysts in deuterium catalysts

racemization, %

0.46% Pd/SiO2 (60.5%D) 38.9 81.1 48.9 56.7 PdsoSi2ometallic glass 10.6 1% Pd/Al:O3 (49.0%D) 30.8 0.4% Pd/Al203 (77.5%D)~ 10. l

symbol + + + *

Pdso-Si20metallic glass

32.2 39.6 1.17% Pd/SiO2 (40%D)@ 33.0

PdsoSi2ometallic glass 62.9 0.46% Pd/SiO2 (60.5%D) 68.0 Pd-black 57.3

,I, x x

% H content in C2 original % final % 19 19 19 19 19 19 19

36.3 50 38.6 41.4 in hydrogen 11.6 in hydrogen 31.9 22.6

15.1 15.1 15.1

33.5 34.4 27.2

6.1 6.1 6.1

35.5 36.1 32.3

poisoned with carbon tetrachloride poisoned with carbon disulfide

The results are easier to see in Figure 1 where representative data are plotted. Dotted lines indicate the change in hydrogen content expected to occur at C2 for a 1,3-bottom shift in a molecule with the particular starting hydrogen content in C2. The solid lines represent the expected absence of change for a 1,3-top shift, and the symbols are those shown in Table 1.

254

60

H content on C2, % 1,3-bottom shift

50 40

..A

t

30 9

.

s

20

1 s

t

1,3-top shift

i

10 0

0

I

1

1

I

1

1

1

1

1

10

20

30

40

50

60

70

80

90

100

racemization, %

Figure 1. Racemization of (1S)-(-)-exo-2,4-dideuteroapopinene on different Pd catalysts.

Although the experiments are imperfect because the percentages of hydrogen at C4 are not known, little evidence exists for the 1,3-top shift. With the exception of the Pds0Si20 metallic glass experiment in hydrogen (fifth experiment from top in Table 1), all data fall near the lines calculated for a 1,3-bottom shift for the respective original percentages of hydrogen at C2. Earlier, with unlabeled apopinenes, we observed that double bond migration appeared to be occurring in deuterium atmospheres without deuterium incorporation into the aUylic position. This led us to postulate an intrarnolecular top shift of hydrogen because a bottom shifting hydrogen should become diluted by surface deuterium if it spends time on the surface [1]. However, these present data suggest a bottom shift is occurring without dilution from surface hydrogen or deuterium and raises several interesting questions. Is the shifting hydrogen or deuterium isolated or protected from the surface hydrogendeuteriurn pool? And how fast does the hydrogen-deuterium pool equilibrate over the surface? To partly answer these questions we consider apopinene. Apopinene is a unique molecule and not representative of most alkenes. Over Pt apopinene undergoes double bond migration approximately eight times faster than cyclohexene [9] and over Pd it undergoes double bond migration as much as seventeen times as fast as addition [Sb]. Assuming the shifting hydrogen or deuterium is not protected from the surface hydrogen-deuterium pool, these results suggest that the rapid double bond migration occurs faster than migration of surface hydrogen or deuterium. It places an upper limit on the rate of migration of surface hydrogen. Apopinene and similar molecules may furnish a unique way to measure rates of migration of hydrogen and deuterium on surfaces under working conditions of catalysts.

255 On the other hand, is the shifting hydrogen protected from incursion by surface hydrogen or deuterium because of some feature of the shift? For example, the shift could occur on the surface but inside the ring such that it is inhibited from migrating outside the reaction sphere (see Scheme 3). In such a case, outside hydrogens or deuteriums would also be inhibited from migrating into the reaction sphere and diluting the shifting hydrogen or deuterium. To answer these and other questions we are continuing to construct likely probe molecules.

4. ACKNOWLEDGEMENT Financial Support from OTKA (1885/91 and 4182/92), NSF INT-8403357, and the USHungarian Joint Fund (No. 177) is gratefully acknowledged. The authors thank Daniel Ostgard for preparing and characterizing some of the catalysts.

REFERENCES

G.V. Smith and D. Desai, Ann. N. Y. Acad. Sci., 214 (1973) 20. G.V. Smith and J. R. Swoap, J. Org. Chem., 31 (1966) 3904. F.D. Mango, Adv. Catal., 20 (1969) 291. F.D. Mango, Coord. Chem. Rev., 15 (1975) 109. A.B. Anderson, J. Chem. Phys., 63 (1975) 4430. A.B. Anderson, D. B. Kang and Y. Kim, J. Am. Chem. Soc., 106 (1984) 6597. B.D. Rihter, Ph.D. Dissertation, Southern Ill. University at Carbondale, 1985. Current address: Aldrich Chemical Co, 230, South Ember Lane, Milwaukee WI. 53233 8. a.G.V. Smith, .~. Molnar, M. M. Khan, D. Ostgard, and T. Yoshida, J. Catal., 98 (1986) 502; b. G. V. Smith, D. Ostgard, M. Bartok, and F. Notheisz, Catalysis of Organic Reactions, (P. N. Rylander, H. Greenfield, and R. L. Augustine, Eds.) Marcel Dekker, Inc., 1988, p 409. 9. G.V. Smith, J. A. Roth, D. S. Desai, and J. L. Kosco, J. Catal.,30, 79 (1973)

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

This Page Intentionally Left Blank On the other hand, is the shifting

J.W. Hightower, W.N. Delgass, E. Iglesia and A.T. Bell (F_xls.) 1 l th International Congress on Catalysis - 40th Anniversary

Studies in Surface Science and Catalysis, Vol. 101 9 1996 Elsevier Science B.V. All rights reserved.

257

F e / M g O catalysts for the selective h y d r o g e n a t i o n of nitriles Gary Bond* and Frank S. Stone School of Chemistry, University of Bath, Bath BA2 7AY, United Kingdom

Fe/MgO catalysts with 5 to 30 mol % Fe have been prepared by impregnation and coprecipitation. Their reducibility has been measured and a comparison made of their Fe ~ surface areas. Catalysts prepared via coprecipitation yielded larger iron areas than those via impregnation. The activity and selectivity of the reduced catalysts for the hydrogenation of propanenitrile at 20-30 bar and 473 K and of ethanenitrile at 1 bar and 508 K have been determined. The most active catalysts are those prepared by coprecipitation and they show high selectivity for primary amines. The activity for ethanenitrile hydrogenation correlates with the iron surface area.

1. INTRODUCTION The catalytic hydrogenation of nitriles is a widely-used method for the industrial production of amines. The reaction, which is catalyzed by a variety of transition metals, also has fundamental interest in that the nature of the metal is known to be a major parameter in influencing the resulting proportions of primary, secondary and tertiary amines. Platinum-based catalysts, for example, direct the hydrogenation in favour of tertiary amines, but among the catalysts most active and selective towards primary amines (often the most desired products industrially) are nickel and cobalt. These metals, and especially nickel, have been frequently studied for the hydrogenation of both low and high molecular weight nitriles il-3], but iron, adjacent to cobalt in the 3d series, has received little attention [4]. Iron, however, has merits in that in comparison with nickel and cobalt it is less prone to catalyze hydrogenolysis and deposit coke. In the industrial context it is also attractive economically. It was accordingly chosen as the catalyst metal for the present investigation. Since the hydrogenation is a metal-catalyzed reaction, it is appropriate to use an oxide support to enhance the dispersion. However, the support, like the metal, needs to be chosen with the desired selectivity in mind. The early view [5] that the selectivity in nitrile hydrogenation is determined largely by the behaviour of the partially-hydrogenated intermediate, the imine R-CH = NH, which can either accept two further hydrogens to form the primary amine or can react with an already-formed amine to start a sequence which * Present address: Department of Chemistry, University of Central Lancashire, Preston, PR1 2HE, United Kingdom

258 leads to the secondary or tertiary amine is still accepted. This latter sequence involves condensation reactions and as such is favoured by acidic conditions. Thus to enhance selectivity towards primary amines, acidic supports are best avoided [2]. Indeed, a basic oxide support is attractive for the added reason that the product amine molecules involved in the formation from the imine of the secondary and tertiary amines, being themselves basic, are less likely to remain adsorbed on the catalyst and be available for onward reaction than would be the case with an acidic oxide. With this in mind, a logical choice for the oxide support for iron in the present work was magnesium oxide, pre-eminent as a basic oxide and readily preparable in high surface area form. Supported iron catalysts are notoriously difficult to reduce [6-8] and thus a substantial fraction of the iron can be expected to remain inactive for the catalysis of hydrogenation. Particular attention has therefore been paid to the preparation of Fe/MgO catalysts by several different methods and examination of their effectiveness in producing metallic iron of adequate specific surface area after reduction in hydrogen. The activity and selectivity for primary amine formation have been determined for the hydrogenation of ethanenitrile (acetonitrile) and propanenitrile.

2. EXPERIMENTAL 2.1 Materials Fe/MgO catalysts with loadings in the range 5-30 mol % Fe were prepared by four methods (a-d) as follows: (a) Impregnation (IP). This method was similar to those used by Boudart et al. [6] and Topsoe et al. [91. A slurry of basic Mg carbonate (Fluka) in 160 ml of distilled water was heated with stirring to 340 K and 160 ml of an aqueous solution of ferric nitrate (BDH) preheated to 340 K was rapidly added. After stirring for 30 min at 340 K the impregnated slurry was filtered, washed and vacuum-dried at 360 K before being calcined in air at 773 K. Different iron loadings were obtained by varying the concentration of the Fe(NO3) 3 solution. The calcined precursor was reduced in a static system (see Sec. 2.2) in hydrogen at ca. 30 Torr for 24 h at 553 K, 24 h at 623 K and 24 h at 693 K, the water vapour produced being frozen out in a 77 K trap. The resulting catalysts are designated as IPx, where x = 100 tool fraction Fe/(Fe + Mg). (b) Carbonate coprecipitation (CCP). Solutions of Fe(NO3) 3 and Mg(NO3) 2 (BDH), each 1M, were mixed in appropriate ratios and heated to 360 K. 1M (NH4)2CO 3 solution was added dropwise until precipitation was complete. After washing and filtering, the precipitate was dried at 373 K and calcined at 773 K. This precursor was reduced as for (a) above. These catalysts are designated CCPx, where x is the mol % Fe as before. (c) Hydroxide coprecipitation (OH). A mixture of Fe(NO3) 3 and Mg(NO3) 2 solutions as in (b) was in this case added dropwise into a stirred vessel containing ammonia solution at 303 K. The pH was maintained constant at pH 11 by further addition of ammonia. The resulting precipitate was washed, dried and calcined as in (b) above, and was reduced as in (a). These catalysts are designated OHx, where x denotes the mol % Fe as previously.

259 (d) Hydrotalcite method (HT).

100 ml 1M Fe(NO3) 3 and 300 ml 1M Mg(NO3)2 was

pumped simultaneously with 400 ml of 0.125M (NH4)2CO3 into a stirred reaction vessel at a rate of 70 ml h -1. The temperature of the vessel was kept constant at 323 K and the pH was maintained constant at 9.8 by the addition of NH4OH. The resulting precipitate of magnesium iron hydrotalcite (pyroaurite, nominally Mg6Fe2(OH)16CO3) was dried, calcined and reduced as in (c) above. This catalyst is designated HT. 2.2 Catalyst characterization The reduction procedure described above was carried out in a static system in order to facilitate determination of the extent of reduction in situ. The reduction was monitored gravimetrically using a microbalance (CI Electronics MK II) and volumetrically by measuring the decrease in hydrogen pressure. Total surface areas were determined by the BET method using nitrogen at 77 K. Carbon monoxide chemisorption was used to estimate the surface area of metallic iron after reduction. The quantity of CO chemisorbed was determined [6] by taking the difference between the volumes adsorbed in two isotherms at 195 K where there had been an intervening evacuation for at least 30 min to remove the physical adsorption. Whilst aware of its arbitrariness, we have followed earlier workers [6,10,11] in assuming a stoichiometry of Fe:CO = 2.1 to estimate and compare the surface areas of metallic iron in our catalysts. As a second index for this comparison we used reactive N20 adsorption, N20(g) ~ N2(g) + O(ads), the method widely applied for supported copper [12]. However, in view of the greater reactivity of iron, measurements were made at ambient temperature and p = 20 Torr, using a static system. Bulk characterization of calcined precursors and reduced catalysts was carried out by X-ray diffractometry using Cu Kcx radiation. Reduced catalysts were first passivated by exposure to N20 as described above. Line-broadening analysis was carried out on the Fe(110) reflection to obtain the iron particle size. Overlap with the MgO(200) reflection limited its usefulness to the more highly-loaded catalysts. 2.3 Catalysis Hydrogenation tests were carried out using an autoclave reactor (ICI, Wilton, UK) in the case of propanenitrile and a flow reactor operating at atmospheric pressure for ethanenitrile. (a) Autoclave reactor experiments. The procedure was as follows. 0.5 g of passivated catalyst was activated by reduction in H 2 at 573 K and after transfer via a nitrogen-purged glove box was installed in the reactor with 10 ml propanenitrile. The reactor was pressurized with H 2 and heated to 473 K. The subsequent fall in pressure from a maximum of about 30 bar (20% nitrile : 80% hydrogen) as hydrogenation proceeded was monitored for 20 hours. The liquid product was subjected to GLC analysis. (b) Flow reactor experiments. For a more precise comparison of the activity and selectivity of the Fe/MgO catalysts, the hydrogenation of ethanenitrile was investigated. A laboratory flow reactor operating at 1 bar pressure was employed. A sample (0.5 g) of passivated catalyst was loaded into the reactor and after activation (as above) the temperature was lowered to 508 K (unless otherwise indicated) and the H 2 flow (at a chosen

260 value in the range 10-50 ml min -1) was diverted through a saturator containing ethanenitrile held normally at 273 K before passing over the catalyst. After achieving steady state conditions the exit gas was sampled and analyzed by GLC. Space velocity, temperature and the respective concentrations of CH3CN and H 2 in the feed were separately varied in order to obtain the rate of of reaction, activation energy and reaction orders in ethanenitrile and hydrogen for individual catalysts.

3. RESULTS 3.1 Characterization of catalyst precursors Four catalysts with nominal compositions of 5, 10, 20 and 30 mol % Fe were prepared by each of the methods (a), (b) and (c) described in Sec. 2.1, together with one by method (d). Actual compositions were close to the nominal ones (see later). in the OH series, two phases were detectable by XRD in the dried precipitate. One was a phase with the pyroaurite structure, carbonate having presumably arisen from atmospheric CO 2, and the other brucite, Mg(OH)2 , to which pyroaurite is closely related structurally. For both the CCP and IP series, the only structure identifiable in the dried precipitate was that of magnesium hydroxycarbonate. X-ray analysis of the calcined precursors showed MgO together with y-Fe20 3 in the case of the OH series and HT, but ~-Fe20 3 with the CCP and IP series. MgFe20 4 spinel was also detectable in some cases. 3.2 Characterization of the reduced catalysts Iron was present as Fe 3+ in the calcined precursors. For all the catalysts the reduction procedure described in Sec. 2.1 resulted in incomplete reduction of the Fe 3+ to metallic iron. This is in agreement with the findings of previous authors [6,11 ]. The individual percentage reductions of Fe 3+ to Fe ~ as determined by the separate gravimetric and volumetric measurements (See. 2.2), are shown in Table 1. The values are calculated on the assumption that all the Fe 3+ is reduced to F e 2 + prior to the onset of reduction to Fe ~ There is good agreement between the two methods. Table 1 also records the actual Fe/(Fe + Mg) ratio in the catalysts as determined by atomic absorption spectroscopy (AAS) on the calcined precursors. We present next the data which relate to the surface area of the reduced catalysts. All the catalysts had a total (BET) surface area in the range 50-100 m2g -1, the lower values being for those with the highest content of iron, but the more important catalyst characterization parameter for the present work is the surface area of the metallic iron, studied here by CO chemisorption and N20 reaction. The CO chemisorption was evaluated using the range of the adsorption isotherm between 75 and 150 Torr, where the difference in volume adsorbed between the first (chemisorption and physical adsorption) and the second (physical adsorption) isotherms was constant. Expressing this volume as a number of molecules and assuming 2 Fe o atoms per chemisorbed CO, followed by conversion to surface area on the basis of 1.2 x 1019 atoms Fe per m 2 (an average for the (110), (100) and (111) faces ofbcc Fe), we obtained the Fe ~ surface areas shown in Table 2. Also shown in this table are

261 values derived from N20 reaction, where the amount of N20 decomposed was determined gravimetrically with the microbalance and where a stoichiometry of 2.5 molecules N20 reacted per surface Fe~ was assumed. This value supposes a lesser oxidation depth than the "approximately 4 monolayers of metallic iron oxidized upon prolonged exposure to N20 at room temperature (and 280 Tort)" reported by Vogler et al. [ 13], but we used a pressure of only 20 Torr to drive the reaction. Fe:N20 stoichiometry of 1:2.5 gives good agreement with the CO data. However, more significant is the agreement of the variation in Fe ~ surface area within the group of 13 catalysts, taking the data from CO chemisorption on the one hand and comparing them against the data from N20 reaction on the other. Table 1 Effect of loading and method of preparation on the reducibility of iron in Fe/MgO catalysts Reduction to Fe ~ (%) Catalyst

Fe/(Fe + Mg) by AAS analysis

Gravimetric determination

Volumetric determination

IP5 IP10 IP20 IP30

4.9 10.0 19.8 29.9

24 32 37 53

24 31 35 49

CCP5 CCP10 CCP20 CCP30

4.8 10.0 18.9 30.8

20 26 59 69

22 25 54 65

OH5 OH10 OH20 OH30

4.8 9.5 18.5 29.2

27 39 54 57

26 37 51 54

HT

22.6

42

40

For the catalysts with highest iron loadings, the average size of the iron crystallites (d) was obtained from the line-broadening of the Fe(ll0) reflection, the instrumental broadening having been determined with a sample of sintered iron powder (Koch-Light). The results (Table 3) show that the OH catalysts have the smallest particle sizes. An estimate of the average particle size can also be made [6] from the relation d(nm) = 0.85/D, where D is the dispersion, the ratio of surface Fe ~ atoms to total Fe ~ atoms, the

262 Table 2 Iron surface areas from CO chemisorption and N20 reaction Fe surface area (m2g-leat) Catalyst by CO chemisorption

by N20 reaction

IP5 IP10 IP20 IP30

1.5 2.1 3.4 5.6

1.3 2.5 3.1 5.5

CCP5 CCP10 CCP20 CCP30

2.9 4.4 5.0 5.3

3.1 4.7 5.1 5.7

OH5 OH10 OH20 OH30

5.0 9.5 10.1 7.2

5.1 9.8 10.2 7.0

7.5

7.4

HT

Table 3 Iron particle size from X-ray line broadening and percentage-reduction/CO-chemisorption

Iron particle size (nm) Catalyst X-ray line broadening

reduction and CO chemisorption

IP20 IP30

47 48

28 33

CCP20 CCP30

45 46

29 47

OH20 OH30

39 36

13 27

263 former being given by the CO chemisorption results and the latter by the loading and percentage reduction data. The d values obtained in this way are also shown in Table 3. 3.3 Hydrogenation of propanenitrile The catalysts were studied for their activity and selectivity in propanenitrile hydrogenation in simple batch experiments at 473 K using the autoclave reactor. A rough comparison of activity was made by taking as a measure of the rate of hydrogenation the slope of the tangent to the curve of H 2 uptake versus time at p = 25 bar. The results showed clearly that the activity of the catalysts prepared by impregnation (IP) was less than that of those prepared by coprecipitation. The selectivity was obtained by carrying out GLC analysis of the products after 20 hours of reaction. It was found that all the catalysts prepared by coprecipitation gave selectivities to the primary amine in excess of 60%, whereas the IP catalysts were again different, giving a selectivity of less than 60% whatever the loading. An informative overview of the activity and selectivity results is obtained by presenting them as a matrix (Fig. 1), which serves also to show the general trend of higher selectivity being associated with higher activity (the arrow at OH10 in Fig.1 is to indicate that it has an off-scale rate, 0.2 bar rain-l). The liquid products other than the primary amine were the secondary and tertiary amines, the latter in very small yield (--1%), 1-propylamino-propene (the most abundant by-product) and 1,1-dipropylamino-propene (~1%). The gas remaining after cooling the reaction mixture was sampled with a gas syringe in a few cases (reactions using OH catalysts) to test whether any hydrogenolysis had occurred. The results indicated that about 4% of the nitrile which had reacted had undergone hydrogenolysis: GC analysis showed that the principal gaseous products from this hydrogenolysis were propane (50%), propene (35 %) and ethane (10%). 0.04

100

8O" ~--

OH5 OH3 0 O t ccP20 ~

. 9

HT 9CCP30

l

OH10

-~" 0 . 0 3

OH20

.,.

E "7 -- 0 . 0 2

. . . . . . C C P 10 .--IP30 > --IP20 o 9

60":e

OH3

0

E

40-

=: 0 . 0 1

20 9 IP5 0

0

9

........ I ......... t 0.05 O. 10

l 0.15

Rate (bar min-1) Fig. 1. C2HsCN hydrogenation at 473 K. Selectivity to C2H5CH2NH 2 vs. activity.

--

0

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

1. . . . . .

5

I

10

Fe ~ area (m2g- 1 ) Fig.2. CH3CN hydrogenation at 508 K. Activity vs. Fe~ surface area.

264 3.4 Hydrogenation of ethanenitrile The aim of the experiments with ethanenitrile (CH3CN) was to study hydrogenation activity more quantitatively than was possible with the autoclave reactor. Catalysts IP30, OH5, OH20, OH30 and HT were selected for investigation. Steady-state conversion to ethylamine was measured at 508 K and atmospheric pressure for various space velocities, the hydrogen flow rate being varied from 10 to 50 ml rain -1. Plots of conversion vs. reciprocal space velocity were linear. The respective gradients, which are measures of the reaction rates for hydrogenation, are plotted in Fig.2 as a function of the iron surface areas from Table 2. The rates correlate with the surface areas: note particularly the low rate for the IP catalyst, in spite of its high loading. The selectivity for hydrogenation to the primary amine was high in all cases, with only minor amounts of secondary amine and no tertiary amine. There was only slight evidence of hydrogenolysis. The activation energy of the reaction, determined for catalysts OH20 and HT over the range of temperature from 450 to 510 K, using a flow rate of 20 ml rain -1 for H 2 and with the CH3CN saturator at 273 K, was 59 kJ mol -] in both cases. The order of reaction with respect to ethanenitrile was measured for catalyst OH5 at 508 K and for HT at 483 K by maintaining the temperature of the saturator at different temperatures between 273 K and 295 K at 20 ml rain -] H 2 flow rate. From plots of log rate vs log [CH3CN ] the order derived was -0.25 + 0.08 for OH5 and 0.28 + 0.04 for HT. The order of reaction for H 2, obtained by diluting the flow of H 2 with He whilst keeping the saturator at 273 K, was measured for HT at 483 K and found to be 1.9 _4- 0.3.

4. DISCUSSION 4.1 Reducibility and metal surface area The characterization results show that the method of preparation of Fe/MgO exerts a significant influence on the amount of iron which can be reduced to Fe ~ (Table 1). The percentage reduction increases with the loading in each series (IP, CCP, OH), but the coprecipitation with NH4OH, which gives brucite-like structure to the precipitate and y-Fe.20 3 on calcination (the OH series) provides more easily reduced precursors than those obtained via precipitation of carbonate-rich hydroxycarbonate. This difference is even more strongly reflected in the results on iron surface areas (Table 2) where the OH series catalysts (and HT) show consistently higher values. As already indicated, the difficulty of reducing supported iron in hydrogen is well-known [6, 8,11 ]. It probably arises from a combination of causes, the two most important of which are a strong interaction with the support 16,8] and reoxidation or inhibition by water vapour in the pores of the oxide [14]. With MgO as support, there is undoubtedly a strong tendency for iron, especially at the F e 2 + stage of reduction, to be present at least in part as FeO-MgO (FexMgl_xO) solid solution [6,8]. This need not be deleterious to the ultimate formation of finely-divided iron, provided the method of preparation has led to a solid solution in which the Fe 2+ ions are well-distributed. The iron particles are limited in size

265 by the extraction and reduction of Fe 2+ ions being restricted to a relatively small depth of solid solution, as in the analogous case of reduction of high surface area NiO-MgO and CoO-MgO [15]. We suggest that this is working to greatest effect with the OH series of catalysts. The average size of the iron crystallites evaluated using the d = 0.85/D formula for OH5 is 3.8 nm as compared with 11.4 nm for IP5. In the IP series, the Fe 2 + ions in the FeO-MgO may be clustered, as noted in previous work [6]. The presence of some very finely-divided iron also in the catalysts with higher loadings is strongly indicated by the comparison between XRD-derived and dispersion-derived sizes in Table 3. The disparity, also evident in previous work [6,11], arises because very fine particles tend to be 'lost' in line-broadening analysis but fully contribute in chemisorption. The OH catalysts again show up favourably, exhibiting the lowest sizes by line-broadening analysis. Significantly, the above disparity is greatest with OH20, suggestive of a high proportion of very small Fe particles. At 30 mol % Fe, the support is less effective in dispersing the iron due to the high loading. m

4.2 Fe-catalyzed hydrogenation of nitriles The results show that iron supported on MgO is an effective catalyst for the selective hydrogenation of nitriles to the primary amines. However, the selectivity is impaired if the catalyst has low activity (Fig. 1). The inference from the experiments with propanenitrile is that the imine intermediate, which one may expect to be strongly chemisorbed and activated on iron, needs to react rapidly with adsorbed hydrogen if condensation reactions to form secondary and tertiary amines are to be avoided. The availability of adsorbed hydrogen may well be enhanced by high dispersion and a plentiful supply of finely-divided iron, conditions which are manifestly weakest in catalysts IP5, IP10 and IP20, and especially so in IP5 (Table 2), and it is these catalysts which are consistently lowest in activity. Catalysts of the CCP series, on the other hand, perform well, but in general the most effective are those of the OH series. These are the catalysts with the highest iron surface areas. Verhaak et al. [2] have shown convincingly that the reactions leading to the secondary and tertiary amines involve the support. The by-product 1-propylaminopropene which we observed is presumably formed via attack of propylamine on propylimine, followed by elimination of NH 3, as shown below: +H 2 CH3-CH2-C-~N ~

CH3-CH2-CH=NH / + NH2-CH2-CH2-CH3 NH 2 ,b I

CH3-CH2-CH2-N-CH-CH2-CH 3 H / - NH3 !

CHB-CH2-CH2-NH-CH =CH-CH 3

266 1-Propylaminopropene is thus on a logical route to dipropylamine, which can form directly by hydrogenation of its C = C bond. This condensation-elimination reaction may indeed involve the support [2] or at least the metal-support interface. The kinetic results with ethanenitrile support the above conclusions. The plot of Fig.2 shows that the hydrogenation to the primary amine is metal-catalyzed with a rate which is proportional to the metal surface area. The fact that the plot does not extrapolate to the origin can be interpreted as indicating that some of the iron measured by exposure to CO or N20 is not accessible to the reactant nitrile. The activation energy (59 kJ mol-1), which is the same for OH20 and HT, is similar to the value of 52-55 kJ mo1-1 reported by Verhaak for the hydrogenation of CH3CN on Ni and Co [16]. The near-zero order in CH3CN confirms strong chemisorption of the nitrile and the imine on the iron surface. The authors thank the Science and Engineering Research Council and also ICI Chemicals and Polymers Ltd for their support of this work.

REFERENCES 1. J. Volf and J. Pa~ek, Stud.Surf.Sci.Catal., 27 (1983) 105. 2. M.J.F.M. Verhaak, A.J. van Dillen and J.W. Geus, Catal.Lett., 26 (1994) 37. 3. F. Medina, P. Salagre, J.E. Sueiras and J.L.G. Fierro, Appl.Catal., A, 92 (1993) 131; 99 (1993) 115. 4. Ullmann's Encyclopaedia of Industrial Chemistry, Vol.A2, VCH, Weinheim, 1985. 5. J. yon Braun, G. Blessing and F. Zobel, Chem.Ber., 36 (1923) 1988. 6. M. Boudart, A. Delbouille, J.A. Dumesic, S. Khammouma and H. Topsae, J.Catal., 37 (1975) 486. 7. A.J.H.M. Kock, H.M. Fortuin and J.W. Geus, J. Catal., 96 (1985) 261. 8. D.E. Stobbe, F.R. van Buren, A.W. Stobbe-Kreemers, A.J. van Dillen and J.W. Geus, J. Chem. Soe., Faraday Trans., 87 (1991) 1631. 9. H. Topsae, J.A. Dumesic, E.G. Derouane, B.S. Clausen, S. Morup, J. Villadsen and N. Topsae, Stud.Surf.Sci.Catal., 3 (1979) 365. 10. H.J. Jung, M.A. Vannice, L.N. Mulay, R.M. Stanfield and W.N. Delgass, J.Catal., 76 (1982) 208. 11. S. Mousty, B.S. Clausen, E.G. Derouane and H. Topsae, Stud.Surf.Sci.Catal., 16 (1983) 385. 12. J.J.F. Scholten and J.A. Konvalinka, Trans.Faraday Soc., 65 (1969) 2465. 13. G.L. Vogler, X.Z. Jiang, J.A. Dumesic and R.J. Marion, J.Catal., 89 (1984) 116. 14. A. Baranski, M. ~agan, A. Pattek, A. Reitzer, L.J. Christiansen and H. Topsae, Stud.Surf.Sci.Catal., 3 (1979) 353. 15. J.G. Highfield, A. Bossi and F.S. Stone, Stud.Surf.Sci.Catal., 16 (1983) 181. 16. M.J.F.M. Verha~, Thesis, University of Utrecht, 1992. ISBN 90-393-0243-X.

J.W. Hightower,W.N. Delgass,E. Iglesia and A.T. Bell (Eds.) 11th International Congress on Catalysis - 40th Anniversary

Studies in Surface Science and Catalysis, Vol. 101 9 1996ElsevierScience B.V. All rights reserved.

267

Selective synthesis of ethylenediamine from ethanolamine over modified H-mordenite catalyst K. Segawa a, S. Mlzuno, " a M. S u g .m r aa, and S. Nakata b aDepartment of Chemistry, Faculty of Science and Technology, Sophia University, 7-1 Kioi-cho Chiyoda-ku, Tokyo 102, Japan* bR & D Center, Chiyoda Corp., 13 Moriya-cho, Kanagawa-ku, Yokohama 221, Japan The synthesis of ethylenediamine (EDA) from ethanolamine (EA) with ammonia over acidic types of zeolite catalyst was investigated. Among the zeolites tested in this study, the protonic form of mordenite catalyst that was treated with EDTA (H-EDTA-MOR) showed the highest activity and selectivity for the formation of EA: at 603 K, W/F=200 g h tool", and NH3/EA=50. The reaction proved to be highly selective for EA over H-EDTA-MOR, with small amounts of ethyleneimine (EI) and piperazine (PA) derivatives as the side products. IR spectroscopic data provide evidence that the protonated EI is the chemical intermediate for the reaction. The reaction for the formation of EDA from EA and ammonia required stronger acidic sites in the mordenite channels for higher yield and selectivity. I. I N T R O D U C T I O N Ethylenediamine (EDA) is used in a wide variety of applications in industry. The applications for EDA include uses as chelating reagents, surfactants, fabric softeners, lubricating oil additives, fungicides, insecticides, and resinous polymers. EDA is made by amminolysis of ethylene dichloride with ammonia [ E D C process], or by reductive amination of ethanolamine (EA) with hydrogen and ammonia fMEA process] on a commercial basis. However, the E D C process h a s numerous drawbacks: the production of by-product sodium chloride and the cost of corrosion-resistant equipment for such a process [1]. On the other hand, the M E A p r o c e s s includes high pressure reactions (10-20 MPa) over transition metal catalysts, and shows lower selectivity for EDA [2]. In view of the wide utility of EDA, there has been substantial work done in the preparation of EDA t h a t involves the reaction of ethylene oxide with ammonia, EA with ammonia, or ethylene glycol with ammonia. Certain zeolite catalysts have been studied in the field ofEA amination to suppress the formation of bulkier by-products, and generate enhanced yields of EDC [3]. The development of a selective catalyst is to reduce utility cost and/or to expand the plant capacity. The process requires a solid acid catalyst; objectives of the process have been to maximize conversion of the

268 organic substrate with ammonia to form EDA and to maximize selectivity. In this study, as an alternative process for EDA synthesis, acid-catalyzed amination of EA under atmospheric pressure, has been studied; such a process is much more environmentally benign than the present industrial processes. A s i ~ i f i cant advantage afforded by the treatment of EDTA in sodium form of mordenite followed by ion-exchange to protonic form is that the process is extremely effective in producing EDA in high conversion while at the same time being highly selective to EDA. This is contrast to many of the subsequent processes [EDC process and MEA process] which sacrificed selectivity for EDA in favor of conversion.

2. E X P E R I M E N T A L

Preparation of catalysts. Zeolites (JRC-Z) and silica-alumina (JRC-SAL-2, Si/Al=5.3) samples were supplied by the Catalysis Society of J a p a n (JRC: J a p a n Reference Catalysts). Three different types of zeolites were studied: Na-FAU (faujasite, JRC-Z-Y5.6, Si/Al=2.8), Na-MOR (mordenite: JRC-HM10, Si/Al=5.0), and Na-lVIFI (ZSM5, JRC-Z5-25, Si/Al=12.5). K-LTL (Linde type L, HSZ500KOA, Si/AI=3.0) and K-CHA (K-chabazite, Si/AI=2), were supplied by TOSOH and by Air Products and Chemicals, Inc., respectively. Some Na-MOR samples were treated with an aqueous solution of H4EDTA (ethylenediamine-tetraacetic acid) under reflux at 383 K for 4 h. After cooling and filtration, the samples were washed with distilled water and dried at 373 K, and were then calcined at 773 K (EDTA-MOR, Si/A1=5.2-11.2). Add-type zeolites were prepared by ion-exchange of Na-form or K-form of zeolite with aqueous solution of NH4NO3; ion-exchanged samples were dried at 373 K for 24 h and then calcined in a furnace at a constant temperature increase (1 K rain i ) from 373 K to 773 K and kept at 773 K for 5 h. Catalytic reactions. The reaction was carried out at 543-643 K by using a flow reaction system with a mixture ofEA, NH3, and N2 in the ratio of 1/50/25 at atmospheric pressure. The flow rate of the mixture gas was 76 cm 3 rain "1. Prior to the reaction, the catalyst was calcined at 773 K under 02 flow for 2 b_ The reaction products were analyzed by an on-line gas chromatograph (FID) which was equipped with a 30-m capillary column (TC1701). Adsorption measurements. Chemisorption of base molecules (NH 3 and EA) on acidic zeolites was confirmed by IR spectroscopy and high-temperature microcalorimetry. A vacuum-tight IR cell with KBr windows was designed to fit an infrared spectrometer (270-30, Hitachi) and to be attached to a vacuum system (10 -4 Pa). The cell was arranged such that the zeolite wafer could be lowered into slots between the optical windows, and withdrawn upward by the action of a magnet into the heated portion for the pretreatment and adsorption of NHs and EA. After evacuation at 773 K for lh, the zeolite sample was cooled to 373 K before adsorption of the base molecules to be studied. IR spectra were obtained at room temperature. High-temperature micro-calorimetry of NH3 on zeolite catalyst was obtained at 473 K by the calorimeter (HAC-450G, Tokyo Rikou). Each sample (1.5 g) was charged to the calorimeter, and it was evacuated at 673 K for 4 h. NH3 (15 mmol g-i) portions were admitted dose by dose at 473 K.

269 3. R E S U L T S AND D I S C U S S I O N The catalytic activity and selectivity of EDA synthesis from EA and N H 3 over various zeolite catalysts are shown in Table 1. Among the various solid adds, the protonic form of mordenite (H-MOR) catalyst and the protonic form of mordenite catalyst that was treated with EDTA (H-EDTA-MOR), showed the higher selectivity for EDA. Over H-EDTA-MOR; the selectivity of EDA was about 72 % at 96 % of EA conversion with the presence of a n excess a m o u n t of ammonia (NHz/EA=50). Small amounts of EI and piperadine derivatives (PA) were formed as by-products. PA derivatives are mainly piperazine, amirtoethylpiperadine, and 1 , 4 - d i a z a b i c y c l o - o ~ e (DABCO). When the reaction t a k e s place over some other solid acid catalyst, such as amorphous silica-alumina, 100 % of EA converted to EI, PA derivatives, and other oligomers including aminoethylaminoethanol (Others: polyamines). On H-CHA (chabazite), H-FAU (faujasite-Y), and H-LTL (L-type), only a small a m o u n t of EDA was formed. The major product was EI and PA, or other polyamine oligomers were formed. On HMFI (ZSM5) catalyst, the major products were PA and other higher polyamines.

Table 1 Catalytic Activities and Selectivities of EDA Synthesis* over Various Zeolite Catalysts Catalyst*** SiO2-AI203 H-CHA H-FAU H-LTL H-MOR H-EDTA-MOR H-MFI

Pore Size

Conversion

Selectivity**/%

Si/A1

/nm

/%

EDA

EI

PA

Others

5.3 2.2 2.8 3.0 5.0 6.9 12.5

--0.38 0.74 0.71 0.70 0.70 0.54

I00 4 6 15 42 96 23

1 4 13 23 77 72 36

7 69 76 35 7 0 2

49 13 6 21 9 18 31

43 14 5 21 7 10 31

*Reaction Conditions: Temperature=603 K, W/F=200 g h mol"1, NH3fEA=50 ** EDA: ethylenediamine, EI: ethyleneimine, PA:piperazine derivatives *** CHA: chabazite, FAU: faujasite (Y), LTL: Linde type L, MOR: mordenite, EDTA-MOR: Namordenite was treated with EDTA and then ion-exchanged to H-form, MFI: ZSM5.

The results (Table 1) suggest that the selective synthesis of EDA from EA and NH 3 requires stronger acidic sites in the mordenite channels to suppress the formation of bulkier PA derivatives and other polyamines (Others in Table 1). The t r e a t m e n t by EDTA on Na-MOR and then ion-exchange to protonic form gives significant improvements for EDA synthesis. Figure 1 shows the catalytic activities and selectivities for EDA synthesis over H-MOR (Figure 1A) and HEDTA-MOR (Figure 1B), as a function of time on stream. Over H-MOR catalyst, the catalytic activity was much lower than t h a t of H-EDTA-MOR catalyst, and some deactivation occurred with increasing time on stream. On the other hand,

270 even if the catalytic activity over H-EDTA-MOR was higher than t h a t of other catalysts, no significant deactivation occurred aIter 3 hours on stream under the whole range of the reaction conditions tested in this study. The selectivity to EDA over H-EDTA-MOR is much higher than t h a t over H-MOR.

"

100

I

I

I

A 80

O

r'-!

60

-

" 40 .9

-

(D

I

I

- 100

EDA

-

._>

I

r-I

~

~

L~

-

I

.

B i.

i

I

I

i

_ Conversion

I

__~_

_

80

I-!

m

W

!--1

LJ

EDA 60

(,D e~

" 0

20

"0.

PA

_i

Others M

El 0

1

2

40

Conversion

-

20 ~

I

3

m

I

4

m

P~~iI~II~

m

I

I

5

6

Others . =

0 7

. 0

1

9 O

_ '~J

_

;,.,,--

~

_A

A

_A

_A

2

3

4

5

6

Time on s t r e a m / h

Time on s t r e a m / h

Figure 1. Catalytic activities and selectivities for EDA synthesis over (A) H-MOR (Si/AI=5.0) and (B) H-EDTA-MOR (Si/AI=6.1). Reaction conditions: temperature=583 K, NH3~A=60, PEA-1.2 kPa.

100

80

Or)

-

I

D

.

.

I

!

- 0 ' "

Conversion

QJb

•

- _

60

r

._o

40

-

o O

==.Others ~ --

e-.

20

-

lib /'%

-

.,,.,,

~.

PA

%./

v

B A,

0 4.0

6.0

8.0

, 10.0

~_ 12.0

Si/AI

Figure 2. The effect of dealumination by EDTA over MOR (Si/AI=5.0) for ethylenediamine synthesis. Reaction conditions: temperature=603 K, NH3/EA=50, PEA=I.4 kPa.

271 Figure 2 shows the catalytic activity and selectivities after treatments by EDTA, as a function of Si/A1 ratios. The treatment of zeolite with H4EDTA has been shown to dealuminate the zeolite framework. The bulk ratio of (Si/A1) in the zeolite was determined by XRF analysis. The reaction of EA with excess amounts of NH 3 over H-MOR provided higher selectivities for EDA (61%) but, activities are low (53 % conversion). In an attempt to improve activity and selectivity by modification of catalyst structure, moderately dealuminated H-MOR catalysts were prepared by EDTA treatments. When the Si/A1 ratios become above 6 from Si/AI=5.0 (H-MOR) of parent mordenite, a large enhancement of the catalytic activities was observed (87-93 % conversion). The selectivity of EDA was not significantly changed over H-EDTA-MOR in comparison with H-MOR, except over higher dealuminated catalysts (Si/AI=ll.2). Due to the channel structure of MOR, EDA is formed over the protonic sites that are located inside the main channels of MOR, and cyclization to PA and polyamines (higher oligomers) might be formed non-selectively on the external surfaces of MOR crystals.

160 L 140

I

I

I .........

.................................. t

t

f

(SI/AI)

120

O H-MOR (5.0)

r-

.g

g

0r "0 0 -r

---! ..........

~00 80

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

i! ...... ......

..........

! .......

oo .......... i ........... 40

....

0.0

0.5

i ..........

O H-EDTA-MOR

(5.9)

9 H-EDTA-MOR

(6.1)

A H-EDTA-MOR

(8.0)

i .......... I t'

i

i

1.0

1.5

2.0

2.5

NH3 adsorbed/mmol g-1

Figure 3. High-temperature micro-calorimetry ofNH 3 on H-MOR (Si/AI=5.0) and H-EDTA- MOR (Si/A1=5.8-8.0): NH 3 adsorbed at 473 K. H-MOR has larger numbers of stronger acidic sites than those of other zeolites [5]. Figure 3 shows high-temperature micro-calorimetry of NH 3 over HMOR (Si/AI=5.0) and H-EDTA-MOR (Si/A1=5.8-8.0). We preferred the temperature at 473 K for calorimetry measurements to collect precisely the information of the stronger acid sites. All micro-calorimetric curves decreased with increasing coverage of NH 3 on zeolites. The results suggest that the acid strength and acid amount of H-MOR are higher than the H-EDTA-MOR. The MOR samples after being dealuminated by EDTA (H-EDTA-MOR) show similar acid site distributions for the range of Si/A1 ratios from 5.8 to 8.0.

272 100. -

.

I

'

I

'

A 80

w

9...@--"

-zl00 !

,...e""'"

-

80

,

40

'

. .

,,'

40

~-

o 20 rO

0

-

U _

0

100

200

,

_A 300

O

,

A

400

Contact time (W/F) /g h tool-1

PAl '

-

hers

20

~

Oon,,or ,on

~

~ 0 500 0

~

P 20

A 40

I

l 60

80

100

NH3/EA

Figure 4. EDA synthesis on H-EDTA-MOR (Si/Al=6.1) as a function of contact time (A) and as a function of partial pressure of NH 3. Reaction conditions: temperature=583 K, NH3/EA=0-80, PEA=I.4 kPa.

Deeba and coworkers [3] reported that the H-MOR dealuminated by acid leaching showed higher selectivity for EDA at lower conversions: about 60 % selectivity at 30 % conversion (NH3~A=16). However, at higher conversions, selectivity for EDA was not as high. In this study, if the reaction conditions included longer contact time (W/F=500 g h tool-Z), conversion exceeds about 95 % of EA with 80 % selectivity to EDA. The time courses of EDA synthesis over HEDTA-MOR catalyst at 583 K (PEA=I.4 kPa, NH3/EA=50) are shown in Figure 4A. The initial product of reaction was EI, and the formation of EDA followed. However, the selectivity of EDA did not exceed 85 % at higher conversion region. The selectivity of PA derivatives and higher polyamines increased with increasing contact time. The catalytic performance over H-EDTA-MOR as a function of partial pressure of N H 3 is also shown in Figure 4B. The activity rises with increasing partial pressure of NH3. Importantly, the formation of higher polyamines and EI is extremely retarded at higher partial pressure of NH3, whereas E D A is the major product. The results suggest that the intramolecular condensation of E A occurred at the initialstage of reaction to produce El intermediate, then adsorbed El was activated by the stronger protonic acid sites to produce EDA. The activation of El over the stronger acidic sitesis the rate-determining step for this reaction; such adsorption species m a y be converted to E D A with excess amounts of N H3. Intermolecular condensation of E A to form El, P A derivatives,and higher polyamines m a y occur at weaker acidic sites,which m a y be located on the external surfaces of mordenite crystals. W h e n E A adsorption occurred under N 2 flow at 583 K on H - E D T A - M O R catalyst without N H 3, El, higher polyamines, and P A were formed (Stage A in Figure 5). But, when an excess amount of N H 3 (NH3/EA=50) was introduced in

273 the reaction system at Stage B, EDA formed selectivity increased as did the conversion. At Stage C (same reaction conditions at Stage A), the formation of EDA was suppressed completely, and EI, higher polyamines, and PA reappeared again.

100

>,

._>

-

80

9

I

A

'

'~

El

"~ 60 (D =o

-

t,.

O

0

-0"

20

-

0

EDA

El

-"

-

"@ . . . .

0,,

Conversion -O ....

e"

'

Conv ers ion

e...

._o 40 O3

i

c

-

O

'

P A' -

e.-

~

Conversion ~

Others

2

4

" --O. . . .

6

@--

8

Time on stream/11

Figure 5. Transient responses for EDA synthesis over ; H-EDTA-MOR 1 (Si/AI=6.1). Reaction conditions: temperature=583 K, W/F=200 g h mol'-, PEA=0.95 kPa, NH3/EA=0 at Stage A and C, NH3/EA=50 at Stage B.

The adsorption studies of EA or ammonia on H-EDTA-MOR (Si/Al=6.1) by IR spectroscopy (Figure 6) suggested that the reaction may proceed through ammonio-ion of EA over protonic acid sites to produce an EI intermediate. When HEDTA-MOR was exposed to 0.3 kPa of EA and evacuated at 473 K (Figure 6A), NH3 T deformation bands build up at 1597 cm* and 1497 cm " together with CH2 deformation band at 1460 cm". Those deformation bands are attributed to the -1 presence of protonated primary amines over the catalysts. At 1372 cm and 1324 cm "1 wave number regions (Figure 6A), OH deformation bands are observed. We previously reported that the major acidic sites were BrCnsted sites on HMOR, which are mainly located in the main channels [5]. The IR spectrum (Figure 6A) suggests that EA is protonated and adsorbed as ammonio-ion of EA (NHs*CH2CH2OH)+ on H-EDTA-MOR, and not adsorbed as an oxonium-ion (NH2CH2CH2OH2-). When the ammonio-ion on H-EDTA-MOR is heated and evacuated at higher temperature (Figure 6B-6D), the adsorption species were changed to secondary amines. The deformation bands of ammonio-ion are shifted to lower wave number: NH2 § deformation bands build up at 1600 cm 1, together with CH2 N§ deformation band at 1442 cm "1. The intensities of OH deformation bands (1370 -1 cm , 1324 cm "1) are decreased with increasing evacuation temperature. The

274 results suggest t h a t the ammonio-ion transformed to protonated EI over the catalyst surface. The intr~molecular condensation of EA occurred at the initial stage of reaction to produce EI. Then EI was activated by the stronger protonic acid sites to produce EDA. The secondary protonated amines are strongly held on the surfaces of H-EDTA-MOR without the presence of ammonia.

1600

1442

o

1600

L

/

.

A, I .... 1750

I 1500

I I 1300 1750

! 1500

I 1300

Wavenumbers/cm 1

Figure 6. IR spectra of adsorbed EA and NH3 on H-EDTA-MOR (Si/AI=6.1): (A) H-EDTA-MOR exposed 0.3 kPa of EA at 473 K and evacuated at 473 K, (B) evacuated at 523 K, (C) evacuated at 573 K, (D) evacuated at 623 K, (E) H-EDTA-MOR exposed 0.3 kPa of Nil3 at 473 K and evacuated at 473 K, (F) after recording of spectrum E, the sample was exposed to 0.3 kPa of EA at 473 K and evacuated at 473 K, (G) evacuated at 523 K, (H) evacuated at 573 K.

When H-EDTA-MOR was exposed to 0.3 kPa of NH3 at 473 K and evacuated at 473 K (Figure 6E), only protonated ammonia (NH4 +) was observed. The NH4 + deformation band builds up at 1443 cm -1 together with a small amount of coordinated NH3 bonds (deformation) t ha t attached to Lewis acid sites at 1624 cm -1.

275 The major acidic sites on H-MOR are Brcnsted sites determined by pyridine adsorption studies: above 80 % of acidic sites are Brcnsted sites and the rest are Lewis acid sites [4, 5]. A f a r adsorption of NH3, 0.3 kPa of EA are admitted on HEDTA-MOR at 473 K (Figure 6F); adsorbed NH3 is easily replaced by EA to produce defo _rma_fion bands of NH3 § (1597 cm -x, 1497 cm-X), CH2 (1460 cm-1). This spectrum is quite the same as the spectrum in Figure 6A. The results suggest that adsorption of EA is much stronger than that of NH3. When adsorbed EA is heated up to 573 K (Figure 6G-6H), the spectra are almost the same as the spectra in Figure 6B and 6C.

HO~NH

2 '---

L

EA

H20

HN \

~NH2

/

NH

\

PA

ll-H2~ N H

El

H+

~ N+ ~ H2

/

]

l EA,EOA -~ NH3

H2N ~

NH2

EDA

Scheme 1. Reaction pathway of EA synthesis from EDA and NH3 over H-EDTAMOR.

4. CONCLUSION The synthesis of EDA from EA with NH3 over acidic types of zeolite catalyst was investigated. Among the zeolites tested in this study, H-EDTA-MOR was the best catalyst. The reaction proved to be highly selective for EA over H-EDTAMOR, with small amounts of EI and PA derivatives as the side products. The catalytic activity for EDA synthesis rises with increasing partial pressure of NH3. The initial product of reaction was EI, and the formation of EDA followed. The reaction pathways for the formation of EDA from EA and NH3 are summarized in Scheme 1. The results suggest that the formation of EDA required stronger acidic sites in the mordenite channels with excess amounts of NH3. The mordenite channels may retard the formation of bulkier PA derivatives and other polyamines. The reactions of EA proceed through ammonio-ions by the addition of protons of H-EDTA-MOR; then intramolecular condensation of ammonio-ions of EA occurred to produce an EI intermediate over the active sites of the catalyst. EI is protonated by the stronger acidic sites with excess amounts of N H 3 to

276 produce EDA. Protonated EI intermediates are quite stable over H-MOR because the presence of stronger acidic sites. The excess amounts of NH s m a y be required to remove the protonated EI to form EDA. This must be the main reason why the dealuminated MOR (H-EDTA-MOR) enhances the catalytic activity for EDA synthesis from EA and NH 3, since the acid strength of H-EDTAMOR is slightly lower than that of H-MOR (see Figure 3), to give relatively easier desorption of protonated EI intermediate from the surfaces to form EDA. REFERENCES

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

S. Kumoi, N. Kubota, and T. Hiroi, Kagaku Keizai, 11 (1985) 56. Ninters, J. R., U.S. Pat., 4 404 405 (1983). Deeba, M., Ford, M. F., and Johnson T. A., European Patent 252 424 (January 13, 1988). M. Deeba, M. E. Ford, T. A. Johnson, and J. E. Premecz, J. Mol. Catal., 60 (1990) 11. K. Segawa, M. Sakaguchi, and Y. Kurusu, Stud. Surlf. Sci. Catal., 36 (1988) 579. K. Segawa and H. Tachibana, J. Catal., 131 (1991).

J.W. Hightower, W.N. Delgass, E. Iglesia and A.T. Bell (Eds.) 11th International Congress on Catalysis - 40th Anniversary

Studies in Surface Science and Catalysis, Vol. 101 9 1996 Elsevier Science B.V. All rights reserved.

277

Competitive Reaction Pathways in Propane Ammoxidation over V-Sb-Oxide Catalysts: an IR and Flow Reactor Study G. Centia, b and F. Marchi ~ a Dip. Chim.Industriale e dei Materiali, University of Bologna, V.le Risorgimento 4, 40136 Bologna, Italy* b Dip. Chimica Industriale, University of Messina, Contrada Papardo, 98100 Messina, Italy The surface t r a n s f o r m a t i o n s of propylene, allyl alcohol and acrylic acid in the presence or absence of NH3 over V-antimonate catalysts were studied by IR spectroscopy. The results show the existence of various possible p a t h w a y s of surface t r a n s f o r m a t i o n in the m e c h a n i s m of propane ammoxidation, depending on the reaction condition and the surface coverage with chemisorbed NH3. A surface reaction n e t w o r k is proposed and used to explain the catalytic behavior observed in flow reactor conditions. 1. I N T R O D U C T I O N V-Sb-oxide based catalysts show interesting catalytic properties in the direct synthesis of acrylonitrile from propane [1,2], a new alternative option to the commercial process s t a r t i n g from propylene. However, further i m p r o v e m e n t of the selectivity to acrylonitrile would strengthen interest in the process. Optimization of the behavior of Sb-V-oxide catalysts requires a thorough analysis of the relationship between structural/surface characteristics and catalytic properties. Various studies have been reported on the analysis of this relationship [3-8] and on the reaction kinetics [9,10], but little attention has been given to the study of the surface reactivity of V-Sb-oxide in the transformation of possible intermediates and on the identification of the surface mechanism of reaction. 2. E X P E R I M E N T A L V-Sb-oxide samples with Sb:V ratios of 1.0 and 3.0 were prepared by the following precipitation-deposition method: VC13 is dissolved in a 0.1N aqueous HC1 solution (A) and SbC15 is dissolved in a 3-5N HC1 aqueous solution (B). The dropwise addition of A to B leads to the formation of a white precipitate (mainly Sb-hydroxide) and a d a r k blue solution due to VO 2§ ions. VO 2§ is then precipitated over the Sb-hydroxide by adding dropwise a concentred aqueous solution of "Fax: +39-51-644.3680;e-mail: [email protected]

278 ammonia up to basic pH. The precipitate is filtered, washed three times with distilled water and then dried at 140~ overnight. The resulting solid is then calcined in a flow of air up to 600~ (3 h) using a constant rate of increase in temperature (50~ The surface areas of the samples after calcination are 17 and 10 m2-g-1 (Sb:V= 1.0 and 3.0, respectively). Prior to r u n n i n g the catalytic tests the samples prepared by this precipitation-deposition method were activated at 500~ (3h) in a flow of propane, a m m o n i a and air. Details on the characterization of these samples have been reported elsewhere [8,11]. The catalytic tests were carried out using an a p p a r a t u s composed of (i) a continuous fixed-bed stainless steel reactor, (ii) a section for the preparation of feed by mixing already calibrated mixtures of the single components in helium, (iii) a system for on-line gas-chromatographic (GC) analysis (two GCs equipped with a flame ionization and thermoconducibility detector, respectively) and (iv) a section for the analysis of NH3 conversion and HCN formation by absorption in appropriate solutions and titration[6,9,10]. Tests were made using 2-4 g of sample with particle dimensions in the 0.1-0.2 m m range and diluted in a 1:5 ratio using an inert support. The axial t e m p e r a t u r e profile was monitored by thermocouples inserted into the catalytic bed. Tests were made using the following feedstock: 7.5% propane and a O2/C3 and NH3/C3 ratio of 1.56 and 1.60, respectively. The gas-space hourly velocity (GHSV) was 2800 h -1. Fourier-transform infrared (IR) spectra (resolution 2 cm -1) were recorded with a Perkin Elmer 1750 i n s t r u m e n t in a quartz cell connected to grease-free evacuation and gas m a n i p u l a t i o n lines. The self-supporting disk technique was used. Before recording the spectra, the samples were treated with 02 at 450~ (lh), then cooled down to r.t. before evacuating the 02. The sample was then evacuated at 400~ Evacuation at higher t e m p e r a t u r e s lead to a drastic cut off of IR trasparency. All reactants were purified prior to the adsorption experiments. Due to the better resolution of the spectra, only results for Sb:V=I.0 are reported here, however the IR data for Sb:V=3.0 were not significantly different.

3. R E S U L T S 3.1 C a t a l y t i c T e s t s The catalytic behavior in propane ammoxidation of Sb:V=I.0 and 3.0 is summarized in Fig. 1. The tests were carried out using a propane concentration of about 8% and oxygen as the limiting reactant, because these experimental condit-ions agree with those indicated as preferable in the patent literature [12] and from the analysis of the reaction kinetics [9,10]. For both catalysts, as the reaction t e m p e r a t u r e increases the the selectivity to acrylonitrile (ACN) passes t h r o u g h a m a x i m u m at about 480~ Maximum selectivity is about 30% and 60% for Sb=l.0 and 3.0, respectively. Increasing the temperature, decreases the selectivity to propylene (C3=) and increases t h a t to carbon oxides (COx). It may be noted t h a t the lower selectivity to ACN in Sb:V= 1.0 is

279

not necessarily associated with a higher formation of carbon oxides, but r a t h e r with a higher side conversion of NH3 to Nz "~which leads to a higher consumption of O2. Sb:V=I.0 is thus less selective to ACN, but more selective to C3=. It is interesting also to note t h a t the formation of acetonitrile does not follow t h a t of ACN, but instead selectivity to AcCN decreases as the temperature increases. The AcCN/ACN ratio thus is m a x i m u m at low temperatures and decreases ~continuously. In Sb:V=I.0 the ~=AcCN/ACN ratio passes from ~0.60 at 440~ to 0.18 at 500~ whereas the ratio is about ~ over Sb:V=3.0 (0.23 and 0.08, respectively).

3.2 Infrared study

Fig. 1 Catalytic behavior in propane ammoxidation of Sb:V = 1.0 and 3.0 (bottom and top, re- Propylene: The IR spectra of praspectively). Symbols: C conversion, S selectivity, pylene in contact at r.t. with C3 propane, C3= propylene, ACN acrylonitrile, Sb:V=I.0 and evacuation at proAcCN acetonitrile, gressively higher t e m p e r a t u r e s are reported in Fig. 2. Well evident bands are noted at 1660, 1468 and 1455 (double), 1388 and 1372 (double), 1328, 1255, 1170, 1132 and 1101 cm -1, the last very intense. The same bands are observed by adsorption of isopropyl alcohol and may be assigned to a mixture of a isopropoxylate species (1101 cm-1; vc-o ) and coordinated acetone (intense band at 1660 cm-1; vc=o); the latter assignment is confirmed by the analysis of the spectrum of acetone adsorbed over the same catalyst. Analogous species were detected by adsorption of propylene over vanadiatitania catalysts [13,14]. Additional weak bands at 1640 and 1070 cm -1 are observed in the spectrum of propylene in contact with the catalyst (Fig. 2a), but disappear by evacuation at r.t. (Fig. 2b). They can be assigned to vc--c and vc-o, respectively, of propylene adsorbed in the form of allyl alcoholate. With increasing evacuation t e m p e r a t u r e (Fig. 2c,d) the bands associated with isopropoxylate progressively decrease and disappear by evacuation at 120~ Likewise, the bands of coordinated acetone increase in relative intensity up to about 160~ and then decrease at higher temperatures. The spectrum of the sample evacuated at 160~ (Fig. 2e) is dominated by an intense band near 1540 cm -1 plus other bands at 1438, 1660, 1420 and 1370 cm-1; the latter three bands

280

are due to acetone which has not been completely transformed. At an evacuation temperature of 200~ (Fig. 2f) the spectrum is dominated by two intense bands at 1532 and 1448 cm -1, typical of carboxylate species. Other weak bands at 1640 and 1375 cm -1 (vcc and 5CH, respectively) indicate the presence of an acrylate species and at 1360 cm -1 (5CH3) of an acetate. The very weak bands at 1860 and 1795 cm -1 suggest the possible formation of a cyclic anhydride (vas and vs of C=O in cyclic anhydrides, respectively), reasonably by dimerization of propylene and oxidative attack to form maleic anhydride.

JD

/; i' :1

"1 .I

".-.

".'l ","1

i

:"

t

9'":

: ~ ' :".

:1

o:

t

"

:

:

....

1

/,, -.. I ~"

~

\A

g

-

.

"-

..... /

',

Propylene thus readily reacts at r.t with V-Sb-oxide to form an isopropoxylate species which then may develop along two possible pathways: (i) the first route (attack from surface hydroxyl groups) gives products of carbon chain degradation (acetone plus CH4) ! ~ '~ e a! * and (ii) the second route occurring at higher ii. temperatures gives rise first to an allyl alcoholate species which then transforms to acrolein and acrylate. The second routes b "' J" t ~ -i ~'~ thus involves nucleophilic oxidative attack by lattice oxygen. Propylene + NH3: The results of propylene and ammonia coadsorption experiments are summarized in Fig. 3. Two series of tests were made. In the first series of tests (procedure A) propylene is put in contact at r.t. with the catalyst for 5 min, then removed by evacuation at r.t.; later ammonia 1800 1600 1400 1200 o m ~ is put in contact for 2 min at r.t. and then Fig. 2 IR spectra of 60 torr propylene removed by evacuation. Only species chemiin contact at r.t with Sb:V=I (a), subsorbed on the catalyst thus remain. Spectra sequent evacuation at r.t. (b) and folwere then recorded at increasing temperalowing evacuations at increasing temperatures: (c) 80, (d) 120, (e) 160 and tures under vacuum. In the second series of experiments (procedure B) propylene and (f) 200~ ammonia together are put in contact with the catalyst and the spectra recorded at increasing temperatures of contact with the mixture of C3H6+NH3. In the first experiments the change in coadsorbed species is studied, whereas in the second type of tests the effect of the presence of propylene and ammonia also in the gas phase is analyzed. ,"

;

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.

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;

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281

n 1~ I I~l

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.

.

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.

.

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...'

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|

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i',,

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.." , - ..........

""

""

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9..

A

2500

2000

1800

1600

1400

1200

2500

2000

1800

1600

1400

1200 cm"

Fig. 3 Left, procedure A: 60 torr C3H6 adsorbed and evacuation at r.t., 2 torr NH3 adsorbed and evacuation at r.t. (a), then evacuation at increasing temperatures: 100 (b), 150 (c), 200 (d) and 250~ (e). Right, procedure B: 30 torr C3H6 + 30 torr NH3 in contact with the catalyst for 20 min at r.t (a) and at 100 (b), 200 (c) and 300~ (d) for 5 min. The two procedures of coadsorption e x p e r i m e n t s give different results as shown in Fig. 3 (procedure A: spectra on the left; procedure B: on the right). The more r e m a r k a b l e evidence is the absence of formation of a w e a k band n e a r 2230 cm -z (VCN) indicating the formation of weakly coordinated acrylonitrile using procedure A, n o t w i t h s t a n d i n g the disappearence of VCH bands. The VCN band forms using procedure B. F u r t h e r m o r e , in the case of procedure A the a m o u n t of ammonia chemisorbed (mainly in the form of a m m o n i u m ion; strong band at about 1430 cm -1) rapidly decreases with increasing t e m p e r a t u r e due to conversion to Ne + H20 with a consequent freeing of the BrOnsted sites blocked by reaction with NH3. The presence of a m m o n i a in the gas phase using procedure B instead inhibits this process. O t h e r bands in the spectra are consistent w i t h those observed after chemisorption of the pure components and indicates t h a t using procedure A surface species develop w i t h formation m a i n l y of chemisorbed acetone and acetic acid a n d using procedure B with formation m a i n l y of chemisorbed acrolein and acrylic acid. This indicates t h a t the presence of a m m o n i a in the gas p h a s e changes the sur-

282 face acidity characteristics of V-Sb-oxide catalysts, thus modifying surface reactivity.

l

.1o

ii i I:

/~':

l il ::

"

..

.-"

"'

II

::

!!ii!

,

~

,

-

I

:"

:"

-

!

: :

:

-

~",

,

- i

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/ :.

;:

:

: .... , l

,, l

::

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/~

.

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, ,,

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.

t~

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gi:

,,

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,,, j a I .....

l

l

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~ "-\

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]'l

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,'\

,

1800 I~X) 1400 1200 C . " Fig. 4 IR spectra of 10 torr allyl acohol in contact at r.t with Sb:V=I (a), subsequent evacuation at r.t. (b) and at increasing temperatures: (c) 80, (d) 110, (e) 150, (f) 200 and (g) 250~

Allyl alcohol: Allyl alcohol in contact with VSb-oxide catalyst at r.t. (Fig. 4a) gives rise to a spectrum characteristic of allyl alcoholate (1645 cm-l: vc=c; 1422 cm-Z: scissoring =CH2; 1445, 1342, 1360 and 1104 cm -1 s c i s s o r i n g CH2-O-, wagging CH2 and vc-o). A progressive increase in the i n t e n s i t y of the band at 1185 cm -1 with increasing t e m p e r a t u r e of evacuation indicates t h a t acrolein, not present initially, forms progressively with a parallel decrease in the relative intensity of the bands of allyl alcoholate. At 200~ the bands of acrolein (1641 cm -1 vc=o, 1625 cm -1 vc=c, 1430 cm -1 scissoring CH2, 1372 and 1282 cm -1 5CH, 1188 cm -1 vcc) disappear with the formation of new b a n d s clearly evident in the spectrum after evacuation at 250~ (Fig. 4g). The l a t t e r s p e c t r u m is characterized by bands at 1636, 1530, 1437, 1375 and 1278 cm 1 which correspond to those observed for acrylate species. The intense band at 1145 cm -1 which develops at higher t e m p e r a t u r e s is instead due to the reduction of the catalyst and formation of oxygen vacancies [5]. The adsorption of allyl alcohol t h u s readily gives rise to an allyl alcoholate species at room t e m p e r a t u r e . At higher t e m p e r a t u r e s this species first t r a n s f o r m s to chemisorbed acrolein and t h e n to an acrylate species. Allyl alcohol + NH3: The coadsorption exp e r i m e n t s were m a d e by first p u t t i n g the catalyst into contact with NH3 at 200~ followed by removal of gaseous NH3 at r.t. and then p u t t i n g the c a t a l y s t with the chemisorbed a m m o n i a into contact w i t h the allyl alcohol at r.t. Using procedures for the coadsorption tests like those used for propylene + NH3, less resolved spectra were obtained. The more r e m a r k a b l e differences in comparison with the analogous spectra for propylene + NH3 (Fig. 3) or acrylic acid + NH3 coadsorption (see below) are t h a t (i) the band in-

283

t

I

!

/''",

'

,

"'

,

:

;

:t

i',,

l

I

,

::

,"

!

'

i

"\

t

i ',..,

l

I

9

'.:. ',,

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._

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|

dicating the formation of a nitrile species (2280 cm-Z) falls at a slightly higher frequency w i t h respect to the former case, due to interaction with Lewis sites, and occurs at lower temp e r a t u r e s (150-200~ i n s t e a d of 250300~ and (ii) the formation of nitrile species also occurs in the absence of a m m o n i a in the gas phase. Bands also are observed at 1550 and 1670 cm -1 which agree w i t h the formation of an imine or oxime species by reaction of acrolein with coordinated NH3. This i n t e r m e d i a t e species t r a n s f o r m s to nitrile species at lower temperatures t h a n those for the i n t e r m e d i a t e species formed in co-adsorption tests of acrylic acid and a m m o n i a .

Acrylic acid: The IR spectra obtained

'

I

J,,

25(m

Fig. 5 IR spectra of 10 torr NH3 in contact at 200~ with Sb:V=I, subsequent evacuation at r.t., contact with allyl alcohol (5 torr, 5 min) (a), and following evacuations at increasing temperatures: (b) 150, (c) 200, (d) 250 and (e) 300~

by adsorption of acrylic acid at r.t. and evacuation at increasing temp e r a t u r e s are shown in Fig. 6. Adsorption and evacuation at r.t. gives rise to a spectrum with bands at 1636 cm -1 (vc=c), 1494 cm 1 (Vas -CO2-), 1436 cm -1 (scissoring CH2 overlap to vs -CO2-), 1376 cm 1 (SCH), 1277 cm 1 (vcc) and 1070 cm -1 (rocking CH2) indicating the rapid formation of an acrylate species. Weak bands are also observed at 1660 and 1600 cm -1 which m a y be a t t r i b u t e d to weakly bonded acid. These bands d i s a p p e a r by evacuation above r.t.. The acrylate species is stable up to an evacuation t e m p e r a t u r e of about 200~ and then decreases in relative intensity at higher t e m p e r a t u r e s . Already in the spectrum obtained by evacuation at 200~ (Fig. 6c) a b r o a d e n i n g of the band at 1494 cm -1 and a shoulder at 1350 cm -1 indicate partial degradation to an acetate species.

Acrylic acid + NH3: Two series of experiments were made. In the first series

284 (procedure A) ammonia is first adsorbed at r.t. followed by removal of gas phase ammonia by evacuation. Later the catalyst with chemisorbed ammonia is put in contact at r.t. with the acrylic acid. Spectra are then recorded at -...: : . . - ..~ increasing temperatures of evacuation. In the second series of experiments (procedure B) acrylic acid is first put into contact with the !." .~ catalyst followed by evacuation. Then the i. "..'/ "" catalyst with chemisorbed acrylic acid is put q-.. into contact with NH3 at increasing temperatures. Analogously to the case of propylene and C ..: a aI 1". ammonia coadsorption, using procedure A only ........._ :I,, , chemisorbed species are formed, whereas in I I I | ii I procedure B ammonia is also present in the gas phase. The two procedures give rise to different results. In both cases acrylic acid, present in the p,( ,I form of acrylate, readily reacts with ammonia / I!' 'I at r.t. forming a species characterized by an inq | i| ~'. I ,,, II I " /ills :l tense band at 1535 cm -z indicating the formaJr I |1 ~ i! 9t i tion of an amide. With increasing reaction b ,:: :; temperature (100~ however, in the case of procedure A the band at 1535 cm -1 shii~s to 1495 cm -1 and a weak band forms at 1720 cm -1. The latter band is characteristic of undissociated and weakly coordinated acrylic acid. This indicates that at 100~ amide dissociates with formation of the free acid. When ammonia is instead present in the gas phase (procedure B), the amide species undergoes transformation to 1800 1600 1400 1 2 0 0 r "1 acrylonitrile with a maximum in the intensity Fig. 6 IR spectra of 1 torr acrylic of the vcs band at 2220 cm -1 at an evacuation acid in contact (5 min) with Sb:V=I temperature of about 300~ and evacuation at r.t (a), and folCoordinated acrylic acid and ammonia thus lowing evacuations at 100 (b) and react faster at r.t. to form acrylamide, but in 200~ (c). the absence of ammonia which inhibits the reactivity of the BrOnsted sites, the amide dissociates at higher temperatures with formation of the free acrylic acid. When the reactivity of the BrOnsted sites is blocked by their transformation to ammonium ions, the amide m a y be dehydrogenated to form the acrylonitrile. ~

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] ---"

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O 0 0

Fig. 7 Surface reaction network in the ammoxidation of propane over V-Sb-oxide. 4. DISCUSSION

Co-adsorption experiments show a complex role of the nature and concentration of chemisorbed ammonia species. Ammonia is not only one of the reactants for the synthesis of acrylonitrile, but also reaction with Br@nsted sites inhibits their reactivity. In particular, IR experiments show that two pathways of reaction are possible from chemisorbed propylene: (i) to acetone via isopropoxylate intermediate or (ii) to acrolein via allyl alcoholate intermediate. The first reaction occurs preferentially at lower temperatures and in the presence of hydroxyl groups. When their reactivity is blocked by the faster reaction with ammonia, the second pathway of reaction becomes preferential. The first pathway of reaction is responsible for a degradative pathway, because acetone further transform to an acetate species with carbon chain breakage. Ammonia as NH4 + reacts faster with acrylate species (formed by transformation of the acrolein intermediate) to give an acrylamide intermediate. At higher temperatures the amide may be transformed to acrylonitrile, but when BrCnsted sites are present, the amide may be hydrolyzed to reform ammonia and free, weakly bonded, acrylic acid. The latter easily decarboxylate forming carbon oxides. Ammonia also reacts with the acrolein intermediate, via the formation of an imine or possibly oxime intermediate which transforms faster to the acrylonitrile than to the acrylamide intermediate. This p a t h w a y of reaction occurs at lower temperatures in comparison to that involving an acrylate intermediate, but its relative importance depends on the competitive reaction of the acrolein intermediate with the ammonia species and with catalyst lattice oxygens. NH3 coordinated on Lewis sites also inhibits the activation of propane differently from that absorbed on Brr sites. The scheme shown in Fig. 7 summarizes the proposed surface reaction network in propane ammoxidation.

286 Although the reaction scheme of Fig. 7 is based mainly on the IR data, it also can be used to discuss various aspects of the surface reactivity found in flow reactor studies (Fig. 1). The higher acetonitrile to acrylonitrile ratio at lower temp e r a t u r e s agrees well with t h a t expected from the effect of the reaction temperature on the relative rates of propylene transformation via isopropoxylate or allyl alcoholate intermediates. The considerable effect of the rate of side a m m o n i a oxidation to N2 on the selectivity to acrylonitrile (compare Fig. 1A and 1B) agrees well with the effect of a m m o n i a on the modification of reaction p a t h w a y s discussed above, in addJt|on to its role as coreactant. The optimization of V-Sb-oxide catalysts thus requires the tailoring of various aspects of the surface reactivity: (i) increasing the rate of the p a t h w a y via allyl alcoholate versus t h a t via isopropoxylate, (ii) increasing the rate of reaction of ammonia with the acrolein intermediate versus conversion of the latter to acrylate and (iii) reducing the rate of conversion of chemisorbed a m m o n i a to N2, the latter a factor which also has a considerable effect on the first two points. The financial support from the Ministero Pubblica Istruzione (60%) is gratefully acknowledged.

REFERENCES 1. Y. Moro-oka, W. Ueda, in Catalysis - Vol. 11, The Royal Society of Chemistry, Cambridge U.K. 1994, p. 223. 2. G. Centi, R.K. Grasselli, F. Trifirb, Catal. Today, 13 (1992) 661. 3. R. Nilsson, T. Lindblad, A. Andersson, C. Song, S. Hansen, New Developments in Selective Oxidation II, V. Cortes Corberan and S. Vic Bellon Eds., Elsevier Pub.: Amsterdam 1994, p. 281. 4. A. Andersson, S.L.T. Andersson, G. Centi, R.K. Grasselli, M. Sanati, F. Trifirb, Appl. Catal. A, 113 (1994) 43. 5. G. Centi, S. Perathoner, Appl. Catal., 124 (1995) 317. 6. G. Centi, R.K. Grasselli, E. Patan~, F. Trifir6, New Developments in Selective Oxidation, G. Centi, F. Trifirb Eds., Elsevier Pub.: Amsterdam 1990, p. 515. 7. G. Centi, E. Foresti, F. Guarneri, New Developments in Selective Oxidation II, V. Cortes Corberan, S. Vic Bellon Eds., Elsevier Pub.: Amsterdam 1994, p. 281. 8. G. Centi, S. Perathoner, Preparation of Catalysts VI, G. Poncelet et al. Eds., Elsevier Science Pub.: Amsterdam 1995, p. 59. 9. R. Catani, G. Centi, F. Trifirb, R.K. Grasselli, Ind. Eng. Chem. Res., 31 (1992) 107. 10. A. Andersson, S.L.T. Andersson, G. Centi, R.K. Grasselli, M. Sanati, F. Trifirb, New Frontiers in Catalysis, L. Guczi et al. Eds., Elsevier Pub: Amsterdam 1993, p. 691. 11. G. Centi, P. Mazzoli, Catal. Today (issue on Fundamentals of Oxide Catalysts), in press. 12. M.A. Toft, J.F. Brazdil, L.C. Glaeser, U.S. Patent, 4,784,979 (1988). 13. V. Sanchez Escribano, G. Busca, V. Lorenzelli, J. Phys. Chem., 94 (1990) 8939. 14. G. Busca, G. Ramis, V. Lorenzelli, J. Molec. Catal., 55 (1989) 1.

J.W. Hightower, W.N. Delgass, E. Iglesia and A.T. Bell (Eds.) 11th International Congress on Catalysis - 40th Anniversary

Studies in Surface Science and Catalysis, Vol. 101 1996 Elsevier Science B.V.

287

'Seeing' The Active Site in Catalysis. STM and Molecular Beam Studies of Surface Reactions Michael Bowker

Reading Catalysis Centre, Department of Chemistry, University of Reading, Whiteknights Park, Reading RG6 6AD. and IRC in Surface Science, University of Liverpool.

1. ABSTRACT The 'Holy Grail' of catalysis has been to identify what Taylor described as the 'active site' that is, that ensemble of atoms which is responsible for the surface reactions involved in catalytic turnover. With the advent of atomically resolving techniques such as scanning tunnelling microscopy it is now possible to identify reaction centres on planar surfaces. This gives a greater insight also into reaction kinetics and mechanisms in catalysis. In this paper two examples of such work are described, namely CO oxidation on a Rh(110) crystal and methanol selective oxidation to formaldehyde on Cu(110).

2. INTRODUCTION Since early in this century the concept of the 'active site' in catalysis [1] has been a focus of attention in this area of chemistry. This was proposed to be that ensemble of surface atoms/reactants which is responsible for the crucial surface reaction step involved in a catalytic conversion. Since those days much work has been done in the area, which cites the concept of the active site. However, no such ensemble has been positively identified due to the lack of availability of techniques which could image such a structure, which is of atomic dimensions. However, in more recent times science has made rapid strides in this direction. It is now possible to use EXAFS in situ during a catalytic reaction to examine the average coordination of metal atoms in the small particles which often exist in precious metal catalysts [2]. High resolution transmission electron microscopy has evolved to the level of atomic resolution, but can only be used ex-situ, or in situ with moderate pressures when special cells are fitted [3].

288 The most recent innovation in this field is scanning tunnelling microscopy, which has the capability of atomic resolution. In the work reported here two surface reactions are examined using this technique. These reactions are of relevance to automobile catalysis (CO oxidation on Rh) and methanol oxidation/synthesis on Cu. It is proposed that active sites are imaged in these reactions and that these active sites can indeexl be extremely dilute on the surface.

3. RESULTS AND DISCUSSION Oxygen adsorption on Rh results in reconstruction of the surface layer and an STM image of an oxygen-dosed surface is shown in fig. 1. The image shows bands of bright regions separated by dark lines. LEED has shown a multiplicity of structures on Rh(110), depending on the exact oxygen coverage. The important common feature of these is a missing row of Rh atoms oriented in the { 110} direction. Different oxygen coverages then produce different numbers of Rh atoms within the bright bands, which are mixed structures of Rh and O; the c(2x6) structure has 2 rows of Rh, while the c(2x8) has 3 [4-7]. Due to a staggering of the atoms in adjacent bands the periodicity in the unit cell is doubled, stretching over two bands in the [001] direction [6,7]. If these structures are exposed to CO, the reaction takes place in quite a surprising manner - the surface shows one dimensional reactivity and fig. 2 is an STM image of the same area before and after such a reaction. The reaction appears to initiate at step edges or defects on the surface and then proceeds in the {011 } direction to eat away at the Rh-O bands. Furthermore, certain structures are more reactive than others; the c(2x8) bands react faster than the c(2x6). It appears, then, that the active sites for CO oxidation are located at the end of these bands of oxidised Rh and are very specific in nature. This explains earlier results in which a low initial reactivity to CO2 formation was observed using molecular beam measurements [8]. The reaction probability was low at high oxygen coverages but increased as the oxygen-coverage diminished, until finally decreasing again as the oxygen becomes dilute on the surface. An example of this is shown in fig. 3a. If we turn to quite a different system, namely methanol oxidation on Cu(110), a similar reactivity pattern can be observed. Fig. 3b compares the reaction of methanol with preadsorbed oxygen at two oxygen coverages - 0.25 and 0.5 monolayers, the latter being saturation of the p(2x 1) structure on the surface. The oxygen saturated layer shows very low reactivity initially, which increases with time, goes through a maximum, and decreases to zero when the oxygen is all used up. When a Cu(110) surface is dosed with a partial layer of oxygen, the absorbate forms long, thin, one-dimensional islands. This structure has been observed by many other workers [9-11] and is quite characteristic of this adsorption system, with oxygen reacting with surface Cu atoms producing an added row structure of alternating oxygen and Cu atoms (a p(2xl) structure and with a surface atomic density of only half that of the original surface. This is thought to be formed by the diffusion of Cu atoms away from step edges to bind to growing CuO chains [ 10,11]. When this surface is reacted with methanol, then methoxy is formed as demonstrated by a variety of techniques, including molecular beam methods [12], TPD [12,13], IRAS [14] and UPS [15]. Fig. 4a shows a layer of oxygen which has partially reacted and formed

289

Figure 2. Showing the beginning of CO reaction with the surface along the long Rh-O islands, which are between the missing rows, as bright streaks.

290

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._.1 uJ or"

300

I

I

350

400

I

450

I

I

I

500

550

600

650

TEMPERATURE (K) Figure 1. Relative rates of methyl chlorosilane formation from a methyl + chlorine monolayer on a Cu3Si surface. The Cu3Si surface was prepared by ion bombardment at 330 K, and a I : 1 ratio of methyl groups and chlorine atoms were reacted in these studies.

carbon. It should be emphasized, however, that carbon deposition likely occurs after the surface CH3 and C1 coverages have b e g u n to decrease as a result of methylchlorosilane formation. By contrast, in steady state catalytic processes, the surface coverages remain constant (probably at or near monolayer saturation). As a result, one should exercise caution in comparing directly the relative yields from these monlayer adsorption studies with catalytic selecfivities. However, trends in rate and selectivity as a function of promoters and surface composition in the monolayer studies are probably relevant for the catalytic process.

311 The reactions of CH3 radicals and C12 alone with Cu3Si have also been investigated. On pure Cu3Si, the dominant silane product from CH3 adsorption is SiH(CH3)3 and the temperature at which the surface is sputtered prior to methyl adsorption has a dramatic effect on the reaction rate (see section 3.3). The C12 reaction gives SIC14 evolution, and the reaction t e m p e r a t u r e is close to that for methylchlorosilane formation. 3.2. Effect of promoters on reaction rates and selectivities

The products from the reaction of CH3 + C1 monolayers on the promoted Cu3Si surface are the same as those for pure Cu3Si, but both the absolute rates and the selectivities are significantly different. In experiments analogous to those described in section 3.1, methylchlorosilanes are evolved from the p r o m o t e d Cu3Si surface between 300 and 450 K. This temperature is 200 K lower than that from the pure Cu3Si surface. This 200 K difference in reaction temperature corresponds to a difference of six orders of magnitude in rate (if the rates are extrapolated to a common reaction temperature of 500 K assuming standard and equivalent pre exponential factors for the reactions on these two surfaces [10]). The reaction selectivity, i.e. the percentage of each methylchlorosilane relative to the total yield of methylchlorosilanes, has been determined from the TPR peak areas of the major cracking fragments using the relative mass spectrometer sensitivity factors determined from studies of authentic samples [11]. These selectivities are compared with those for a typical commercial catalyst [12] in Figure 2. As shown, the 84% selectivity to (CH3)2SIC12 on promoted Cu3Si is close to the commercial value of 90%. On pure Cu3Si, however, the selectivity to (CH3)2SIC12 is only 58%. This decrease in selectivity is not unexpected for a Cu3Si sample without promoters on the basis of catalytic results [1]. For reference, it should be noted that the equilibrium concentration of dimethyldichlorosilane in a methylchlorosilane mixture containing equal numbers of methyl and chlorine is 88% [13] as a result of reversible reactions such as that shown below: 2(CH3)2SIC12 ~

(CH3)SiCI 3 + (CH3)3SiCI

The addition of promoters (and excess Si) to Cu3Si also has a significant effect on the reactivity of pure C1 monolayers formed on this surface by the adsorption of C12. We find that on the promoted samples, C1 monolayers produce SIC13 rather than SIC14. The identification of SIC13 as the chlorosilane product is based on the detection of SIC13+ and the absence of SIC14+ or higher chlorosilane ions in mass spectrometric studies. A second effect of promotion is to lower the chlorosilane evolution temperature. Specifically, on promoted samples the SIC13+ peak m a x i m u m in TPR studies is 390 K vs. --500 K for u n p r o m o t e d samples. Thus, the rates of both chlorosilane and methylchlorosilane formation are dramatically enhanced on the promoted vs. unpromoted samples. By contrast, the rate of trimethylsilane formation when methyls alone are adsorbed on the surface is relatively unaffected by the presence or absence of promoters. In this regard, it should be noted that Falconer and c o w o r k e r s h a v e p r e v i o u s l y s u g g e s t e d that the active surface sites for methylchlorosilane formation involve surface Si-C1 species [14].

312

Figure 2. Comparison of the relative yields of methylchlorosilanes from methyl + chlorine monolayers on Cu3Si and doped Cu3Si samples with typical yields for the catalytic Rochow process.

3.3. Effect of surface composition on reaction activity In all of the studies described above, the Cu3Si samples were prepared by ion bombardment at 330 K followed by cooling of the surface to 180 K before adsorbing the methyl radicals and chlorine. AES studies as well as ion scattering results in the literature [7, 15] show that this procedure produces a surface that is enriched in silicon compared with the Cu3Si bulk stoichiometry. We have found that surfaces with less Si enrichment (possibly even copper enriched relative to the bulk stoichiometry) can be prepared by ion bombardment at temperatures below 300 K. Specifically, Cu(60 eV)/Si(92 eV) Auger peak ratios of 1.2 - 1.7 compared with a ratio of 0.5 at 400 K can be obtained by sputtering at 180 K. In the case of unpromoted Cu3Si surfaces, the effect of this copper enrichment on methylchlorosilane formation appears to be relatively minor. The majority of methylchlorosilanes are evolved at 400 - 650 K as on the unpromoted surface. There is, however, a small yield of methylchlorosilanes with a peak temperature of -370 K. By contrast, for trimethylsilane formation from pure methyl monolayers, copper enrichment by low temperature sputtering shifts the dominant product peak from

313 580 K to 280 K. This dramatic effect is shown in Figure 3 for the case of samples sputtered at 330 K vs. 160 K, and in these studies the evolution of trimethylsilane is monitored via m / e = 73 as shown. This 580 to 280 K decrease in reaction t e m p e r a t u r e for trimethylsilane formation corresponds to a decrease in the reaction activation energy from -33 k c a l / m o l to -16 kcal/mol. Alternatively, if the rates of these processes could be measured at a common temperature, they would differ by more than 5 orders of magnitude. The fact that trimethylsilane is evolved at either 580 K or 280 K and not at temperatures in between suggests that there are distinctly different active sites for forming this product. The ratio of these active sites is a function of the temperature at which the surface is ion bombarded, and the transition from high to low temperature product evolution correlates directly with a factor of 1.5 to 2 increase in the Cu/Si

Sputtered at 330 K

Sputtered at 160 K

580 K

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280 K

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200

300

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TEMPERATURE (K)

Figure 3. Temperature-programmed reaction spectra monitoring m / e = 73 (a cracking fragment of trimethylsilane) after adsorbing a saturation coverage of methyl groups on Cu3Si surfaces prepared by ion bombardment at 160 K (dotted curve) and 330 K (solid

curve).

314 Auger peak ratio. Although the effects of differing surface roughness for these different s p u t t e r i n g temperatures remain to be addressed, the 580 K peak temperature correlates both with the extent of silicon enrichment and also with the temperature to which methyl groups are found to be stable on a Si(100) surface [16]. Furthermore, the 280 K temperature for trimethylsilane formation on the copperenriched surfaces is consistent with the finding the methyl groups are thermally stable to -400 K on copper surfaces [9]. Finally, it should also be noted that early studies [5] have suggested that an important catalytic role for copper is to break up the Si lattice thereby making the Si more available to CH3 and C1 since Cu-Si bonds are weaker than Si-Si bonds. The results here demonstrating the enhanced reactivity for formation of trimethylsilane on copper-enriched surfaces are consistent with this interpretation. 4. CONCLUSIONS The studies presented here show that if a monolayer of methyl groups and chlorine atoms is produced on the surface of a Cu3Si wafer (free of oxygen, carbon and other impurities) then methylchlorosilanes can be formed in 10-20% yield by heating the monolayer to temperatures above 450 K. For a 1:1 ratio of methyls to chlorine, the selectivity for formation of dimethyldichlorosilane is --60%. The ratedetermining step in the evolution of these methylchlorosilanes is a surface-mediated coupling reaction as opposed to product desorption. Both the rate and the selectivity of the process are significantly increased if the Cu3Si samples contain excess Si as well as Sn, Zn, and A1. These promotional effects are consistent with those found for similar additives in the catalytic Rochow process for synthesizing methylchlorosilanes from methyl chloride and silicon. The monolayer studies here suggest that the increase in rate reflects an effect of the promoters on the surface chlorine as opposed to the surface methyl groups. By contrast, it is found that changes in the surface C u / S i ratio have a more significant effect on the reactivity of the surface methyl groups than the surface chlorine. Studies are in progress to further characterize the active sites for these reactions, and we believe that combination of a controlled monolayer approach and UHV surface analysis techniques will lead to even more definitive results in planned future studies. 5. ACKNOWLEDGMENTS Financial support from Dow Coming Corporation and from the American Chemical Society (Grant #CHE 93-18625) is gratefully acknowledged. BEB also gratefully acknowledges support from the Camille and Henry Dreyfus Foundation in the form of a Teacher-Scholar Award.

315

REFERENCES

1. 2. 3. 4. 5.

6. 7. 8. 9. 10. 11. 12. 13. 14. 15. 16.

K.M. Lewis and D. G. Rethwisch (eds.), Catalyzed Direct Reactions Of Silicon, Studies in Organic Chemistry 49, Elsevier, Amsterdam, 1993 and articles therein. W.J. Ward, A. Ritzer, K. M. Carroll and J. W. Flock, J. Catal., 100 (1986) 240. K J. H. Voorhoeve, Organosilanes: Precursors to Silicones, Elsevier, Amsterdam, 1967. T.C. Frank, K. B. Kester and J. L. Falconer, J. Catal., 91 (1985) 44. K.M. Lewis, D. McLeod, B. Kanner, J. L. Falconer and T. C. Frank, in: Catalyzed Direct Reactions of Silicon, edited by K. M. Lewis and D. G. Rethwisch, Elsevier, Amsterdam, 1993, 333, and references there in. C.-M. Chiang, T. H. Wentzlaff and B. E. Bent, J. Phys. Chem., 96 (1992) 1836. T.C. Frank and J. L. Falconer, Appl. Surf. Sci, 14 (1982-83) 359. G.H. Smudde, Jr., X. D. Peng, IL Viswanathan and P. C. Stair, J. Vac. Sci. Technol., A 9 (3) (1991). C.-M. Chiang and B. E. Bent, Surf. Sci., 279 (1992) 79. P. A. Redhead, Vacuum, 12 (1962) 203. D.-H. Sun, L. J. Kaufman, B. E. Bent, A. P. Wright, and B. M. Naasz, manuscript in preparation. B. Kanner and K. M. Lewis, in: Catalyzed Direct Reactions of Silicon, edited by K. M. Lewis and D. G. Rethwisch, Elsevier, Amsterdam, 1993, 1. D. R. Weyenberg, L. G. Mahone and W. H. Atwell, Annals of the New York Academy of Sciences, 159 (1969) 38. T. C. Frank and J. L. Falconer, Langmuir, 1 (1985) 104. Y. Samson, J. L. Rousset, G. Bergeret, B. Tardy and J. G. Bertolini, Appl. Surf. Sci., 72 (1993) 373. H. Gutleben, S. R. Lucas, C. C. Cheng, W. J. Choyke and J. T. Yates, Jr., Surf. Sci., 257 (1991) 146.

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J.W. Hightower, W.N. Delgass, E. Iglesia and A.T. Bell (Eds.) 11th International Congress on Catalysis - 40th Anniversary

Studies in Surface Science and Catalysis, Vol. 101 9 1996 Elsevier Science B.V. All rights reserved.

317

Ruthenium as Catalyst for Ammonia Synthesis M. Muhler*, E Rosowski, O. Hinrichsen, A. Hornung and G. Ertl Fritz-Haber-Institut der Max-Planck-GeseUschaft Faradayweg 4-6, D- 14195 Berlin (Dahlem), Germany Five Ru-based catalysts were prepared to study the effect of the support and the role of the alkali promoter in NH8 synthesis: Ru/AlzOa, Ru/MgO, Cs-Ru/A1203 and Cs(K)-Ru/MgO. The catalysts were characterized by N2 physisorption, H2 chemisorption and XPS. The absence of chlorine- and sulphur containing compounds turned out to be important for the preparation of highly active catalysts. Power law expressions were derived from conversion measurements at atmospheric pressure and at 20 bar. For all catalysts, the reaction order for H2 was found to be negative suggesting that a Pr% / PrI2 ratio in the feed gas higher than 1 / 3 would be favourable for industrial NHs synthesis at high pressure. The microkinetic analysis of the temperature-programmed desorption and adsorption of N2 and of the kinetics of isotopic exchange demonstrated the enhancing influence of the Cs promoter on the rate of N2 dissociation and recombination. XPS measurements after thorough reduction revealed a shift of the Ru 3d5/2 peak to lower binding energy by about 1 eV in the presence of Cs suggesting an electronic promoter effect. 1. Introduction

Alkali-promoted Ru-based catalysts are expected to become the second generation NHs synthesis catalysts [ 1]. In 1992 the 600 ton/day Ocelot Ammonia Plant started to produce NHa with promoted Ru catalysts supported on carbon based on the Kellogg Advanced Ammonia Process (KAAP) [2]. The Ru-based catalysts permit milder operating conditions compared with the magnetite-based systems, such as low synthesis pressure (70 - 105 bars compared with 150 - 300 bars) and lower synthesis temperatures, while maintaining higher conversion than a conventional system [3]. In spite of the industrial importance, relatively few studies in the catalytic literature deal with the kinetics of NH3 synthesis over supported Ru catalysts [4-8]. These studies focus on the inhibiting influences of PH2 and PNH8 on the rate of NH3 formation. Alkali promotion was found to decrease the inhibition by PNH8 significantly thus causing an increase in the inhibiting ' effect of PH2 [5-7]. The reaction order for N2 was found to be essentially unity indicating that the dissociative chemisorption of N2 is the rate-determining step (rds) in the overall mechanism. On multiply promoted Fe catalysts, H2 has a positive reaction order due to high coverages of adsorbed atomic nitrogen (N-.) [9]. In our laboratory a systematic study is in progress aiming at a detailed understanding of the "Corresponding author

318 catalytic phenomena involved in the synthesis of ammonia on ruthenium. The following catalysts were prepared from Rua(CO)x2 and high-purity supports: Ruthenium supported on MgO (Ru/MgO) and AlzOa (Ru/A12Oa), potassium promoted Ru/MgO (K-Ru/MgO) and cesium promoted Ru/MgO (Cs-Ru/MgO) and Ru/AIzOa (Cs-Ru/AlzOa). The catalysts were characterized by N2 physisorption (BET area), H2 chemisorption and X-ray photoelectron spectroscopy (XPS). This study presents kinetic data obtained with a microreactor set-up both at atmospheric pressure and at high pressures up to 50 bar as a function of temperature and of the partial pressures from which pov~er-law expressions and apparent activation energies are derived. An additional microreactor set-up equipped with a calibrated mass spectrometer was used for the isotopic exchange reaction (IER) 2aN 2 + a~ 2 = 2 29N2 and the transient kinetic experiments. The transient experiments comprised the temperature-programmed desorption (TPD) of N2 and H2. Furthermore, the interaction of N2 with Ru surfaces was monitored by means of temperature-programmed adsorption (TPA) using a dilute mixture of N2 in He. The kinetic data set is intended to serve as basis for a detailed microkinetic analysis of NH3 synthesis kinetics [ 10] following the concepts by Dumesic et al. [ 11].

2. Experimental The catalysts were prepared from high purity A12Os (99.99%, Johnson Matthey) or MgO (Puratronic, 99.996% metals basis, Johnson Matthey) and Rua(CO)I2 (Johnson Matthey) by wet impregnation in a rotary evaporator and subsequent heating in high vacuum following the procedures in ref. [12-14]. Details of the preparation are given in ref. [15]. The achieved metal loading was 5 wt. % Ru. The cesium-promoted catalysts Cs-Ru/MgO and Cs-Ru/A12Oz were obtained by impregnating the Ru/MgO or Ru/Al2Oa catalysts subsequent to heating in vacuum to 723 K with an aqueous solution of CsNOa (99.99 %, Strem). For the preparation of K-Ru/MgO, an aqueous solution of KNOa (99.997 %, Johnson Matthey) was used. The atomic ratios were Cs(K) / Ru - 1 / 1 for Cs(K)-Ru/MgO and Cs / Ru - 3 / 1 for Cs-Ru/A12Os. The Ru metal area was determined by volumetric H2 chemisorption in the quartz U-tube of an Autosorb 1-C set-up (Quantachrome) following the procedure described in ref. [ 16]. Prior to chemisorption, the catalysts were activated by passing 80 Nml/min high-purity synthesis gas (PN2 / Pn2 -" 1 / 3) from a connected feed system through the U-tube and heating to 673 K for alkali-promoted catalysts or to 773 K for alkali-free catalysts with a heating rate of 1 K/rain. The BET area was measured by static N2 physisorption in the same set-up. The kinetic experiments were carried out in an all stainless steel microreactor system with three gas lines which could be operated at pressures up to 100 bar. The gases used had the following purities: Ar 99.9993%, N2 99.9993%, H2 99.9993%. The feed gas was further purified by means of a self-designed guard reactor [17]. Gas analysis was performed by a non-dispersive infrared detector (BINOS, Fisher-Rosemount) which was calibrated by using a reference gas mixture (Linde). 138 mg of the 250/zm-800/zm sieve fraction were used for the kinetic experiments resulting in bed heights of less than 15 mm which prevented limitations by heat or mass transport. The reduction was carried out in synthesis gas using 40 Nml/min with a heating ramp of 1 K/min up to 673 K. The spectroscopic investigations were carried out in a modified LHS 12 MCD system. For the XPS measurements (Mg Kct 1253.6 eV, 240 W power) a fixed analyser pass energy of 108 eV

319 was used resulting in a resolution of 1.1 eV FWHM of the Ag 3d5/2 peak. The binding energy scale was calibrated using EB(AU 4fr/2) - 84.0 eV. The samples were activated in a directly attached preparation chamber (base pressure < 10-Smbar) from which the sample could be transferred into the UHV analysis chamber (base pressure 1.10 - l ~ mbar) within 1 min. The reduction was carried out in 1000 mbar synthesis gas by heating with 2 K/min to 673 K followed by 3 h NHa synthesis at this temperature. The synthesis gas mixture was replaced several times during the reduction. Charging was corrected using Mg 2s at 88.1 eV as internal standard [14]. Quantitative data analysis was performed by subtracting stepped backgrounds and using empirical cross sections [ 18]. 3. Results and Discussion

As expected, 3,-A12Os (BET area 110 mS/g) turned out to be the more stable support with a higher surface area than MgO (BET area 52 mS/g). The BET area of Ru/AlsOa was found to be 104 m2/g after NHa synthesis at 773 K which decreased significantly to 70 mS/g as a result of cesium impregnation. After NHa synthesis at 773 K, the specific area of Ru/MgO was observed to be 25 mS/g compared with 52 mS/g found for the MgO support. Cesium impregnation caused a further decrease in specific area to 23 m2/g. Table 1 Results of the H2 chemisorption measurements after NHa synthesis based on H/Ru = 1/1. NH3 synthesis was run at 773 K with Ru/MgO and Ru/AlsOa, and at 673 K with all alkali-promoted catalysts. The mean particle size was calculated assuming spherical particles. Catalyst H2 monolayer Metal area Dispersion Particle size / ~mol Hs/g / mS/g /% /nm Ru/MgO 130 12.9 53 1.9 Ru/AI203 118 11.7 48 2.1 Cs-Ru/MgO 69 6.8 28 3.6 Cs-Ru/AlsOs 100 9.9 41 2.5

The H2 chemisorption results are summarized in table 1. The Hs monolayer capacities were used to derive Ru metal dispersions and mean particle sizes assuming spherical particles. On both MgO and AlsOa, the impregnation with Rua(CO)12 resulted in mean particle sizes of about 2 nm after NHa synthesis at 773 K. It is remarkable that about the same Ru metal areas were obtained on MgO and A12Oa in spite of the largely differing BET areas of the supports. For the Cs-Ru/MgO catalyst, the amount of chemisorbed hydrogen was found to be reduced by about a factor of two. XRD measurements and TEM images revealed that the decrease in metal area is indeed due to sintering of the Ru metal particles [15]. The Ru/AlsOa catalyst was not significantly affected by the impregnation with CsNOa as shown by the increase in particle size from 2.1 nm to 2.5 nm derived from H2 chemisorption. The results of the conversion measurements at atmospheric pressure using 138 mg catalyst are shown in fig. 1. The following sequence in catalytic activity was observed: Cs-Ru/MgO > Ru/MgO > Cs-Ru/AlsOa > Ru/AlsOa. It is noteworthy that the catalytic activity of the Cs-Ru/MgO catalyst exceeds significantly the catalytic activity of a multiply promoted iron

320 catalyst (trace D in fig. 1A). A recent kinetic study provides evidence that the MgO support acts as alkaline earth promoter creating promoted sites at the interface [19]. The influence of the akali promoter is shown in fig. lB. The traces were obtained with somewhat deactivated catalysts after several weeks of NHa synthesis. Cesium ttmaed out to be a better promoter than potassium in agreement with the results obtained by Aika et al. [14]. The same authors furthermore suggest that Cs promotion is less efficient on A12Os since CsOH interacts mainly with the acidic support whereas on the basic MgO support more CsOH should be in contact with the Ru metal particles [ 14].

0.8

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-A

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Z~Cs - Ru / MgO

0.6

e'-

._o r tO ~ 0.4 0 0

-r e-

~

%%

0.2

%%%

w

,

550

l

,

600

,

,

l

l

,

650

, ,

,

i

,

700

Teml:~mture / K

,

,

l

l

,

750

I

4,

1

,

,

#

550

,

,

9

,

|

600

,

,

,

1

|

,

650

J

,

1

l

,

l

700

Temperature / K

Figure 1. NHa concentration in the reactor effluent gas using a total flow of 40 Nml/min with

PN2 / PI-I2 - 1 / 3 at atmospheric pressure. Traces A-E in fig.lA (from bottom to top) were obtained with Ru/AlzOa, Cs-Ru/AlzOa, Ru/MgO, a multiply promoted iron-based catalyst, and Cs-Ru/MgO. The corresponding NI-Ia equilibrium concentration is displayed as dashed line. Traces A-C in fig.1B (from bottom to top) were obtained with Ru/MgO, K-Ru/MgO, and Cs-Ru/MgO.

It is known that chlorine acts as severe poison for NHa synthesis [20,21]. Hence recent kinetic studies used chlorine-free Ru precursors like Rus(COh2 [8,22] or Ru(NO)(NOs)a [7]. In addition to chlorine, the presence of sulphur was found to poison Ru catalysts. Fig. 2A demonstrates that both poisons may originate from the Ru precursor. The binding energies for the C1 2p peak and of the S 2p peak observed for Ru prepared form RuO2 are typical for chloride and sulfide anions, respectively [23]. Ru prepared from Rua(CO)x2 was found to have a significantly higher purity. As shown in fig. 2B, sulphur and chlorine impurities can also originate from the support. The XPS data of MgO with a purity of 98 % reveal the presence

321 of chloride and sulphate anions which are essentially absent in MgO with a purity of 99.999 %. Hence it is mandatory to use high-purity Ru precursors and supports to prepare poison-free catalysts.

A

I XPS

C, 2 p

,- ~t~l(

Ru from

d

Ru from

I 198.6 e V RU3(00)12 , .... , . . , //L... 200 195 170

J .... 205

162.1 e V , i i , , , i .... 165 160

, 155

Binding Energy / eV

!

MgO

98%

~

MgO 99.999% ,

,

.

I

200

,

,

,

,

i

.

190

.

.

.

i

.

.

.

.

180

i

170

!

i

i

!

I

!

'

'

160

Binding Energy / eV

Figure 2. XPS C1 2p and S 2p data obtained with RuO2 and Rua(CO)x2 after reduction in synthesis gas at 1 bar and 673 K and transfer in UHV (Fig. 2A, upper half) and with two different MgO supports with purities of 98% and 99.999%, respectively (fig. 2B, lower half).

The quantitative XPS results obtained after reduction in synthesis gas are summarized in table 2. The observed ratios of O / Mg = 1.1 / 1 and O / A1 = 3.4 / 2 are in reasonable agreement with the stoichiometric ratios. The Ru concentration determined by XPS is somewhat higher for Ru/MgO than for Ru/A1203 in agreement with the results obtained with Hz chemisorption. Cs impregnation leads to a stronger decrease in the amount of Ru observed by XPS for Cs-Ru/MgO than for Cs-Ru/A12Oz due to the sintering of the Ru metal particles on MgO. The decrease of the Ru / support ratio observed for both catalysts may also be due to the shielding of the

322 Table 2 Surface composition determined by XPS after reduction in synthesis gas at 1 bar at 673 K and transfer in UHV. Catalyst Cs 3d O ls Ru 3p Mg 2s A1 2s Ru/MgO 50.4 3.3 46.4 Ru/A12Os 61.9 2.1 36.0 Cs-Ru/MgO 5.7 57.9 1.2 35.2 Cs-Ru/Al2Os 6.3 58.8 1.6 33.2

underlying Ru metal particle by a CsOH overlayer. Although a Cs / Ru ratio of 3 / 1 was used for the preparation of Cs-Ru/Al~Os, the ratio observed by XPS is similar to the ratio found for Cs-Ru/MgO. The latter catalyst was prepared with a Cs / Ru ratio of 1 / 1. This result indicates that the excess of Cs used for Cs-Ru/Al2Os has reacted with the bulk forming ternary phases with A12Os. Table 3 Power law exponents r - l~m,-~tt8 .PaN2.P~2 and apparent activation energies as a function of the total pressure determined in the given temperature range. The accuracy of the determination of the power law exponents and of the apparent activation energy is about q-0.1 and -4- 5 kJ/mol, respectively. Catalyst Pressure Temperature range c~(NHa) j3(N2) 7(H2) E,, / bar /K / kJ/mol Ru/MgO 1 513 - 603 -0.3 0.8 -0.3 69 20 573 - 663 -0.3 1.0 -0.5 78 Ru/AlzOs 1 593 - 663 -0.4 0.9 -0.1 70 20 573 - 688 -0.5 0.9 -0.3 76 Cs-Ru/MgO 1 498- 570 0.0 0.7 -0.7 96 20 550 - 630 0.0 0.8 -0.9 109 Cs-Ru/AlzO3 1 543 - 608 0.0 0.7 -0.6 103 20 573 - 663 0.0 0.9 -0.6 101

The results of the conversion measurements are summarized in table 3. The power law exponents and the apparent activation energies were derived following the analysis given in ref. [5]. The reaction orders of NHa and the reaction orders of N~ and H~ were determined by varying the synthesis gas flow between 40 Nml/min and 160 Nml/min and by varying the N2 / H2 ratio between 3 / 1 and 1 / 3 using a total flow of 120 Nml/min, respectively. Both determinations were carded out in the temperature range specified in table 3 ensuring the measurements to be in the kineticaUy controlled regime far from equilibrium. From the data shown in table 3, it is evident that the effect of Cs promotion on the power law kinetics is twofold: First, the reaction order for NHs is changed to essentially zero, and secondly, the apparent activation energy is higher by more than 20 kJ/mol in the presence of Cs. Contrary to the results obtained by Aika et al. [5], the reaction order for H2 was negative for all catalysts investigated. The positive reaction order for H~ reported by Aika et al. [5] for

323 Ru/Al2Oa and Ru/MgO may be due to the presence of chlorine originating from RuCla used for catalyst preparation. Fig. 3A shows the effluent NHa concentration observed for Ru/MgO as a function of reaction temperature for three different PN2 / Pn2 / P A r ratios at 20 bar total pressure. It is obvious that the reaction orders for N2 and H2 have opposite signs. Fig. 3B illustrates that the reaction orders for N2 and H2 partly compensate each other in the kinetically controlled temperature regime. Hence an increase in total pressure with a constant PN2 / Pn2 = 1 / 3 ratio does not lead to a significant increase in conversion at lower temperatures. For the application of alkali-promoted Ru catalysts under industrial synthesis conditions, it is necessary to find a compromise between kinetics and thermodynamics by increasing the PN2 / PH= ratio. The optimum observed for Cs-Ru/MgO prepared from Cs2COa at 50 bar is at about PN2 / Pn2 = 40 / 60 [15]. The high NHa concentration of about 8 % obtained with 0.138 g catalyst using a total flow of 100 Nml/min clearly shows that Ru catalysts have indeed the potential to replace Fe-based catalysts in industrial synthesis [ 15].

,, ,,

1.8 1.6 0

>

1

A

, ,,,,,,,,,,,,,,,,,,,,,

_

, ,,

Ru / M g O

/P /

1220b~rl/min

cO

B -

Ru / M g O 40 N m l / m i n

/

1.4

~ 1.2

4.5

3.5 _

N2 : H2 :Ar

//~t~

C

/

\

3

_

o ta v o

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m 0.8 -r 7 -,-. 0.6 C :~

IJJ

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.

/ 7

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p = 1bar p -- 9bar p = 20bar p = 50bar

1.5

"

0.4

0.5

0.2 0

i

550

600

650

700

Temperature / K

750

I

I

l

s

550

I

1

1

[

600

l

1

1

1

650

1

I

l

I

700

Temperature / K

1

~

I

II

750

Figure 3. Dependence of the NHa effluent concentrations observed for Ru/MgO on the feed gas composition (fig. 3A, left figure) and on the total pressure. In fig. 3A, trace A was obtained with PN2 / Pn2 / P A r - 1 / 1 / 2, trace B with PN2 / Pn2 / P A r = 1 / 3 / 0, trace C with PN2 / Pn2 / P A r - 3 / 1 / 0, respectively, using a total flow of 120 Nml/min at 20 bar. In fig. 3B, traces A-D (from bottom to top) were obtained at 1 bar, 9 bar, 20 bar and 50 bar, respectively, using a total flow of 40 Nml/min with PN~ / Pn2 - 1 / 3.

The reaction orders for N2 observed for all catalysts were close to 1.0 indicating that the

324 dissociative chemisorption of Nz is the rate-determining step in NHs synthesis. The kinetics of the interaction of Nz with Ru/MgO, Ru/AlzOa and Cs-Ru/MgO have been studied recently by performing N2 TPD and N2 TPA experiments and by determining the rate of isotopic exchange

28N2 + a~

- 2

2aN2 [ 2 4 ] .

Table 4 Rate constants ki - Ai- exp(-Ei / RT) for N2 + 2 , - 2 N - - , . Units of A/are ( torr- s) -~ for the forward reaction and s-1 for the reverse reaction forward rate constant reverse rate constant catalyst preexponential activation energy preexponential activation energy factor (kJ/mol) factor (kJ/tool) Cs-Ru/MgO 7.4.10 ~ 33.0 2.0.10 x~ 137.0 Ru/MgO 7.4.10 ~ 48.0 1.5.101~ 158.0 Ru/AI20s 7.4.10 ~ 60.6 1.5.101~ 158.0 .

.

.

.

.

.

Table 4 summarizes the rate constants ki - - A i 9 exp(-E//RT) for the forward and the reverse reaction derived from our microkinetic analysis of the steady-state and transient experiments with the three catalysts, i.e. Cs-Ru/MgO, Ru/MgO, and Ru/A12Os catalyst [24]. The rate constants in table 4 for Ru/A12Os should be considered as initial rate constants since it was not possible to achieve a higher coverage of N - - , than 0.25. Furthermore, it was not possible to detect TPA peaks for Ru/AlzOs within the experimental detection limit of about 20 ppm. Ru/MgO is a heterogeneous system with respect to the adsorption and desorption of N2 due to the presence of promoted active sites which dominate under NHa synthesis conditions. The rate constant of desorption given in table 4 for Ru/MgO refers to the unpromoted sites [ 19]. The N~ TPD, N2 TPA and IER results thus demonstrate the enhancing influence of the alkali promoter on the rate of N2 dissociation and recombination as expected based on the principle of microscopic reversibility. Adding alkali renders the Ru metal surfaces more uniform towards the interaction with N2. On polycrystalline Ru samples, IR measurements by Aika and Tamaru [25] revealed the influence of the alkali promoter on the stretching frequency of N2 -- * which was interpreted in the frame of a charge transfer mechanism. XPS should be the appropriate technique to detect charge transfer from the promoter to the Ru metal clusters. The Ru 3d spectrum of the Ru/MgO precursor after heating in vacuum to 723 K in order to decompose the adsorbed Ru carbonyl compounds is shown as lower trace in fig.4. The binding energy of the Ru 3ds/~ peak indicates that Ru is not yet reduced to the metallic state. Furthermore, the intensity ratio of the Ru 3da/2 and Ru 3d5/2 peaks shows that significant amounts of carbon compounds are present giving rise to overlapping C ls peaks at about 285 - 290 eV. After reduction (trace in the middle), the binding energy of the Ru 3d5/2 peak was found to be 280.0 eV indicating complete reduction to Ru metal. After reduction of the Cs-Ru/MgO catalyst, the Ru 3d speaks were observed to be shifted by 1 eV to lower binding energy (top trace in fig.4). It has to be noted that the Mg 2s peak had to be used as internal standard (EB(Mg 2s) - 88.1 eV) to correct for charging. However, the MgO bulk should not be affected by cesium impregnation. The XPS shift is influenced by many factors like the extraatomic relaxation energy which might change due to the presence of a Cs+O coadsorbate layer resulting from the decomposition of presumably CsOH as mobile

325

1 eV "~/~

XPS Ru3d

8000

r r

6OOO

r r-

-'-" C: 4000

2000

0 295

I

290

285

' x_..,

I ,

280

275

Binding Energy / eV

Figure 4. XPS Ru 3d data observed for the Ru/MgO catalysts. The Ru 3d spectra (from bottom to top) were obtained with the precursor after heating in high vacuum to 773 K , after reduction in 1 bar synthesis gas up to 773 K, and after impregnation with aqueous CsNOa solution and subsequent reduction in synthesis gas up to 673 K.

species. Since this shift was only observed after thorough reduction in the directly attached preparation chamber with rapid transfer in UHV, it seems plausible to assume that the treatment at 673 K in 750 mbar H2 (purity 99.9999 %) caused a partial reduction of the Cs+O coadsorbate layer thus creating oxygen vacancies which might serve as electron-donating adsorption sites for N2. Further studies are in progress to clarify this hypothesis. 4. Conclusions The preparation of Ru-based catalysts from high-purity supports using Rus(CO)xz followed by impregnation with aqueous Cs solution was shown to result in stable and active NHs synthesis catalysts. Cs-Ru/MgO was found to have a higher catalytic activity at atmospheric pressure than a multiply promoted Fe-based catalyst. Power law expressions were derived from conversion measurements at atmospheric pressure and at 20 bar. For all catalysts, the reaction order for H2 was found to be negative suggesting that a higher PN2 / Px2 ratio in the feed gas than 1 / 3 would be favourable for industrial NHa synthesis at high pressure. Studying the kinetics of the interaction of N2 with the Ru catalysts revealed that the Cs promoter enhances both the rate of dissociative chemisorption and the rate of recombinative desorption. Ru catalysts were found to be rather inactive for NHa synthesis without alkali

326 promotion. Ru/MgO turned out to be a heterogeneous system with respect to the adsorption and desorption of N2 due to the presence of promoted active sites which dominate under NHa synthesis conditions. Adding alkali renders the Ru metal surfaces more uniform towards the interaction with N2. XPS results provide evidence for an electronic promoter effect. REFERENCES 1. S.R. Tennison, in Catalytic Ammonia Synthesis, Plenum Press, New York, (Ed. J.R. Jennings), 1st. ed. (1991 ) 303. 2. EJ. Shires, J.R. Cassata, B.G. Mandelik, C.E van Dijk, U.S. Patent, 4479925 (1984) Oct. 30. 3. T.A. Czuppon, S.A. Knez, R.V. Schneider IN, G. Worobets, Chem. Engineering, March

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

18. 19. 20. 21. 22. 23. 24. 25.

1993, presented at the 1993 AICHE Ammonia Safety Symposium, Sept. 1993, Orlando, Florida, 100, No.3 (1993) 19. ER. Holzman, W.K. Shiflett, J.A. Dumesic, J. Catal., 62 (1980) 167. K. Aika, M. Kumasaka, T. Oma, O. Kato, H. Matsuda, N. Watanabe, K. Yamazaki, A. Ozaki, T. Onishi, Appl. Catal., 28 (1986) 57. H. Bails, M. Glinski, J. Kijenski, A. Wokaun, A. Baiker, Appl. Catal., 28 (1986) 295. J.U. Nwalor, J.G. Goodwin Jr., Topics in Catal., 1 (1994) 285. E Moggi, G. Albanesi, G. Predieri, G. Spoto, Appl. Catal. A, 123 (1995) 145. L.M. Aparicio, J.A. Dumesic, Topics in Catal., 1 (1994) 233. O. Hinrichsen, E Rosowski, M. Muhler, G. Ertl, Chem. Eng. Sci., Proc. of the 14th Int. Symposium on Chem. React. Engineering, (1996) accepted. J.A. Dumesic, D.E Rudd, L.M. Aparicio, J.E. Rekoske, A.A. Trevino, The Microkinetics of Heterogeneous Catalysis, ACS Professional Reference Book, Washington, DC, (1993). H. Kn6zinger, Y. Zhao, B. Tesche, R. Barth, R. Epstein, B.C. Gates, J.E Scott, Faraday Discuss. Chem. Soc., 72 (1982) 53. E Moggi, G. Predieri, G. Albanesi, S. Papadopulos, E. Sappa, Appl. Catal., 53 (1989) L 1. K. Aika, T. Takano, S. Murata, J. Catal., 136 (1992) 126. E Rosowski, A. Hornung, O. Him'ichsen, D. Herein, M. Muhler, G. Ertl, Appl. Catal., (1996) in preparation. R.A. Dalla Betta, J. Catal., 34 (1974) 57. B. Fastrup, H.N. Nielsen, Catal. Lett., 14 (1992) 233. Practical Surface Analysis, John Wiley, Chichester, (Ed. D. Briggs, M.E Seah) , 2nd ed. (1994). E Rosowski, O. Hinrichsen, M. Muhler, G. Ertl, Catal. Lett., (1996) accepted. W.K. Shiflett, J.A. Dumesic, Ind. Eng. Chem. Fundam., 20 (1981 ) 246. S. Murata und K.-I. Aika, Appl. Catal. A, 82 (1992) 1. Y. Kadowaki, S. Murat~ K.-I. Aika, Stud. Surf. Sci. Catal., Elsevier Science Publishers, (Ed. L. Guczi, E Solymosi, E Tetenyi), 75 (1993) 2055. J.E Moulder, W.E Stickle, EE. Sobol, K.D. Bomben, Handbook of X-ray Photoelectron Spectroscopy, Perkin-Elmer, (1992). O. Hinrichsen, E Rosowski, A. Hornung, M. Muhler, G. Ertl, J. Catal., (1996) submitted. K. Aika, K. Tamaru, in Ammonia: Catalysis and Manufacture, Springer-Verlag, Berlin, (Ed. A. Nielsen), 1st ed. (1995).

J.W. Hightower, W.N. Delgass, E. Iglesia and A.T. Bell (Eds.) 11th International Congress on Catalysis - 40th Anniversary

Studies in Surface Science and Catalysis, Vol. 101 9 1996 Elsevier Science B.V. All rights reserved.

327

SURFACE-S~UCTURE-DEPENDE~ REACTION PATHWAYS OF M E T H Y L G R O U P S O N NI(100) and NI(111) S U R F A C E S Richard B. Hall, Miguel Castro*, Chang Min Kim*, Charles A. Mimst Exxon Research and Engineering, Rt. 22E, Annandale, N.J. 08801 *Department of Chemistry, University of Puerto Rico, Mayaguez, Puerto Rico *Department of Chemistry, Kyung Pook National University, Daegu-Si South Korea "tDepartment of Chemical Engineering, University of Toronto, Toronto, Canada 1. Introduction-

Reactions of hydrocarbon fragments such as methyl groups on transition metal surfaces play an important role in a number of catalytic processes, including oxidative coupling of methane, methanation of syngas, steam reforming, and partial oxidation of methane to syngas. Although important, our knowledge of the fundamental mechanisms and kinetics of their reactions is limited. Because of the significance of these reactions, there are an increasing number of investigations of the fundamental reaction steps of small alkyl fragments on well characterized surfaces. Several methods have been used for the preparation of adsorbed alkyl radicals and for the study of their reactions. Most frequently, surface-bound alkyl fragments are produced by decomposing alkyl halides on the metal surface. This approach has been quite successful in characterizing reaction pathways of C1 and C2 fragment reactions on Ni [ 1-3], Pt [ 1] and Cu [46] surfaces. However, it intrinsically involves co-adsorbed halide atoms. Significantly, the co-adsorbed halide limits the coverage (i.e. concentration) of alkyl groups that can be adsorbed on the surface. As we will show, the dominant reaction pathway, at least on nickel surfaces, is a function of surface coverage. High coverages are likely to be more representative of reactions under high pressure, commercial process conditions, and the use of alkyl halide precursors preclude studies of high coverages. Surface methyls have also been synthesized by the collision-induced dissociation of methane physisorbed on Ni(111) surfaces.[7, 8] This approach avoids the effects of coadsorbates other than hydrogen, and a number of aspects of the reaction and decomposition of CH3 and CH fragments on Ni(111) have been determined.J9] However, the method is relatively complex and best suited for study of low coverages. Recently, Stair and coworkers [10, 11] developed a method to produce gas-phase methyl radicals, and used this to study reactions of methyl groups on Pt surfaces [12] and on molybdenum oxide thin films [ 13]. In this approach, methyl radicals are produced by pyrolysis of azomethane in a tubular reactor located inside an ultrahigh vacuum chamber. This method avoids the complications of co-adsorbed halide atoms, it allows higher coverages to be reached, and it allows the study of reactions on oxide and other surfaces that do not dissociate methyl halides effectively. We have adopted this last approach. We report here results from studies of the reaction of methyl groups on two different nickel single-crystal surfaces, Ni (100) and Ni(111). Nickel is

328 of particular interest because it is used as a commercial catalyst in methanation [14], steam reforming [15], and partial oxidation of methane to syngas.[16] These processes have many fundamental reaction steps in comrr~n, and our goal is to understand these fundamental reaction steps and to develop detailed kinetic models of these processes. We report here on some of the mechanistic trends and describe an unexpected dependence of the reaction pathway on metal surface structure. A more quantitative kinetic analysis will be presented in a subsequent paper.

2. ExperimentalThe experiments were carried out in a modified Leybold-Heraeus stainless steel ultrahigh vacuum chamber which has been described in detail elsewhere [ 17]. Briefly, it has a Balzers mass spectrometer for residual gas analysis (RGA) and temperature programmed desorption (TPD) measurements, a hemispherical electron energy analyzer and electron gun for AES measurements, a Varian LEED spectrometer, and a differentially pumped ion gun for sputtering and low energy ion scattering (LEIS). Background pressures below 1 x 10-10 torr are achieved by a combination of turbo molecular and liquid nitrogen cooled titanium sublimation pumps. The single-crystal nickel sample (either (100) or (111) orientation) was mounted on a liquid nitrogen cooled sample manipulator via two tantalum wires spot-welded to the side of the crystal. The substrate temperature was measured using a chromel-alumel thermocouple spotwelded to the side of the crystal. The crystal could be cooled to 110 K and resistively heated to 1200 K. The crystal was cleaned by cycles of Ne+ ion bombardment (primary energy of 1.0 KeV) for 5 minutes and annealing to 1000 K for 10 minutes until most impurities were removed, as detected by AES. Residual carbon was removed by exposing the surface at 300 K to 1 langmuir (L) of oxygen, followed by heating to 800 K. Trace oxygen was removed by reducing with hydrogen (lx10-7 Torr) at 800 K Several cycles were required before no carbon was detected by AES. Methyl radicals were produced by pyrolysis of azomethane (CH3N2CH3). Azomethane was synthesized as described earlier [ 18]. It was purified periodically by freeze-pump cycles at 77 K, and the gas purity verified by RGA. The methyl radical source was similar to that developed by Stair and coworkers.[ 10, 11] The source was made of a quartz tube with 3 mm OD and 1 ram ID, resistive heating was supplied by means of a 0.25 mm diameter tantalum wire wrapped outside the quartz tube. The length of the heating zone was 4 era, recessed from the end of the tube by 1 era. An alumina tube around the outside of the heating zone served as a radiation shield. Azomethane was admitted to the hot tube at a pressure of l x10-8 to l x10-7 Torr via a high-vacuum precision leak valve. The pyrolysis tube was maintained at about 1200 K, adequate to decrease the major peaks in the mass spectrum of the parent azomethane at 58 and 43 ainu by at least a factor of 100. Ideally, pyrolysis of azomethane produces only methyl radicals and N2 through cleavage of the relatively weak methyl-N2 bonds. However, secondary reactions of the methyl radicals are unavoidable. These reactions produce ethane, methane and hydrogen in addition to the methyl and nitrogen. The product distribution was 16% CH3, 16% CI-14, 13% C2H6, 39% N2, and 17% H2 (by mole %), based on mass specmun intensities, corrected for sensitivity factors of the individual species. The distribution varied only slightly over the range of pressures used in this work, however it changes significantly at higher pressures. In this paper, the dose of methyl is reported as Langmuirs (L) of exposure to the total product gas pressure. Surface temperature was held at 120 K during dosing. At this temperature, methane, ethane and nitrogen are too weakly bound to adsorb on the surface. Control exper~maents confLrmed that pressures of methane, ethane and nitrogen 10 times higher than produced in the pyrolysis source do not interfere with methyl adsorption or reaction. Hydrogen, on the other hand, does adsorb during dosing of the methyl radicals. It is possible that the hydrogen concentration is comparable to the methyl concentration in experiments on the Ni(100) surface because the sticking probabilities are comparable. In experiments on the Ni(111) surface, it is

329 likely that the hydrogen concentration is less than 10% of the methyl concentration because the sticking probability of hydrogen on this surface is only around 0.01. In experiments on the surfaces with chemisorbed oxygen or multilayer oxide films, hydrogen does not adsorb. Calibration of the amount of desorbing hydrogen was carried out by comparing the hydrogen TPD peak area with that measured for a saturation dose of hydrogen, which is known to give 1 H atom per surface Ni atom.[19] This is defined as 1 monolayer (ML) coverage and corresponds a surface density of 1.9 x 1015 atoms/era2 for N i ( l l l ) and 1.6 x 1015 atoms/era2 for Ni(100). The coverages of all species are referenced to these values. Calibration of the amount of desorbing CH4 was carried out by comparing the methane TPD peak area to that measured for a saturation exposure of CO (.67 ML) scaled by the ratio of the mass spectrometer sensitivities, and UHV-chamber residence times for these two gases. Calibration of the amount of residual carbon was carried out by comparing the (2(272 eV)/Ni(848 eV) AES peak to peak ratio with the one measured after decomposition of a saturation amount of ethylene on Ni(111), which gives a surface carbon coverage of I/4 monolayer.[20] Self-consistent mass balances were achieved using these calibrations in experiments in which CH3 and CO were coadsorbed in varying amounts.

3. Results and Discussion3.1 CH3 on Ni(100) A representative temperature programmed desorption (TPD) specmma of products resulting from reactions of methyl on Ni(100) is shown in Fig. 1. The only gas-phase products detected were methane (16 and 15 ainu curves) and hydrogen (2 ainu curve). No C2 or higher hydrocarbons were observed. A portion of the adsorbed methyl groups decomposed on the surface by reactions R1 through R3, CH3 CH2 CH

--> CH2 + H(S) --> CH + H(S) --> C(S) + H(S)

(R1) (R2) (R3)

where (S) refers to surface-bound species. (Because only surface-bound hydrocarbon fragment groups are discussed here, this designation for surface species is omitted). The m o u n t of CH3 decomposed is determined by measuring the amount of residual surface carbon by Auger spectroscopy. The peak in the methane TPD curve occurs at 230 K. The appearance of gas-phase methane marks the temperature at which methane is formed on the surface because the reaction temperature is well above the desorption temperature for methane. As will be discussed below, the dominant mechanism for methane formation on Ni(100) is: CH3 + H(S) --> CI-I4(g)

(R4)

Hence, the appearance of gas-phase methane marks the occurrence of a C-H bond-breaking step. In the hydrogen TPD ~ m m a , the lower-temperature peak occurs at 355 K, characteristic of desorption of surface hydrogen from Ni(100). Hence it is defined by the desorption kinetics and does not provide information on C-H bond-breaking reactions that might occur at lower temperatures. The higher-temperature desorption peak at 390K on the other hand occurs above the temperature for hydrogen desorption and is believed to be related to decomposition of surface the most stable fragment, CH, or a dimer thereof to carbon and hydrogen. The spectra shown in fig. 1 are for a methyl surface coverage of about 0.1 ML. The results obtained at this coverage are in excellent agreement with earlier experiments involving decomposition of methyl iodide to produce surface-bound methyl groups.[ 1, 3, 21] The methyl coverage as a function of dose is shown in fig. 2. The methyl coverage is

330 1 E 3

v

~. 0.5 ec-

15 AMU _(CH_4)

m

0

i

i

i

,

,

i

i

i

100

i

I

t

.

.

.

.

.

.

300

.

.

i

.

.

.

.

.

.

500

.

.

.

i

i

.

i

.

700

Temperature, K Figure 1 Mass spectrometer (MS) intensities versus surface temperature in a representative temperature programmed desorption (TPD) profile of methyl on a clean Ni(100) surface. Total exposure was 1 L . Adsorption temperature was 105 K and heating rate was 3 K/s. Methane is produced near 225 K. Hydrogen desorbs in a wide temperature range, 260 - 420 K. taken to be the sum of the methane observed in TPD and the residual surface carbon measured after annealing to 600 K. The maximum methyl coverage is 0.42 ML (+ 10%). This is about 2 times higher than can be achieved via the decomposition of methyl iodide [ 1]. This qualitatively seems reasonable since methyl groups and I atoms have roughly the same diameters on the surface, and there is necessarily 1 iodine atom per methyl group. The uptake of methyl roughly follows langmuirian adsorption kinetics with a sticking probability per collision with the clean metal surface of about 0.1. The sticking probability determined from these experiments has an uncertainty of about a factor of 2, due primarily in the uncertainty in determining the absolute partial pressure of methyl groups coming from the pyrolysis source.

0.5'

I

I

I

I

~'~ 0.4 r

I

I

"="

~'

0.3 0>

ro -r-

0.2 0.1 0.0 0

2

4

I

I

I

I

6

8

10

12

C H3 Exposure (L)

--

Figure 2 Methyl coverage in monolayers (ML) as a function of methyl exposure, in langmuirs (L) of mixture of gases from pyrolysis source. Symbols are data showing the sum of methane formed plus residual carbon. Consistent values are obtained by summing hydrogen appearing in CH4 and H2 gas-phase products. Solid line is a guide to the eye.

331 The dependence of methane formation on methyl coverage is shown in fig. 3. The yield of methane increases as the coverage increases. At coverages below 0.05 ML, little methane is formed. Nearly all of the methyl groups decompose. At saturation coverage, about 0.14 ML of the methyl groups form methane, and roughly 0.27 ML decompose An important trend to note in fig.3 is that the peak in the TPD spectrum shifts to higher temperatures as the coverage increases. If the rate determining step is the rupture of the CH2-H bond, and this step is f'trst order, the peak in the TPD curves should occur at the same temperature. The shift to higher temperatures is at least in part due to the fact that higher temperatures are required for the CH2-H bond breaking reaction as the surface becomes more crowded. We find that preadsorbing carbon in various amounts up to 0.3 ML leads to a progressive shift to higher temperatures. The temperature shift observed with 0.3 ML of preadsorbed carbon is similar to that exhibited by curve h in fig. 3. The effects of st~ace crowding and co-adsorbed species will be reported in more detail elsewhere.

CH 4 / OH 3 / Ni(100) "S"

g :E

1

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

1O0

i

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

|

200 300 Temperature, K

....

.........

400

Figure 3 CI-h TPD from Ni(100) at various initial coverages of CH3. Curve a=0.025; b-----0.06; c--0.10; d=0.18; e---0.33; f=0.37; g=0.40 ML.

3.2 CH3 on N i ( l l l ) Methane formation from the reactions of methyl groups on a Ni(111) surface at various methyl coverages is shown in fig. 4. The saturation coverage we observe for methyl on Ni(111) is 0.32 M'J., (+ 10%). Within experimental error, this is the same surface density of methyl groups as the saturation coverage found for Ni(100) (because there are more surface Ni atoms per ML on Ni(111)). As with Ni(100), we fred that methane yield increases faster than linearly in methyl coverage. This is illustrated in fig. 5. In the lower panel, the yield (total amount formed) is plotted, and in the upper panel, the selectivity (fraction of adsorbed methyls going to methane) is shown. Again we see that below about 0.05 ML that little or no methane is formed. This is consistent with experiments of Ceyer and coworkers who observe little methane formation from low coverages of methyl groups on Ni(111), even in the presence of considerable co-adsorbed hydrogen.[7] At coverages above 0.1 ML, about 15% of the methyl groups desorb as methane, increasing to about 30% at coverages of 0.3 ML.

332

_>, i-

1.0 A

E

-r0

0

0.5

o,p 2" 0.0 u ~ - ~

0.08

0.0

2

-

I

!

-

-

-I-

=~

~ 0.04 200

300

400

Temperature (K) Figure 4 CH4 TPD from Ni(111) at various initial coverages of methyl. Curve a--0; b=0.03; c=0.08; d--0.11; e=0.13; f=0.24; g=0.26; h--0.29; i----0.31 ML.

0.00 m -=r" 0.0

J

J

I

0.1

0.2

0.3

Initial Methyl Coverage (ML) Figure 5 CH4 yield (bottom plot) and selectivity (top plot) versus initial CH3 coverage on N i ( l l l ) . Yield is the absolute amount of CI-I4 formed, selectivity is the fraction of adsorbed CH3 that goes to CI-I4

The shape of the TPD spectra are different than they were for Ni(100), and the peak occurs at slightly higher temperatures, about 245 K. Also, in contrast to the trend observed on Ni(100), on Ni(111) the peak in the methane TPD curve shifts to lower temperatures as the surface coverage is increased. This trend is frequently associated with a reaction step that is higher than first order in reactant coverage. It is an important indication that the dominant reaction pathway is probably different than it is on the (100) surface. Further evidence for different reaction pathways is obtained in isotope labelling experiments, illustrated in fig.6. Here we present a comparison of the effects of preadsorbing various amounts of deuterium with methyl groups on the two metal surfaces. For a reaction network consisting of reactions R1 through R4, the following trend would be expected. At low deuterium coverages, most of the hydrogen required for methane formation must come from decomposition of some fraction of the methyl groups. As the deuterium coverages increases, it will compete with increasing effectiveness for undecomposed methyl groups. The balance between these two pathways can be determined from the relative amounts of CH4 and CH3]) in the TPD spectra. (No methane with more than 1 deuterium atom is observed, even at the highest deuterium coverages. This is clear evidence that at least reaction R1 is irreversible. It also suggests that reactions R2 and R3 are irreversible. Further evidence for this is presented below.) The results obtained for Ni(100) are consistent with this trend. As the deuterium coverage increases from 0.1 ML to 0.7 NIL, CH3D yields increase, and CI-I4 yields decrease, reflecting an increasing contribution from reaction of methyl with surface deuterium, and decreasing contribution from reaction with hydrogen supplied from other methyl groups on the surface. Note especially that the TPD curves shift in opposite directions for each product. As the deuterium coverages increases, the peak in the CH3D spectrum shifts to lower temperature and the peak in the CH4 spectrum shifts to higher temperature. This is the qualitative behavior

333

~ II

04

_>, "~~

_>, "~~

I

I

I

I

)D(ML)

r~=-----.-.~.-,~--~-~- ~ - - . L - . ~ 0 . 1 5 -

=

,.

o.o o.o

150

200

250

300

350

Temperature (K)

400

Figure 6 CI-I~D (dotted curve) and ~

150

200

250

300

350

Temperature (K)

_

400

(solid curve) formed on Ni(]00) (left hand

figure) and Ni(111) (fight hand figure) at a fixed CHa initial coverage of 0.2 ML, and various amounts of preadsorbed deuterium (indicated in the figures, units are monolayers of deuterium atoms) expected for reactions R 1 through R4. In fact, a simple kinetic model that includes only these reactions can quantitatively reproduce the dependence on deuterium coverage of both the relative yields of CHaD and CH4 and the peak desorption temperatures. The details of this model will be published separately. The results obtained for a Ni(111) surface are significantly different. Most striking is that even with a 3 fold excess of deuterium relative to methyl, very little CHaD is produced. Reaction R4 appears to be a minor channel on Ni(111). Furthermore, because the deuterium coverage is nearly invariant over the entire temperature range over which methane is produced, it is unlikely that the CH4 that is formed comes methyl reaction with surface hydrogen generated by methyl decomposition. One would expect more CHaD than CI-I4 because the D coverage is at all times equal to or greater than the H coverage. The possibility that the relatively slow reaction of CHa with D(S) might be due to an unusually large deuterium kinetic isotope effect was tested by performing experiments with preadsorbed H(S). We observe no significant increase in the amount of methane formed in the presence of preadsorbed H atoms on Ni(111), indicating that D(S) and H(S) have similar reaction rates. CH4 must therefore come from a disproportionation or other reaction of two hydrocarbon fragments, and not by a reaction that involves surface-bound hydrogen atoms. Similar experiments have been done with 0.25 ML of deuterium and varying amounts of methyl. The CH4 yield as a function of methyl coverage at fixed D coverage is very similar to that found in the absence of coadsorbed deuterium. This is not surprising in that the results illustrated in fig. 6 show that the main peak in the methane TPD spectrum is not significantly affected by coadsorbed deuterium. Results obtained with 0.25 ML of coadsorbed deuterium are very similar to those shown in figs. 4 and 5. The CH4 yield has a higher than first order dependence on coverage. This cannot be attributed to the coverage dependence of reaction R4, since it makes only a minor contribution. It is more likely to be due to a bimolecular reaction

334 between two hydrocarbon fragments, which would exhibit a greater than f'wst order dependence on coverage. The mechanism for a non-surface-mediated hydrogen-tranfer from one hydrocarbon fragment to another cannot be identified in ambiguously from the current data alone. Possible reaction pathways include: CH3 + CH3 --> [C2H6] --> CH4(g) + CH2

(disproportionation)

R5

where, [C2H6]is a short-lived complex involving one or more bonds to the metal surface, and (g) designates a gas-phase product; CH3 + CH3 --> CI-h(g) + CH2

(abstraction)

R6

i.e., a prompt reaction in which an H atom is stripped from a surface bound CH3 group by one that is not; CH2 + CH3 --> [-CH2CH3] --> CH4(g) + CH

(methylene insertion)

R7

(methyne insertion)

R8

where [-CH2CH3]is a surface bound ethyl group; CH + CH3 --> [=CHCH3] --> CH4(g) + C and, CH2 + CH2---> [C2H4] --> CH4(g) + C

R9

We discount the likelihood of reaction R5 because two metal-bonded methyl groups have no molecular orbitals readily available for reaction. These groups are sp3 hybridized with all 4 orbitals fully occupied. The lowest lying unoccupied orbital is so much higher in energy that it is inaccessible, and there are no obvious interactions that might enable a rehybridization to lower this energy. An abstraction mechanism, R6, might proceed on an available orbital basis, but seems unlikely based on the following. A free, or nearly free, methyl group desorbing from the surface has an orbital available for forming a fourth C-H bond. The abstraction of a hydrogen from CHa(g) by CH3(g) is well known; the heat of reaction is zero, and the energy barrier is only about 50 kJhnol. Since the C-H bond strength of CH3 is less than that in CH4 by about 50 kJ/mol, this energy barrier could be relatively low. However, it is also necessary to break, or nearly break, the metal-methyl bond to provide the free methyl group. The chemisorption energy of methyl on Ni(111) has been estimated to be 160 kJ/mol.[22] This would make the overall reaction endothermic by about this amount, much too high to expect it to occur at 250 IC Furthermore, if the metal-methyl bond were weak enough for this process to occur, we would expect to see at least some free methyl groups in the TPD specmun. We were unable to detect any desorption of free methyl groups in these experiments, although it has been observed from Cu [6] and NiO [23] surfaces. The CH2 and CH groups generated by reactions R1 or R2 do have orbitals readily available for reaction, and the energy barriers are likely to be quite low. We postulate reactions R7 through R9 based on analogy with known organometallic or metal surface chemistry. The formation of ethyl groups from CH3 and CH2 fragments (the first step in reaction R7) has been observed on Cu surfaces in experiments in which the fragments were created by the dissociation of coadsorbed CH3I and CH212 respectively.[5] However, the dominant reaction pathway of ethyl groups is to undergo a 13-hydrogen elimination to give ethylene, not methane.[5] We would expect to detect at least some formation of ethylene in conjunction with reaction R7.

335 Since we do not observe any evidence for ethylene formation, we rule out reaction R7. Similarly, we rule out reaction R9, which should also produce at least some ethylene. We believe reaction R8 is the most likely route by which the methane is formed on Ni(111). CH has been observed from the decomposition of methyl at low coverages on Ni(111) [9, 24], and, in the absence of methyl groups, CH can dimerize to give C2H2.[9] In the presence of methyl groups, other C-C bond forming reactions should be expected. Support for this comes from the fact that ethylidyne (CCH3) has been observed in reactions of CH3 on P t ( l l l ) under conditions where both CH and CH3 are likely to present.[12] Ethylidyne formation is unusually favorable on Pt(111), and it is reasonable to expect that related C2 adducts formed on Ni(111) might rearrange to give methane according to reaction R8, rather than ethylidyne. Lastly, there is some evidence in the shape of the TPD curves shown in fig. 6 for the formation of a short-lived reaction intermediate that is a precursor to the methane formation on Ni(111). Note that the CH3D curve drops sharply at the same time (temperanwe) that CH4 production is at a maximum. A reaction network consisting of reactions R1 through R4 is not able to reproduce this behavior. We can envision only two ways in which CH3D production might decline while the CH4 formation rate is increasing. The fast is that there are two different kinds of CH3 on the surface, a minority of CH3's (e.g. those at a special site) that get used up in reacting with surface D, but cannot form CH4, and the rest of the CH3's that can form CH4. The second is that all of the CH3 groups react to form a C2 complex from which CH4 is formed, leaving no CH3 groups to react with surface D. We believe that the latter is a more reasonable explanation, and that reaction R8 is the pathway by which this is occurring. A more detailed description of the kinetic modeling of the reactions of methyl on Ni(111) and (100) will be presented in a subsequent paper. 4.0 Conclusions: We have characterized the reaction pathways of CH3 groups as a function of coverage on Ni(100)andNi(lll)surfaces. Relatively high coverages, up to 0.4 ML, have been investigated for the first time. On both surfaces, methane and hydrogen are the only gas-phase products detected. Methyl C-H bonds begin breaking at around 220 K, and methyl decomposition occurs in parallel with methane formation over a limited temperature range. At high coverage, up to 1/3 of the methyl groups form methane, and the remainder decompose, ultimately producing surface carbon and hydrogen gas. The dependence of the methane yield on methyl coverage is similar on the two surfaces, but the mechanism by which methane formation occurs is different. On Ni(100), the dominant mechanism is the reaction of CH3 and surface hydrogen. On Ni(111), the reaction of CH3 with surface hydrogen is relatively slow, and the dominant mechanism involves a hydrogen transfer between two hydrocarbon fragment groups. Experimental results suggest that CI-h is formed from a C2 reaction intermediate (CHCH3) from the reaction of CH with CH3. References: 1. F. Zaera, Ace. Chem. Res, 25 (1992): 260. 2. S. Tjandra and F. Zaera, J. Catal., 157 (1994): 598. 3. S. Tjandra and F. Zaera, Langmuir, 8 (1992): 2090. 4. C.-M. Chiang, T.H. Wentzlaff, C.J. Jenks, and B.E. Bent, J.Vac. Sci. Technol., A10 (1992): 2185. 5. C.-M. Chiang, T.H. Wentzlaff, and B.E. Bent, J. Phys. Chem., 96 (1992): 1836. 6. J.-L. Lin and B.E. Bent, J. Vac. Sci. Technol., A10 (1992): 2202. 7. A.D. Johnson, S.P. Daley, A.L. Utz, and S.T. Ceyer, Science, 257 (1992): 223. 8. M.B. Lee, Q.Y. Yang, S.L. Tang, and S.T. Ceyer, J. Chem. Phys, 85 (1986): 1693.

336 9. Q.Y. Yang, K.J. Maynard, A~D. Johnson, and S.T. Ceyer, J. Chem. Phys., 102, no. 19 (1995): 7734. 10. X.D. Peng, R. Viswanathan, G.H.J. Smudde, and P.C. Stair, Rev. Sci. Instrum., 63 (1992): 3930. 11. G.H.J. Smudde, X.D. Peng, R. Viswanathan, and P.C. Stair, J. Vac. Sci. Technol., A9 (1991): 1885. 12. D.H. Fairbrother, X.D. Peng, R. Viswanathan, P.C. Stair, M. Trenary, and J. Fan, Surf. Sci, 285 (1993): LA55. 13. G.H. Smudde, Jr., X.D. Peng, R. Viswanathan, and P.C. Stair, J. Am. Chem. So Pd/AI, Pd/Zr. At higher temperatures the NO conversion is enhanced for the Pd/AI-Zr-Ba and Pd/AI-Zr solids but the Pt-Rh/AI203 solid remains the best one. At the maximum of conversion, the activities obey the order: Pt-Rh > PdAIZrBa, PdAIZr > PdZr > PdAI Table 2 Light-off temperature, T50 (K), for NO in the presence of the CO-NO-O2-C3H6 and CO-NOO2-C3H6-H20 mixtures (s = 1.03) (flow rate 22 l/h) (4568 vpm CO, 623 vpm NO, 5611 vpm 0 2, 799 vpm C3H6, 10 vol. % H20 ). T50(NO)(K ) Pd/AI20~ Pd/ZrO2 Pd/AI20~- ZrO 2 Pd/AI203- ZrO2-BaO pt-Rh/Al203 ,,,

CO-NO-O2,C3H6 650 660 610 585 575

CO-NO-O2-C~FI6-H20 690 620 665 .. 590 5~/5

In the presence of water, with the exception of Pt-Rh/AI203, the best Pd-based catalysts contain ZrO 2 (Table 2, Figure 2), either alone (Pd/ZrO2) or associated with Ba (Pd/Al203-ZrOE-BaO). Above 673 K, Pd/ZrO 2 is the best catalyst for the NO reduction but the Pd/Al2Oa-ZrO2-BaO solid has the highest performances in the overall range of temperature. Near T50 the NO conversions obey the sequence: PtRh >Pd/AIZrBa> Pd/Zr > Pd/AIZr > Pd/Al.

349 At the maximum the sequence becomes: PtRh, Pd/Zr > Pd/AIZrBa > Pd/AIZr > Pd/AI It is noteworthy that the Pd/ZrO 2 and Pd/A1203-ZrO2-BaO activities are close to that of PtRh/AI203

.

100

.

.

.

.

.

.

.

.

.

.

.

.

.

.

.

.

.

.

.

.

.

.

.

.

.

.

.

, c~-."~ . ~ c , ~

A

qr~JP ~

9

~A |Ik,,

Q g=

o o

0 Z

50-

9

.

~ e

9

e" %* . , , . . - . - - * ~ ~. ........ 473,"" " 573 6 } 3 "" Temperature (K)

7;'73

Figure 2. NO conversion with the CO-NO-O2-C3H6-H20 (10 vol. %) (s = 1.03) mixture on: Pd/AI203, A Pd/ZrO2, m Pd/AI203-ZrO2 , ~ PtRh/AI203 and 9 Pd/AI203-ZrO2-BaO. Let us also notice that: i) the changes in the activity sequences with the temperature show the necessity to compare the activities at various temperatures, not only near the light-off temperature ii) since the activity sequences depend also on the mixture, for example on the presence or on the absence of H20, it is necessary to study complex realistic mixtures. Furthermore whatever the solid the N20 quantity remains moderate. 3.2. Influence of the support on the Pd electronic state

In order to control the effect of ZrO 2 on the electronic properties of Pd, an infrared study (using the CO probe molecule) and an XPS determination of the Pd binding energies have been performed. The infrared spectrum of CO irreversibly adsorbed at room temperature on the Pd/AI203 solid previously reduced at 673 K and evacuated at 623 K shows (Figure 3) the classic v CO bands assigned to lineary bonded CO (2065 cmq) and to multi-bonded CO (1960 and 1925 cm-l) on metallic Pd [2-7]. The low frequency bands are more precisely attributed to bridged species adsorbed on the (100) or (110) (1960 cmq) and (111) (1925cm q) faces of Pd crystaUites. The position of the bands is coverage-dependent with a downwards frequency shift as the CO coverage decreases upon evacuations at increasing temperatures. Upon evacuation, the linear and multibonded species are no longer seen at 423 K and 573 K, respectively. The spectra of CO irreversibly adsorbed at 298 K on Pd/ZrO 2 (Figure 3), Pd/AI203ZrO 2 and Pd/AI203-ZrO2-BaO reduced at 673 or 773 K are similar to the spectrum of CO on

350 Pd/A]20 3. Linear CO species (band at 2070 cm-1) and multi-bonded CO species (band between 1970 and 1945 cm -1 and shoulder between 1925 and 1920 cm -1) on Pd ~ are observed. The variations in wavenumbers are not strong enough to indicate electronic effects rather than changes in CO coverage.

1960 1925

. m , m ~

2065

t ~

1119'0

'

..2..2 0 0

'| ............2 "i 'l -O00

....... 1 8 0..... 0 i ' --'cm-1

Wavenumber

Figure 3. Infrared spectra of CO irreversibly adsorbed at 298 K on Pd/AI20 3 (a) and Pd/ZrO 2 (b) reduced at 673 K and evacuated at 623 K

/

Zr

31~/,

................. 349; .... ; Binding

Pd3d5/2

I

......341; ............... ; ............ 333'

; ..... eV

Energy

Figure 4. Original and decomposed XPS spectra of the Pd/AI203-ZrO 2BaO reduced "in situ" Pd 3d and Zr 3p photopeaks.

This has been corroborated by the complememary XPS study in which the spectra were recorded without air contacting the reduced samples. As already discussed, the binding energies are referred to an internal standard, ie., the AI 2p line of AI203 (73.8 eV) or the Zr 3d line of ZrO 2 (182.2 eV). The binding energies 0~d 3d5/2 and Pd 3d3/2 levels) are characteristic of non-modified metallic palladium ffigure 4) [8-10]. There is no detectable difference between the spectra of Pd/AI20 3 and the deconvoluted spectra Pd/Al2OyZrO2-BaO. In the case of Pd/ZrO 2, the low Pd content with respect to ZrO 2 and the closeness of the binding energies of Pd 3d and Zr 3p levels do not permit reliable deconvolution. We can conclude however that the electronic state of Pd ~ is not noticeably modified whatever the support.

351 3.3. Characterization of the supports From the previous results, it has been proven that the nature of the support, although it has no significant influence on the Pd electronic properties, modifies the catalytic properties of the solids. To permit a better understanding of these supports effects, the surface properties of the supports (in the presence of the metal) have been studied, in particular the acidic properties and the oxygen mobilities. The AI203 and ZrO 2 supports have been mainly onsidered.

As already described (1) the inhibition of the NO reduction by CO due to carbon deposits in mixtures containing hydrocarbons depends on the support, with the most acidic supports leading to higher amounts of carbon deposits. The nature (Bronsted or Lewis centers), the number, and the strength of the acidic sites of the Pd/AI203 and Pd/ZrO 2 solids have been checked using infrared spectroscopy of adsorbed pyridine and thermoprogrammed desorption of ammonia. The infrared spectra were recorded after equilibrating the reduced and evacuated solids with an excess of pyridine vapor and further evacuation at various temperatures. After evacuation at 423 K there is no more physically adsorbed pyridine. There is no characteristic band of pyridine adsorbed on Bronsted acid sites (no appearance of the 19b vibration at 154045 cm-1) [11,12]. The OH groups observed on the solids are thus non acidic. The existence of Lewis acid centers (coordinatively unsatured AP + or Zr 4+) is proven by the presence of the 19b vibration at 1440-50 cm-1 and of the 8a vibration at 1610-1620 cm-1. The absorbances of the 1440-50 cm -1 band show that the acidity difference between the Pd/AI203 and Pd/ZrO 2 solids is not significant. This is corroborated by a thermodesorption study of NH 3 adsorbed at 373 K, which shows that the quantity of adsorbed NH 3 is only slightly higher on Pd/AI203 (4.7 * 10-4 mol NH3/g solid) than on Pd/ZrO 2 (3.3 * 10-4 mol NH3/g solid). The different behavior between Pd/AI203 and Pc[/ZrO2 cannot therefore be explained only by difference in the acidic properties. Nevertheless after reduction of Pd/ZrO 2 and Pd/Al20-ZrO2-BaO , IR spectroscopy of adsorbed CO allows to detect, aside from the bands due to CO adsorbed on Pd ~ (2000-1850 cm-1), other absorption bands in the range 1200-1700 cm-1 characteristic of "carbonatecarboxylate-formate" structures [3, 13-16] located on the supports. These "carbonatecarboxylate" structures imply the formation of CO 2 arising either from the participation of surface oxygen species of the supports (since Pd has been totally reduced) according to the surface reaction COads + O s --> CO 2 + ~ (oxygen vacancy) or from the CO disproportionation. Such species (carbonates, carboxylates, formates ...) are not formed after CO chemisorption on Pd/AI20 3. With Pd/ZrO 2 previously reduced (no ionic Pd), the actions of C3H6 at 473 K or 673 K and of C3H8 at 673 K lead also to the formation of hydrogenocarbonates (1615, 1450, 1220 cm-1) due probably to the presence of"oxygen species" located on the support. To confirm this formation of reactive surface oxygen species and oxygen vacancies, thermoprogrammed oxidations and reductions have been performed on Pd/ZrO 2. On this Pd/ZrO 2, calcined under 0 2 at 723 K, reduced in flowing H 2 at 773 K and evacuated at 773 K, the oxidation is performed with a He-l% 02 mixture. At 298 K, the 0 2 consumption is weak (0.2 mmol.for 14.8 mmol. total Pd). This corresponds to the oxidation of

352 the superficial Pd according to the reaction: Pd s + 8902 --> PdsO. Upon heating from 298 K to 773 K (10 K/rnin), the 02 consumption reaches 10.3 retool, with the maximum of the peak at 658 K. The total 02 consumption (10.5 mmol. 02 for 14.8 mmol. Pd) exceeds the amount of O 2 needed for the Pd reoxidation into PdO. It can be assumed that this 02 excess (3.1 mmol.) is trapped by the support according to the reaction : ZrO2. x + x/2 02 --> ZrO 2. Then, the solid is cooled under the He-1% 02 mixture and an At-1% H 2 mixture is introduced at 298 K. The H 2 consumption which reaches 12.3 retool. (for 14.8 retool, total Pd) can arise from two main phenomena : the reduction of PdO and the adsorption of H 2. PdO + H 2 --> Pd ~ + H20 pd s + 1/~H2 ._> Pdsi_I (s: superficial) The H 2 absorption can be neglected since the H 2 pressure is low (17). During the heating from 298 K to 773 K (10 K/min) the H 2 consumption reaches 4.75 mmol. Thus, the total H 2 comsumption reaches 17 retool, for 14.8 retool, total Pd and it can be concluded that this excess o f H 2 corresponds to a partial ZrO 2 reduction: ZrO 2 + X H 2 --> ZrO2.x + H20. After a purge under Ar at 773 K a second temperature-programmed oxidation leads also to a consumption of 0 2 in excess with respect to the Pd amount. In the absence of palladium, the ZrO 2 support (after treatment under H 2 at 773 K in order to remove possible impurities) does not absorb 02 or H 2 during successive temperatureprogrammed oxidations and reductions until 773 K. Moreover, with the Pd/Al203 solid the 02 or H 2 consumption never exceed the amounts corresponding to Pd. In the presence of Pd, the ZrO 2 surface presents therefore some redox properties and, assuming an amount of 10 oxygen atoms per (angstrom) E, 2 or 3% of the superficial oxygen atoms would be reduced, leading to the formation of oxygen vacancies on ZrO 2. 4. DISCUSSION AND CONCLUSION Supports effects do not drastically modify Pd, this is shown by XPS and by IR spectroscopy of adsorbed CO. Nevertheless, the catalytic performances of the materials have been significantly improved with the supports containing zirconia. For instance, the activity of Pd/AI203-BaO-ZrO 2 nearly reaches the activity of the Pt-Rh/AI20 3 reference.

The role of the metal in the NO reduction by CO is clearly shown by CO, NO and CONO (1:1 mixture) adsorptions at room temperature and at reaction temperatures. i) At 298 K, the adsorption of NO and CO takes place on the same Pd centers and a competitive adsorption is observed. ii) Upon heating, NO is dissociatively adsorbed on Pd ~ according to the reaction : NOads --> Nads + Oads and Pd ~ is oxidized into Pd n+ as shown by the IR bands at 1815-1810 crn-l assigned to NO on Pd n+ and observed after heating under NO at 473 K and evacuation at 298 K. iii) The number of such Pd n+ sites is drastically decreased when the samples are put into contact, at 473 K, with CO or with a CO-NO (l:l) mixture. Thus CO 2 is detected in the gaseous phase. The Oads species created by the NO dissociation react with adsorbed CO according to the reaction COads + Oads --> CO 2 . This reaction cleanses the metal surface of excess adsorbed oxygen, permitting a new dissociative NO adsorption. These two reaction steps are generally invoked in the CO+NO mechanism [ 18-21 ] and, as can be concluded from the above results, occur mainly on Pd.

353

For the role of the support, two effects have to be considered i) An intrinsic effect: for instance the acidity or the stabilizing effect toward thermal sintering of Pd. The influence of the support acidity has been clearly illustrated by the hydrocarbon poisoning which is strongly decreased in the presence of less acidic supports [ 1]. ii) The effect of an oxygen-ions mobility which is connected to the redox properties of the support but needs the presence of the metal. In the present work, thermoprogrammed oxidations and reductions have shown the presence of "mobile and active oxygen atoms" on ZrO 2. With such solids, CO adsorbed on the metal can migrate to and react with these reactive oxygens giving CO 2 and an oxygen vacancy (in close proximity to the metal) according to the scheme : COads(Pd) +O- Zr-O --> CO 2 + [3- Zr-O. These oxygen vacancies can be directly involved in the water gas shitt reaction (CO + 1-120 --> CO 2 + 1-12)or in the NO dissociation or in a trapping of 0 2. In the first case the oxygen vacancy would be filled by 1-120, leading to the formation o f H 2 [22-25] according to the reaction: 1-120 + O- Zr- [3 --> H 2 +O- Zr-O In fact the water gas shitt reaction is much more promoted by the ZrO 2 support than by the AI203 one. While the CO conversion reaches 100 % at 623 K on Pd/ZrO 2 , it does not exceed 50 % at 773 K on Pd/AI20 3. Such a H 2 formation via the water gas shift reaction (or the steam reforming reaction which is also greatly enhanced on Pd/ZrO 2 ) implies the participation of both the metal and the support. H 2 is generally considered as a better reducing agent for NO than CO [18] and therefore, this H 2 formation via the water gas shitt reaction could explain the enhancement of activity observed in the presence of 1-120 for ZrO 2containing solids which possess "mobile oxygen atoms". Therefore, a bifunctional mechanistic scheme, including the participation of both the metal (via the adsorption of CO) and the support (via the formation of "oxygen vacancies" which are active sites for the H20 dissociative adsorption) seems quite relevant to explain the specific behavior, for the NO reduction in the presence of water, of samples containing ZrO2. Such active sites would be located at the metal-support interface and are linked to the redox properties of the support. In conclusion, a "specific role" of these dual "Metal-Support" sites at the metal-support interface has to be considered in addition of the "intrinsic roles" of the metal and the support and a bifunctional mechanism can be reasonably proposed. ACKNOWLEDGEMENTS This work has been carried out with the financial support of the "Groupement de Recherches catalyseurs d'epuration des gaz d'echappement automobile" funded by the "Centre National de la Recherche Scientifique, the "Institut Fran~ais du Petrole" and the PIRSEM (Programme Interdisciplinaire de Recherches Scientifiques pour rEnergie et les Matieres Premieres).

354 REFERENCES

.

5. .

7. 8. .

10. 11. 12. 13.

14. 15. 16. 17. 18. 19. 20. 21. 22. 23. 24. 25.

A. Lemake, J. Massardier, H. Praliaud, G. Mabilon and M. Prigent, Catalysis and Automotive Pollution Control Ill, Elsevier, Amsterdam, (1995), in press. J.L. Duplan and H. Praliaud, Appl. Catal., 67 (1991) 325 A. A. Davydov, Infrared Spectroscopy of Adsorbed Species on the Surface of Transition Metal Oxides, (J. Wiley ed.), Rochester, (1984) A. M. Bradshaw and F. Hoffmann, Surf. Sci., 72 (1978) 513 D. Tessier, A. Rakai and F. Bozon-Verduraz, J. Chem. Soc. Farad. Trans., 88 (1992) 741 P. Gelin, A. R. Siedle and J.T. Yates, J. Phys. Chem., 88 (1984) 2978 A. Palazov, G. Kadinov, Ch; Bonev and D. Shopov, Surf. Sci., 188 (1987) 505 J. Z. Shyu, K. Otto, W.L.H. Watkins, G. W. Graham, R.K. Blitz and H.S. Gandhi, J. Catal., 124 (1988) 2. V. Pitchon, M. Guenin and H. Praliaud, Appl. Catal., 63 (1990) 333 Tran Thanh Phuong, J. Massardier and P. Gallezot, J. Catal. 102 (1986) 456 G.M. Alikina, A.A Davydov, I.S. Sazonova and V.V. Popovskii, Kinetik. i Katal. 28 (1987) 418 J. W. Ward, J. Catal., 10 (1968) 34 and 11 (1968) 271 A. Laachir, V. Perrichon, A. Badri, J. Lamotte, E. Catherine, J.C. Lavalley, J.E1Fallah, L. Hilake, F. Le Normand, E. Quemere, G.N. Sauvion and O. Touret, J. Chem. Soc. Farad. Trans.,87 (1991) 1601 Can Li, F. Domen, K.J. Maruya and T. Onishi, J. Catal., 141 (1993) 540 C. Binet, A. Badri, M. Boutonnet- Kizling and J.C. Lavalley, J. Chem. Soc. Farad. Trans., 90 (1994) 1023 E. Guglielminotti and F. Boccuzzi, J. Catal., 141 (1993) 486 E. Wicke and H. Brodowsky, Topics in Applied Physics, 29 (1978) 73 W. F. Egelhoff, The Chemical Physics of solid surfaces and Heterogeneous Catalysis,(D. A. King and D. P. Woodruff, eds.) Vol. 4 (1984) 397 D' Arc3, Lorimer and A.T. Bell, J. Catal. 59 (1979) 223 W. C. Hecker and A.T. Bell, J. Catal., 84 (1983) 200 S. H. Oh, G. B. Fisher, J. E. Carpenter and D. W. Goodman, J. Catal. 100 (1986) 360 B. Harrison, A. F. Diwell and C. Hallet, Platinum Metals Rev., 32 (1988) 73 J. C. Schlatter and P. J. Mitchell, Ind. Eng. Chem. Prod. Res. Dev., 19 (1980) 288 J. Barbier, Jr. and D. Duprez, Applied. Catal. B., 4 (1994) 105 M. Weibel, F. Garin, P. Bernhardt, G. Make, and M. Prigent, Catalysis and Automotive Pollution Control II, Elsevier, Amsterdam, 71 (1991) 195

J.W. Hightower, W.N. Delgass, E. Iglesia and A.T. Bell (Eds.) 11th International Congress on Catalysis - 40th Anniversary

Studies in Surface Science and Catalysis, Vol. 101 9 1996 Elsevier Science B.V. All rights reserved.

355

Characterization of Pd-based Automotive Catalysts R. W. McCabe and R. K. Usmen Ford Research Laboratory, MD 3179, SRL, Dearborn, MI 48121-2053 1. ABSTRACT Characterization studies were undertaken to determine the cause of large differences in activity between various commercial automotive catalysts after aging for 75 or 120 h on an accelerated engine-dynamometer cycle. In all, a set of nine catalysts was examined, comprised of both fresh and aged Pal-only, Pd/Rh, and Pt/Rh catalysts. Catalyst activity, as measured by CO/NOx crossover efficiencies in dynamometer airfuel sweep tests, showed no correlation with either noble metal dispersion or noble metal surface area. The amount of stored oxygen required to obtain 100% CO conversion in the A/F-modulated dynamometer sweep experiments was estimated at 15 p-mol O/g-cat.. Bench reactor experiments involving both titration of pre-oxidized catalysts with CO and cyclic CO oxidation confirmed that the 15 p-mol O/g-cat. storage requirement represents a threshold level separating high- and low-activity catalysts. Formation of bulk PdO is the main oxygen storage mode in the 120 h dynamometeraged Pd-based catalysts. Pd loading is important: the higher the Pd loading, the greater the capacity for oxygen storage via PdO. Dispersion of Pd in the dynamometer-aged catalysts (2-6%) was too low to account for significant oxygen storage via chemisorbed oxygen. Furthermore, temperature-programmed reduction experiments and comparisons of oxygen uptakes on catalysts with and without rare earth oxides both indicated that the rare earth oxides play little role in oxygen storage after 120 h dynamometer aging. Not only does the formation of bulk PdO account for the quantities of oxygen stored in the aged catalysts, but the observed rates of oxygen uptake are consistent with bulk Pd oxidation kinetics reported by Remillard et al [J. Appl. Phys. 71 (9), 1992, pp. 4515-4522]. 2. INTRODUCTION Three-way automotive catalysts based on palladium, rather than the more expensive metals platinum and rhodium, have long been desired. However, Pd is more sensitive than Pt to poisoning by lead (Pb) compounds [1-4]. Consequently, widespread commercial use of Pd-based automotive three-way catalysts (TWC) was delayed in the U.S. until the early 1990s, by which time residual Pb concentrations in unleaded gasoline had decreased to negligible levels. The past five years have witnessed

356

considerable research and development of various types of Pd-based TWCs including Pd-only [5-11 ], Pd/Rh [12-15], and Pt/Pd/Rh '~rimetal" catalysts [16,17]. Aside from its historically low price compared to Pt and Rh, Pd has distinct catalytic properties which make it a desirable component of today's three-way catalysts. Chief among these is thermal durability, particularly Pd's ability to maintain activity under high-temperature lean (i.e. excess O2) conditions. Pd also has excellent light-off characteristics, especially when deployed at higher concentrations than traditional Pt/Rh catalysts. Highly loaded Pd-based catalysts are thus an obvious choice for socalled close-coupled or starter catalysts mounted close to the exhaust manifold [5,18,19]. Such catalysts reach operating temperatures much faster than underbody catalysts but also experience higher warmed-up operating temperatures and greater risk of thermal deactivation. The present study was initiated to understand the causes of large differences in performance of various catalyst formulations after accelerated thermal aging on an engine dynamometer. In particular, we wished to determine whether performance characteristics were related to noble metal dispersion (i.e. noble metal surface area), as previous studies have suggested that the thermal durability of alumina-supported Pd catalysts is due to high-temperature spreading or re-dispersion of Pd particles [20-

25].

Catalyst performance (as evaluated in dynamometer sweep evaluations) did not correlate with noble metal particle dispersion. Instead, bench reactor experiments involving titration of preadsorbed oxygen with CO showed that total Pd load in@ rather than Pd surface area is the key factor affecting performance. Pd serves as its own oxygen storage agent through formation of bulk PdO, and the amount of PdO formed depends primarily on the amount of Pd available, not the surface area of the Pd. 3. EXPERIMENTAL 3.1. Catalysts Table 1 lists characteristics of the catalysts. Those labeled "TWC" are commercial formulations from Ford's catalyst suppliers, each letter designating a different formulation. The commercial catalysts all contained various rare earth and alkaline earth oxide promoters and stabilizers in addition to the noble metals. Two simple Pd-on-alumina reference catalysts (A and B) were prepared in our laboratory by impregnating alumina-coated ceramic monoliths with aqueous solutions of Pd nitrate. The laboratory-prepared catalysts were dried and calcined in air at 550 ~ for 5 h prior to evaluation. Noble metal concentrations were determined by x-ray fluorescence. Some of the commercial formulations were aged on an engine dynamometer for 75 or 120 h according to a standard 4-mode aging procedure with an inlet exhaust gas temperature of 760 ~ (peak catalyst temperature ca. 900 ~ [6]. The dynamometer aging cycle simulates vehicle aging of catalysts - 75 h for 50,000 miles and 120 h for 100,000 miles.

357

Table 1 Catalyst Description Catalyst

Source

Aging

A) Pd/AI203 B) Pd/AI203 C1) Pd-only TWC C2) Pd-only TWC C3) Pd-only TWC D) Pd-only TWC E) Pd/Rh TWC

Lab Lab Supplier Supplier Supplier Supplier Supplier

Fresh Fresh Fresh 75 h 120 h Fresh 120h

F) Pd/Rh TWC

Supplier

120h

G) Pt/Rh TWC

Supplier

120h

Pd concentration (%) 0.30 0.79 0.66 0.66 0.66 0.60 0.30 (0.026 Rh) 0.33 (0.036 Rh) (0.21 Pt) (0.042 Rh)

Oxygen storage component (~) No No Yes Yes Yes No Yes Yes Yes

(1) rare earth oxide component

32. ~ meUxx~ Conversion efficiencies of the dynamometer-aged catalysts were measured in a standard A/F sweep test on an engine dynamometer [6]. The sweep experiments were carried out at 450 ~ and 85,000 h"1 space velocity (volumetric basis; standard conditions). The sweep ranged from 0.5 A/F lean of stoichiometry to 0.5 A/F rich of stoichiometry with imposed A/F perturbations of +_0.5A/F at 1 Hz. After sweep evaluation, small samples of catalyst were removed from the front region of the brick for chemisorption and flow reactor experiments. Chemisorption measurements employed the CO-methanation technique of Komai et al [26] as modified in our laboratory [27] for application to automotive catalysts. In particular, modifications were made to the pretreatment to avoid the formation of Pd hydrides [27]. We have found the method well-suited to automotive catalysts, both because of its high sensitivity and apparent freedom from complications due to adsorption of CO on sites other than noble metal sites. In previous studies involving a series of Pd/Rh and Pt/Rh TWCs aged on vehicles, we obtained good correlation between apparent dispersions (and noble metal surface areas) measured by the CO methanation technique and those determined from both x-ray diffraction and transmission electron microscopy [28,29]. As with other chemisorption methods, the CO-methanation technique does not distinguish between different noble metals. Thus, Rh was treated equivalently to Pt or Pd, and a standard adsorption stoichiometry of 1 CO molecule per surface metal atom was assumed for all three noble metals. In general, contributions from Rh are expected to be small due to the low concentrations employed (one-tenth the Pd loading in the case of the Pd/Rh catalysts). Thus, the Rh concentration was simply added to the Pd or Pt loading and treated as Pd or Pt.

358

T~ration of pre-dosed oxygen by CO was carded out at 500 ~ in a 1" o.d. quartz reactor tube housing a 3/4" diameter by 1/2" long catalyst button. The reactor contained two solenoid-controlled three-way valves th= were used to inject alternating pulses of secondary feed streams (each 0.1 I.Jmin) into a main feed of N2 carder gas (2.9 L/min). Two types of experiments were conducted. One involved pre-dosing the catalyst for 120 s with 0.87% 02 in the main feed. The O2 flow was then stopped, the system was purged for 120 s, and the adsorbed oxygen was titrated from the catalyst by injecting alternating 15 s pulses of CO and N2 into the N2 carder stream (the CO concentration was 0.3% after dilution with the carder N2 stream). The second type of experiment exposed the catalyst to alternating pulses of CO (0.3%) and 02 (0.185%) for durations of 15 s for 02 and variable times between 1 and 7 s for CO. Conversion of CO and formation of CO2 were monitored by non-dispersive infrared analyzers. 4. RESULTS

DyrBmonteter ev uaSons

Figure 1 shows sweep data for two of the dynamometer aged catalysts: the 120 h aged Pd-only ((33) catalyst (Fig. 1A), and the 120 h aged Pd/Rh (E) catalyst (Fig. 1B). Despite equivalent aging, the Pd-only catalyst gave much higher conversions, especially around the stoichiometric point. CO/NOx cross-over efficiencies of the other dynamometer-aged catalysts are reported in Table 2. The Pd-only catalyst stands out, showing crossover efficiencies in excess of 95% after both 75 and 120 h aging. In contrast, the formulations containing Pd/Rh or PURh have crossover efficiencies between 50 and 54%.

Figure 1. Engine dynamometer sweep plots of 120-hr aged (A) Pd-only (C3) and (B) Pd/Rh (E) catalysts at 450~ and modulations of +0.5 A/F at 1Hz.

359

4.2. Noble metal dispersions and surface areas Table 2 lists the apparent dispersions obtained from the CO methanation technique. No correlation is observed between dispersion and catalyst performance as measured by the CO/NOx crossover efficiencies. The C2 and C3 Pd-only TWCs, despite their extremely high CO/NOx crossover efficiencies, gave apparent dispersions of 3.5 and 3.0% after 75 and 120 h aging versus higher values of 5.9% for the Pd/Rh catalyst (E) and 4.3% for the Pt/Rh catalyst (G), both of which displayed low CO/NOx crossover efficiencies. Even between the two Pd/Rh catalysts, catalyst E has an apparent dispersion more than four times that of catalyst F, yet the two are nearly identical in their CO/NOx crossover efficiencies. Table 2 Catalyst Properties Catalyst

CO/NOx crossover efficiency

(%)

A) Pd/AI203 B) Pd/AI203 C1) Pd-only TWC C2) Pd-only TWC C3) Pd-only TWC D) Pd-only TWC E) Pd/Rh TWC F) Pd/Rh TWC G) Pt/Rh TWC

NA NA NA 96 98 NA 50 52 54

Dispersion (%)

10.1 9.5 10.8 3.5 3.0 6.6 5.9 1.6 4.3

NM Surface Area (m2/g-cat.)

0.17 0.42 0.40 0.13 0.11 0.22 0.11 0.03 0.04

Table 2 also lists the noble metal surface areas normalized to the total mass of the catalyst. The surface areas were calculated directly from the dispersion data taking into account the different mass of noble metal in each catalyst and assuming a constant site density of 1x1019/m2. As with dispersion, no clear correlation exists between mass-specific noble metal surface areas and CO/NOx cross-over efficiencies. 4.3. Oxygen tJtr"~n expedrnents The engine dynamometer sweep evaluations were carried out under modulated air-fuel conditions of + 0.5 A/F ratio at 1 Hz. To achieve high conversions under these conditions, the catalyst must store oxygen during lean excursions in order to convert CO and HC under rich excursions. Likewise, some of the oxygen related inhibition of NOx reduction on the lean side is mitigated by replenishment of oxygen to the storage agent. The lack of a correlation between the CO/NOx crossover frequencies of these catalysts and either noble metal dispersion or mass-specific surface area suggests, in turn, that oxygen storage in aged catalysts is not strongly dependent on either noble metal dispersion or noble metal surface area.

360

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Figure 2. Trtration of pre-adsorbed oxygen by pulses of CO (0.3% CO in N2) at 500 ~

Experiments involving titration of preadsorbed oxygen with pulses of CO in nitrogen were carded out to assess the oxygen storage capacity of both the fresh and dynamometer aged catalysts. Figure 2 shows the pulse profiles for the two catalysts with the highest (CI) and lowest (G) oxygen uptakes. The areas under the CO pulses were integrated to determine the amount of oxygen pre-adsorbed during the 120 s O2 pretreatment at 500 ~ W'rth the exception of the fresh CI Pd-only catalyst, all of the other catalysts produced CO breakthrough in excess of 90% after the third pulse. The cumulative O-atom uptakes, expressed as /~-mol O/g-cat., are summarized in Table 3. Both the first-pulse uptakes and the cumula'dve uptakes corresponding to the first three pulses are listed. Table 3 also contains the theoretical oxygen uptakes associated with both chemisorbed oxygen (i.e. assuming one chemisorbed O-atom per surface noble metal atom) and formation of bulk noble metal oxides (i.e. PdO or PtO).

Table 3 Oxy.qe.n........capac...ities(#-mol O/Q-~.)... Catalyst

- Theoretical O uptakesSurface O ~ Bulk O 2

- O titrated by CO 1st pulse 1st 3 pulses

A) Pd/AI20 3 B) Pd/AI20 s C1) Pd-only TWC C2) Pd-only TWC C3) Pd-only TWC D) Pd-only TWC E) Pd/Rh "IWC F) Pd/Rh TWC G) PURh TWC

1.0 1.7 6.8 2.7 1.8 3.4 1.9 0.4 0.6

11.5 24.8 35.5 27.2 21.6 20.9 9.2 12.7 5.2

28.2 74.2 62.0 62.0 62.0 56.4 30.6 34.4 14.0

1~ ...assumes i'" O-atom per surface""or""buJk noble metal atom. '..........

15.6 32.0 66.8 39.0 32.0 31.0 14.3 23.0 7.2

361

Two key observations can be drawn from the data in Table 3: 1) the oxygen consumed in either the first pulse or the first three pulses is much greater than that which can be attributed to oxygen adsorbed on the surface of the noble metal particles, and 2) with the exception of ~ fresh (CI) Pd-only IWC, none of the other catalysts have either first-pulse or cumulat~e CO uptakes that exceed the theoretical oxygen uptake associated with formation of bulk noble me'~ oxide. Even though all of the aged commercial TWCs contain oxygen storage agents, the quantities of oxygen taken up after 120 h dynamometer aging are not sufficient to require storage via those agents. 40

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10

4.4. Role of bulk PdO The importance of bulk PdO as an oxygen storage component is illustrated in Figure 3 which plots the amount of oxygen consumed in the first CO pulse versus the theoretical bulk oxygen capacity of each Pd-containing catalyst (expressed as/~-mol PdO/g-cat.). The solid curve is fit to the data for the three catalysts which do not contain rare earth oxygen storage agents plus an additional point at the origin reflecting the experimental observation that negligible oxygen is consumed on a blank alumina catalyst. Given the absence of rare earth oxides, and recognizing that the amount of stored oxygen is far greater than that available from chemisorption (as shown by the data points denoted by diamonds in Fig. 3), the only other source of oxygen is from reduction of bulk PdOo Thus the solid curve can be taken as representative of bulk oxidation of Pd during the 120 s exposure to O2 at 500 ~ Note that all of the 12=3h dynamometer aged catalysts have oxygen uptakes on or below the curve defined by the catalysts which do not contain rare earth oxygen storage agents. This suggests that rare earth oxides do not contribute significantly to oxygen uptakes after 120 h dynamometer aging. The only catalysts showing greater oxygen

362

storage are the fresh and 75 h aged Pd-only TWCs (C1 and C2), and it is likely that rare earth oxides do contribute to oxygen uptakes in those catalysts. Interestingly, the C1 and C2 catalysts are the only pair which show a correlation between oxygen uptake and noble metal dispersion (i.e. the oxygen titrated by the first CO pulse drops from 35.5 to 27.2/~-mol O/g-cat. as the dispersion drops from 10.8% (C1) to 3.5%

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Figure 4. 1-12TPR traces of fresh (CI) and b) 120 h aged (C3) Pd-only TWCs (30 ~ heating rate in 9% i"La/Ar). Inset: TPR traces of a 0.32% Pd/15% CeO2/AI=O3 catalyst fresh and after aging for 1 h at 900 ~ in a laboratory flow reactor. 4.5. TPR The absence of a strong contribution from rare earl~ oxygen storage agents is confirmed by hydrogen temperature-programmed reduction (TPR) experiments. Figure 4 shows hydrogen TPR traces for the fresh and 120 h aged Pd-only catalysts (C1 and C3). The fresh catalyst contains a broad peak centered around 100 ~ which has an area of 127/~-mol 1-12/g-cat. (corresponding to twice the theoretical I-Iz uptake of 62 /~-mol I-Iz/g-cat. required to reduce PdO to Pd metal). The "extra" area is attributed to the reduction of surface oxygen associated with ceria in intimate contact with Pd as shown by the close similarity to the TPR trace of a fresh 0.32=,(, Pd/15% CeO2/AI203

363

catalyst prepared in our laboratory (Fig. 4 inset). After 120 h dynamometer aging, the low-temperature feature of the C3 catalyst peaks near 40 ~ and decreases in area to 23/~-mol H2/g-cat. (Fig. 4b). Again, the changes upon aging are similar to those of the reference 0.32% Pd/15% CeO2/AI20~ catalyst after aging for 1 h at 900 ~ in a laboratory reactor (Fig. 4 inset). Changes in TPR spectra of the type shown by both the Pal-only catalyst and the model Pd/CeO2/AI20 s catalyst are characteristic of large reductions in contact area between Pd and rare earth oxides effected by thermal aging [30]. Pd in intimate contact with ceria [30,31] or lanthana [32] gives a TPR feature between 100 and 200 ~ whereas Pd on alumina has its chief TPR feature at temperatures between 0 and 60 ~ [23,32,33]. In fact, partial reduction of PdO may be occurring in our experiments prior to the start of the temperature ramp from 0 ~ The TPR data thus support the oxygen uptake experiments in suggesting that rare earth oxides contribute little to the uptakes observed for the 120 h dynamometer-aged catalysts. 5. DISCUSSION Automotive catalysts have traditionally been designed with the objective of maximizing dispersion of the noble metals. Some exceptions exist, such as the well-documented observation that oxidation rates of saturated hydrocarbons increase with increasing noble metal particle size [34,35]. However, one of the main objectives in designing automotive catalysts is to stabilize the dispersion of the noble metals against thermal sintering. Pd undergoes complex changes in oxidation state and structure in response to variations in temperature and gas environment [20-25,33,36-40]. Spreading and/or redispersion of Pd (as a cationic form of Pd) has been reported in oxidizing environments at temperatures between 700 and 900 oC [20-25]. At first glance, such a mechanism would appear to offer a plausible explanation for the thermal durability of Pd-basecl catalysts, particularly in aging cycles of the type reported here involving lean (i.e., oxidizing) modes. The data of this study, however, indicate that dispersions of Pd-based catalysts after thermal aging are quite low, and no greater than comparably aged PURh catalysts. Catalytic activ'rty, as measured in modulated NF sweep experiments, neither correlates with noble metal dispersion nor with noble metal surface area. As shown in Fig. 3, oxygen uptakes (as reflected in amounts of CO oxidized) show an increasing trend with the bulk oxygen capacities of the catalysts (expressed as the theoretical amount of bulk PdO which can be formed). Comparing the oxygen uptake data of Fig. 3 to the CO/NOx crossover efficiencies of Table 2, a threshold level of oxygen uptake is suggested, between 13 and 20/~-mol O/g-cat., above which the catalyst can store enough oxygen to ensure high efficiency during the A/F perturbations encountered in the dynamometer sweep test. This range is consistent with a stored oxygen demand of 15.3/~-mol O/g-cat. which we have estimated as required to ensure stoichiometric oxidation of all reducing species during the 0.5 s, -0.5 delta NF half cycle of the dynamometer sweep (centered at the stoichiometric A/F ratio) [41]. To a first approximation, the high activity of catalyst C3 results simply from its having an oxygen uptake capacity above the threshold requirement, whereas the other

364 dynamometer-aged catalysts have oxygen uptakes below the threshold requirement. The oxygen uptakes shown in Fig. 3 reflect much longer oxygen exposures (120 s) and larger CO titers (48 #-mol CO/g-cat.) than charactedstic of the dynamometer sweep cycle. Therefore, additional bench reactor CO oxidation experiments were carried out at shorter time intervals cycling between 15 s pulses of O2 (0.185 mol%) and 2 to 7 s pulses of CO (0.3 mol%). Results are summarized in Table 4, with the catalysts listed in order of highest to lowest CO conversions. The most active catalyst (C1 Pd-only TWC) stored sufficient oxygen to convert a 7 s CO pulse with 98% efficiency. At the other extreme, the PURh catalyst (G) gave only 48% conversion of a 7 s CO pulse and reached only 9'2% conversion with a 2 s pulse. The most pertinent condition for comparing the bench reactor and dynamometer results is at 5 s, since the bench reactor CO dose at 5 s (16/~-mol CO/g-cat.) is close to the dynamometer stored oxygen demand (15.3/~-mol O/g-cat.) Significantly, we find close quantitative agreement between the CO conversions at 5 s CO pulse length and the CO/NOx crossover efficiencies reported in Table 2. In both cases, catalysts E, F, and G give conversions in the 50-55% range whereas the C2 and C3 catalysts give conversions in excess of 85%. The good quantitative agreement between the engine dynamometer data and the laboratory pulsed CO oxidation experiments supports our interpretation that high CO/NOx crossover efficiencies in dynamometer sweep evaluations reflect oxygen storage above a threshold level. Table 4 Per cent conversion of CO pulses of various len.qths Catalyst

C1) Pd-only TWC C2) Pd-only TWC D) Pd-only TWC C3) Pd-only TWC B) Pd-AI20 3 A) Pd-AI20 s E) Pd/Rh TWC F) Pd/Rh TWC G) Pt/Rh TWC

7 98 NM NM NM 62 NM 42 NM 48

length of CO pulse (s) ............. 5 4 3 >99 95 89 86 78 67 51 51 NM

>99 >99 95 96 89 NM 67 NM 59

>99 >99 >99 > 99 >99 92 82 NM 71

>99 >99 >99 >99 >99 >99 >99 93 92

Given the connection between oxygen uptakes and catalyst sweep performance, the lack of a correlation between noble metal dispersion (and surface area) and catalyst performance implies that oxygen uptake does not depend on noble metal dispersion or surface area. The lack of a correlation is not surprising given that, 1) the quantities of oxygen involved, 2) the similarities in oxygen uptakes between catalysts with and without oxygen storage agents, and 3) the absence of TPR features characteristic of noble metal-rare earth oxide interactions all point to bulk oxidation of PdO as the

365

dominant oxygen storage mechanism in the 120 h dynamometer-aged Pd-based catalysts. For bulk oxidation of Pd to be the dominant oxygen storage mechanism, Pd oxidation kinetics must be rapid enough to account for the quantities of oxygen stored. Data reported by Remillard et al [38] on the air oxidation of thin Pd films indicate that the rates are indeed fast enough to account for the oxygen uptakes observed for the aged catalysts in these experiments. The 500 ~ 120 s pre-oxidation employed in the CO titration experiments would produce an oxide film thickness of 89 Angstroms using the growth expression reported in Ref. 38 (assuming equivalent kinetics in air and in the 0.87% O2/N2 mixture of our experiments). Taking the C3 catalyst, for example, with a dispersion of 3% (corresponding to a mean particle diameter of 373 Angstroms [42]), the formation of an outer PdO band of 89 Angstrom thickness would result in shrinkage of the metallic core to 274 Angstroms and growth of the overall diameter to 452 Angstroms. The increase in particle diameter owes to both the lower density of PdO compared to Pd and the presence of about 21% void volume in the PdO band [38]. Approximately 61% of the Pd in the particle is oxidized. Experimentally, 120 s oxidation of the C3 catalyst yielded a 3-pulse CO consumption corresponding to oxidation of about 52% of the Pd in the catalyst. Thus, the bulk oxidation kinetics are fast enough to account for the experimental observations. Even at the shorter 15 s oxygen exposure of the cyclic CO oxidation experiments, the kinetics of Remillard et al predict a 35 Angstrom thick oxide film. This corresponds to 25% oxidation of the Pd (or 15.2/~-mol O/g-cat.), just slightly less than the uptake of 16/~-mol O/g-cat. required for complete conversion of the 5 s CO pulse in the cyclic CO oxidation experiments (Table 4). Note that the oxidation kinetics of Remillard et al were obtained on 1-/~m-thick sputtered films of Pd on quartz, indicating that Pd oxidation kinetics are rapid enough, even for large-grain Pd particles, to account for oxygen uptakes required in both the dynamometer and bench reactor activity evaluations. The conclusions reached in the present study, namely that the catalytic activity of aged Pd-based catalysts depends primarily on Pd loading, begs the question '~hat is the function of the various additives such as ceria?". One obvious answer is that much effort has gone into stabilizing support components against loss of surface area and associated occlusion of noble metal particles. Both rare earth and alkaline earth oxides are important in this regard. In addition, we have addressed catalytic activity under very limited conditions -- modulated A/F sweep experiments carried out with lowsulfur fuel on an engine dynamometer. Factors such as noble metal dispersion and effects of promoters/stabilizers may be more important under other conditions, e.g. during light-off, at different temperatures and space velocities, with high-sulfur fuel, or in different A/F regimes during aging and evaluation. Also, we have focussed exclusively on CO oxidation, whereas different effects of dispersion and promoter/stabilizers may obtain for HC oxidation and for NOx reduction. Our study does suggest, however, that at least part of the Pd should be deployed in the form of large particles to obtain good CO and NOx conversions under modulated A/F conditions. The par~cles will consist of a metallic core with a roughened outer band capable of facile interconversion between Pd oxide and Pd metal.

366

6. SUMMARY Chemisorption measurements, combined with oxygen uptake, TPR, and pulsed CO02 experiments were employed to determine the source of large differences in dynamometer sweep performance of a series of Pt/Rh, Pd/Rh and Pd-only TWCs after dynamometer aging. The following observations have been made: 1) Apparent noble metal dispersions of 75 and 120 h dynamometer-aged TWCs range from about 2 to 6%. 2) CO/NOx cross-over efficiencies of aged catalysts in dynamometer sweep experiments do not correlate with either noble metal dispersion or noble metal surface area. 3) Oxygen uptakes in both dynamometer and bench reactor experiments at 500 ~ are much too great to attribute to oxygen chemisorption. 4) After 120 s oxygen exposure (500 *C), all of the dynamometer-aged Palbased catalysts gave oxygen uptakes that could be accounted for by the formation of bulk PdO. 5) Dynamometer-aged (120 h) catalysts showed no evidence for oxygen storage via rare earth oxides. 6) Formation of bulk PdO is the primary oxygen storage mechanism in the dynamometer-aged Pd-based catalysts. 7) A threshold level of oxygen storage (via bulk PdO) is required to reach high CO/NOx conversion levels in dynamometer sweep tests; Pd loading, rather than dispersion or surface area, is the most important factor affecting oxygen uptakes. 8) Rates of oxygen uptake in the dynamometer-aged catalysts are consistent with published oxidation kinetics of 1-/~m-thick Pd films. 7. ACKNOWI.EDGMENTS We thank K. S. Patel and D. M. DiCicco for providing the dynamometer-aged catalysts and sweep evaluation data. E. Gulari and C. Sze (U. of Michigan) assisted with the design of the pulsed reactor system. REFERENCES 1. M. Shelef, K. Otto, and N.C. Otto, Adv. in. Catal. 27 (1978) 311-365. 2. H.S. Gandhi, W.B. Williamson, E.M. Logothetis, J. Tabcock, C. Peters, M.D. Hurley, and M. Shelef, Surf. & Interface Anal. 6(4) (1984) 149. 3. W.B. Williamson, D. Lewis, J. Perry, and H.S. Gandhi, Ind. Eng. Chem. Prod. Res. Dev., 23 (1984) 531. 4. R.L Klimisch, J.C. Summers, and J.C. Schlatter, Amer. Chem. Soc. Adv. Chem. Ser. 143 (1975) 103. 5. J.C. Summers, J.F. Skowron, and M.J. Miller, Soc. of Automotive Eng., Paper

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930386 (1993). 6. J.S. Hepburn, K.S. Patel, M.G. Meneghel, H.S. Gandhi, and Engelhard and Johnson Matthey Three Way Catalyst Development Teams, Soc. of Automotive Eng., Paper 941058 (1994). 7. M. Harkonen, M. Kivioja, P. Lappi, P. Mannila, T. Maunula, and T. Slotte, Soc. of Automotive Eng., Paper 940935 (1994). 8. J.C. Summers, J.J. White, and W.B. Williamson, Soc. of Automotive Eng., Paper 890794 (1989). 9. H. Muraki, Soc. of Automotive Eng., Paper 910842 (1991). 10. H. Tanaka, H. Fujikawa, and I. Takahashi, Soc. of Automotive Eng., Paper 930251 (1993). 11. T. Yamada, K. Kayano, and M. Funabiki, Soc. of Automotive Eng., Paper 930253 (1993). 12. Y.-K. Lui and J.C. Dettling, Soc. of Automotive Eng., Paper 930249 (1993). 13. J.K. Hochmuth and J.J. Mooney, Soc. of Automotive Eng., Paper 930219 (1993). 14. J.C. Summers, W.B. Williamson, and J.A. Scaparo, Soc. of Automotive Eng., Paper 900495 (1990). 15. H. Muraki, H. Sobukawa, M. Kimura, and A. Isogai, Soc. of Automotive Eng., Paper 900610 (1990). 16. B.H. Engler, E.S. Lox, I~ Ostgathe, T. Ohata, I~ Tsuchitani, S. Ichihara, H. Onoda, G.T. Garr, and D. Psaras, Soc. of Automotive Eng., Paper 940928 (1994). 17. A. Punke, U. Dahle, S.J. Tauster, and H.N. Rabinowitz, Soc. of Automotive Eng., Paper 950255 (1995). 18. D. Ball, Soc. of Automotive Eng., Paper 922338 (1992). 19. Z. Hu and R.M. Heck, Soc. of Automotive Eng., Paper 950254 (1995). 20. E. Ruckenstein and J.J. Chen, J. Catal. 70 (1981) 233. 21. J.J. Chen and E. Ruckenstein, J. Phys. Chem. 85 (1981) 1606. 22. E. Ruckenstein and J.J. Chen, J. Colloid Interface Sci. 86 (1982) 1. 23. H. Lieske and J. Volter, J. Phys. Chem. 89 (1985) 1841. 24. J.W.M. Jacobs and D. Schryvers, J. Catal. 103 (1987) 436. 25. J.G. McCarty and Y-F. Chang, Scripta Metal. et Mater. 31 (1994) 1115. 26. S. Komai, T. Hattori, and Y. Murakami, J. Catal. 120 (1989) 370. 27. R.K. Usmen, R.W. McCabe, and M. Shelef, "Proceedings of the Third Congress on Automotive Pollution Control," Elsevier, Brussels, Belgium, in press. 28. M.H. Yao, D.R. Liu, R.J. Baird, R.K. Usmen, and R.W. McCabe, ext. abstr., "Proc. 52rid Ann. Mtg. Microscopy Soc. of Amer.," C.W. Bailey and A~J. Garratt-Reed, eds., San Francisco Press, Inc., p.776, 1994. 29. M.H. Yao, D.R. Liu, R.J. Baird, R.K. Usmen, and R.W. McCabe, submitted to J. Catal., Nov., 1995. 30. R.K. Usmen, R.W. McCabe, G.W. Graham, W.H. Weber, C.R. Peters, and H.S. Gandhi, Soc. of Automotive Eng., Paper 922336 (1992). 31. B. Harrison, A.F. Diwell, and C. Hallett, Plat. Met. Rev. 32 (1988) 73. 32. S. Subramanian, R.J. Kudla, C.R. Peters, and M.S. Chattha, Catal. Letts. 16 (1992) 323. 33. T.R. Baldwin and R. Burch, Appl. Catal. 66 (1990) 359. 34. T. Tokoro, I~ Hori, T. Nagira, T. Uchyima, and Y. Yoneda, Nippon Kagaku Kaishi

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12 (1979) 1646. H. Yao, Y. Yao, and K. Otto, J. Catal. 56 (1979) 21. R.F. Hicks, Q. Haihua, M.L. Young, and R.G. Lee, J. Catal. 122 (1990) 295. P. Briot and M. Primet, Appl. Catal. 68 (1991) 301. J.T. RemiUard, W.H. Weber, J.R. McBride, and R.E. Soltis, J. Appl. Phys. 71 (1992) 4515. 39. D. Konig, W.H. Weber, B.D. Poindexter, J.R. McBride, G.W. Graham, and I~ Otto, Catal. Letts. 29 (1994) 329. 40. R. Burch and F.J. Urbano, Appl. Catal. A: Gen. 124 (1995) 121. 41. The stored oxygen demand was estimated from the difference between the actual conversions measured at the stoichiometric point with +0.5 A/F perturbations at 1 Hz and steady state conversions measured at the +0.5 and -0.5 A/F extremes. 42. F.H. Ribeiro, M. Chow, and R.A. Dalla Betta, J. Catal. 146 (1994) 537. 35. 36. 37. 38.

j.w. Hightower, W.N. Delgass, E. Iglesia and A.T. Bell (Eds.) 11th International Congress on Catalysis - 40th Anniversary

Studies in Surface Science and Catalysis, Vol. 101 9 1996 Elsevier Science B.V. All rights reserved.

369

P r o c e s s d e v e l o p m e n t for the selective h y d r o g e n o l y s i s o f CC12F2 ( C F C - 1 2 ) into CH2F2 ( H F C - 3 2 ) A. Wiersma l, E.J.A.X. van de Sandt 2, M. Makkee 1, H. van Bekkum 2 and J.A. Moulijn I Department of Chemical Process Technology, Section Industrial Catalysis 2 Department of Organic Chemistry and Catalysis Delft University of Technology, Julianalaan 136, 2628 BL, Delft, The Netherlands

A palladium on activated carbon catalyst is a suitable catalyst for the selective hydrogenolysis of CC12F2 into CH2F2. Its stability is good: no deactivation occurred after 800 hours of operation. Even after addition of possible recycle components methane and CHC1F2, only minor deactivation was observed during 1600 hours of total operation. The catalyst performance and stability strongly depend on the hydrogen to CC12F2 feed ratio. A ratio of 6 leads to higher conversion, higher selectivity to CH2F2 and higher stability. At low H2 to CC12F2 ratios the mechanism of deactivation is presumably coke deposition. However, the ratio cannot be raised too high because then sintering of the palladium particles causes deactivation. Addition of methane to the feed leads to additional deactivation, and, therefore, methane is only allowed in a recycle stream when high hydrogen to CC12F2 feed ratios are applied. Addition of CHC1F2 to the feed does not lead to a higher CH2F2 yield. The reaction pathway to CHCIF2 and CH2F2 is different from the route to methane. The selectivity to methane depends on the adsorption mode of CC12F2 and is independent of both the CC12F2 and the hydrogen concentration. The ratio between CHC1F2 and CH2F2 is mainly determined by the CC12F2 concentration. A CFC-destruction process, which can destruct about 90% of the CFCs in use, based on this catalyst is both technically and economically feasible. The main features of this process are a liquid cooled, multi-tube fixed bed reactor, excess hydrogen, 100% CC12F2 conversion and a hydrogen recycle, in which methane is allowed.

1. INTRODUCTION The term CFCs is a general abbreviation for ChloroFluoroCarbons. They have been extensively used since their discovery in the thirties, mainly as refrigerant, foam blowing agent, or solvent because of their unique properties (non toxic, non flammable, cheap). However, after the first warning of Rowland and Molina [ 1] in 1974 that CFCs could destroy the protective ozone layer, the world has moved rapidly towards a phase-out of CFCs. Because the destruction of stratospheric ozone would lead to an increase of harmful UV-B radiation reaching the earth's surface, the production and use of CFCs is prohibited (since January 1, 1995 in the European Union and since January 1, 1996 worldwide). However, the depletion of stratospheric ozone will continue in spite of this prohibition. This is caused by slow diffusion of CFCs present in the troposphere to the stratosphere and

370 the eventual emission of the CFCs still in use [2]. It is, therefore, of utmost importance to prevent the banked CFCs (45% CC12F2 as refrigerant, 45% CCI3F as foam blowing agent) from being emitted into the atmosphere. Commercially available CFC destruction processes are incineration and pyrolysis, but these processes only have a limited capacity. The long term replacements for CFCs in their application as refrigerant are most probably HFCs. Especially HFC-134a (C2H2F4) is often mentioned, but the total number of possible replacements still increase. CH2F2 (HFC-32) or a mixture thereof are good replacements in heavy duty cooling applications, ' because CH2F2 has excellent cooling properties and in addition a lower global warming potential than HFC-134a. Combining the fact that there is not enough destruction capacity for the banked CFCs and that the market for CFCreplacements such as CH2F2 is growing, a challenging task is to convert the waste CFCs into valuable HFCs. At Delft University of Technology a catalytic process is under development in which the harmful CC12F2 is converted into the valuable, ozone friendly CH2F2. With this process both the waste materials CC13F, which can be converted into CC12F2 by use of HF, and CC12F2 can be converted. Because of the limited time available to develop this process, the catalyst development is strongly related to a future process design. Palladium on a purified activated carbon support has been selected as a very suitable catalyst for the reaction. We have reported that the performance of this catalyst looks very promising and that a CFC hydrogenolysis plant based on this catalyst is both technically and economically feasible [3-5]. This paper deals with the stability of the selected catalyst, the long term influence of the hydrogen to CC12F2 feed ratio on the catalyst performance and the influence of the possible recycle components methane and CHC1F2 on the performance of the catalyst.

2. E X P E R I M E N T A L 2.1. Preparation of catalyst Catalysts were prepared by incipient wetness impregnation of palladium chloride dissolved in HCI (CI/pd2+=10) on an activated carbon support. 50 Gram of the gas activated, peat based carbon extrudates (Norit RB1, d=l mm, 1=3-5 mm, BET=1060 m2/g) were consecutively washed with 0.5 M caustic hydroxide, water, 0.5 M hydrochloric acid, and water in a flow set-up prior to the impregnation. The catalysts were dried overnight at 373 K and subsequently activated at 623 K for 1 hour in nitrogen. Reduction of catalyst was done insitu with hydrogen at room temperature and subsequently the temperature was raised to the desired level under flowing hydrogen. The CC12F2 was introduced and after 10 hours of catalyst activation, the temperature was set at the experimental temperature.

2.2. Testing of catalyst The activity tests of the catalyst were carried out in a microflow reactor set-up in which all the high temperature parts are constructed of hastelloy-C and monel. The reactor effluent was analyzed by an on-line gas chromatograph with an Ultimetal Q column (75 m x 0.53 mm), a flame ionization detector, and a thermal conductivity detector. The composition of the feed to the reactor can be varied, besides the temperature, pressure, and space velocity. The influence of the recycle components CHC1F2 and methane was tested by adding these components to the feed. In total five stability experiments of over 1600 hours were performed. In each

371

experiment temperature, pressure, and WHSV were the same. The H2 to CC12F2 feed ratio ranged from 1.5 (hydrogen limiting) to 10 (excess hydrogen). The experimental set-up was the same for each stability test: After 830 hours on stream methane was added to the feed as to give a CC12F2 to methane feed ratio of about 1. The methane was removed from the feed after 1000 hours. After 1150 hours CHC1F2 was added to give a ratio of CHCIF2 to CC12F2 of 4 and was removed after 1300 hours. After 1650 hours on stream the experiments were ended. In the experiment with the hydrogen to CC12F2 feed ratio of 10, this ratio was doubled after 680 hours of operation.

3. R E S U L T S The main products of the CC12F2 hydrogenolysis are CH2F2, CHC1F2, and methane. The initial performance of the catalyst, after activation, is strongly dependent on the HE to CCI2F2 feed ratio. The influence of the H2 to CC12F2 feed ratio on the initial catalyst performance is depicted in Figure 1. Both the conversion and the selectivity to the desired product CHEF2 increase with increasing H2 to CC12F2 feed ratio, while the selectivity to CHC1F2 increases. The selectivity to methane is not influenced by the HE to CC12F2 feed ratio. These differences become even more pronounced as a function of time on stream. This can be seen in Figure 2 in which the influence of the HE to CC12F2 feed ratio on the catalyst performance after 700 hours of operation is depicted. 100

25

conv. CC12F2

I00

=

O

~176176

20

>~ 80 7, 60

>~ 80 ~ .~_ 15 >~ ~ 60

\ sel. CHCIF2 t:t,

40

~ 2o

35 30-25-" 20 "7-~,

. . . . . . . .

~176176

9 sel. CHCIF2 ~ -

-I% , +, +- - - "~ -~--~ = .------.-+ sel. metnanqmp, cl

I....

40

-ql sel. CH2F2 9

5

O

~ 20

-15

lO

- - - +

~ ~

O

o

a.....

0

l

i

o

1

2 4 6 8 10 Feed ratio hydrogen/CCl~F2

Figure 1. Initial catalyst performance.

12

o

I

0

I

I

i

I

2 4 6 8 10 Feed ratio hydrogen/CCl2F2

12

Figure 2. Catalyst performance after 700 hours.

Thus the catalyst performances change as a function of time on stream. The extend of these changes depends on the H2 to CC12F2 feed ratio. In Figures 3a and 3b the catalyst performance during the first 800 hours of operation is depicted for two different H2 to CC12F2 feed ratios. The results show that not only the catalyst performance is better at higher H2 to CC12F2 feed ratios, but also the stability of the catalyst is better. At a ratio of 1.5 both the catalyst activity and selectivity keep declining during the 800 hour experiment. At the ratio of 6, however, no selectivity change is observed after 300 hours of operation and after 450 hours also the activity of the catalyst remains unchanged.

372

100 e

~

~

~' S ~'~ 80 "; ~ 60 ~ "~ 40 ~

g

20

~

1/40

100

..... CH2F2 -ql sel.

35 30 I 25 20 t 15

~ 80 ~ ~ ~ 60

cony. CC12F2

5

0

0 0

200 400 600 Time on stream [hi

800

~, >., "F~

_ "7,

seI. CHCIF211~

>o 20

sel. methaneu~-

0 0

20 15

~ 40

~

25

.~_ 10 ~

' '~. . . . . ' 200 400 600 Time on stream [hi

9 800

(a) H2 / CC12F2 = 1.5. (b) H2 / CC12F2 = 6. Figure 3a, b. Performance of catalysts as a function of time on stream. The effect of further increase of the high H2 to CC12F2 ratio on the catalyst performance is depicted in Figure 4. In this experiment the ratio was doubled to 20 after 680 hours. The steady-state catalyst performance is not only more enhanced at the ratio of 10, but it is also reached far more quickly in comparison to the results in Figure 3b. However, an increase of the ratio to 20 does not lead to a further increase in catalyst performance and at this ratio the catalyst also tends to deactivate. 100

9

~" 80

conv. CCI2F2

ratio doubled

o 60 o 40 e,,,

o 20

25 20 15-, ._>. 10oo

(D

t__.

2F2

(D r~

I 0

100

200

300

400 500 600 Time on stream [h]

700

800

0 900

Figure 4. Performance of catalyst as a function of time on stream (after 680 hours H2 to CC12F2 feed ratio doubled from lO to 20). Table 1 summarizes the performances of the catalysts for different H2 to CC12F2 feed ratios after about 700 hours of operation. A low HE to CCl2F2 feed ratio is poor choice for three reasons: low conversion, low selectivity, and low stability, while a high ratio leads to a stable catalyst, with a superior performance. A stable catalyst performance has only been reached with both the HE to CC12F2 feed ratios of 6 and 10. When the ratio of l0 is doubled, the catalyst also tends to deactivate.

373

Table 1 Catal~,st performance after 700 hours ratio H2/CC12F2 1.5 2.2 3 6 10 20 ,,

,,, ,,,

conv. CC12F2 50 60 67 79 88 84

sel. CH2F2 53 67 75 85 88 88

sel. CHC1F2 37 24 16 7 4 3

sel. CH4 7 7 7 6 6 7

%

deact. 17 14 11 5 3 -

The effect of addition of methane to the feed is depicted in Figures 5a and b. During methane addition the amount of methane produced cannot be exactly determined and, therefore, the selectivity to methane is assumed to be the same as prior to the addition of methane. The influence of methane is larger at lower H2 to CC12F2 feed ratios. The addition causes dilution and will result in both a lower CC12F2 and hydrogen concentration. Thus, methane addition leads to a lower CC12F2 conversion. An unexpected effect is the higher selectivity to CH2F2, which coincides with a lower selectivity to CHC1F2. This selectivity change was observed in all experiments, but with higher H2 to CC12F2 feed ratio, the change in selectivity was less pronounced. After removal of methane from the feed, the conversion increases again, but does not return to its original value. As can be seen from Figure 5a, during the addition of methane additional deactivation of the catalyst was observed, especially at lower hydrogen to CC12F2 feed ratios.

~9 '~ "~ .,-

= 0

~

t._.

o

I00 90 80 7o 60 50 40 30 20 10 0 600

methane -.9 sel. CH2F2

se

2~"-

sel. methane 700

50 45 40 35 30 25 20 15 ~0 5 0

~" >, "~ ~ ~ ,___,

800 900 1000 1100 Time on stream [hi

(a) H2 [ CCI2F2 = 2.2. Figures 5a, b. Influence of addition of methane.

methane ' 100 25 I..91 sel CH2F2' ~' 90 Ll_ "~ 80 20 -~ 70 ~ 60 15 ~ 50 "~40 sel CHCIF2 10 "~ 30 ~ 20 sel methane I~5 ~ 10 ~ 0 600 700 800 900 1000 1100 Time on stream [hi I

l

I

>

'1

(b) H2 / CC12F2 = 6.

No influence was observed in the selectivities of addition of CHC1F2 to the feed even at low H2 to CC12F2 feed ratios. The influence of addition of CHC1F2 on the conversion is comparable to the influence of methane. Figure 6 gives a comparison of the CC12F2 conversions in all experiments as a function of time on stream. Both in the presence of methane and of CHCIF2 additional catalyst deactivation was observed, especially at the lower hydrogen to CC12F2 feed ratios. None of the catalysts restores its initial activity completely after the methane is switched off. After the CHC1F2 is switched off at the ratio of 6 the

374

catalyst stops to deactivate but does not restore its initial activity. Therefore, the catalyst deactivation is irreversible under the conditions applied. methane

I0080

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

CHC1F2

' .............. i

..........

H2

,.__..

g 60 40

r

;'2

20 [

ti

0

0

250

500

750

I000

1250

1500

1750

Time on stream [h]

Figure 6. CC12F2 conversion as a function of time for several hydrogen to CCI2F2 feed ratios. The main by-products of the reaction were ethane, CH3F, CH3CI, and propane. The selectivities as a function of time are depicted in Figure 7a and b. At low hydrogen to CC12F2 feed ratios the selectivities to ethane and propane are higher. Addition of methane or CHC1F2 to the feed also leads to an increase in the selectivities to ethane and propane. methane CHCIF2

3

methane

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

CHC1F2

0

I sel CH3F

.

,

,

i ~

0 0

250

500

250

750 1000 1250 1500 1750

H2 / CC12F2 = 1.5.

,

~

,~

set, CH3C1

....-

750 1000 1250 1500 1750

Time on stream [hi

Time on stream [hi

(a)

500

.,

(b)

H2 / CC12F2 = 6.

Figure 7 a, b. Selectivities to by-products as a function of time on stream.

4. D I S C U S S I O N 4.1. Catalyst stability The catalyst activity, selectivity, and stability are simultaneously influenced by the H2 to CCI2F2 feed ratio. Also the time needed to reach a steady-state catalyst performance depends

375

on this ratio. Apparently it takes some time before the adsorption/desorption equilibrium between all the reactants has been established. The catalyst deactivation at lower H2 to CC12F2 ratio can be tentatively explained by coke deposition on the catalyst surface. As the concentration of hydrogen is increased, coke precursors are hydrogenated from the catalyst surface. The amount of ethane and propane produced, which can be an indication for coke formation, is proportional to the deactivation and the H2 to CC12F2 feed ratio. In separate studies with a model catalyst evidence for formation of a palladium carbide phase was found also indicating coke formation on the catalyst surface[6]. This suggests that high hydrogen concentrations should be used to avoid the coke deposition. It can, however, be expected that the hydrogen concentration cannot be increased infinitely because palladium is expected to sinter easy in a 100% hydrogen atmosphere at reaction temperatures [ 11 ]. On the other hand CFCs can be expected to regenerate a deactivated palladium on activated carbon catalyst, as has been claimed by Kellner [12]. Therefore, catalyst stability can be obtained when the hydrogen concentration is high enough to prevent coke formation and the CC12F2 concentration is high enough for in-situ stabilization of the palladium dispersion. This explains why only two stable operations, at hydrogen to CC12F2 ratio of 6 and 10, were observed. When the H2 to CC12F2 feed ratio of 10 is doubled to 20 the catalyst starts to deactivate. This is probably caused by sintering of the metallic palladium, but also the presence of an optimum in activity and selectivity as a function of disperion could explain the observed phenomena. Thus the determination of the dispersion under reaction conditions is important. The addition of methane leads to a lower CC12F2 concentration and a lower hydrogen concentration. Thus the conversion decreases by addition of methane. The addition of methane also leads to a higher selectivity to ethane and thus coke formation. This effect can best seen at the ratio of 6, where the stable catalyst starts to deactivate after addition of methane. At the ratio of 20 no influence of methane on the selectivities and no additional deactivation were observed. Therefore, the mechanism of deactivation is better explained by sintering at this H2 to CCI2F2 ratio. It is not clear whether the added methane interferes with the catalyst surface and has an influence on the desorption of methane from the surface or that the observed additional ethane is caused by a lower hydrogen concentration. There are, however, indications that the observed additional deactivation is caused only by a lower hydrogen concentration. This can be seen in the influence of addition of CHC1F2 to the feed. Although there is little influence of addition of CHC1F2 on the selectivities, a comparable influence on the catalyst stability as addition of methane is observed. The catalyst deactivation increases by the addition of CHC1F2, and this deactivation coincides with a higher selectivity to ethane. Thus, the most probable explanation for the additional deactivation is the lower hydrogen concentration, and, therefore, higher coke formation during addition of both CHC1F2 and methane. This explains why the effects of extra deactivation are less at higher ratios and not observed at the high ratio of 20.

4.2. Catalyst performance The catalyst performance depends on the H2 to CC12F2 feed ratio. The selectivities to CHEF2 and CHC1F2 are influenced by the HE to CCIEF2 feed ratio, while the selectivity to methane is independent of this ratio. We have previously proposed a reaction mechanism with serial reactions on the catalyst surface and minor readsorption of the intermediate products, which is depicted in figure 8 [4,5]. Thus the kinetics of the reaction follows mainly parallel reaction pathways, in which the selectivities are not influenced by the conversion, and a

376 remarkable low change in selectivities as a function of temperature is observed. Such a behaviour has also been observed by other authors [7-9]. A mechanism with a serial reaction on the catalyst surface can explain the observed changes in selectivity of CHCIF2 and CH2F2 as a function of H2 to CC12F2 feed ratio. A higher hydrogen concentration would lead to more hydrogen adsorbed on the catalyst surface thus favouring the adsorption steps in which more hydrogen are involved. However this mechanism does not fully explain the constant selectivity to methane. An explanation for this phenomenon would be that the methane is formed via a different route on the catalyst surface. This could be caused by either different sites on the catalyst surface or a difference in the adsorption of CC12F2. It is not likely that the methane is formed on different sites because the selectivity to methane remains unchanged in spite of deactivation of the catalyst. Therefore, the conclusion would be that the 'methane' and the 'non-methane' sites deactivate at the same rate. A difference in the adsorption of CC12F2 would be a more simple explanation. Figure 9 shows part of a reaction scheme for the CC12F2 hydrogenolysis reaction. In this scheme all the hydrogenolysis reactions starting from CC12F2 are depicted.

HO,F

7,F

[CCI2F'2 ~

C ,F OH,

CHCI2F "

12

/ [CHCIF2} 22

= CH2CIF 31

\/ l

Figure 8. Reaction mechanism with apparent parallel reaction kinetics.

= CH2CI2

-

32

=

= CH3CI

\/

40

',,

CHeF - - 41

CH, 50

Figure 9. Reaction scheme of CFCs starting from CC12F2.

If during the adsorption of CC12F2 a carbon-chlorine bond is broken the products CHC1F2 and CH2F2 are preferentially formed. The sequential reaction to methane via CH3F will be of minor importance. This can be expected because the carbon fluorine bond is stronger than the carbon chlorine bond. If on the other hand a carbon-fluorine bond is broken during adsorption, the reaction follows a different path. In this case intermediates, which can be expected to be more reactive, are formed, which results in a high selectivity to methane. The ratio between the selectivity to methane and the sum of the selectivities to CHCIF2 and CH2F2 depends on the adsorption mode of CCI2F2, which is independent of the hydrogen concentration. This explanation is consistent with the observation that the sum of the selectivities to CHC1F2 and CHEF2 is the same during each experiment. The lower selectivity to CHCIF2 during addition of methane is remarkable. Because of the lower hydrogen concentration during addition a higher selectivity to CHC1F2 can be expected. A tentative explanation would be that the selectivity to CHCIF2 is proportional to the CCI2F2 concentration. The higher selectivity to CHCIF2 is then caused by coupled desorption of CHC1F2 and adsorption of CC12F2. CHC1F2 is removed from the surface by adsorption of CCIEF2, and, therefore, the selectivity to CHCIF2 is proportional to the CCIEF2 concentration. This also explains the fact that CHC1F2 has no influence on the catalyst performance in spite of the fact that the reactivity of CHCIF2 is 5% of the CCl2F2 reactivity [4,5]. CHC1F2 does not adsorb on the catalyst surface in the presence of CC12F2.

377 Thus it is clear from both the remarkable constant selectivity to methane and the influence of addition of methane on the selectivities to CHCIF2 and CH2F2, that the serial reaction mechanism has to be modified in order to explain the observed phenomena. Therefore, more measurements will be performed in order to determine the adsorption of CFCs on the metal surface and the influence of other products, such as CH2Fz and HC1 on the catalyst performance [ 11 ].

4.3. Process design The research is strongly related to a future plant operation. From these stability tests a preliminary layout of a future process can be determined. Obviously a high hydrogen to CC12F2 feed ratio is needed. This leads to a higher conversion, higher selectivity, and higher stability. The excess hydrogen cannot be purged and has to be recycled. It can be expected that some light products such as methane and ethane will accumulate in the recycle. From the methane addition experiments it can be concluded that some build-up of methane in the hydrogen recycle can be tolerated without a negative influence on the catalyst performance. The additional methane has to be compensated for with a higher hydrogen to CC12F2 feed ratio to avoid coke formation. However, the hydrogen concentration cannot be raised too high and therefore, the amount of methane in the recycle is limited. The methane to CC12F2 feed ratio is estimated to have a maximum of 1. The amount of methane in the recycle can be controlled with a purge. A scheme of a plant layout is depicted in Figure 10. recycle

D.. purge

Hydrogen .._

product H2F~

oil water

Feed pretreatment

waste

acids

Reactor

Acid removal

Lights separation

Distillation

Figure 10. Process lay out. In the feed pretreatment section oil and water are removed from the recovered or converted CC12F2. The reactor type will be a multi-tubular fixed bed reactor because of the exothermic reaction (standard heat of reaction -150 kJ/mol). After the reactor the acids are selectively removed and collected as products of the reaction. In the light removal section the CFCs are condensed and the excess hydrogen is separated and recycled. The product CH2F2 is separated from the waste such as other CFCs produced and unconverted CC12F2. The waste will be catalytically converted or incinerated. A preliminary process design has shown that such a CFC-destruction process would be both technically and economically feasible.

378 5. CONCLUSIONS Palladium on activated carbon is a very suitable and stable catalyst for the selective conversion of CC12F2 into CH2F2. The performance and stability of the catalyst strongly depend on the H2 to CC12F2 feed ratio. At low H2 to CC12F2 feed ratio, the mechanism of deactivation is presumably coke deposition, while at high ratios sintering of palladium causes catalyst deactivation. An optimum in both catalyst performance and stability as a function of the hydrogen to CC12F2 feed between 6 and 20 is found. The reaction follows mainly parallel pathways, rather than the expected serial reaction. Methane is formed via a reaction pathway different from the formation of CHCIF2 and CH2F2. The ratio between the selectivity to CHC1F2 and CH2F2 is proportional to the concentration of CC12F2. A process for the conversion of CC12F2 into CH2F2 would include a multi-tube fixed bed reactor with a hydrogen recycle in which a limited amount of methane is allowed. This process would be both technically and economically feasible. REFERENCES 1. 2.

M.J. Molina and F.S. Rowland, Nature, 249 (1974) 810. United Nations Environmental Programme, "Report of the Ad-Hoc Technical Advisory Committee on ODS Destruction Technologies", (May 1992). 3. A. Wiersma, E.J.A.X. van de Sandt, M. Makkee, H. van Bekkum, and J.A. Moulijn, Dutch Patent Application, NL 94.01574, (1994). 4. A. Wiersma, E.J.A.X. van de Sandt, M. Makkee, H. van Bekkum and J.A. Moulijn, in Proceedings of the 1st World Congress Environmental Catalysis, Pisa, Italy, G. Centi, C Cristiani, P. Forzatti and S. Perathoner (Eds.),(May 1-5 1995) 171. 5. A. Wiersma, E.J.A.X. van de Sandt, M. Makkee, C.P. Luteijn, H.van Bekkum and J.A. Moulijn, Catal. Today, 27 (1996) 257. 6. E.J.A.X. van de Sandt, A. Wiersma, M. Makkee, H. van Bekkum and J.A. Moulijn, submitted for publication in catalysis today. 7. B. Coq, J.M. Cognion, F. Figueras and D. Tournigant, J. Catal. 141 (1993) 21. 8. G. Moore and J. O'Kelly, European Patent Application 508660, Imperial Chemical Industries (1992). 9. R. Ohnishi, W.-L. Wang and H. Ichikawa, Stud. Surf. Sc. Catal., 90 (1994), 258 10. H. Jin, S.Y. Jeong, B.S. Kim, J.M. Lee and S.K. Ryu, in Proceeding of the 6th carbon conference, Granda, Spain, (July 3-8 1994), 352. 11. E.J.A.X. van de Sandt, A. Wiersma, M. Makkee, H. van Bekkum, and J.A. Moulijn, to be published. 12. C.S. Kellner, J.J. Lerou, V. Rao and K.G. Wuttke, WO patent 91/04097, Du Pont de Nemours, (1991).

Acknowledgement: R. Artan, R.A.M. Avontuur, B.J. Bezemer, J.H.A. van den Hoogen, N.B.G. Nijhuis, Y. Verbeek, and C.P. Luteijn are greatfully acknowledged for their work in the preliminary process design. Sponsors:

AKZO-Nobel, Dutch Ministry of Housing and Environmental Affairs, AlliedSignal Fluorocarbon Europe B.V., Johnson Matthey Plc., KTI b.v., and the European Union.

J.W. Hightower, W.N. Delgass, E. Iglesia and A.T. Bell (Eds.) 11th International Congress on Catalysis- 40th Anniversary

Studies in Surface Science and Catalysis, Vol. 101 9 1996 Elsevier Science B.V. All rights reserved.

379

Catalytic fluorination o v e r c h r o m i u m oxides. Preparation of hydrofluorocarbons

S. Brunet, B. Boussand and J. Barrault

Laboratoire de Catalyse - URA CNRS 3 50 - Eeole Sup6rieure d'Ingenieurs de Poitiers 40 Avenue du Recteur Pineau - 86022 POITIERS Cedex - FRANCE

Abstract

The fluorination of CF3CH2CI into CF3CH2F over chromium oxides is accompanied by a dehydrofluorination reaction (formation mainly of CF2=CHCI). This dehydrofluorination is responsible for the deactivation of the catalyst. A study of the dehydrofluorination reaction of CF3CH2CI proves that the reaction is favoured when the degree of fluorination of chromium oxide increases. Consequently it would be favoured on strong acid sites. Adding nickel to chromium oxide decreases the formation of alkenes and increases the selectivity for fluorination while the total activity decreases. Two kinds of active sites would be present at the catalyst surface. The one would be active for both the reactions of dehydrofluorination and of fluorination, the other only for the fluorination reaction.

1. INTRODUCTION The F / CI exchange in chloroalkanes is a route to HFCs. For example, different routes can be possible for the synthesis of CF3CH2F [ 1,2 ]. Our focus is on its preparation from CF3CH2CI and HF with chromium (III) oxide as a catalyst. This fluorination is accompanied by a dehydrofluorination which produces chloroalkenes (mainly CF2=CHCI ) resulting in a deactivation of the catalyst. Indeed this haloalkene could polymerise and thus lead to coke formation. The reactions involved are : CF3CH2C1 + HF

.~

~ CF3CH2F + HCI

CF3CH2CI

-.-

~

CF2 = CHCI + HF

The catalytic activity for F/CI exchange depends on the amount of reversibly oxidized sites with a linear relationship between the activity and the number of such sites [3,4].

380 The aim of our work is to study, under adequate operating conditions, the dehydrofluorination reaction of CF3CH2CI so as to determine the nature of the sites involved in the fluorination and the dehydrofluorination of CF3CH2CI. Thus a selective poisoning of dehydrofluorination sites would allow to increase the selectivity for the fluorination reactions. With this in view, we studied the development of the dehydrofluorination reaction of CF3CH2CI as function of the degree of fluorination of chromium oxide. Moreover, nickel and chromium oxide catalysts were prepared and tested for the dehydrofluorination reaction. Nickel oxide, a basic compound [5], could poison selectively the sites involved du~ng the dehydrofluorination reaction.

2. EXPERIMENTAL

2.1. Catalyst preparation The chromium oxides were prepared according to the following procedure [6-8]. Chromium oxide resulted from the dehydration of chromium hydroxide obtained by the addition of an ammonia solution (5M) to a solution of chromium nitrate (0.5M). The final pH was equal to 7.5 and the hydroxide formed was kept constantly stired and heated at 80~ for 1 h so as to obtain complete precipitation. This solid was filtered and washed three times with hot distilled water and dried for 16 h in an oven at 90~ It was then submitted to a dynamic thermal treatment under nitrogen at 380~ for 8 h. The chromium oxide formed was cooled down under the same vector gas. The mixed chromium and nickel (5% and 10% Ni atomic) catalysts were prepared by dehydration of mixed chromium and nickel hydroxides prepared by adding an ammonia solution to a solution of chromium and nickel nitrate in order to maintain a pH = 7+_1. The final pH was equal to 7.5. The hydroxyde treatment was the same as the one already described for chromium oxide.

2.2. Catalytic fluorination and dehydrofluorination The fluorination of CF3CH2CI was carried out at 380~ under atmospheric pressure in a fixed bed dynamic reactor. In a first operation, the catalyst was fluorinated "in tim" for 2 or 70 hours by HF mixed with nitrogen (ratio N 2 : HF: = 5 : 4) at 380~ CF3CH2CI was then injected into the reactor in the presence of HF. The operating conditions of this fluorination of CF3CH2CI were: temperature = 380~ catalyst weight = 50 mgo contact time = O.Ols, HF : CF3CH2CI : N 2 = 4 : 1 : 5. The same operation was repeated with a nonfluorinated catalyst. The catalytic activity for the CF3CH2F formation was measured both over the prefluorinated and the non-fluorinated catalyst after a 2 hour reaction with CF3CH2CI. The products resulting from the reaction were injected with an automatic sampling valve into a GIRA GC 181 gas phase chromatograph and analyzed with a flame ionization detector. The separation was carried out in a capillary column DB5 (J and W Scientific).

381 The catalytic activity for the dehydrofluorination of CF3CH2CI was measured at 320~ under atmospheric pressure, in a pulse flow reactor [9]. Pulses of pure CF3CH2CI were injected into a helium stream every ten minutes. The amount of ehloroeompound was adjusted (40.9 pmol) in order to obtain a conversion ea 10 %. The products resulting from the reaction were injected into a Varian 3400 gas phase chromatograph and analyzed with a flame ionization detector. The separation was made in a capillary column BP5 (SGE). The catalytic activity of the catalyst was measured at~er a 5 hour reaction with CF3CH2CI (by the amount of ehloroalkene formed : CF2=CHCI , CFCI=CHCI (Z and E)). 2.3. Fluorine titration

The fluorine titration of chromium oxide was carried out at the Elf-Atoehem Research Center, Pierre-B6nite.The catalyst mineralization was carried out in a Parr bomb by reaction with sodium peroxyde. Fluorine ions were then titrated by a potentiometric method with a specific fluoride electrode. 3. RESULTS 3.1. Transformation of CF3CH2C! over a non fluorinated chromium oxide

The transformation of CF3CH2CI was studied at 320~ in a pulse flow reactor. Indeed, in a dynamic reactor, the significant alkene formation leads to a rapid deactivation of the catalyst. The reaction is carried out in absence of I-IF in order to favour the dehydrofluorination reaction. Products distribution is shown in Fig. 1. 1,50

E

],oo

00000000

0

E

.~0,50 r

- ~

,/~

.

r 0 550

0 600

650 T/K

700

750

550

600

650 T/K

700

750

Fig. 1 Oxidation of CH4 in the presence and absence of H2 over Fe0.sAl0.sPO4 catalyst. (A) CH4 conversion; (O) in the absence of H2, (Q) in the presence of 50.7 kPa Hz; (B) Product selectivities; black symbols: in the presence of 50.7 kPa He; white symbols: in the absence of H2. ( S ) and (C)) CH3OH, (O) and ( ~ ) HCHO, (A) and (/X) CO, ( I ) and (7-1) CO2. Conditions: P(CH4) = 33.8 kPa, P(O2)= 8.4 kPa, W/F= 0.028 g h dm-3

The AIPO 4 prepared by a similar method to that used for the preparation of the Fe0.sAlo sPO 4 catalyst was also tested for the oxidation of CH 4 by 02 both in the presence and absence of H 2. No reaction occurred at 573-723K, suggesting that the iron was an indispensable active component for the conversion of C H 4. AS compared with the FePO 4 catalyst reported previously [13], the space time yield for the sum of CH3OH and HCHO (per unit gram)at 723 K increased for about 10 times for the Fe05Alo 5PO4 catalyst, i.e., from 1.36 mmol h 1 g-I for FePO 4 to 14.1 mmol h a g.1. This must be ascribed to the high specific surface area of the catalyst used in this study.

400 3 . 2 . Characterization of the Feo.sAio.sPO4 catalyst The results of the XRD measurement showed that the Fe0.sA10.sPO4 catalyst was "~most in amorphous state. Only a very broad peak at 20 of ca. 23 degree was observed. The M6ssbauer spectroscopic study on this catalyst showed one doublet of iron with the isomeric shift of 0.31 mm s a (ot-Fe was used as the reference) and the quadrupole splitting of 0.62 man s 1. These parameters are very close to those observed for FePO 4 [13, 14], suggesting that the iron cation in the catalyst is tetrahedrally coordinated with oxygen and isolated by four PO4 tetrahedral units. Such coordination circumstance was suggested to be a key factor for the iron site effective for the oxidation of CH4 to CH3OH by H2-O 2 gas mixture [15]. After the reaction for 5 h in a reactant stream of CH,, O z and H 2 (P(CH4)= 33.7, P(O2)= 8.4 and P(H2)= 50.7 kPa), the catalyst was analyzed 200 by XRD, M6ssbauer and XPS studies. As regarding the XRD and M6ssbauer spectroscopic measurements, obvious changes were not observed before and after the reaction. On the other hand, a marked change was observed in the XPS spectrum of the catalyst after the reaction. As r shown in Fig. 2, besides the peak at 57.7 eV, N which was the only peak of Fe3p obtained for the sample before the reaction and was ascribed to Fe(III), a clear shoulder at 56.1 eV was observed after the reaction. This can be ascribed to the ! I I l Fe(lI) on the catalyst surface. The same 64 60 56 52 phenomenon has been reported for FePO4 catalyst [13]. Such observations suggest the occurrence of Binding Energy/eV the redox of iron between Fe(llI) and Fe(ll) during Fig. 2 XPS spectra for Fe 3p. the reaction. We believe that this redox plays a (a) before the reaction, (b) after the key role in the formation of a new active center and reaction in a reactant stream of CH4, thus is important in the selective oxidation of CH4 H2 and 02. to CH3OH as will be discussed later.

I

3 . 3 . A d s o r b e d s p e c i e s on the catalyst in H2-O 2 gas mixture In situ FTIR spectroscopy was used to study the adsorbed species generated on the catalyst surface in the presence of H 2 and 02. Before the experiment, the catalyst wafer was pretreated by 02 (5.3 kPa) at 723 K for 1 h followed by evacuation at the same temperature in vacuum (ca. 6x10 3 Pa) for 2 h. After the pretreatment, the temperature was decreased to a desired one in vacuum and IR spectrum was recorded at that temperature. The spectra of the catalyst wafer recorded at different temperatures were used as the background ones for the adsorption studies described below. Fig. 3 shows the IR spectra of the adsorbed species generated on the surface in the presence of both H 2 and 02 at different temperatures. No obvious absorption band due to the adsorbed species was observed at 298 K. When the temperature was increased to 473 K, two weak bands at 3740 and 3670 cm 1 assigned to the stretching vibrations of a non-acidic and acidic OH groups, respectively, were observed. These two bands were also observed in H 2

401 alone at >473 K. Moreover, the isotope substitution of 1602 with ISoz in the H2-O 2 mixture did not affect the position of these OH stretching bands. Therefore, the OH groups must be formed due to the interaction of H z with the lattice oxygen of the catalyst surface. Particularly, the acidic OH at 3670 cm 1 may be generated as a result of the reduction of the surface Fe(IlI) according to the following reaction, -Fe(IlI)- O-P-O- + 1/2H z ---, -Fe(II) + HO-P-O(1) Besides the absorption bands attributed to OH groups, another weak band at 1075 cm 1 was observed at 473 K. This band can tentatively be assigned to an 02- species. The formation of 02- species with an absorption band at 1050-1200 cm a was reported for many oxides [1618]. Here, the O.,- species must be formed on the reduced iron site through one electron transfer from Fe(II) to the adsorbed Oz. However, this 02- species was very unstable because it disappeared when the temperature was increased above 473 K, probably due to the desorption or to the transformation to u% another type of oxygen species. As shown in Fig. 3, when the 0.01 O.Ol temperature was increased to 573 K, an absorption band at 895 crn 1 appeared and the bands of OH groups became stronger. The r162 intensity of the absorption band at 895 cm ~ 673K increased with a rise in temperature to 623 and 673 K. The stability of the species giving the absorption band at 895 cm 1 was examined under vacuum at different 623K temperatures. Only a slight decrease in the intensity was observed when the system was degassed at 623 K. These results indicate that the absorption at 895 cm 1 was ascribed to a chemisorbed species peculiarly 298K generated and stable in the presence of H 2 and O2. It should be noted that the band at 895 cm 1 has never been observed in the | ! presence of O., alone. 1400 800 ;00 3 4 0 0 3 0 0 0 2 6 0 0 In order to identify the new adsorbed Wavenumber/cmq species of the band at 895 cm a, the effect of the isotope substitution of 160z with 1sOz Fig. 3 IR spectra of adsorbed species on was examined. The spectra shown in Fig. 4 the catalyst in H2-O2 at different were recorded for the catalyst in the presence temperatures. of H 2 and aSO2 in the wavenumber region of P(H2)= 13.3 kPa, P(O2)- 1.33 kPa. 700-950 cm -1 with 100 scans. The results show that the absorption band at 849 cm 1

I

402 appears when the temperature is higher than 573 K. This means that the absorption band at 895 cm -1 is shifted to 849 cm 1 when 1602 is replaced by 180,_. This result strongly suggests that the absorption at 895 cm 1 is ascribed to an adsorbed oxygen species on the catalyst surface. However, the results obtained above do not allow us to distinguish whether this adsorbed oxygen species is a diatomic type oxygen such as peroxide species (022) or a monoatomic type oxygen such as iron oxo species (Fe=O). Generally, both the adsorbed peroxide and M=O type oxygen species (where M is a metal) give IR absorption band at 800-950 cm 1 [ 17-21]. We can calculate the isotope shift in IR absorption due to the substitution of 1602 with 1802 by assuming the adsorbed species to be an O, 2 or a Fe=O species. If the absorption is ascribed to the O-O stretching mode of a peroxide species, the ratio of the wavenumbe~ for the 180-180 and 160-160 stretching modes should roughly be given by the ratio of the square root of the reduced masses, viz. [(1/18 + 1/18)/(1/16 + 1/16)] 1/2 or 0.943. Thus, the band at 895 cm -1 should be shifted to 844 cm 1 by the substitution of 1602 with 1802 in this case. If the adsorbed species is of Fe=O type, a similar calculation suggests that the substitution of 1602 with 1802 will shift the peak at 895 crn ~ to 855 cm -1. As described above, the experimental results showed that the band at 895 cm 1 was shifted to 849 cm 1 due to the replacement of 1602 with 1802 in the H2-O 2 gas

[0.005

473K

i 900

i 800

700

Wavenumber/cm-I

Fig. 4 IR spectra of the adsorbed species on the catalyst in Hz-18(b_ at different temperatures. P(H2)= 9.3 kPa, P(1802)= 0.93 kPa.

1 900

I 800

700

Wavenumber/cm-1

Fig. 5 IR spectra of the adsorbed species on the catalyst in the presence of H2 and a mixture of 1602, 16Oa80 and 1802_ at different temperatures.

403

mixture. Therefore, we still cannot identify the oxygen species, viz. peroxide or iron-oxo species? Thus, in this context, further experiments using an isotopic mixture of 1602, ~60180 and 180z were carried out. IR spectra were recorded under the mixture of H a, 1602, 160~80 and 180, at different temperatures. The initial pressures of H e and O2 at 298 K were 8.7 and 0.93 kPa, respectively. The analytical result of the composition of the isotopic oxygen by a quadrupole mass spectrometer was 2.0:0.5:1.0 for the 1602"160180:1802. The IR spectra obtained in the region of 950-700 crn 1 were shown in Fig. 5. Three absorption bands at 895, 870 and 849 crn a were observed at 573 K and their intensities were increased further with the increase in temperature to 623 K. These observations can only be explained by assuming that these absorption bands are ascribed to the O-O stretching mode of the adsorbed peroxide. The calculation shows that if the band at 895 cm 1 is ascribed to the stretching of 160-160, the wavenumbers for the stretching vibrations of 160-180 and ~80-180 should shift to 870 and 844 cm 1. Therefore, the observed bands at 895, 870 and 849 cm 1 can be assigned to the O-O stretching vibrations of 160-160, 160-180 and 180-180, respectively. In conclusion, an adsorbed peroxide species is formed on the catalyst surface in the presence of H,_ and O,. at e573 K.

3 . 4 . The reaction of the adsorbed oxygen species with methane As described above, the temperature needed for the formation of the adsorbed peroxide was ca. 573 K, which was almost the same as that for the initiation of the oxidation of CH 4 to CH3OH in the presence of H, and O2. This peroxide may directly be related to the selective oxidation of CH 4 to CH3OH. Therefore, the reactivity of this adsorbed peroxide with CH4 was investigated by in situ FTIR spectroscopy as follows. The adsorbed peroxide species for the reaction was prepared in a gas mixture of H a and O~ (P(H2)= 13.3 kPa, P(O2)= 1.33 kPa)at 673 K. Then, the temperature was rapidly lowered to 373 K and the gas phase was evacuated at this temperature. The intensity of the band at 895 cm ~ attributed to the peroxide was not affected by evacuation at this temperature. After the evacuation, CH4 gas (5.33 kPa) was introduced to the cell. IR spectra were recorded at different temperatures in the atmosphere of CH 4. Fig. 6 shows the IR spectra where the absorption due to the gaseous CH 4 at the corresponding temperatures has been subtracted. No obvious change in the intensity of the band at 895 cm x occurred after CH4 was introduced at 373 K. When the temperature was increased to 473 K, the intensity of the band at 895 cm 1 decreased markedly and three new bands at 2936, 2870 and 1050 cm -1 appeared simultaneously. The former two bands can be assigned to the anti-symmetric and symmetric stretching vibrations of C H 3 groups and the latter to the C-O stretching vibration for the methoxide species [22]. The reverse band at ca. 3010 cm ~ arose due to the decrease in the amount of the gaseous CH4 because the gaseous CH,, shows a very strong absorption band at 3010 cm ~. Accompanying the formation of the methoxide, the intensity of the band at 3668 cm I, assigned to the acidic OH group on the catalyst, increased obviously. In order to confirm that the absorption bands appeared at 3668, 2936 and 2870 cm -~ are derived from the reaction of CH4 with the adsorbed peroxide, CD 4 was used instead of CH 4. The IR spectra recorded in the regions of 4000--2000 cm 1 and of 1500-700 cm ~ are shown in Fig. 7. The experimental procedures for the results of Fig. 7 were the ~ m e as those for Fig. 6. When the temperature was increased to 473 K in the presence of CD4, the bands at 895 cm 1, ascribed to the adsorbed peroxide, decreased remarkably. The new bands at 2692, 2208 and 2124 cm 1 were observed simultaneously. Obviously, these are the isotopic bands of those

404 of 3668, 2936 and 2870 cm 1 observed in the case of C H 4 a s the reactant in Fig. 6. The ratios of wavenumbers for the corresponding pairs of the bands were 1.375, 1.330 and 1.351, respectively. All these values are reasonable ones for O-H/O-D, C-H/C-D (symmetric) and CH/C-D (anti-symmetric)stretching vibration modes. In conclusion, the adsorbed iron peroxide generated in the gas mixture of H 2 and 02 reacts with C H 4 at >473 K, forming adsorbed OH and methoxide species on the catalyst surface. Both species must be the intermediates in the formation of CH3OH.

I e~

I

O.Ol

,~,~

I

O.O1

II

tt%

' a2

I

O.Ol

I

0.01

!

J

r~l r~l .--I ~ r {'xll

~I~

573K ~ 573K

o

< 173K

473K

;73K 173K !

4000 3800

3400

3000

2600

1400

800

I

I

3000

2000

1400

1

I

801

Wavenumber/cmq

Wavenumber/cm-I

Fig. 6 IR spectra of the adsorbed species on the catalyst after the reaction of CH4 with the oxygen species generated in a H2-O2 gas mixture at different temperatures.

Fig. 7 IR spectra of the adsorbed species on the catalyst after the reaction of CD4 with the oxygen species generated in a H2-O2 gas mixture at different temperatures.

The effects of the partial pressures of H 2 and 02 on the formation of the adsorbed peroxide species were examined. These results have been compared with the kinetic results for the conversion of CH, by using the flow system. As shown in Fig. 8 (A), the surface concentration of the peroxide increased roughly linearly with a rise in the partial pressure of H~.. On the other hand, it was saturated at a low partial pressure of O2 (Fig. 8 (B)). Very similar trends were observed for the kinetic measurements for the conversion rate of CH 4 as functions of the partial pressures of H2 and 02 as shown in Fig. 9. These observations further support that the peroxide species is responsible for the partial oxidation of CH4.

405

~/~

0.03 _ (A)

0.03

0.02

0.02

.= < 0.01

0.01

t~ t.... O

I 1 1 0 10 15 20 0 1 2 3 P(H2)/kPa P(O2)/kPa Fig. 8 The intensity of the band at 895 cm -1 vs. the pressures of H2 (A) and 02 (B). T= 673 K; (A) P(O2)= 1.3 kPa; (B) P(H2)= 13.3 kPa.

5

"7, .m

_~ 0.6

(B)

(A) 0.4-

o4

o= "7.~ 0.2 w

= rj

~

0_

0

O.2-

658 K

20

40 P(H2)/kPa

60

00

1

1

1

1

1

2

3

4

5

P(O2)/kPa

Fig. 9 The conversion rate of CH4 vs. the pressures of H2 (A) and 02 (B). (A) P(O2)= 8.4 kPa; (B) P(H2)= 50.7 kPa.

3.5.

Reaction

mechanism

As described earlier, it is suggested that H 2 reduces the iron sites on the catalyst surface, forming water and OH groups. Thus, the cofeeding of H 2 must guarantee a high concentration of the active iron site, viz. Fe(II), as indicated by the XPS studies. Oxygen must be adsorbed on this reduced iron site by accepting electrons, initially forming O~.- species (absorption band at 1075 cm 1) as demonstrated in Fig. 3. This oxygen species itself may be very unstable because the band at 1075 cm 1 disappeared at >473 K. It was probably further reduced into 02- species by electrons trapped in Fe(ll) sites. The in situ FTIR results have shown that an adsorbed peroxide is generated on the catalyst surface in a H2-O: gas mixture at a573 K. Although we cannot exclude the possibility that the true active oxygen species could be an atomic oxygen intermediate (O-) resulted from the splitting of the adsorbed peroxide, the results of Figs. 6-9 strongly suggest that the peroxide species itself is responsible for the activation of CH 4. The results obtained in this study suggest the reaction mechanism in Fig. 10. As can be seen in Fig. 10, besides the redoxing iron sites, the proton accepting and donating sites, viz. phosphate groups, are also important for the generation of active oxygen species as well as for the accomplishment of the catalytic cycle. We believe that the phosphate groups directly bonded to the iron site play these specific roles [15].

406

CH3OHy

O~ . 0 ~ i )...,.Fe(II 185nm .OH =

892 +,OH

(2)

CH3~ + H20

(3)

H20

CI"I3OH +

(4)

=

892

SCHEME 1. Photochemical Conversion of Methane. Reference 1.

L.a/V~

hV

L >410nm

=

e~+

+

hvB

(1)

e ~ + MV 2+

=- MV, +

(2)

+ + H20 hvB

=

(3)

MV'++ H +

=

2.0H-----~ t-1202 ~

H + + ,OH 892 + MV 2+

(4)

H20

(5)

890 2 +

SCHEME 2. Catalytic Photolysis of Water. Reference 5. e-c~n= electron in Jr. ~.,x.J conduction band, h VB = positive hole in valance band, MV - methylviologen.

initial production of hydroxyl radical through photolysis of water. This radical may then react with a methane molecule to produce methyl radical. In the preferred reaction, the methyl radical then reacts with another water molecule to produce methanol and hydrogen. Catalytic photolysis of water to hydrogen and oxygen occurs during hxa~ation of liquid water with visible light (at wavelengths longer than 410 nm) in the presence of a solid photocztalyst suspended in the solution (Scheme 2) [5]. The photolysis sequence of interest initially produces a hydroxyl radical through the reaction of water in the presence of a doped tungsten oxide photocatalyst and an electron transfer molecule, methyl viologen dichloride hydrate (1, l'dimethyl-4,4'-bipyridinium dichloride). The proposed mechanism invokes the coupling of two hydroxyl radicals to form hydrogen peroxide, which decomposes to water and oxygen. By combining these reactions, hydroxyl radicals, generated with the photocatalyst and the electron transfer reagent, should react with methane to produce methyl radicals. In our

409

proposed reaction pathway (Scheme 3), methyl radicals react with an additional water molecule to form methanol and hydrogen, identical to Scheme 1.

hV

LaNkO3

-

Z. > 4 1 0 n m -

h

4-

ec8+ hvB

(1)

el~B + MV2+

= MVo +

(2)

+ , H20 hvB

.~ H * + , O H

(3)

MV,++ H +

=

(4)

CH4 + "OH

"

CH~ + H20

= CH3OH +

892 + MV 2+ CH3~ * H20

(5) 892

(6)

SCHEME 3. Photocatalytic Conversion of Methane e-Z 2 B = electron in 9 9 .. conduction band, h-VB = posmve hole in valance band, MV = methylviologen.

FIGURE 1. Schematic of photocatalytic reactor.

Previous research by our group [6] has confirmed literature reports [1,2] that it is possible to photolyze methane, saturated with water vapor, to produce methanol and hydrogen. In a modification of the above experiment, we were also able to photolyze methane sparged through a photochemical reactor filled with water. Recently, we began investigating the photocatalytic conversion of methane in water. 3. EXPERIMENTAL The reactor, a commercially supplied quartz photochemical reaction vessel, was fitted to meet the needs of this research (Figure 1). This included use of a Teflon-coated magnetic stirring bar in the reactor, a fritted glass sparger, a nitrogen line used to cool the UV lamp, and an injection port. Deionized water was distilled prior to use. The semiconductor photocatalysts were synthesized following a modification of the proc~ure in the literate [4]. Four dopants, copper, lanthanum, platinum, and a mixture of copper and lanthanmrt, were selected for study on the ttmgsten oxide catalyst base. In a typical experiment, 1.0000 grams of the sintered catalyst is suspended, by mechanical stirring, in water (--750 mL) containing an electron-transfer reagent, methyl viologen dichloride. A mixture of methane (5 ml_Jmin) and helium (16 mL/min) is sparged through the photocatalytic reactor. The helium is an intemal standard for on-line analysis of the reactor effluent. The reaction t ~ t u r e is maintained at --94~ by circulation of heated (-q 20~ silicone oil in the outer jacket of the reactor. A high-pressure mercury-vapor quartz lamp is used as the light source. The spectral characteristics and energy output of the lamp, supplied by the lamp's manufacturer, show that --46% of the radiated energy of the UV lamp used in this study is in the visible region. The outer surface of the lamp is cooled by a stream of nitrogen gas, while the lamp's immersion well is cc~led by a flow of tap water. The gaseous products of reaction

410 are analyzed on line and in real time by a quadrupole mass spectrometer. Liquid products are condensed from the gas stream at 0~ and analyzed by HP 5890A capillary gas chromatography. The chromatograph was equipped with a 60 m x 0.25 mm i.d. fused silica column coated with a 0.25 ~rn film of a crossed-linked version of SP-1000~ (commercially known as Nukol| Helium was used as the carrier gas with an average linear velocity of 30 cm/sec at 50~ After injection, the column was held at 50~ for ten minutes, then temperature programmed from 50~ to 200~ at 4 ~ per minutes with a five minute hold at 200~ The injector t ~ e was 200~ and the FID temperature was 250~ The injection volume was --2~tL. 4. RI~ULTS

The first series of experiments was conducted with no photocatalyst in the reactor. ~ g these experiments, a temperature d ~ d a n c e of the reaction was observed; photoconversion of methane decreased sharply with decreased temperature and was not observed below -~70~ This observation implies that a non-photochemical process is part of the reaction sequence. Several experiments were performed where the t ~ of the reactor was allowed to cycle between 60~ and 95~ In all experiments, as the t e ~ e of the reactor decreased, conversions of methane and the production of methanol decreased and were not observed below ~70~ The effect was reversible; when the reactor temperature increased above -~70~ conversion of methane and the production of methanol resumed and increased with temperature. Figure 2 shows the result of an experiment without photocatalyst where the reactor t ~ e was maintained at 97~ &ring the run. Note that conversions of methane remain relatively constant at -4% and production of hydrogen, methanol, FIGURE 2. Typical results of noncztalytic oxygen, and c.zrtx>n monoxide remain methane photoconversion. Observed flow of constant during the experiment. The products in mlJmin of CH3OH ( )'a~ large oscillations in the conversion of (O----O), H2 (t~---O), arid CD (D C3~ methane and the production of methanol were not observed during percent conversion of CH4 ( ............. ). Reaction conditions: He = 16 mL/min, CH4 = 5.0 this experiment. The four doped tungsten oxide mUmin, reactor temperature = 97~ catalysts (noted above) were synthesized and used in this study. The catalysts were analyzed by scanning electron microscopy (SEM), energy dispersive x-ray spectroscopy (EDS), x-ray diffraction (XRD), and electron

411 spectroscopy for chemical analysis (ESCA). For all catalysts except the platinum-doped tungsten oxide, these techniques were not able to detect any differences between the tungsten oxide as received and the unsintered-doped oxide because the level of doping, ~ 4 atom percent, is below the detection limits of these instruments. The sintering process produced differences that were detectable by SEM and XRD. After sintering, XRD dam showed the doped tungsten oxides to be more crystalline than the tmsintered materials as evidenced by the separation of a broad diffraction peak into two separate peaks having 2-theta values of 28.8 ~ and 42.0 ~ (Figure 3). Analysis of the sintered, doped tungsten oxides by SEM revealed that the sintered materials contained larger crystaUites with smoother edges. SEM and EDS analysis of the platinum-doped tungsten oxide photocatalyst after sintering showed the presence of platinum particles on the surface of the tungsten oxide. Figure 4 is a back-scattered electron image of the sintered platinum-doped tungsten oxide photocatalyst (the bright spheres are platinum). Analysis of the sintered platinum-doped tungsten oxide by ESCA revealed that the platinum on the surface is Pt". The catalysts were tested for their ability to catalytically photolyze water prior to their me in the methane conversion experiments. We were able to reproduce photolysis results reported in the literature [4] using these catalysts under similar conditions. All of the following FIGURE 4. Backscatter SEM Ofg~h3hr~t experiments were conducted with the photocamlyst. Bright circles in micro are photocatalyst and electron transfer platinum as identified by EDS. agent in the reactor. Figure 5 shows the results of a typical photocatalytic methane conversion experiment. Methane conversions are --4% with hydrogen and methanol as the main products of reaction. Note that after the UV lamp is turned off, the detected flow of methanol

412 decreases slowly to zero (over --2 hours). It was hypothesized that this behavior was due to stripping of methanol from the water in the reactor by the reactant gases. 0 . 0 s ~ ~ loo To confima this, methanol was injected t /I I ='~~__~;"4"x] into the reactor, previously filled with 0.04} // ON ON 9": ~ _ J ~ u ~ x , q l /-O- -F F I G" 750 mL at the operating I/ I L/ 1% temperature,w and er the concentration of 0.03~// I /~ ~/ ] ~n~' methanol in the gas flow from the => t./t_.. I ~ I u~F~ I t 6~ reactor was measured. Behavior :: similar to the reaction experiment was

!/I

o.ol

t~

I~

~:~~_~c=~

Iff

observed.

~

~J / ~ ~ ~ , ~ ' _ _ 4 TIME (HOURS)

6

~_140 820

FIGURE 5. Typical results of photocatalytic methane conversion. Observed flow of products hamtSn~ of~aoH ( ~ ch ( o - - o ~ I-h ( o - - - e ) , and CO ( m ~ m ) , ar~ reactor water temperature ( .............). WO3kLa photocatalyst.

Gas chromatographic analysis (Figure 6) of the liquid product that had condensed at O~ revealed the presence of methanol and acetic acid. Fm'ther analysis to identify other components by C~-MS was not possible due to the low concentration of products in the trap. The products were diluted by water carried over from the reactor in the flow of helium that is used as an intemal standard. As noted previously, the proposed reaction sequence of interest initially produces a hydroxyl radical,

which then reacts with methane to produce methanol. To test the validity of this hypothesis, a 30% solution of hydrogen peroxide, a good source of hydroxyl radicals, was injected into the reactor &wing p h o t ~ y t i c methane conversion. Figure 7 (a different photocatalyst preparation than Figtm~ 5) shows the results typical of the 30 peroxide injection experiments. After peroxide injection, 25 conversion of methane increases fi'om --4% to--10% methanol production increases 17 fold, and caz~n dioxide increases 5 fold, along with modest CH3COOH / 6.990 increases in hydrogen and carbon CH3OH monoxide. Introduction of hydroxyl ___.k radicals to the reactor leads to a 0 5 10 15 20 25 30 35 40 45 greater fraction of product going to TIME (MINUTES) methanol as evidenced by methane conversion increasing 2.5 times, whereas methanol production increases FIGURE 6. GC of condensed liquid product. 17 times. The increase in carbon dioxide is from "deep" oxidation of

I

413

0.20 ,..,.

200 pL 30% H202/ ADDED il~"~l.. \,

9

1

~9

i

10 2"

d!

8~E

N o.o5 -.................... tu

r~

]

~

.... w ~ ~O .

a.6

3.5

.

.

.

.

.

.

.

.

.

.

.

.

3.7

3.8

3.9

0

4.0

TIME (HOURS)

FIGURE 7. Results of addition of 200 laL of 30% hydrogen peroxide solution to a steadystate photocatalytic ~ i o n . Observed flow of products in ml_/min of CH3OH ( .... ), 0 2 !o-----o~ I-h ( e - - - e ~ c ~ e - - e ~ and O3 ( [ 3 ~ D ) , and percent conversion of CH4 ( ............. ). WO 3kLa photocatalyst.

,,,., 0.010 ._= E -J E 0.008

CH4 FLOW STARTEDf

~ 0.006 U. a LU

0.004

w

i

"'

L

a 0.001

1

TIME (HOURS)

FIGURE 8. Methanol production from various doped WO 3 photocxaalysts. WO3 dopants of La ( ~ ) , Pt ( O ~ O ) , La\Cu 50/50 (e-----el r~ c a a ~ (e- - - e l and Cu (D- - ~ l

[

methane and/or fiarther oxidation of the carbon containing products. Note the drop in methane conversion to zero for approximately 12 minutes after injection of the hydrogen peroxide. Prior to injection of hydrogen peroxide, a steady-state condition existed between the methane dissolving in the water and methane being consumed. It is likely that the introduction of excess hydroxyl radicals depleted the dissolved methane. This resulted in little methane available for conversion until steady-state conditions could be reestablished. Four doped tungsten oxide photocatalysts were prepared, platinmn, lanthanmn, copper, and a 50/50 mixture of copper and lanthanum. These catalyst were tested for their catalytic activity against a blank experiment, where all reaction conditions were identic~ except, that in the blank, no caA_alystwas present. Figtres 8-11 display the results for these experiments. After steady state conditions for the reactions were established (the reactor was at operating temperature and was being irradiated by the UV lamp), methane flow was started. In Figure 8 the production of methanol as a fimction of time is displayed. As shown in this figure, the lanthanum doped catalyst exhibits an increase in methanol production over the n o n ~ y d c reaction. The platinum and lanthanum/~ doped catalysts exhibit approximately the same production of methanol as the noncatalytic reaction. The presence of copper on tungsten oxide inhibits the production of methanol. The effects on hydrogen production for the various catalysts is shown in Figure 9.

414 Prior to the inEoduction of methane to the reactor, the n o n ~ y t i c reaction produces the most hydrogen. The only catalyzed system that produces hydrogen in significant quantities is the c/~4 0.14 one doped with lanthanum/eo~. FLOW .=_ A~er methane flow begins, the E 0.12 _J lanthanmn catalyst exhibits a large g O.lO increase in hydrogen production. A -0.08 modest increase is also exhibited by tl a the non-catalyzed and ,,, 0.06 t-c~ lanth~um\copper doped reactions. t.u 0.04 ]'he results of the addition of u.! o 0.02 l 2001LtL of 30% hydrogen peroxide solt~ion to the reactor converting oJ ~ 15 ~2 0 25 0 0.5 methane under steady-state conditions TIME (HOURS) are shown in Figtres 10 and 11. All of the calalysts exhibited larger peak production of methanol than the noncatalytic reaction after injection of FIGURE 9. Hydrogen production from various hydrogen peroxide. The lanthanum doped WO 3 photocaIalysts. WO3 dopants of doped catalyst showed the largest La ( 3 , ~ (O O), La\Cu 50/50 increase of all the catalyst with the copped doped catalyst exhibiting the least. The effect on the production of hydrogen was not as dramatic as those 0.20 ...., ... observed for methanol production with --200pL[/%, .=_ hydrogen peroxide injection. All the E H202 catalysts exhibited a slight increase in ""1 0.15 g hydrogen production while the noncatalytic reaction exhibited a decrease .J u. 0.10 in hydrogen production until the O LU I-hydrogen peroxide was consumed. LLI Analysis of the copper doped ~LU 0.05 E~ photocatalysts was conducted using scanning electron microscopy in an attempt to relate catalyst morphology d~)~ L.J ~LO.~ 0.4"" ~ ~ 0 . 6 0.8 1 1.2 characteristics to the low conversion TIME (HOURS) and selectivity of this material. Some areas of the surface of the copper FIGURE 10. Effects of addition of 200 ~tL of doped catalyst ~ e d smoother than 30% hydrogen peroxide solution on methanol that of the undoped WO3 ffigtre 12). production. WO dopants of La ( L ), Pt The grain structure of the tungsten ( O ~ O ) , La\C~ 50/50 mixture ( O ~ O ) , no oxide was obscured by this apparent catalyst ( 1 1 ~ 1 1 ) , and Cu ( ~ ~ [ 1 3 ) . coating. The smooth layer was examined using energy dispersive x-

,•o•1

415

FIGURE 12. photocatalyst.

SEM of Sintered WO3'\Cu

ray spectroscopy (EDS) in an attempt to detect differences in composition that might explain the change in morphology fi~om the extx~ed. . WO3 grain structm'e. However, if mcreas~ levels of copper were present in the smooth areas, they were below the limits of detection of EDS. This result is not surprising, since the coating was estimated to be around 0.1~tm thick, while the analysis depth was arotmd l~trn. Fm~er examination of the catalyst using X-ray photoelectron SlXCtroscopy (XPS) did indicate the presence of coptxr, but the signal was too small to accmmely determine its chemical state. The XPS copper peak was not distinct since the coming was not continuous over all catalyst particles, and the X-ray probe for XPS was of millimeter size. This resulted in contributions to the coptxr s ~ from both, coptxr in solid solution in the tungsten oxide, and from copper in the coating. It is possible that under the processing conditions, a thin layer of copper tungstate formed at the surface of the catalyst that effectively formed a barrier to the catalytic WO3 surface. The smooth appearance of ihis layer can be explained by the fact that copper tungstate melts around 1000~ and at 900~ (0.9 TH), one would expect appreciable atormc mobility and d i ~ i o n of the material in the coating layer, explaining its smooth appearance. Finally, more active WO photocatalysts did not exhibit tfi3 smooth surface appearance of the coptx~ doped material.

5. C(~CTUSIONS We have reproduced the results reported in the literature for both methane photolysis and catalytic photolysis of water. In experiments that combine elements of both systems,

416 methane and water have been converted to methanol, hydrogen, and acetic acid by a doped semiconductor p h o t ~ y s t at t ~ of--94~ and atmospheric pressure. Conversion of methane and the production of methanol are augmented by the addition of hydrogen peroxide, consistent with the postulated mechanism that proposes hydroxyl radical as an intermediate in the reaction sequence. The use of a UV filter or alternate visible light source is the next step to ensure that no UV induced reactions are occtrfing in parallel with the photocatalytic process. 6. ACKNOWI EDGMENT We would like to acknowledge the technical assistance of Richard R. Anderson, John Balmas, J. Rodney Diehl, E I ~ A. Frommell, Neil Johnson, and Joseph P. Tamilia. 7. DISCLAIMER Reference in this report to any specific commercial product, process, or service is to facilitate understanding and does not necessarily imply its endorsement or favoring by the United States ~ e n t of Energy. ~ C E S

.

K. O g t ~ M. Kataoka, J Mol. Ca., 43 (1988) 371-379. K. Ogtwa, C.T. Migita, M. Fujita, Ind Eng. Chem. Res., 27 (1988) 1387-1390. M. Ashokkumm', P. Marmhamuthu, j Mat. Sci. Lea., 24 (1988) 2135-2139. P. Maruthamuthu, M. Ashokkaamar, Int. J Hya~ogen Energy, 14 (4) (1989) 275-277. P. Mmaathamuthu, M. Ashokkaamm',K. Cmrtmath~ E. Subranmnian, M.V.C. Sastri, Int. J Hy&~ggenEnergy, 14 (8)(1989) 525-528. C.E. Taylor, R.P. Noceti, J.1L IYEste, Proceedings of the 15th Coal Liquefaction Gas Conversion Contractor's Review Meeting, Pittsburgh, PA, (1994) 777-783.

J.W. Hightower, W.N. Delgass, E. Iglesia and A.T. Bell (Eds.) 11th International Congress on Catalysis - 40th Anniversary

Studies in Surface Science and Catalysis, Vol. 101 1996 Elsevier Science B.V.

417

Role of A- a n d B - C a t i o n s in C a t a l y t i c P r o p e r t y of S u b s t i t u t e d H e x a a l u m i n a t e (ABA111019_~) for H i g h T e m p e r a t u r e C o m b u s t i o n Koichi Eguchi, Hiroshi Inoue, Koshi Sekizawa, and Hiromichi. Arai Graduate School of Engineering Sciences, Kyushu University 6-1 Kasugakoen, Kasuga, Fukuoka 816 Japan

Cation-substituted hexaaluminate compounds, ABAlllO19.a, w e r e investigated for application to high temperature catalytic combustion. Two series of modifications of the compounds was made by cation substitution; substitution of large cations in the mirror plane with lanthanides ions, and substitution of transition metals for A1 site in the spinel block. In a series of AMnAlllO19-a (A = La, Pr, Sin, and Nd), surface area and catalytic activity increased with an increase in ionic radius of lanthanides. La 3§ is superior as the large cation in the mirror plane of the hexaaluminate to other tri-valent cations with small ionic radii. The catalytic activities of LaBAlllO19.a (B = Cr, Mn, Fe, Co, Ni, and Cu) were enhanced when Mn and Cu were employed as the B-site substituents. Although Mn and Cu were also effective substituents for enhancing catalytic activity in Ba-based hexaaluminate compounds, their activity was low as compared with the La-based catalysts. These results indicate that the redox cycle of transition metal in hexaaluminate lattice and catalytic activity appears to be affected sensitively with the electronic or structural effect of large cation in the mirror plane.

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

Catalytic combustion has attracted attention from the viewpoint of the protection of atmospheric environment [1-3]. As compared with conventional flame combustion method, emission of nitrogen oxides and unburned hydrocarbon can be significantly diminished and high energy efficiency can be achieved by using catalytic combustion method. For application of catalysts to high temperature combustion system above 1000 ~ it is desirable for the development of thermally stable catalyst materials. Conventional heterogeneous catalysts are immediately deteriorated at high operation temperatures (> 1200~ Further difficulty for the combustion catalyst is the wide range of reaction temperatures exposed to the catalysts. The

418 catalysts are requested to cover several different kinetic processes and/or their transient regions. Combustion is initiated as kinetically controlled surface reaction, since the catalyst temperature is low in the inlet zone. Mass transfer process controls the reaction rate when exothermic surface reaction reached to a certain conversion level. A further increase in t e m p e r a t u r e initiates gas-phase reaction and catalyst temperature reaches the maximum (>1200~ Some hexaaluminate compounds have been reported as a new material for high temperature catalytic combustion, since they possess excellent thermal stability in m a i n t a i n i n g large surface areas above 1300~ [4-6]. Partial substitution of some transition metals for A1 significantly promotes catalytic reaction due to high reduction-oxidation activity of transition metals, and Sr0.sLa0.2MnA111019~ is one successful design for the catalyst of high t e m p e r a t u r e combustor [6]. It m e a n s t h a t catalytic and thermal behavior of hexaaluminate is greatly influenced by composing cationic species. In this study, the cationic composition was optimized to attain high catalytic activity and thermal resistance by clarifying roles of A- and B- cations in substituted hexaaluminate catalysts (ABA111019-a).

2. E x p e r i m e n t a l

ABAlllO19.a (A = Sr, Ba, La, Pr, Nd, Sin, and Gd; B = Cr, Mn, Fe, Co, Ni, and Cu) samples were prepared by hydrolysis of metal alkoxides. Calculated amounts of Sr or Ba metal and Al(OC3H7)3 were vigorously stirred in 2-propanol at 80~ for 3 h. After dissolution was complete, an aqueous solution of other metal n i t r a t e s was added to this alcoholic solution. The resultant gel was dried and calcined at 1200~ for 5 h in air. Crystalline phase and surface area of calcined samples were characterized by XRD (RIGAKU, RINT-1400) and BET method. An analytical electron microscope (JEOL, JEM-2000FX) was used for estimating crystal morphology and composition. Catalytic combustion of methane over hexaaluminate catalysts was examined in a conventional flow reactor at atmospheric pressure. A gaseous mixture of CH4 (1 vol%) and air (99 vol%) was supplied at a space velocity of 48000 h -1. Methane conversion in the effluent gas was analyzed by on-line gas chromatography.

3. R e s u l t s a n d D i s c u s s i o n 3.1. Effect of A- c a t i o n in s u b s t i t u t e d h e x a a l u m i n a t e c o m p o u n d s AMnAlllO19-a (A = La, Pr, Nd, Sm, and Gd) samples were used to investigate the dependence of the catalytic property on A cation, which occupies the large cationic

419 site in the m i r r o r p l a n e of m a g n e t o p l u m b i t e - t y p e c r y s t a l s t r u c t u r e . F o r the AMnA111019.~ samples, except for A = Gd, the hexaaluminate phase was formed as a m a i n crystal phase, but a small a m o u n t of AA103 was formed as an i m p u r i t y phase. The amount of AA103 increased with a decrease in ionic radius of A. The cationic size of Gd in GdMnAlnO19_a sample was too small to lead to h e x a a l u m i n a t e formation, but resulted in phase separation of GdA1Os and a-A12Os. Surface area of AMnAlnO19.a (A = La, Pr, Nd, and Sin) samples increased monotonously with an increase in ionic radius of A (Figure 1). T10% is the temp e r a t u r e at which m e t h a n e conversion reaches 10%, and was lowered with an increase in ionic radius of A (Figure 1). This result suggests t h a t the sintering resistance of samples are enhanced as the ionic radius of A becomes large. It is, therefore, a p p a r e n t t h a t La s§ is superior as the large cation in the mirror plane of the hexaaluminate in achieving both high t h e r m a l resistance and high activity to other tri-valent A cations with small ionic radii. Every hexaaluminate compounds crystallizes in ~-alumina or magnetoplumbite structure (Figure 2). Both of these s t r u c t u r e s consist of alternative stacking of a spinel block and a mirror plane along the c direction [7]. Closed packing layer of oxygen is located in the spinel block; however, the packing in the mirror plane and along the c direction is relatively loose. These two crystal structures, being different in the coordination in the mirror plane, are regarded as hexagonal layered structures. The anisotropic nature of this crystal strongly affects the kinetics of crystal growth and diffusion. Iyi et al. studied the h e x a a l u m i n a t e

Fig. 2 Crystal structures of hexaaluminate compounds.

420 Table 1 Surface areas, crystal phases, and catalytic activities of Srz.xAxMnAlzlO19-a (A = La, Ce, Pr, Nd, Sin, and Gd). 0

0.2

0.4

0.6

1.0

-

-

H, P

0.98

19.7 b)

La

H c)

520 d) 755 e) 13.4

10.7

-

-

0.93

20.9 H,P 20.8 520 780 H,A, SA2. 17.8 555 780 H, P 520 780

14.9 H,P 520 770

17.7 H,P,A 530 770

12.5 H,P,A 520 810

0.91

18.6 H, P 500 770

16.9 H, P,A 520 780

11.1 H, P,A 520 770

0.91

Sm

17.6 H,P 520 770

16.8 H,P 520 780

13.7 H,P,A 510 780

6.5 H,P,A 570 870

0.89

Gd

17.2 H, P 530 820

15.2 H, P 540 780

13.9 H, P,A 520 780

5.3 Ia, A - -

0.87

Nd

H, F

485 760

H, F,A

Ce

9

17.5

rMa)

a) Ratio of ionic radii between Sr and A in Srl.xAxMnAI11019. b) Surface area (rn2 g-l) after calcination at 1200~ c) Crystal phase; H= hexaaluminate, P= perovskite, A= a-alumina, F= fluorite, SA2= SrAI407. d,e) Temperatures at which conversions of methane are 10% and 90%, respectively. Reaction condition; CH4, 1 vol%; air, 99 vol%; spece velocity, 48 000 h -1.

crystal structure in detail [8]. The thickness of mirror plane in the hexaaluminate is determined by the ionic radius, valence number, and number of elements in the mirror plane. The thickness of spinel blocks is also affected by the concentration of Frenkel defect due to excess charge in the mirror plane. In the previous study, we found t h a t Sr0.sLa0.2MnAlllO19-a calcined at 1300~ possesses the prominent catalytic activity and thermal stability as a result of partial substitution [6]. Therefore, the cationic substitution by lanthanide ions was further investigated for SrMnAlllO19-based h e x a a l u m i n a t e s for enhancing catalytic activity and/or thermal stability. Crystal phases, surface areas, and catalytic activities of Srl.xAxMnAlllO19~ (A = La, Ce, Pr, Nd, Sin, and Gd) catalysts are summarized in Table 1. La, Pr, and Nd are excellent dopants for enhancing catalytic activity and thermal stability as compared with SrMnAl11019-a catalyst. In Srl-xCexMnA111019~, CeO2 was deposited as impurity crystalline phase at every Ce concentration, since Ce is stable at the quad-valent state. Although, the ions with smaller ionic radius, e.g., Gd, lead to the formation of GdA103 and a-A1203, hexaaluminate structure was obtained at the range of 0<x_~0.6. These i m p u r i t y phases deteriorate the .,~ ; ~ 9 oi ~ o catalytic activity due to reduction of suro i eo foe 9 face areas. To elucidate the effect of l a n t h a n i d e ions on the catalytic properties of hexaI _ Ill II o ,t o aluminate, Srl-xPrxMnAl11Ols~ system was chosen as optimum substituent due 20 30 40 to its ionic radius and valence of Pr. The 2e / deg crystal structure of Srl.xPrxMnAlnO19-~ calcined at 1400~ are shown in Figure 3. T h e s a m p l e s consisted of the two Fig. 3 X-ray diffraction patterns of Srl.xPrxMnAlllO19-a calcined at 1400~ hexaaluminate phases at 0 E

0

,~

o

L_

; ;

+

.

0

80

+'+ #+

-

m 60

=

..4

--L

v

~O

o~

,~. ~..

o~'80

100

+ o

O e-eE

w

,

1 012

.......

1

~m O

conversion ~. rate I

013

,

,,jl,,,I

1 0

I

TM

3

t,,,,r~

10

9

1 015

Dark adsorption ( m o l e c u l e c m 2

) cat

Figure 2a: Maximum single pass conversion and initial rate in the absence of TCE for each compound vs. dark adsorption. (+ conversion, o, * rates; filled circles i n d i c a t e compounds for which conversion is less than 95%).

Figure 2b" M a x i m u m single pass conversion and initial rate in the presence of TCE for each compound vs. dark adsorption. (Same symbols as (2a)).

440

3.3. Rate constant

trends

(Rideal-Eley)

Figure 3a presents the maximtlm conversion and the initial rate of each compound versus the literature second order rate constant for contaminant reaction with hydroxyl radicals. The conversion increases with increasing values of koH, but with considerable scatter. For chlorine enhanced conversion and initial rate vs. the second order chlorine rate constant (presence of TCE) (Figure 3b), the trends are clear: conversion and initial rate increase with increasing kcl values. For values ofkcl above 2e-l l cm 3 molecule -t s -1, the enhanced conversion for all corresponding compounds is greater than 80% per pass. 100

~80

-

§ §

v

-

60

..,,.

+

,--80

..~,,.

-

§

++~:.~

100

4 " 1 0 1 2

o. 4 0

+ 9

0 toO_

+

9

0

0

0 0

-

20

9

%

I1)

c

+

o

9

--z

o

+

3

0

-

0

o0. .

:t - ' 2 0

0

O 0

"1o 0

o

0

e-

, , J,..l

s i .J,,,J

, , ,,-J

OH

(cm 3 molecule" 1 s 1)

Figure 3a: Maximum conversion in a single pass reactor and the initial rate (no TCE) vs. literature second order h y d r o x y l r a t e c o n s t a n t s . (Same symbols as (2a)). 3.4. (Rate constant Hinshelwood)

* specific

_

m_.,.

-

..,

3o

t101~

.~=

. -

+ o

conversion ~,=

i rate= ,2-J 10 g

.......

(D c~ 0

~

,

3

' ~ .L

1 0 "13 1 0 "121 0 "111 0 "1~

10 "14 10 "1310 "1210 "11 k

5" .,.,.

.,...,,

L ,,,d . . . . . . . d . . . . . . .

4 " 1 0 "9

Q-

:"

:

.

_+

r

+

-

W

, , ,,,t.[

%

+

0 IID

0

- 1 0 1 2

~40

3

conversion - - - e - - - rate

W ::~

o60

0

0

~,.

+"

.

~. o

0

1 013 m -

k

ci

(cm 3

molecule

1 s " 1)

F i g u r e 3b: Maximum enhanced conversion in a single pass and the enhanced initial rate (TCE added) vs. literature second order chlorine rate constants. (Same symbols as (2a)). dark

adsorption)

trends

(Langmuir-

The conversion and initial rate in the presence and absence of TCE versus the product of the second order rate constant and the dark adsorption appear in Figure 4a,b. Figure 4a shows considerable scatter in the data, revealing only broad, general trends between conversion or rate and hydroxyl second order rate constant. However, the plots of enhanced conversion and initial rate vs. the corresponding chlorine second order rate constant multiplied by the dark adsorption data are smoother (Figure 4b).

441

,- : ~ - 4 -

100 §

o~" 8 0

+

v

+

.

.=.

+ + "

c'-

6O

+

W

o60

o~_

20

:: +

+

~1013

m

"

W

. .

..~

+

0

c40

~0

9~ ~ _ o

+

0

+

@-

?o

o_'

1011

.

9

~"

3

.

>e - 2 0

0

o

.

e,-

>

0

"~

0

~-

+

101~

_

o

0

0

+

. _

o

conversion =

* Dark

"1

101

adsorption

+

e-e'-

3

9

,r

(cm 3 cm 2

cat

s "1)

Figure 4a: M a x i m u m conversion in a single pass and the initial rate in the absence of TCE vs. the product of second order hydroxyl rate constants time the dark adsorption. (Same symbols as (2a)).

k

CI

* dark

=

rate

"

~i,=.i ~,,,"J .... ,,I .... ,,,I , ,,,,,I , ,,'~

1 O" 1

103

"

conversion

o

w

rate

...... ~ ..... .1 ,, ....J ..... .! .... .,I ,,, 4 " 1 0 10

OH

o~ ,

~ 0

e-

k

~,-

_ 1012 _.~

0

++

~

0

o~ L_

~80

3

0

o

_., _,~.. W

O)

+

c, 4 0

+~++

100

1012

1 01

adsorption

1

0 3

1

109 05

-

(cm 3 cm 2

~. v

cat

s

-1)

F i g u r e 4b" Enhanced maximum conversion in a single pass and the enhanced initial rate (TCE added) vs. the product of second order chlorine rate constant time the dark adsorption. (Same symbols as (2a)).

4. D I S C U S S I O N 4.1 In t h e a b s e n c e o f (CI) TCE In the absence of TCE and chlorine, the possible active species are holes (h§ anion vacancies, or anions ( 0 2 ) , and hydroxyl radicals (OH 9 At constant illumination and oxygen concentration, we may expect h § and O2 concentrations to be approximately constant, and the dark adsorption to be a dominant variable. Ifkh+, or ko2- does not vary appreciably with the c o n ~ m i n a n t structure, the rate would depend clearly on the contaminant coverage as shown in Figure 2a, and the reaction would therefore occur via Langmuir-Hinshelwood m e c h a n i s m . (Note: only rates with conversions below 95% are correlated here (filled circles), as the 100% conversion data contains no kinetic information). This rate vs. r LH plot is smoother than those for koH or koH @T, suggesting that non-OH species (holes, anion vacancies, or O2) are the active species reacting with an adsorbed c o n ~ m i n a n t . 4.2 In t h e p r e s e n c e o f (CI) TCE The rate of gas phase reaction of pollutants with chlorine radicals is 10 to 100 times faster than with hydroxyl radicals, and we [1, 2] have proposed C1 9 as the surface active species responsible for the rate enhancement observed on addition

442 of TCE, PCE or TCP. The rate vs. kcl or kc1 (]~Tcorrelations developed in Figures 3b and 4b, respectively verify this involvement, but a choice between the LH and RE mechanism is not clearly evident from the available data. 5. C O N C L U S I O N 9 We can correlate our experimental conversions and rates with the extent of (dark) cont3minant adsorption (~), and the literature homogeneous second order rate constants for OH 9 and C1 9 radicals, and the product of the rate constant times coverage. 9 In the absence of chlorine atoms, our results suggest that hydroxyl radical is not the reactive site; other alternatives not tested here would be a hole (h§ an oxygen vacancy, or a dioxygen anion ( 0 2 ) as recommended by Fox [9] and Yates [4] respectively. 9 In the presence of chlorine atoms, the chlorine radical appears to be the active surface species. It is not possible from our limited data to establish whether most reaction occur via Langmuir-Hinshelwood or Rideal-Eley mechanisms. 9 However, none of the correlations a t t e m p t e d display r a t e v a r i a t i o n proportional to reactant coverage to the first power. F u r t h e r work under conditions of lower conversion per pass will allow inclusion in these correlations of all compounds examined here and should assist in resolving current mechanistic uncertainties. REFERENCES

1. Y. Luo and D.F. Ollis, J. Catal., (1995) Submitted. 2. M.L. Sauer, M.A. Hale, and D.F. Ollis, J. Photochem. Photobiol. A: Chem., 88 (1995) 169. 3. O. d'Hennezel and D.F. Ollis, J. Cat~., (1996) Submitted. 4. G. Lu, A. Linsebigler, and J. J. T. Yates, J. Phys. Chem., 99 (1995) 7626. 5. M.R. Nimlos, W.A. Jacoby, D.M. Blake, and T.A. Milne, Environ. Sci. Technol., 27 (1993) 732. 6. E. Berman and J. Dong, in "The Third International Symposium Chemical Oxidation: Technology for the Nineties" (W.W. Eckenfelder, A.R. Bowers, and J.A. Roth, Eds.), p. 183. Technomic Publishers, Chicago, 1993. 7. R. Miller and R. Fox, in "Photocatalytic Purification and Treatment of Water and Air" (D.F. Ollis, and H. A1-Ekabi, Eds.), p. 573. Elsevier, Amsterdam, London, New York, Tokyo, 1993. 8. M.G. White, "Heterogeneous Catalysis" (N.R. Amundson, Ed.), p. 202, Prentice Hall, Englewood Cliffs, 1990. 9. R.B. Draper and M.A. Fox, Langmuir, 6 (1990) 1396. 10. M. Formenti, F. Juillet, P. Merideau, and S.J. Teichner, Chem Tech 1, Nov ( 1971) 680. 11. S.M. Aschmann and R. Atkinson, Int. J. Chem. Kinet., 27 (1995) 613. 12. L. Nelson, O. Rattigan, R. Neavyn, and H. Sidebottom, Int. J. Chem. Kinet., 22 (1990) 1111. 13. R. Atkinson and S.M. Aschmann, Int. J. Chem. Kinet., 19 (1987) 1097. 14. T.J. Wallington, L.M. Skewes, and W.O. Siegl, J. Phys. Chem., 93 (1989) 3649.

j.w. Hightower, W.N. Delgass, E. Iglesia and A.T. Bell (Eds.) 11th International Congress on Catalysis - 40th Anniversary

Studies in Surface Science and Catalysis, Vol. 101 9 1996 Elsevier Science B.V. All rights reserved.

443

Partial o x i d a t i o n o f m e t h a n e to synthesis gas o v e r R u / T i O 2 catalysts

Y. Boucouvalas, Z.L. Zhang, A. M. Efstathiou and X. E. Verykios Department of Chemical Engineering

and Institute of Chemical Engineering and High

Temperature Processes, University of Patras, GR-26500 Patras, GREECE The catalytic partial oxidation of methane to synthesis gas is investigated over Group VIII metal catalysts. It is shown that while all catalysts promote methane combustion followed by reforming with H 2 0 and CO 2, the Ru/TiO 2 catalyst, to a large extent, promotes the direct formation of synthesis gas. The existence of the direct reaction route is probed by steady-state isotopic transient experiments. The extent of the direct route is found to be sensitive to modifications of the TiO 2 carrier. FTIR and XANES studies indicate that the unique performance of the Ru/TiO 2 catalyst is related to its high resistance to oxidation, which renders high selectivity to synthesis gas in the presence of oxygen. 1. INTRODUCTION The catalytic partial oxidation of methane for the production of synthesis gas is an interesting alternative to steam reforming which is currently practiced in industry [1]. Significant research efforts have been exerted worldwide in recent years to develop a viable process based on the partial oxidation route [2-9]. This process would offer many advantages over steam reforming, namely: (a) the formation of a suitable H2/CO ratio for use in the Fischer-Tropsch synthesis network, (b) the requirement of less energy input due to its exothermic nature, (c) high activity and selectivity for synthesis gas formation. Concerning the reaction pathway, two routes have been proposed: the sequence of total oxidation of methane, followed by reforming of the unconverted methane with CO 2 and H20 (designated as indirect scheme), and the direct partial oxidation of methane to synthesis gas without the experience of CO 2 and H20 as reaction intermediates. The results obtained by Schmidt and his co-workers [4, 5] indicate that the direct reaction scheme may be followed in a monolith reactor when an extremely short contact time is employed at temperatures in the neighborhood of 1000~ However, the majority of previous studies over numerous types of catalysts show that the partial oxidation of methane follows the indirect reaction scheme, which is supported by the observation that a sharp temperature spike occurs near the entrance of the catalyst bed, and that essentially zero CO and H 2 selectivity is obtained at low methane conversions ( 90%) in the temperature range 250-300 ~ where conversion is up to 30 %. In these conditions the main byproduct is propene, produced by propanol dehydration. Above 300 ~ (the temperature at which acetone, when feeded, strats to be oxidized) the selectivity to acetone from 2-propanol falls to below 10 % and is zero at 350 ~ when conversion approaches 100 %. Thus MgCr204 can give rise to partial oxidation of organic compounds like alcohols at relatively low temperature.

486 However, in the same temperature range and 0 2 partial pressure total oxidation of acrolein and propene largely predominates. This can be taken as a further support that on transition metal oxide catalysts the same oxygen species (lattice oxygen) are involved in both partial and total oxidation. It seems valuable to remark the similarity of the behavior of the chromite catalysts, whose surface is covered by chromate species in oxidizing conditions, with that of chromate ions and chromic acid in solution, that are choice oxidizing agents for the very selective conversion of alcohols into carbonylic compounds [19]. The same reagents under more drastic conditions cause C-C bond oxidative cleavage and the deep oxidative destruction of organics. This allows to count the methyls of organic compounds because they, after oxidative chain clevage, give rise to acetic acid [20]. This corresponds to the formation of acetates that are detected as the most stable adsorbed species on chromite catalysts after alkane oxidation [10-12].

100

'

100

90

90

~9 70 60 =- 50 .~ 40

~. 30 10 0 1~

250

3~ T

~ ~

450

5~

150

250

350

450

550

T ~ ~

Fig. 2. Left: catalytic oxidation of C3 organic compounds over MgCr204. Conversion of [] propane; A acetone; X acrolein; 9 propene. Right: catalytic oxidation of 2-propanol over MgCr204. 9 conversion of 2-propanol; selectivities to II acetone; A propene; X COx. 3.3. Surface oxidation pathways: the activation at the C-H bonds and the nature of "surface alkoxide species". The interaction of alkanes with the oxidized surfaces of combustion catalysts gives rise to oxygenate species. For the less reactive hydrocarbons (i.e. methane and ethane) the reaction is detected only at rather high temperature, and carboxylate species (acetates from ethane) or carbonate species are found as the adsorbed products. With more reactive hydrocarbons like propane and butane, carbonyl compounds (acetone and methyl-ethyl-ketone) have been observed at lower temperature on MgCrEO,~x. If the reactant is even more active, features typical of alkoxide species have been observed. This is shown in Fig. 3 where the spectrum of the adsorbed species arising from isobutane interaction with MgCrEO,~x at 423 K are compared with those obeserved after interaction of tert-butanol over the same catalyst in the same conditions. The bands observed in the region 1300-1000 cm ~ are typical of the C-O / CC stretchings of tert-butoxide species. The detection of these species parallels the detection of allyl-alkoxides from propene, over the same catalyst [10,11]. The rate of the oxidative cleavage of C-H bonds over this catalyst nearly follows the order: benzyl --- allyl > methine >

487 methylene > methyl > methane, and has been related either to the dissociation energy of the C-H bonds (inversely) [11] or to the electrostatic potential charge (the more negative the charge, the higher the cleavage rate) on hydrogen [21]. In any case, this trend is the same previously observed for the combustion rate on copper oxide catalysts [2,22], for which the first C-H cleavage step is thought to be rate-determining. Again we remark that the activity of chromite catalysts (whose surface is covered by chromates) closely corresponds to the activity of chromates in acid solution, that are able to oxidize tertiary alkanes to the corresponding tertialy alcohols and secondary alkanes to the corresponding ketones [20]. Interestingly, the oxidation of chiral tertiary alkanes by chromic acid in solution gives rise to the corresponding alcohols stereospecifically, with retention of the configuration [23]. ~.~-. J

i , .o.,o~.,.

!

.....

~ ....

i

(

..... i

- i v

a) ~

!"X,~.,.,/

;;

"'../"

/i

,

i"

i

!\

i,, .oj2s-. X i

.o,~o~

.~r-"

0~."

i k;

.0. 090 ".

'

~,oi

[',,

o.,. i

\\

~

ol

t",

i

,

"

',

'

!

i ..i

\

-I

/

/I

t.cs

/

~, ', ',\ ","',

;"

/

)

',,,

',\

,,

y,\

/

\

9 ''!'

k.,'

w~'e~e~

(d>

A,-

i ../ ",~ \

',~1

,/

,

,~ "","~ "","~ "' .;,i; "' ;'~" T~"; ,',,;"; ;~; "; ,',,;""; i;o""; ;~" ";~ ";;,;;' ";~;' "";~'" " ............

//%

/ I

', " ',\

IX// I

....~ I

o,i/

'v-"/~

~,

I! ~i

I!

r,.

~

"..

i"

i

,,

..

o.q tc)

s

Co) ',

",.

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

02~ ....

I /-f'

ii

, ..... ~.... -,'r .... ,v,-,',',

I~iOo 121~) I ~

12~

..... c---,-, .....

I ~p'~Io 12(]o 1 1 ~ .

.

.

.

.

.

.

.

.

.

.

.

.

11~

v,'---I .... 11~

wml.~.:!l

r,-,,- v,-,-- v,',-,-r

11~i~) 11oo .

.

.

.

.

.

.

.

.

.

.

.

.

i~ .

.

.

.

.

.

.

I~ .

.

.

.

.

.

.

.

.

..... , .....

1o4o .

.

.

.

.

.

.

.

.

.

.

io~c~ .

.

Fig. 3. FT-IR spectra of the adsorbed species arising from the interaction of (a) tert-butanol and (b) isobutane over a combustion catalyst (MgCr204) at 423 K, and from tert-butanol (373 K, c), isobutene (300 K, d) and isobutane (380 K, e) on a selective oxidation catalyst. In Fig. 3, on the left, the spectra of the adsorbed species arising from isobutane interaction with a molybdena-based catalyst that behaves quite selectively for the production of methacrylic acid from isobutane are compared with those arise from isobutene and tertbutanol adsorption on the same surface. In all three cases the typical features of tert-butoxy species are observed, that are stable up to near 400 K. According to our data, isobutane is activated by C-H bond oxidative cleavage giving rise to tert-butoxide that can decompose to isobutene and an OH group. Isobutene can later undergo aUylic oxidation on selective catalysts, while it can undergo oxidative C-C bond breaking to give one acetate and two formate species on the combustion catalysts. These data show that the C-H activation mode is similar on both partial and total oxidation catalysts. Starting from n-butane, 2-butoxides that rapidly convert to 2-butanone are found over MgCr204 [24]. However, the further oxidation of adsorbed 2-butanone only gives rise to the acetate species, while starting from n-butane, formate species are also observed. This can be explained assuming that sec-butoxides can partly isomerize to tert-butoxides before further oxidation. This implies that the C-O bond formed is partly ionic and the alkyl moiety has the

488 character of a carbenium ion. Similarly, the allyloxy species we observed to be formed by oxidation of propene can have the pronounced nature of a an ally1 cation, stabilized by the 1,3 charge delocalization. It is well known that symmetric aUyl species are involved in propene selective oxidation and ammoxidation to acrolein and acrylonitrile, and that no C-O bonds is formed irreversibly between the first and the second hydrogen abstraction in acrolein synthesis [ 1-3,25]. The picture we propose supports the idea that the adsorbed allylic species involved in propene oxidation is cationic in nature. A further support to the cationic nature of allyl species observed on the surface of chromite catalysts is given again by the comparison with the behavior of chromates in solution, that also give rise to aUylic oxidation and where transpositions of intermediate aUylic carbenium ions to more stable tertiary carbocations have been reported [20]. As for selective catalysts for acrolein production, the allyl species intermediates have been proposed to be radical-like, although allyloxy- species are assumed to be formed later [3,25]. In effect desorption of gas-phase allyl radicals has been observed from BiEMoO6 (a very selective catalyst) and, mostly, from Bi203, that is not a selective catalyst for acrolien production [26]. Over BiE(MoO4)3 allyl radicals were not formed while MoO3 (that allows acrolein production) acts a a sink for allyl radicals [26]. It seems quite reasonable to think that the allyl radicals formed on bismuth centers are intermediates favoring the formation of allyl carbeniurn ions / allyloxy- species on molybdenum. This picture also agrees with the detection of "isolated" benzyl species (with no C-O bonds) upon toluene and o-xylene selective oxidation over vanadia-titania selective oxidation catalysts [27], according to the very strong delocalization of the cationic charge over the aromatic ring. Due to this stability, the rate detrmining step in toluene oxidation is apparently shifted to the successive step, the reaction of benzyl species with surface oxide to give benzaldehyde [27]. In conclusion, the carbocationic character of the intermediate and its stability should decrease and the alkoxide character of the same species should increase following the sequence benzyl, allyl, saturated tertiary, secondary, primary.

3.4. Surface oxidation pathways: C3 hydrocarbons. IR studies have shown that propane and propene oxidation over a chromite catalyst follow two different paths constituted by partial oxidation steps that give rise to adsorbed carboxylate species; these species burn totally before leaving the surface [ 10]. The spectra of the adsorbed species arising from propene oxidation at r.t. over oxidized Co304, an even more active catalyst [ 12], are very similar to those arise from aUyl alcohol, acrolein and acrylic acid (see Fig. 1). They show that on Co304 the allylic methyl group of propene is attacked and oxidized to give the carboxylate ion acrylate. Also on MgCr204 acrylate species are formed from propene, although at higher temperature according to the weaker activity of this catalyst. Propane oxidation gives mainly rise to adsorbed acetone at low temperature on MgCr204, while on Coat4 propane gives rise to several different species, i.e. acetate, acrylate and propanoate species. On both surfaces, at the temperature at which the catalytic hydrocarbon conversion is sufficiently fast, the surface species disappear or are substituted by carbonate species, and CO2 gas is observed. In the following scheme, an oxidation pathway for propane and propene is proposed. This mechanism, that could be generalized to different transition metal oxide catalysts, implies that propene oxidation can follow the allylic oxidation way, or alternatively, the oxidation way at C2, through acetone. The latter easily gives rise to combustion, because it can give rise to enolization and C-C bond oxidative breaking. This is believed to be the main combustion way for propene over some catalysts, while for other catalysts acrolein overoxidation could

489 predominate. Consequently, propene combustion can be either successive or competitive to propene partial allylic oxidation. On the other hand, starting from propane the alkoxide intermediate can either decompose to propene or be oxidized to acetone and overoxidized later. So, the alkoxide evolution step can play the role of a "selectivity determining step" in the oxy-dehydrogenation of propane to propene, using an expression proposed by Kung and coworkers [28]. CH3CH2CH 3

CH3CH=CH2

~'4&

C~x~ /CH3 S

OH

c.

I

I

,c-2

.~~

"-

O

o 1

acetone

CH /+\ CH 2 CH 2

CH2=CH

C~x~ /CH3

CH31

C II O

oS C...~CH2

-..

f

O"

I

I

CH2=CH

~

~

CH ii 0

acrolein

,

I

acetic acid

,~

CH 3 C, O'O I

~

~

C Ox

.~~

CH2=CH ,C, OO ~

acrilic acid

Scheme I. A generalized pathway for C3 organics partial and total oxidation. 3.$. Surface oxidation pathways: C4 linear hydrocarbons. Similar experiments allowed us to propose a reaction pathway for n-butane oxidative conversion apparently common to selective oxidation catalysts like (VOhP2OT, that allows the production of maleic anhydride from butane; VEOs-TiO2 and MoO3-TiO2, that allow the production of acetic acid from butane and of maleic anhydride from butene; Mg3(VO4h that allows the oxy-dehydrogenation of butane to butene and butadiene; FeCrO3 that allows the oxy-dehydrogenation of butene to butadiene; and combustion catalysts [29]. The reaction network is substantially the same in all cases, but the different behavior of the catalysts is explained by the different rates of some alternative steps on the different surfaces. Again, a key step concerns the dehydration or the oxy-dehydrogenation of the 2-butoxide intermediate, playing the role of "selectivity determining step", where the oxy-dehydrogenation route diverges from the main combustion route. The low activity in oxidation and the significant basicity with absence of BrOnsted acidity typical of Mg vanadate are likely main factors favoring the production of butene from the 2-butoxy- species. On the contrary, the weak BrCnsted acidity of oxidized Mg-chromite is likely responsible for blocking of the 2-butoxide

490 species at the surface whilethe high oxidizing power of Cr 6+ causes its rapid oxydehydrogenation to the ketone and the following C-C bond breaking. The same mechanism proposed for the combustion catalyst Mg-chromite apply also to catalysts that allow significant yields in acetic acid from n-butane, like vanadia-titania, that accordingly also show a medium-high BrCnsted acidity. Being acetate ions intermediates in the combustion way, it is easily rationalized that the production of acetic acid is favored by the addition of steam in the reactant mixture and by adjusting the reaction conditions. The catalysts that allow the production of maleic anhydride from n-butane with high selectivity, like (VO)2P2OT, are characterized by a strong acidity, that, like a strong basicity, favors the decomposition of alkoxides to give the olefin and the diene. The catalysts that allow the production of maleic anhydride, either from n-butane or from butenes and butadiene, necessarily have particular sites that allow the insertion of oxygen atoms in the 1,4-position of butadiene. These sites are definitely absent on combustion catalysts. . . . . . . . . . . . .

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

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

o.o8

//

0,06

'\

!,'

0,04

\

..

,',

o;

0.02 0.00 = - 0 . 0 2 -~

/

-0.04

.~ ~""~"~--~ .... - 0 . 1 0 ~-"~"~-'~.N'~J'~...'-~-J'V,,'~

/

-o.,.~ ..... "---"'. . ..,,,__:

-

,

/

...

.I

/ .," _ ._/I ," ,]

- o . o 6 -Z

-0.14

i

-

~

.....

s "-4

x..

~

;;

\

:x.x

_ i ,, J

-0.18

\

! ' , ..

\d

,,_._; ...

"-..... <J I

-0.20

/,,,.,,'-\ ....

-..\ ..j' .,.,/,,,i

',,I::!

.....',., (c) ""\

f~ '

(d)

\\

\ " ',.

[~

/

/ ""

:' 'i ,,

, #

/\ .,h,,,-"~..\

,,

'""

1

t"')""'~..., ~'-'''''F " ----v-,, /-,,,~,--,,-~..

'

t/ """

"'f

'

I

-0.~ ~

. . . . ~, .. , ~ , ,

-0.24 ~ 2000

/\ :,,,"

1950

1900

1850

1800

1750

1700

1650

1600

1550

i

1500

Wavenumberl

1450 1 4 0 0 (ore-1)

1350

1300

1250

1200

~,7

f

,-.'' 1150

\1

1100

10.50

Fig. 4. Ft-IR spectra of the adsorbed species arising from interaction of 1-butene (a,b) and butadiene (c,d) on MgCr2Oa+x at r.t. (a,c) and at 373 K (b,d). Accordingly, the interaction of butadiene with (VO)2P207, V2Os-TiO2, and MoO3-TiO2 is strong and reactive at r.t. (giving rise to furan-like species) and, at higher temperatures, gives rise to the typical bands of cyclic anhydrides in the region 1900-1700 cm l [29,30]. This reactivity is lacking on some very reactive catalysts like MgCr204+x where both butene and butadiene give rise to molecular adsorbed species at r.t. (Fig. 4, a and c), while at slightly higher temperatures methyl vinyl ketone is formed from butenes (typical doublet at 1670, 1640 cm q, vC=O and vC=C, and band at 1180 cm l, Fig. 4,b) and a mixture of carbonyl compounds and carboxylates from butadiene (Fig. 4,d). As a conclusion of our experiments the following generalized pathway is proposed for n-C4 oxidations:

491

CH3COOH

O

CO, CO 2

O

HOOC

o,

0

,o

O

>

o

o--~~-o 0

OO

-I~

~~I~OOH

o---~~-o 0

Scheme II. Reaction pathways for butane oxidation. 4. CONCLUSIONS The results of the experiments we carded out and the comparison with the data arising from homogeneous oxidation chemistry allow us to propose and generalize the following conclusions that in part contrast previous literature findings: i) activation of hydrocarbons occurs over transition metal oxides by abstraction of two electrons by the oxidized cationic center. ii) nucleophilic oxygen species (lattice oxygen) contribute to the C-H activation step by interacting with the resulting carbenium ion and proton species, giving rise to an alkoxide and to an hydroxide, respectively. iii) the resulting alkoxide has a partial carbocationic nature, the more important the higher is the charge delocalization (i.e. for benzyl and allyl species). iv) so, the C-O bond of the alkoxide intermediate is not irreversibly formed, and its reversible rupture justifies the symmetry of the aUyl species in propene oxidation. v) the alkoxides so formed can be further oxidized to carbonylic compounds. vi) if the surface and the carbonyl compounds are both sufficiently weakly reactive, desorption can occur and partial oxidation takes place.

492 vii) if the surface is very reactive and the carbonyl compounds has hydrogen in the ~position, enolization occurs that finally gives rise to C-C bond breaking and either to combustion or to the production of acetic acid. viii) total combustion pathways should involve nucleophilic (lattice) oxygen species and a sequence of partial oxidation reactions that finally give overoxidation of the products to COx. The authors acknowledge MURST (Rome, Italy) and NATO for financial support. REFERENCES

1. .

3. 4. .

6. .

.

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

C.N. Satterfield, Heterogeneous Catalysis in Industrial Practice, 2nd ed., McGraw Hill, New York, 1991. A. Bielanski and J. Haber, Oxygen in Catalysis, Dekker, New York, 1991. V.D. Sokolovskii, Catl. Rev. Sci. Eng. 32 (1990) 1. M.F.M. Zwinkels, S.G. Jaras, P.G. Menon and T.A. Griffin, Catal. Rev. Sci. Eng. 35 (1993) 319 P. Mars and D.W. van Krevelen, Chem. Eng. Sci. 3 (special supplement) (1954) 41. J.J. Spivey, Ind. Eng. Chem. Res., 26 (1987) 2165; J.J. Spivey, in "Catalysis", vol 8, The Royal Society of Chemistry, Cambridge, 1989, pag.158. J.E. Germain and L. Laugier, Bull. Soc. Chim. France, (1972) 541 and 2910;. J.E. Germain and R. Perez, Bull. Soc. Chim. France, (1972) 2042 and 4683. E. Garbowski, M. Guenin, M.-C. Marion and M. Primet, Appl. Catal. 1990, 64, 209 L. Ya. Margolis, Advan. Catal. Relat. Subj. 14 (1963) 493 E. Finocchio, G. Busca, V. Lorenzelli, R.J. Willey, J. Chem. Soc. Faraday trans., 90 (1994) 3347. E. Finocchio, G. Busca, V. Lorenzelli and R.J. Willey, J. Catal. 151 (1995) 204. E. Finocchio, G. Busca and V. Lorenzelli, J. Chem. Soc. Faraday trans., submitted G. Busca, M. Daturi, E. Kotur, G. Oliveri and R.J. Willey, in "Preparation of Catalysts VI", G. Poncelet et al. eds., Elsevier, Amsterdam, 1995, p. 667 K. Hadjiivanov and G. Busca, Langmuir, 10 (1994) 4534 G. Busca, R. Guidetti and V. Lorenzelli, J. Chem. Soc. Faraday Trans. 86 (1990) 989. G. Tulyev and S. Angelov, Appl. Surface Sci. 32 (1988) 381; B. Marcus-Saubat, J.P. Beaufils and Y. Barbaux, J. Chim. Phys. 83 (1986) 317. G. Busca, G. Ramis and V. Lorenzelli, J. Mol. Catal. 50 (1989) 231 G. Busca, Catalysis today, in press A.H. Haines, Methods for the Oxidation of Organic Compounds, Academic Press, 1988. G.A. Olah and A. Molnar, Hydrocarbon Chemistry, Wiley, New York, 1995. G. Busca, E. Finocchio, G. Ricchiardi and G. Ramis, submitted paper G.K. Boreskov, in Catalysis Science and Technology, J.R. Anderson and M. Boudart eds., Vol. 3, Springer Verlag, New York, 1982, p. 39. K.B. Wiberg and G. Foster, J. Am. Chem. Soc. 83 (1961) 423. E. Finocchio, G. Busca, V. LorenzeUi and R.J. Willey, submitted paper T.P. Snyder and C.G. Hill, Jr., Catal. Rev. Sci. Eng. 31 (1989)43 D.J. Driscoll, K.D. Campbell and J.H. Lunsford, Advan. Catal. 35 (1987) 139. G. Busca, J. Chem. Soc. Faraday trans. , 89 (1993) 753; G. Busca, in "Catalytic Selective Oxidation", S.T. Oyama and J.Hightower eds., ACS, Washington, 1993,p. 168 H.C. Kung and M. Kung, Advan. Catal. Relat. Subj. 40 (1994) G. Busca, G. Ramis, V. Lorenzelli and G. Oliveri, in "New Developments in Selective Oxidation II", V. Cortes Corberan and V. Bellon ed., Elsevier, 1994, p. 253, G. Ramis and G. Busca and V. Lorenzelli, J.Mol. Catal. 55 (1989)

J.W. Hightower, W.N. Delgass, E. Iglesia and A.T. Bell (Eds.) 11th International Congress on Catalysis - 40th Anniversary

Studies in Surface Science and Catalysis, Vol. 101 9 1996 Elsevier Science B.V. All rights reserved.

493

Biomimetic oxidation on Fe complexes in zeolites Gennady I. Panov, Vladimir I. Sobolev, Konstantin A. Dubkov and Alexander S. Kharitonov Boreskov Institute of Catalysis, Novosibirsk, 630090 Russia

One-step hydroxylation of aromatic nucleus with nitrous oxide (N20) is among recently discovered organic reactions. A high efficiency of FeZSM-5 zeolites in this reaction relates to a pronounced biomimetic-type activity of iron complexes stabilized in ZSM-5 matrix. N20 decomposition on these complexes produces particular atomic oxygen form (a-oxygen), whose chemistry is similar to that performed by the active oxygen of enzyme monooxygenases. Room temperature oxidation reactions of a-oxygen as well as the data on the kinetic isotope effect and Moessbauer spectroscopy show FeZSM-5 zeolite to be a successful biomimetic model.

1. INTRODUCTION A perfect operation of enzyme monooxygenases provides a wide variety of fine oxidation reactions occurring in the living nature at ambient conditions. This remarkable phenomenon relates to a unique ability of monooxygenases (MO) to activate dioxygen and to produce very reactive oxygen species. These species coordinated on the iron active sites can selectively add to organic molecules to yield oxygenated products [ 1]. Man-made catalysts are not so perfect as enzymes. One often needs to perform several stages under harsh conditions to accomplish a simple chemical reaction. As an example we may mention two oxidation reactions, which will be discussed further, i.e. benzene oxidation to phenol via the three stage cumin process, and methane oxidation to methanol via the intermediate CO and 1-12production. MOs elegantly do these reactions by a single move just adding one oxygen atom to a starting molecule. Researchers take strong efforts to understand the mechanism of MO oxidation and to apply biomimetic strategy in the practice. Progress in this field is strongly related to the development of successful models capable to mimic MO activity. In this paper we for the first time consider biomimetic aspects of a recent approach in the catalytic oxidation dealing with the iron-containing zeolites as catalysts and nitrous oxide (N20) as a source of oxygen supply [2]. The paper will refer rather frequently to the previous results obtained in this laboratory, since during several years our research group was being intensively involved in the N20 oxidation study. Further we shall show that analyzing the published data and new results concerning the room temperature methane and ethane oxidation we have found Fe complexes in ZSM-5 matrix to be a promising biomimetic model.

Acknowledgment. Research described in this paper was made possible in part by Grant No JDB 100 from International Science Foundation and Russian Government.

494 2. AROMATICS

HYDROXYLATION WITH NITROUS OXIDE

2.1. Discovering zeolite catalysts Partial oxidation reactions are usually carried out over transition metal oxides capable of changing their valent state during their interaction with reacting molecules. Naturally, zeolites with their alumina-silicate composition did not prove themselves as good oxidation catalysts. They failed also to serve as efficient catalyst supporters, since transition metals being introduced into the zeolite matrix lose their ability to activate dioxygen [3,4]. In the 1980's zeolites attracted a renewed attention. They were shown to be rather promising catalysts if, instead of 02, a chemically pre-modified oxygen entering the oxygen-containing molecules is used. The most known example is an excellent catalytic performance of titanium silicalites in the liquid phase oxidations with H202 [5]. A gas phase oxidation with nitrous oxide is another approach in this field being intensively developed in the last years [2]. Nitrous oxide as an efficient oxygen donor was noticed when used in such a delicate reaction as the direct oxidation of benzene to phenol: C6H6 + N20

=

C6HsOH + N2

(1)

This reaction was first demonstrated over V, Mo and W oxides [6]. At 823 K vanadium oxide provided phenol selectivity up to 71%, which was much higher than it had been ever achieved with 02. This result stimulated further efforts in searching for more efficient catalytic systems. As a result, in 1988 three groups of researchers [7-9] have independently discovered ZSM-5 zeolites to be the most efficient catalysts. They allowed the reaction to proceed at much lower temperature (573-623 K) with nearly a 100% selectivity. Later, more complex aromatic compounds were also hydroxylated in this way [2]. These results being quite untypical for zeolites give rise to a number of fundamental questions: i) what makes the zeolite to function as an active catalyst; ii) what makes N20 to function as a selective oxidant; iii) what is the reaction mechanism. We shall shortly discuss the situation with these issues because of their importance for our further consideration.

2.2. Role of iron A strong zeolite acidity had been first suggested to explain the catalytic activity origin [7]. Later, it was not supported experimentally. The activity was shown to relate to Fe admixture always presented in zeolites in a small amount. The role of iron is clearly seen in Fig 1. The later presents the data obtained with two ZSM-5 zeolite systems of Fe-Si [ 10] and Fe-AI-Si [ 11] composition with a widely varied Fe concentration. In fact, the rate of benzene oxidation to phenol over the purest samples (0.003 and 0.004 wt % Fe2Oa) is very low and strongly increases with the increasing iron content. The catalytic activity of Fe proved to be very high (especially in the AI-Si matrix) so that an admixture of 0.01 wt % Fe2Oa may cause a noticeable benzene conversion. This discovery was quite unexpected, since iron oxide has been never reported as an active catalyst in either partial or full oxidation. The studies of two simplest reactions, i.e. 02 isotopic exchange and N20 decomposition, revealed a dramatic change of Fe properties in the ZSM-5 matrix compared to Fe203 [4]. Fe atoms lose their ability to activate 02 but gain remarkably in their ability to activate N20. It gives rise to a great effect of the oxidant nature in the reaction of benzene oxidation over the FeZSM-5 zeolite (Table 1). Thus, in the presence of N20 benzene conversion is 27% at 623 K, while in the presence of 02 it is only 0.3% at 773 K. And what is more, there is a perfect change of the reaction route. Instead of selective phenol formation with

495 N20 (S=98%), their are only full oxidation products with 02. Over Fe203 none of the oxidants produces phenol.

~ 30 o

623 K -

O I

30[

O

--

=

20~

o

20

~

623K

"

o 573 K

O

o 10 [ Fe-Si ] 0

I

I

I

I

2

3

4

5

0

Fe concentration ( wt %Fe203)

0.2

0.4

0.6

Fe concentration ( wt % Fe203)

Figure 1. Effect of iron concentration on FeZSM-5 catalytic activity. Beside iron, the catalytic properties of many other transition metals (V, Mo, Cr, Mn, Co, Ni, Cu, Ti, Zn, Pd, Pt) have been tested. These metals exhibited no activity in phenol production [7,11 ]. This means that Fe might be the one of few particular elements or even the only one, which can effeciently catalyze this reaction. Table 1 Effect of oxidant nature on benzene oxidation (X-benzene conversion, S-selectivity to phenol) Sample

Oxidant N20

Fe-AI-Si (0.08 wt % FEE03) Fe203

Oxidant 02

T (K)

X (%)

S (%)

T (K)

X (%)

623

27

98

773

0.3

0.0

623

24.5

0.0

623

5.5

0.0

S (%)

2.3. Particular features of nitrous oxide In order to understand the reason for such a beneficial N20 oxidizing effect, a detailed mechanism of its decomposition as a stage supplying oxygen to the surface has been studied [4,12]. This study revealed a special type of iron active sites in ZSM-5 matrix (called a-sites), which decompose N20 producing a new oxygen form (a-form): N20 + ( )or

-(O)ct

+N2

(2)

At temperatures lower than 573 K a-oxygen is termally stable and reaction (2) selectively occurs with no oxygen evolution into the gas phase (Fig 2). Note that this phenomenon is not the result of oxygen consumption for the surface reoxidation, since our experimental conditions completely exclude a reduction of the sample alter its treatment in O2 [ 12]. When all ix-sites are occupied, the reaction terminates. By measuring the amount of N2 produced (or that one of

496

12~

N20 + ( )~, --) N2 + (O)~'

90

I

N~ o

60

Figure 2. Kinetics of N20 decomposition at 523 K followed by (x-oxygen loading on FeZSM-5 zeolite surface.

30

0

4

8 12 Time (min)

16

N20 consumed) one can determine an amount of (x-oxygen loaded and the density of (x-sites. To verify these results, the reaction of isotope exchange can be additionally used to measure (xoxygen amount [ 12]"

~SO2 + (~60)~

=

~602 + (~SO)a

(3)

(x-Site density can be regulated by the iron content and procedure of FeZSM-5 activation. High density of (x-sites provides a good experimental opportunity to study properties of (xoxygen, what has been done in a number of papers [4,12-15]. A very low energy of bonding with the surface and a very high reactivity are the most remarkable features of (x-oxygen. At room temperature, it participates in the isotopic 02 exchange as well as in oxidation of various organic molecules. (x-Oxygen has been probably first observed when studying the mechanism of N20 decomposition over FeM zeolite [16], though its concentration was too low for its clear manifestation. 2.4. Reaction mechanism

After discovering (x-oxygen formation as a particular feature of N20 decomposition, an important question concerning the reaction mechanism arises: does (x-oxygen participate or not participate in the oxidation of benzene to phenol? According to a generally accepted view [ 17], the surface oxygen providing partial oxidation must not have a low energy of bonding to the surface, and must not exhibit a high reactivity, which is in a conspicuous contrast to (x-oxygen properties. Therefore, the idea to relate (x-oxygen to phenol formation needs a strong experimental support. Attempts to use for this purpose some spectroscopic techniques 0IL NMR) have not provided reliable data because of their low sensitivities. Another approach was found to be successful. If the interaction of benzene with (x-oxygen actually produces phenol, and if this phenol can be extracted from the surface, its amount will be sufficient for reliable chromatographic analysis. Based on this idea, experiments were carried out according to the following three stage scheme [18]:

497

1. N20 + ( )or 2. Cd-I6 + ( O )or 3. (Cd-I~OH)ct

520K

)

(O)tx + N2

.~

(Cd-IsOH)ot

;

Cd-IsOH

298 K 298 K

+

()~

The scheme includes a-oxygen loading to the surface (stage 1), its interaction with benzene at room temperature (stage 2), and product extraction from the surface (stage 3). Results obtained with the Fe-AI-Si sample are given in Table 2. Phenol was found to be the only reaction product, and its average amount corresponded to 90% of the a-oxygen loaded. A somewhat underestimated yield is probably related to the incomplete phenol extraction. Similar results were obtained with ZSM-5 sample of Fe-Si composition [ 18]. Table 2 Room temperature benzene oxidation with a-oxyBen Sample

FeZSM-5 (0.07 wt % Fe203)

Run

(O)a (~mol/g)

C6I-~OH

Phenol yield

(p.mol/g)

(%) 93 85 93

1

6.0

5.6

2 3

5.5 5.5 0.0

4.7 5.1 0.0

4 (blank run)

-

These experiments clearly showed that it is a-oxygen participation that provides FeZSM-5 zeolites with such a remarkable catalytic performance in the reaction of benzene to phenol oxidation. Equations (1-3) written above are the main stages of the reaction mechanism.

3. BIOMIMETIC FEATURES OF FeZSM-5 ZEOLITE 3.1. cz-Oxygen m e t h a n e oxidation Results on the room temperature oxidation of benzene to phenol are rather suggestive with respect to the monooxygenase mimicking, since the hydroxylation of aromatics is one of typical reactions catalyzed by these enzymes [19]. Understanding the mechanism of MO oxygen activation is a difficult problem and is a subject of many studies and debates. Using the complexes of iron and other metals, including those encapsulated into zeolites [20-22], some successful models have been developed to simulate the MO function of cytochrome P-450. But all attempts to mimic methane monooxygenase (MMO) have failed because of not enough powerful oxygen activation. The MMO dinuclear iron sites are capable to produce oxygen species of much superior activity compared to other MOs. Therefore, in addition to aromatics and various other compounds, MMO exhibits a unique ability to hydroxylate methane producing methanol [23]. That is why, it is of special interest to test the a-oxygen oxidizing potential with respect to methane.

498

J

120 I00

Temperature ( K )

Temperature ( K )

300 400 500 600 700 800

300 400 500 600 700 800

i

AB ~ ,

.

|

J

I

-

i

i

1

i

120 ~

,

~

CH4

100

a)

80

80

~

60

60

o

40

40

~

co

20

,

|

~_

2'0

4'0

6'0

T i m e ( m i n)

8'0

i

.

i

,

1

CH4

.

|

.

b)

CO

20 0

I~

.

AB -]

,

0

20

40

.

,

.

60

,

.

gO

T i m e (m i n)

Figure3. Temperature-programmed reaction of methane with FeZSM-5 surface before a-oxygen loading (a) and after a-oxygen loading (b). A - time moment of opening the microreactor; B - time moment of switching on the programmed heating (6 K/s). For this purpose we studied a temperature-programmed interaction of CI-L with a-oxygen. Experiments were carried out in a static setup with FeZSM-5 zeolite catalyst containing 0.80 wt % Fe203. The setup was equipped with an on-line mass-spectrometer and a microreactor which can be easily isolated from the rest part of the reaction volume. The sample pretreatment procedure was as follows. After heating in dioxygen at 823 K FeZSM-5 cooled down to 523 K. At this temperature, N20 decomposition was performed at 108 Pa to provide the a-oxygen deposition on the surface. After evacuation, the reactor was cooled down to the room temperature, and CI-I4 was fed into the reaction volume at 108 Pa. Fig 3 shows the results of two temperature-programmed experiments. In the first (blank) experiment CH4 reacts with a "bare" FeZSM-5 zeolite, while in the second one it reacts with the zeolite after a-oxygen loading on its surface. Obviously, the bare surface is quite inert towards methane (Fig 3a): after reactor opening a weak CI-I4 adsorption occurs at room temperature. A slight heating results in a complete recovery of the CI-L pressure. There is quite a different picture in the a-oxygen presence (Fig 3b). Just after reactor opening a large and irreversible CH4 consumption occurs now, which evidences its chemical reaction with a-oxygen to take place at room temperature. The reaction product is strongly bound to the surface, and can not desorb into the gas phase under heating without destruction, which is accompanied by the CO evolution. In order to identify the product, we used a procedure of its extraction from the surface similar to that used in the case of a-oxygen benzene oxidation [18]. For this purpose, a number of single-turn-over runs in the room temperature methane oxidation were carried out according to the following scheme: 1) a-Oxygen loading to FeZSM-5 surface by N20 decomposition at 523 K; 2) Its interaction with CH4 (108 Pa) at room temperature; 3) Sample unloading and product extraction by an acetonitrile-water mixture with its further NMR and chromatographic analysis. Results are shown in Table 3. In all runs methanol has been detected as a sole product and its amount within the experiment error correlates with that of CH4 consumed. It is of interest that methane reaction with a-oxygen:

499 CH4 + (O)(x

=

(CH3OH)a

(4)

occurs surprisingly fast. We failed to measure its rate not only at room temperature but even after sample cooling down to 243 K. Similar experiments on a-oxygen oxidation were also carried out with ethane resulting in a selective formation of ethanol. Table 3 Room temperature ,methan,e. oxidation with a-oxygen Run CI-h reacted CH3OH formed No (gmol/g) (pmol/g) 20 23 18 20

1 2 3 4

CH3OH yield

(%)

19 21 18 18

95 91 100 90

3.2. Kinetic isotope effect A rate-determining step of MO alkane oxidation involves the cleavage of C-H bonds, which brings about high values of kinetic isotope effects (KIE = ka/kD) [24]. We measured KIE for methane oxidation according to reaction (4) using the intramolecular competition of C-H and C-D bonds of CH2D2 molecules in their reaction with a-oxygen. Product methanol was extracted with a deuterated mixture of acetonitrile and water, and its isotope composition was analyzed by the M R spectroscopy. Unlike monooxygenases, FeZSM-5 zeolite is a robust system and may be studied under wide conditions. In a temperature range of 223 - 373 K KIE value for methane oxidation varies from 5.5 to 1.9 (Table 4). It corresponds to an "activation energy" of 5.0 kJ/mol, which is in a good agreement with the difference of the zero point energies of C-H and C-D bonds. Thus, similar to MMO with its KIE value of 5 [25], the rate determining step of CI-L oxidation by a-oxygen also involves the cleavage of C-H bond.

Table 4 NMR analysis of product methanol isotops and KIE values of CH2D2 methane oxidation with a-oxygen (8 is a chemical shift, IH and ID are signal intensities) CHD2OD

CH2DOD

Run

T

KIE = 2IH/ID

No

(K)

8 (ppm)

IH (rel.unit)

8 (ppm)

ID (rel.unit)

1 2 3

223 293 373

3.30 3.30 3.30

63 68 56

3.32 3.32 3.32

23 42 58

5.5 3.2 1.9

3.3. Active state of iron The distribution of iron and other metals in zeolites has been studied by many authors. In general, Fe can occupy three positions in ZSM-5 matrix [26]" (1) as isolated ions in the tetrahedral lattice positions; (2) as isolated ions or small complexes outside the lattice but inside

500 the intracrystalline micropore space; (3) as clusters and finely dispersed oxide particles on the outer surface of zeolite crystals. Which of these positions occupy a-sites is an important question for our consideration. The first position can be safely excluded since a high temperature calcination, causing the removal of Fe atoms from the lattice, remarkably increases the a-site concentration [27]. Besides, a-sites can be prepared via the impregnation of a ready zeolite matrix [28], when the probability for Fe atoms to incorporate into the lattice is very low. a-Sites do not occupy also the 3rd type position: deactivation of the outer zeolite surface by its coveting with an inert SiO2 layer affects neither catalytic activity no a-site concentration [29]. Thus, we may deduce that the active iron occupies the second type position in ZSM-5 matrix and is either isolated Fe ions or small complexes inside the mieropore zeolite space. According to the following evidences, a-sites are most probably di-iron complexes similar to the di-iron active sites of methane monooxygenase: a) The ratio between the number of a-sites and that of Fe atoms in the FeZSM-5 zeolites sometimes achieves but never exceeds 2. b) The active iron is invisible in the ESR spectra [12], which is consistent with the formation of dinuclear complexes. c) Below 110 K, the magnetic susceptibility of FeZSM-5 decreases with the decreasing temperature, which is typical for the antiferromagneticaUy bound iron complexes [30]. d) Experimentally observed a-oxygen features are well interpreted within a dinuclear quantum-chemical model of a-sites [31 ]. Recent results obtained with Moessbauer spectroscopy [ 13] provided additional arguments in favor of MMO similarity. According to [32] iron atoms in MMO, depending on conditions, exist m both oxidized and reduced states. Two different quadrupol doublets correspond to each of these states with main parameters given in Table 5. The Moessbauer spectra of FeZSM-5 also reveal two states of iron, which are represented by two quadrupol doublets ffig 4). The narrow doublet parameters correspond to the Fe(III) state, the broad doublet parameters correspond to the Fe0I) state. One can see an excellent agreement between the spectral characteristics of both reduced and oxid~ed iron complexes in MMO and in the ZSM-5 matrix. An agreement of this quality is a rather difficult situation to achieve. Beside FeZSM-5, Table 5 includes also Moessbauer data for the FeY zeolite and for some model compounds specially I

L i

Fe-'+ ,.."

9

"."

,

I

I

/"

9

.~.~,-~

Figure 4. Moessbauer spectrum of 57Fe enriched FeZSM-5 recorded at 90 K Solid line is the best fit obtained with the two quadrupole doublets. The sample was calcined at 820 K, excess of inactive iron was removed by oxalate extraction.

i l

.

-2

0

i

4 .................

Velocity (mm/s) prepared to simulate dinuclear iron centers in biology. One can see that their spectral parameters do not correlate so well with MMO, as FeZSM-5 parameters do

501 Table 5 Moessbauer spectra parameters of Fe-containing systems Sample .

.

.

.

.

.

.

Fe state . . .

Isomer shif~ (mm/s)

Quadrupol splitting AE(mm/s)

Reference

MMO

Fe(II)-Fe~) Fe(III)-Fe(III)

1.30 0.50

3.014 1.05

32

FeZSM-5

F~ Fe(I~

1.34 0.47

3.09 0.99

this work

[Fe2(OH)(OAc)2(Mes TACN)] CIO4

Fe(I~Fe(ID

1.16

2.83

33

[Fe2(BPMP)(OPr)2](BPh4)

Fe(II)-Fe(II)

1.20

2.72

34

[Fe20(OAc)2I-IB(pz)s)2]

Fe(RI)-Fe(III)

0.52

1.60

35

[Fe20(O2CH)4(BiPhMe)2]

Fe(III)-Fe(IlI)

0.54

1.81

36

FeY

Fe(II) Fe(III)

1.17 0.34

2.14 1.22

37

4. CONCLUSION Results discussed above show in several lines a distinct biomimetic-type activity of iron complexes stabilized in the ZSM-5 matrix. The most important feature is their unique ability to coordinate a very reactive c~-oxygen form which is similar to the active oxygen species of M/dO. At room temperature or-oxygen provides various oxidation reactions including selective hydroxylation of methane to methanol. Like in biological oxidation, the rate determining step of this reaction involves the cleavage of C-H bond. These data allow to consider FeZSM-5 zeolite as a new successful monooxygenase model. It is worth noticing the important role of the zeolite structure for constructing such an inorganic model, though in this paper we did not have room to discuss this topic. A remarkable advantage of this robust model is an opportunity to study the monooxygenase-like oxygen as well as the active state of iron under easily controlled and reproducible conditions. It opens new possibilities in getting more knowledge on the mechanism of enzymatic oxidation. REFERENCES

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

A.E. Shilov, in Activation and Functionalization of Alkanes (ed. C.L. Hill), Wiley, New York, 1989, p. 13. G.I. Panov, A.S. Kharitonov and V.I. Sobolev, Appl. Catal., 98 (1993) 1. D.B. Tagiyev and Kh.M. Minachev, Stud. Surf. Sci. Catal., 28 (1986) 981. G.I. Panov, V.I. Sobolev and A.S. Kharitonov, J. Mol. Catal., 61 (1990) 85. G. Belussi, A. Carati, M. Clerici, and R. Millini, J. Catal., 133 (1992) 220. M. Ivamoto, K. Matsukami and S. Kagawa, J. Phys. Chem., 87 (1983) 903. E. Suzuki, K. Nakashiro and Y. Ono, Chem. Lett., (1988) 953.

502 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. 34. 35. 36. 37.

M. Gubelmann and Ph. Tirel, Preparation of Phenol by Direct Hydroxylation of Benzene, Fr. Patent No.2 630 735 (1989). A. Kharitonov, L. Vostrikova, K. Ione and G. Panov, Phenol Preparation, Rus.Patent No. 1805127 (1989). A.S. Kharitonov, G.A. Sheveleva, G.I. Panov, Ye.A. Paukshtis and V.N.Romannikov, Appl. Catal. A, 98 (1993) 33. G.I. Panov, G.A. Sheveleva, A.S. Kharitonov, V.N. Romannikov and L.A.Vostrikova, Appl. Catal. A, 82 (1992) 31. V.I. Sobolev, G.I. Panov, A.S. Kharitonov, V.N. Romannikov, A.M. Volodin and K.G. Ione,. J. Catal., 139 (1993) 435. G.I. Panov, V.I. Sobolev, K.A. Dubkov, V.N. Parmon, N. Ovanesyan, A.Ye. Shilov and A.A. Shteinman, J. Amer. Chem. Soc., submitted for publication. V.I. Sobolev, O.N. Kovalenko, A.S. Kharitonov, Yu.D. Pankrat'ev and G.I. Panov Mendeleev Commun., 1 (1991) 29. V. Zholobenko, L. Kustov and V. Kazansky, in R.Ballmoos and M. Treacy (Eds.), Proc. 9th Intern. Zeolite Conf., Butterworth-Heinemann, Boston 1992, vol. 2, p.299. J. Leglise, J.O. Petunnchi and W.K. Hall, J.Catal., 86 (1984) 392. G.K. Boreskov, Catalysis: Science and Technology, 3 (1982) 40. V.Sobolev, A.Kharitonov, E.Paukshtis and G. Panov, J. Mol. Catal., 84 (1993) 117. Y. Moro--oka, Stud. Surf. Sci. Catal., 54 (1990) 53. B.V. Romanovsky, Micromol. Symp., 80 (1994) 185. C.A. Tolman, J.D. Druliner, M.J. Nippa and N. Herron, ref. 2, Chapter 10. R.E. Parton, I.F.J. Vankelecom, M.J.A. Casselman, C.P. Bezoukhanova, J.B. Uytterhoeven and P.A. Jacobs, Nature, 370 (1994) 541. H. Dalton, Catal. Today, 13 (1992) 455. J.T. Groves and G.A. McClusky, Biochim. Biophys. Res. Commun., 81 (1978) 154. A.M. Khenkin and A.E. Shilov, New J. Chem., 13 (1989) 659. P. Kamasamy and R. Cumar, Catal. Today, 9 (1991) 328. V.I. Sobolev, K.A. Dubkov, Ye.A. Paukshtis, L.V. Pirutko, M.A. Rodkin, A.S. Kharitonov and G.I. Panov, Appl. Catal., submitted for publication. L. Pirutko, A. Kharitonov and V. Buchtiyarov, Kinet. Katal., to be published. L.Pirutko, O.Parenago, E.Lunina, A.Kharitonov, L.Okkel and G.Panov, React. Kinet. Katal. Lett., 52 (1994) 275. E. Smimov, L. Makarshin, V. Sobolev, V. Parmon and G. Panov, in preparation. M. Filatov, A. Pelmenschikov and G. Zhidomirov, J. Mol. Catal., 80 (1993) 243. J.G. Dewitt, J.G. Bentsen, A.C. Rosenzweig, B. Hedman, J. Green, S. Pilkington, G.C. Papaefthymion, H. Dalton, K.O. Hodgson and S.J. Lippard, J. Am. Chem. Soc., 113 (1991)9219. J. Hartman, R. Rardin, P. Chaudhuri, K. Pohl, K. Wieghardt, B. Nuber, J. Weiss G.Papaefthymion, R.Frankel and S.Lippard, J. Am. Chem. Soc., 109 (1987) 7387. A.S. Borovik and L.Jr. Que, J. Am. Chem. Soc. 110 (1988) 2345. W.H. Armstrong and S.J. Lippard, J. Am. Chem. Soc., 106 (1984) 4632. W. Tolman, S. Lin, J. Bentsen and S. Lippard, J. Am. Chem. Soc., 113 (1991) 152. Aparicio. J.A. Dumesic, S.M. Fang, M.A. Long, M.A. Ulla, W.S. Millman and W.K. Hall, J. Catal., 104 (1987) 381.

J.W. Hightower, W.N. Delgass, E. Iglesia and A.T. Bell (Eds.) I I th International Congress on Catalysis - 40th Anniversary

Studies in Surface Science and Catalysis, Vol. 101 9 1996 Elsevier Science B.V. All rights reserved.

503

A Molecular Approach to Synergy Generation in Co-Mo Binary Sulfide Catalysts for Hydrodesulfurization Yasuaki Okamoto and Hiromoto Katsuyama Department of Chemical Engineering, Faculty of Engineering Science, Osaka University, Toyonaka, Osaka 560, Japan The mechanism of synergy generation between Co and Mo sulfides for Hydrodesulfurization was studied for highly dispersed Co-Mo binary sulfide clusters encaged in a NaY zeolite. It was shown by EXAFS, XRD, and HREM that highly dispersed Mo sulfides were prepared in the zeolite cages by sulfiding adsorbed Mo(CO)6 . Highly dispersed Co-Mo binary mixed sulfides were synthesized in the zeolite cavities by using Mo(CO)6 and Co(CO)3NO as precursors. It is concluded that thiophene hydrodesulfurization and butadiene hydrogenation take place on the Co sites of the Co-Mo binary sulfide clusters and that chemicai bondings of Co-S-Mo are required for the synergy generation. Coordinatively unsaturated sites were detected by NO adsorption only on the Co sites of the Co-Mo mixed sulfides. It is proposed that the catalytic activities of the Co sites are promoted by electronic modifications induced by the Mo sulfides. 1. INTRODUCTION Hydrotreatmgs of petroleum feedstocks have recently become more and more crucial not only for protecting environments but also for efficient utilization of limited natural resources. Developments of highly active and selective hydrotreating catalysts, in particular hydrodesulfurization (HDS) catalysts, are one of the most urgent problems in petroleum industries. Sulfided Co-Mo or Ni-Mo based catalysts have been used in industry for HDS reactions. It is well known that strong catalytic synergies generate between Co(Ni) and Mo sulfides [ 1-4]. The mechanism of synergy generation in Co-Mo binary sulfide catalysts has been extensively studied for HDS but still remains controversial at present [1-4]. Two main synergy models are under discussion now: a CoMoS model [1,5] and a contact synergy model [3]. In the CoMoS model, the formation of atomically dispersed Co sulfide species anchored on the edge sites of MoS 2 is claimed to be the origin of the synergy generation. The local structures of Co and Mo atoms for supported Co-Mo catalysts have been proposed on the basis of EXAFS results [6-9]. In the contact synergy model, the promotional effects of Co are explained in terms of a remote control mechanism. It is believed that the HDS activity of Mo sulfides is enhanced by spillover hydrogen originally generated on highly dispersed Co sulfides in contact with or in the proximity of the Mo sulfides. Preparations of supported Co-Mo binary sulfide clusters well defined in the structure on a molecular level are considered to be a promising approach to clarify the mechanism of synergy generation. Supported and unsupported metal sulfide clusters have been examined for the purpose up to now [ 10,11 ]. In the present study, we synthesized in zeolite cavities Co-Mo binary sulfide clusters by using Co and Mo carbonyls and characterized the clusters by extended X-ray absorption fine structure (EXAFS), X-ray photoelectron spectroscopy (XPS), Fourier transform infrared spectroscopy (F'FIR), and high resolution electron microscopy (HREM). The mechanism of catalytic synergy generation in HDS is discussed.

504

2. EXPERIMENTAL 2.1. Catalyst Preparation A NaY zeolite (A1/Si atomic ratio; 0.41) was supplied by Shokubai Kasei Kogyo Ltd. After an evacuation at 673 K for 1 h (lx 10 .3 Pa), the zeolite powder was exposed to a vapor of Mo(CO) 6 or Co(CO)3NO at room temperature, followed by an evacuation at room temperature for 10 rain to remove physisorbed metal carbonyl molecules on the external surface of the zeolite. Mo(CO)6/NaY or Co(CO)~NO/NaY was sulfided in a stream of an atmospheric pressure of 10% H2S/H 2 (0.2 dm 3 rain1). The sulfidation temperature was increased from room temperature to 373 K at a rate of 2 K rain ~ and kept at the temperature for 1 h. Subsequently, the temperature was increased up to 673 K at a rate of 5 K min ~ and kept at 673 K for 1.5 h. After the sulfidation, the sample was cooled in the H2S/H 2 stream to room temperature. The Mo and Co sulfide catalysts thus prepared are denoted MoSx/NaY and CoSx/NaY, respectively. Mo sulfide catalysts, MoS2/NaY , were also prepared by a conventional impregnation method by using ammonium heptamolybdate, for comparison. A series of CoSx-MoSx/NaY catalysts was synthesized by intoducing Co(CO)~NO into MoSx/NaY evacuated at 673 K for 1 h, followed by second programmed sulfidation procedures. MoSx-CoSx/NaY catalysts were prepared in the reversed order of the metal sulfide accommodations into the zeolite cavities. When Co2(CO)8 was used as the Co precursor, MoSx/NaY was impregnated with Co2(CO)8 dispersed in n-hexane, followed by evacuation at room temperature to remove the solvent. Co2(CO)s/MoSx/NaY was subsequently sulfided at 673 K to give CoSx/MoSx/NaY. The catalyst composition was determined by AAS and ICP.

2.2. Reaction Procedures The sulfided catalysts were evacuated for 1 h before catalytic reactions. The reactions were carried out under mild conditions by using a circulation system (0.2 dm 3) made of glass. The HDS of thiophene was conducted at 623 K and an initial pressure of 20 kPa (H2/CnH4S = 36). The thiophene pressure was kept constant (0.54 kPa) during the reaction by holding a small amount of liquid thiophene kept at 273 K in the bottom of a U-tube in the reaction system. The products were analyzed by gas chromatography. The HDS activity was calculated from the amount of H2S produced during the reaction. The hydrogenation (HYD) of butadiene was subsequently conducted at 473 K over the catalyst, that had been used for the HDS of thiophene for 1 h, after evacuation at 673 K for 1 h. The initial pressure was 14 kPa (H2/C4H6 = 2). The HYD products were butenes with a small amount of butane. The catalytic activity was calculated on the basis of the total amount of the reaction products.

2.3. Catalyst Characterization The Mo K-edge EXAFS spectra for the catalysts and reference compounds (1k/loS2 and ]'qa2MoO4) were measured on the BL-10B instruments of the Photon Factory at the National Laboratory for High Energy Physics by using a synchrotron radiation. The EXAFS spectra were obtained at room temperature without exposing the sample to air by using an in situ EXAFS cell with Kapton windows [12]. Data analysis was carried out assuming a plane wave approximation. The XPS spectra of the freshly sulfided Co-Mo/NaY catalysts were measured on an XPS-7000 photoelectron spectrometer (Rigaku, A1 anode; 1486.6 eV). The sample mounted on a holder was transferred from a glove bag into a pretreatment chamber attached to the spectrometer as possible as carefully not to be contacted with air. The binding energies (BE) were referenced to the Si2p band at 103.0 eV for the NaY zeolite, which had been determined by the Cls reference level at 285.0 eV due to adventitious carbon. The FTIR spectra of the NO molecules adsorbed on the catalyst were measured on a JEOL 3IR-100 spectrophotometer in a diffuse reflectance mode. After a back ground spectrum was measured in situ for a freshly sulfided catalyst, NO was introduced to the catalyst as 10 % NO/He pulses (5.1 cm3). The FTIR spectra were recorded after an introduction of 5 pulses.

505

3. RESULTS and DISCUSSION 3.1. Mo and Co Sulfide Catalysts The catalytic activities of MoSx/NaY, CoSx/NaY, and MoS2/NaY are shown as a function of the metal atoms per supercage (SC) for the HDS of thiophene and for the HYD of butadiene in Figs. 1 and 2, respectively. MoSx/NaY catalysts exhibited higher catalytic activities by a factor of 3-4 than the corresponding impregnation catalysts, MoSJNaY, as previously reported [13]. The HDS activity of MoSx/NaY increased linearly with the Mo content up to 2Mo atoms/SC, suggesting a formation of uniform Mo sulfide species in this concentration range. The activity leveled off, however, at a higher Mo content, indicating an agglomeration of the Mo sulfide species. Similar loading effects were observed for the HYD of butadiene over MoSx/NaY and MoS 2/NAY as shown in Fig.2. It is noteworthy that CoSx/NaY showed a considerably high HDS activity, being comparable with that of MoSx/NaY. In contrast to relatively low HDS activities of the Co sulfide catalysts supported on A1203, the Co sulfide species supported on activated carbon have been reported to show even higher HDS activities than Mo sulfide catalysts [ 14,15]. This is attributed to an extremely high dispersion of the Co sulfide species on activated carbon. The high HDS activity of CoSx/NaY suggests a high dispersion of the Co sulfide species. With the HYD of butadiene, CoSx/NaY showed a much lower activity than MoSx/NaY. The local structure and dispersion of the Mo sulfide species were examined by EXAFS techniques. The Fourier transforms of the Mo K-edge EXAFS for MoSx/NaY showed two peaks ascribed to Mo-S and Mo-Mo bondings (vide infra). The coordination number of the Mo-Mo bondings for MoSx/NaY was calculated to be close to unity, as summarized in Table 1, indicating the formation of highly dispersed Mo sulfide species, possibly Mo dimer species, in the supercage of the zeolite. High NO adsorption capacities of MoSx/NaY [ 13] are consistent with the high dispersion of the Mo sulfide species. Mo sulfide catalysts were also synthesized by sulfiding Mo oxide dimer species encaged in the NaY zeolite, (MoO3)2/NaY, which had been prepared by a mild oxidation of Mo(CO)6/NaY by using molecular oxygen [16]. The structural parameters derived from EXAFS analysis were identical with those for MoSx/NaY in Table 1.

6ot

80 120 ,.o0 100 x

=-->.- 8O z

o

60

40

"'o-

._Z. 40

20

..~>_ f,O


C O 2 + 2

(1)

2 N O + C O ---> C O 2 + N 2 0

For the NO+propene reaction the rN2 values were obtained directly from experiment by gas chromatographic measurement of the N2 production.

515 Catalytic rate measurements under potentiostatic or galvanostatic conditions were carried out using a galvanostat-potentiostat (Amel type 553). The reactant gas mixture was delivered at total flowrates of 1-2 x 10 -4 mol sec -1, with partial pressures PNO, Pco, Ppropene varied between 0 - 6.5 k Pa, 0 - 1.5 k Pa, 0 - 0.4 k Pa, respectively with PHe, bringing the total pressure to 1 atmosphere in every case. Conversion of the reactants was typically =15%. Control experiments confirmed that the Au reference and counter electrodes were catalytically inert under all conditions. The sample used for the XPS measurements was first tested in the EP reactor to ensure that it exhibited the same catalytic behaviour as that of the samples used to acquire the reactor data. XPS measurements were carried out in a VG ADES 400 UHV spectrometer system. The sample was mounted on a molybdenum block resistively heated by imbedded, electrically insulated tungsten filaments. XP spectra were acquired with Mg Ka radiation with the working (Pt) electrode always at ground potential; appropriate electrochemical potentials (Vwc) were applied between the working (Pt) electrode and the Au counter electrode by applying voltage bias to the latter. The potential of the working electrode with respect to the Au reference electrode (Vwr) was also measured. Quoted binding energies are referenced to the Au 4t"7/2 emission at 83.8 eV; the Au reference spectra were provided by a grounded Au film deposited on the outer face of the quartz sample holder, as illustrated in the inset to Figure 6. 3. RESULTS 3. 1. NO Reduction by CO Steady state measurements of NO decomposition in the absence of CO under potentiostatic conditions gave the expected result, namely rapid self-poisoning of the system by chemisorbed oxygen: addition of CO resulted immediately in a finite reaction rate which varied reversibly and reproducibly with changes in catalyst potential (VwR) and reactant partial pressures. Figure 1 shows steady state (potentiostatic) rate data for CO2, N2 and N20 production as a function of VWR at 621 K for a constant inlet pressures (P~ P~ of NO and CO of 0.75 k Pa. Also shown is the VWR dependence of N2 selectivity where the latter quantity is defined as SN ~ =

rN; rN2 + rN:O

(2)

Note that there is a sharp increase in activity as VWR is reduced below --0.2 V (Na pumping towards the Pt catalyst) with a concomitant threefold enhancement in the selectivity to N2. As it can be seen from Figure 1, the highest selectivities to nitrogen production always occur in the presence of the highest Na loading (most negative potential). The N20 rate actually goes through a maximum at VWR=-0.2 V. Control experiments were carried out in which the total flowrate was varied by a factor of 2 in order to check that the observed step change in

516 activity is due to a true increase in catalyst activity and is not influenced by mass transfer limitations or by multiple steady states in the reactor. As in previous EP studies with 13" alumina [20,25] for any fixed gaseous composition the promotional effect of Na is fully controlled by VWR and the resulting Na coverage. Upon current interruption the promoted rates remain practically constant over periods of many min. Potentiostatic imposition of the initial potential is necessary to restore the initial (unpromoted) rate [20]. 10

1

P NO= Pco = 0.75 k P~z, T = 621 K

'7r I,,.

r9 8

CO

0 E

'o

2

Io

~e (D

0

>"

oo

.~

selectivity

Z

V wr

0

0 -200mV

9 +200mV

.e..,

Z

04

9 +lO00mV

(clean Pt)

04

0

N2

T=621 K, PNO=0.52 kPa &

V wR / volts

-2

-'1

zx 4 ~

zx

()

zx

zx

1

Figure 1- NO+CO. Effect of VWR on CO s, N 2, N O rates and N 2 selectivity.

0

0.5

P co / kPa

1

1.5

Figure 2: NO+CO. Effect of Pco on N=, rate as a function of catalyst potential.

Figure 2 depicts the dependence of N2 rate on Pco at fixed PNO= 0.52 k Pa for three different values of the catalyst potential. VWR=+1000 mV corresponds to the clean Pt surface (unpromoted rate) and VWR=- 200 mV corresponds to a sodium promoted surface. Both CO2 and N2 rates exhibit Langmuir-Hinshelwood behaviour and as can be seen from Figure 2 for N2 rate, increased levels of Na result in a systematic increase in the CO partial pressure (P'co) necessary for inhibition. The N20 rate also exhibits Langmuir-Hinshelwood kinetics, but the effect of increased Na is somewhat different: in particular, high levels of Na tend to suppress the N20 rate and there is no systematic shift in P'co. 3. 2. NO Reduction by Propene In the case of NO reduction by propene, the only detectable reaction products were CO2, N2, N20 and H20. The overall mass balance was found to close within 5% as observed by a combination of GC and mass spectroscopic analyses. Figure 3 shows the effect of varying the catalyst potential on the rate of production of CO2, N2, N20 and on the selectivity towards nitrogen formation, SN2. As can be seen from this figure, both the CO2 and N2

517 reaction rates exhibit volcano-type behaviour: the rate of production is low at both high positive and high negative catalyst potentials, corresponding to zero and high coverage of sodium species, respectively. The region of strong electrochemical promotion corresponds to intermediate values of VWR, where a sharp asymmetric peak in rate is observed. In this promotion regime (+ 100 to -350 mV), the activity shows an exponential dependence on VWR. The behaviour of N20 production is similar, though the peak is not so sharp, and the rate at high negative potentials is higher than at high positive ones by approximately a factor of two. The selectivity towards N2 formation v e r s u s N20 production is thus slightly lower at high negative Vw-Rthan at high positive VWR, with a maximum (0.8) at -0.3 V. 12 10

P No=1.3 kPa, P =0.6 kPa t propene T--648 K &

":

"...

8

i i

_r ~

4

T=648 K, PNO=1.4 kPa

._>

10

-0.5

300mV

0

9

9

O

zx

0.5

Figure 3: NO+propene. Effect of V WR CO 2,N2' NO rates and N2 selectivity.

9

8

3 0) resulted in complete disappearance of vacuum-deposited Na from within the XPS samping depth. This behaviour clearly shows that the electropumped Na and the vacuum-deposited Na behave identically on the surface of the Pt electrode. Further studies to quantify the transport and morphological properties of Na on the catalyst surface are in progress.

Figure 5: NO+propene. Dependence of apparent activation energy for CO 2, N 2, NO rates on VwR.

Figure 6: XPS of Pt/beta alumina acquired under electrochemical bias. A) cleaned surface, B) Na-promoted surface.

4. DISCUSSION It is important to point out that in the discussion that follows the term "Na coverage" is used; this does not imply that the promoter is thought to be present in the form of chemisorbed metallic sodium as it would be in vacuum. The reactive gas atmosphere is expected to lead to the formation of surface compounds of Na, and single crystal data indicate

519 that stable Na-CO complexes [23] or Na carbonates [24] can be formed, depending on the composition of the ambient gas. Adsorbed polar alkali compounds lead to large decreases in work function, of the same order as those produced by the alkali metal itself, so the general theory of electrochemical promotion [1 ] is nevertheless applicable. The NO+CO reaction exhibits strong electrochemical promotion under Na pumping to the catalyst when the partial pressures of both reactants are similar. This and the EP behaviour as a function of CO pressure may be rationahsed as follows. At low partial pressures of CO the rate is low due to the restricted availability of CO~a~; CO+O is the rate limiting step and so the EP effect is small. At intermediate PCO the rate determining step becomes NO dissociation, as the coverages of NO and CO are similar. Thus a strong electrochemical promotion effect is observed: supply of sodium to the catalyst causes enhanced NO dissociation and hence a dramatic rate increase (Figure 1). At high CO partial pressures the EP effect is again attenuated as a result of limited NO coverage and site blocking by CO-Na surface complexes [20]. The selectivity towards N2 formation versus N20 production increases with increasing Na coverage; this follows from increased availability of N(a ) + O(a ) versus NO(a ), again strongly supporting the idea of Na-enhanced NO dissociation. The maximum p values obtained are 13 and 1.5 for N2 and N20 respectively, at a gas composition of P0t~o = P0co = 0.75 k Pa. For P0co = 0.75 k Pa, Na pumping to the catalyst leads to an increase in SN2 by up to a factor of three. Sharp changes in activity which occur as a function of Pt~o at fixed catalyst potential resemble the behaviour observed for the CO+O2 reaction over Pt/I]"alumina [20] and, as in that case, are ascribed to a surface phase transition. At low PNo the surface is dominated by islands of CO; reaction occurs only at the peripheries of these islands, resulting in a low rate. At sufficiently high PNo the CO islands are disrupted by NO chemisorption and the rate rises sharply as intermixing of the reactants occurs. This model is strongly supported by the observed effects of Na promotion. The higher the Na coverage (more negative catalyst potentials) the lower the value of PNO at which the phase transition occurs, reflecting the increased strength of NO chemisorption relative to that of CO. In the case of NO reduction by propene, very large changes in activity also occur under the influence of Na pumping to the Pt catalyst. The overall activity towards formation of carbon dioxide is strongly enhanced when Na is pumped to the Pt catalyst (Figure 3). Activity towards formation of the other products shows a similar dependence on catalyst potential. The maximum gain in rate over the clean surface rate is of the order of a factor of 10 in the case of nitrogen and CO2 production. In the region of strong electrochemical promotion (+100 to - 350 mV), the activity towards formation of all products shows an exponential dependence on catalyst potential, as predicted by theory [25]. It is evident from Figure 3 that there is a precipitous fall in rates as the coverage of Na species increases beyond a critical valued, i.e. the regime of electrochemical promotion (~ 0 -- -350 mV) is followed by a regime of strong poisoning. Results obtained with related systems, including single crystal/electron spectroscopy data obtained with model planar

520 catalysts, indicate that the poisoning behaviour is associated with the formation of stable surface compounds of Na which serve to block active sites at high Na loadings. In this regard, there is a striking difference between the two reactions at very high negative catalyst potentials, NO+propene exhibits a much greater sensitivity to poisoning while the reaction of NO+CO is relativily resistant to poisoning. To put it another way, promotion of NO+propene occurs over a much narrower range of Na coverage. Logically, one should consider this in terms of possible differences in the surface chemical behaviour of CO and propene on a Nadoped Pt surface in the presence of chemisorbed oxygen, in addition to other adsorbed species. Two factors may contribute to this, although further spectroscopic data are needed in order to provide a definitive answer. Factor 1" single crystal studies show that the surface compounds formed in the NO+CO system under reaction conditions undergo significant agglomeration to form 3-dimensional crystallites. In other words, pumping the equivalent of one monolayer of Na to the Pt does n o t lead to the formation of a monolayer of Na compound which would completely block the surface and presumably completely poison the system. Propene has a much greater tendency to deposit carbon on Pt than does CO [26,27,28]. The presence of strongly adsorbed carbon atoms could inhibit agglomeration of Na surface compounds, thereby maximising the number of Pt sites that are affected by a given amount of Na. Factor 2: the chemisorption of both CO and NO is strengthened by coadsorbed Na [29]; the chemisorption bond of propene should be weakened by Na. Therefore at high levels of Na the coverage of propene should be strongly attenuated, with a corresponding large decrease in reaction rate. The observed dependence of the CO2 rate on propene partial pressure for fixed NO partial pressure and for a range of different catalyst potentials (Figure 4) demonstrates that the system exhibits classical Langmuir-Hinshelwood behaviour - a characteristic rate maximum reflecting competitive adsorption of the two reactants. Under all conditions of partial pressure, there was an overall increase in activity with increased pumping of Na to the catalyst: we again associate this with dissociation of NO induced by the Na promoter. The rate maxima shift systematically to higher propene partial pressures as the sodium coverage is increased, reflecting the increase in the binding of NO relative to propene with increasing Na coverage. This kind of behaviour is exactly what one would expect in the case of an electropositive promoter: the chemisorption strength of electron donors (propene) should be decreased whereas the chemisorption of electron acceptors (NO and its dissociation products) should be enhanced. For both reactions studied, NO+CO and NO+propene, the effect of electrochemically pumped Na in increasing the extent of NO dissociation is large and significant. This is because unpromoted low index planes of Pt, Pt(111), are relatively inert towards NO dissociation and we adscribe the NO dissociation as the key reaction-initiating step. Such dissociation of diatomic molecules in the field of coadsorbed cations has been discussed in detail by Lang et al [29]. The rates of production of CO2, N2 and N20 all depend on

521 dissociation of NO for their formation, as it can be analised from the following proposed elementary steps: CO(a ) 4" O(a ) "~ CO 2 N( a ) + N( a ) ===)N 2

(3)

N(a) + NO(a ) --=) N 2 0

The observed increase in the selectivity towards N2 is a consequence of NO dissociation, i.e. a decreased amount of molecular NO, and increased amount of atomic N on the surface, both factors favours the second of the above reactions over the last one. This dissociative mechanism is the generally accepted pathway under ultra high vacuum conditions [7, 8, 30]. However, a recent study by Klein et al [9] has questioned the validity of the dissociative mechanism under atmospheric pressure conditions in favour of a non-dissociative mechanism. A particular difficulty with the non-dissociative mechanism is that it cannot readily account for the lack of reactivity of low index planes of Pt. Our EP results strongly suggest that the dissociative mechanism holds, even in the high pressure regime. The catalyst film consist of large polycrystalline Pt particles whose surfaces are dominated by low index planes that are inactive for NO dissociation. The low rates observed at high positive catalyst potentials (Na-free system) may be ascribed to defects and high index planes that are inevitably present at crystallite edges. Both N2 and N20 are produced in this region as there is a mixture of molecular NO plus atomic N and O. Na supplied to the Pt surface strongly enhances the overall activity by inducing NO dissociation on the otherwise ineffective low index planes in accord with both theory and experiment. For the reduction of NO with propene, the catalyst potential dependence of the apparent activation energies does not show a step change and is much less pronounced than it is for the CO+O2 and NO+CO systems. There is persuasive evidence [20] that the step change is associated with a surface phase transition - the formation or disruption of islands of CO. It is reasonable to assume that this phenomenon cannot occur in the NO+propene case, since there is no reason to expect that large amounts of chemisorbed CO can be present under a n y conditions. That there should be a difference in this respect between CO+O2/CO+NO on the one hand, and NO+propene on the other hand, is therefore understandable; however, the chemical complexity of the adsorbed layer in the NO+propene precludes any detailed analysis of the Ea(VwR) effect. The central assumption underlying all of the preceding discussion is that under EP conditions, reversible changes in VWR correspond to the reversible pumping of Na to/from the Pt from/to the solid electrolyte. Our XPS data clearly show that such reversible transport of Na between 15"-alumina and the surface of the Pt film does indeed occur under the conditions of voltage and temperature that were used for the reactor studies; increasingly negative potentials corresponding to increasing amounts of Na. Furthermore, we have demonstrated the equivalence of vacuum-deposited and electrochemically pumped Na on the catalyst

522 surface. These are important observations that serve to underpin the fundamental theory developed by Vayenas et al [1]. ACKNOWLEDGEMENTS Support under grant GR/J00632 from the UK EPSRC is gratefully acknowledged. MST acknowledges support from British Gas plc. AP and RML acknowledge additional support under a grant from the British Council and Fundaci6n Antorchas. REFERENCES

3 4 5 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

C.G. Vayenas, S. Bebelis, I.V. Yentekakis, and Lintz, H. G., Catalysis Today, Vol. 11, No 3, p. 303. Elsevier, Amsterdam, 1992. S.G. Neophytides, D. Tsiplakides, P. Stonehart, M.M. Jaksic and C.G. Vayenas, Nature 370 (1994) 45. C. Pliangos, I.V. Yentekakis, X. Verykios and C.G. Vayenas, J. Catal. 154 (1995) 124. J. A. Rodriguez and D. W. Goodman, Surf. Sci. Reports 14 (1991) 1. Y. O. Park, W. F. Banholzer and R. I. Masel, Surf. Sci. 155 (1985) 341. W. F. Banholzer, R. E. Parise and R. I. Masel, Surf. Sci. 155 (1985) 653. D'Arcy Lorimer and A. T. Bell, J. CataL 59 (1979) 223. B. A. Banse, D. T. Wickham and B. E. Koel, J. Catal. 119 (1989) 238. R. L. Klein, S. Schwartz and L. D. Schimdt, J. Phys. Chem. 89 (1985) 4908. D. N. Belton and S. J. Scmieg, J. Catal. 138 (1992) 70. D. N. Belton and S. J. Scmieg, J. Catal. 144 (1993) 9. S. E. Oh., G. B. Fischer, J. E. Carpenter and D. W. Goodman, J. Catal. 100 (1986) 360. A. Obuchi, A. Ohi, M. Nakamura, A. Ogata, K. Mizuno and H. Ohuchi, Appl. Catal. B: Environmental 2 (1993) 71. J.R. Hardee and J.W. Hightower, J. Catal. 86 (1984) 137. S. Naito and M. Tanimoto, Chem. Lett. 1993 1935. R. Burch, P.J. Millington, A.P. Walker, Appl. Catal. B: Environmental 4 (1994) 65. T. Miyadera and K. Yoshida, Chem. Len. 1993 p. 1483. H. Hamada, Y. Kintaichi, M.Sasaki, T. Ito and M. Tabata, Appl. Catal. 75 (1991) L 1. C. G. Vayenas, S. Bebelis, S. Neophytides and I.V. Yentekakis, Appl. Phys. A 49 (1989) 95. I.V.Yentekakis, G.D. Moggridge, C.G.Vayenas and R.M.Lambert, J.Catal. 146 (1994) 292. I.V. Yentekakis, S. Neophytides and C.G. Vayenas, J. Catal. 111 (1988) 152. I.V. Yentekakis and S. Bebelis, J. Catal. 137 (1992) 278. J.C. Bertolini, P. Delichere and J. Massardier, Surf. Sci. 160 (1985) 531. I.R. Harkness and R.M. Lambert, J. CataL 152 1995 211. C.G.Vayenas, S. Bebelis and S Ladas, Nature 343 (1990) 625. N.R.Avery, N.S.Sheppard, Proc.Roy.Soc.Lond. A 405 (1986) 1 R.J.Koestner, J.C.Frost, P.C.Stair, M.A.Van-Hove, G.A.Somorjai, Su~Sci. 116 (1982) 85. M.Salmeron, G.A.Somorjai, J.Phys.Chem. 86 (1982) 341. N.D. Lang, S. Holloway and J.K. Norskov, Surface Science 150, 24 (1985). G. Pirug and H.P. Bonzel, J. Catal. 50 (1977) 64.

J.W. Hightower, W.N. Delgass, E. Iglesia and A.T. Bell (Eds.) 11th International Congress on Catalysis - 40th Anniversary

Studies in Surface Science and Catalysis, Vol. 101 9 1996 Elsevier Science B.V. All rights reserved.

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P r o m o t i o n of m o l e c u l a r h y d r o g e n on solid acid c r a c k i n g a c t i v i t y T. Shishido, T. Nagase, K. Higo, J. Tsuji and H. Hattori Center for Advanced Research of Energy Technology, Hokkaido University, Sapporo 060, Japan IR spectroscopy of adsorbed pyridine, and temperature-progrRmmed desorption (TPD) of hydrogen and deuterium were studied for several solid acid catalysts to examine whether the formation of protonic acid sites from hydrogen molecules occurs and promotes acid-catalyzed reactions. The catalysts examined were Co.Mo/SiO2-A1203, Co.Mo/A1203, SIO2-A1203, H-ZSM-5, physical mixture of Pt/SiO2 and H-ZSM-5, and Pt/SO42"-ZrO2, and the reaction chosen was cumene cracking. For all catalysts, the formation of protonic acid site was observed by heating in the presence of hydrogen, and the amount of hydrogen (or deuterium) adsorbed increased as the adsorption temperature was raised. The promotion effects of hydrogen on the catalytic activity for cumene cracking were observed for the catalysts containing metallic components. The concept "molecular hydrogen-originated protonic acid site" is proposed. 1. I N T R O D U C T I O N It is established that the active sites of solid acid catalysts for most of the acid-catalyzed reactions are protons located on the surfaces. Although a proton is formed by heterolytic dissociation of hydrogen molecule, promotion effects of hydrogen on the acid-catalyzed reactions have scarcely been reported. Most of the papers studying hydrogen effects on the acid-catalyzed reactions reported suppression effects of hydrogen; the reactions were retarded in the presence of hydrogen[ 1-5]. In recent years, several papers have reported promotion effects of hydrogen. Baba et al. reported that the promotion effects of hydrogen on ethylbenzene disproportionation over Ag-Y zeolite are caused by the formation of protons from hydrogen molecules accompanied by reduction of Ag § ions to Ag o metals in the presence of hydrogen[6-8]. The protons thus formed are eliminated by the reverse reaction of Ag o to Ag + if the gaseous hydrogen is removed from the reaction mixture. Although the formation of a proton by reduction of metal ion with hydrogen is generally observable, the unique feature with Ag-Y is the occurrence of the reverse reaction as the gas phase hydrogen is removed. Prominent effect of hydrogen was reported by Sachtler et al. with Pd-Y zeolite for isomerization of methylcyclopentane in the presence of hydrogen[9-11]. They proposed "electron-deficient"Pd-H + adducts as the active sites in the presence of hydrogen. Hosoi et al. reported that Pt/SO42--ZrO2 persists a high activity for a long period in alkane skeletal isomerization when the reaction is carried out in the

524 presence of hydrogen[12]. They ascribed a high and stable activity in the presence of hydrogen to a removal of carbonaceous residues deposited on the catalyst by hydrogenation. We have studied the hydrogen effects on the catalytic acyivity of Pt/SO42-ZrO2 for alkane skeletal isomerization, and concluded that the promotion effect of hydrogen is caused by the formation of protons from molecular hydrogen[1317]. Hydrogen molecule is dissociated on the platinum species to form hydrogen atoms which spiUover onto the support. The spiltover hydrogen atom migrates on the support to reach Lewis acid site where hydrogen atom loses an electron to form a proton. The proton is stabilized on the oxygen atom nearby the Lewis acid site. The electron trapped at Lewis acid site reacts with a second hydrogen atom to form a bond of Lewis acid-H-. We proposed the protons originating from molecular hydrogen act as active sites on Pt/SO42--ZrO2 . Recently, the formation of proton from molecular hydrogen was proposed for the physical mixture of NiS and USY zeolite[18]. Skeletal isomerization of alkanes possibly proceeds by two mechanisms; metal-acid bifunctional mechanism and acid-catalyzed monofunctional mechanism. Since Pt/SO42"-ZrO2 contains metallic component, there still remains a possibility that the promotion effect of hydrogen on skeletal isomerization is caused by some interaction of hydrogen with the surface, but not by the formation of protons. It would be more clearly demonstrated that the protons formed from hydrogen molecules act as active sites for acid-catalyzedreactions if the reaction catalyzed only by acid sites are promoted in the presence of hydrogen over Pt/SO42"-ZrO2. In addition, it would be important to examine whether the concept "molecular hydrogen-originated protonic acid site" can be extended to the catalysts other than Pt/SO42--ZrO2 . The present paper aims to clarify the following points. 1 Does the generation of protonic acid sites originating from molecular hydrogen occur for the catalysts other than Pt/SO42"-ZrO2? 2 Are the promotion effects of hydrogen on the typycal acid-catalyzed reaction observable for Pt/SO42"-ZrO2 and the other catalysts? As a pure acid-catalyzed reaction, cumene cracking was chosen, and hydrogen effects were examined for Co.Mo/SiO2-AI203, Co.Mo/A1203, SIO2Al203, H-ZSM-5, a physical mixture of Pt/SiO2 and H-ZSM-5, and Pt/SO42-ZrO2. 2. E X P E R I M E N T A L

2.1 Catalyst preparation The Pt/SO42--ZrO2 was prepared as follows. The sulfated ion treated Zr(OH)4 was prepared by impregnation of Zr(OH)4 with 1N H2SO4 aq. solution followed by filtration and drying at 383K. The Zr(OH)4 was obtained by the hydrolysis of ZrOC128H20 with aqueous ammonia. The obtained gel was washed with deionized water until no C1- ions could be detected. The Pt/SO42--ZrO2 sample (0.5 wt%Pt) was prepared by impregnation of SO42--ZrO2 with 1% H2PtC16 aq. solution followed by drying at 383K and calcination at 873K in air. The amount of S remained in the resulting catalyst was 1.5 wt% determined by XRF. The catalyst was pretreated in a hydrogen stream at 623K for 2 h before

525 use for reaction. For IR and TPD measurements, the catalyst was treated further in a vacuum at 623K. Pt/SiO2 was prepared by impregnation of silica supplied from the Catalysis Society, Japan (JRC-SIO-1) with H2PtC16 aq. solution followed by drying, and calcining at 773K for 5 h. The content of Pt was adjusted to be 2 wt%. H-ZSM-5 was prepared by ten times ion exchange of Na-ZSM-5 supplied from TOSOH with 1N NH4C1 aq. solution. The SIO2/A1203 molar ratio was 23.3. The catalyst was finally calcined at 803K for 3 h. The physical mixture of the Pt/SiO2 and HZSM-5 was prepared by mixing the two components in the ratio 1:10 (Pt/SiO2 : H-ZSM-5) in an agate mortar. Co.Mo/SiO2-A1203 was prepared as follows. Into a solution of cobalt nitrate, aqueous ammonia was added to form aqueous solution of cobalt ammoniun complex. To the solution, aqueous molybdenum ammoniun was addes to make a mixed solution containing cobalt and molybdenum. Silica-alumina supplied from the Catalysis Society, Japan(JRC-SAL-2) was impregnated with the mixed solution followed by drying at 353K for 36 h, crushing into a powder, and calcining at 773K for 16h. The contents of Co and Mo were 4 and 8 wt% as CoO and MOO3, respectively. The Co.Mo/A1203 catalyst used as a reference catalyst was supplied from Cyanamid Co. Ltd. (HDS-20).

2.2 IR study of adsorbed p y r i d i n e A self-supported wafer of the sample was placed in an in situ IR cell with CaF2 windows, and pretreated. The pretreatment procedures varied with the catalyst samples. For Co.Mo/SiO2-A1203 and Co.Mo/A1203, the calcined sample was pretreated in a hydrogen flow at 673K for 2h followed by outgassing at 773K. For presulfided sample, the calcined sample was treated in a 3%H2S in H2 at 673K for 2h followed by outgassing at 773K for 17 h. For the physical mixture of Pt/SiO2 and H-ZSM-5, the calcined sample was outgassed at 773K and cooled to 673K at which the sample was treated in a hydrogen flow for 2 h followed by outgassing at 773K for 17h. For H-ZSM-5, the calcined sample was outgassed at 773K for 17h. After the pretreatment, pyridine was introduced into the cell at a pressure of ca. 2 Torr for 15 rain at 423K, and then outgassed normally at 673K for 20 rain. In the case of Co.Mo/SiO2-AI203, the final outgassing was done at 423K. To examine the effects of hydrogen on the adsorbed pyridine, the pyridine-covered sample was exposed to 500Torr of hydrogen at room temperature, and heated stepwise from room temperature to 573K or 673K by 50K increments. Following the processes of heating in the presence of hydrogen, the sample was outgassed stepwise from 373K to 673K by 50K increments. All spectra were recorded on an FT/IR-5300 infrared spectrometer(JASCO) at room temperature. For determination of the number of protonic sites and Lewis acid sites on the surface, the integrated absorbances of the bands at 1450 cm -1 (due to pyridine chemisorbed on Lewis acid sites, L-Py) and 1490 cm "1 (due to both the L-Py and pyridine chemisorbed on protonic acid sites, B-Py) were used with the tangent background for all samples. The values obtained were normalized to the weight of the sample wafer. To obtain the apparent absorption coefficients of the bands, a known amount of pyridine was adsorbed on the sample, and the absorbance of each band was measured. Then, a small quantity of water which is sufficient to convert all Lewis acid sites into protonic acid sites was introduced into the IR-

526 cell. The apparent absorption coefficients were calculated from the changes of the absorbances at 1450 and 1490 cm "1 on exposure to water vapor. As the apparent absorption coefficient for a certain band is different for a different wafer, the calibration of the apparent absorption coefficient was performed for each wafer.

2.3 T e m p e r a t u r e - P r o g r a m m e d Desorption of h y d r o g e n and/or deuterium The p r e t r e a t m e n t conditions were the same as those for IR study of adsorbed pyridine. The sample pretreated in a vacuum was exposed to ca. 300Torr D2 at different t e m p e r a t u r e s for 1 h. After cooling to room temperature, the sample was outgassed for 10 rain prior to TPD run. TPD was run at a heating r a t e of 10K rain "1, and the desorbed gases were analyzed by mass spectrometry. 2.4 R e a c t i o n p r o c e d u r e s for c u m e n e c r a c k i n g A pressurized flow reactor and an atmospheric pulse reactor were employed for carrying out cumene cracking. For the reaction over Co.Mo/SiO2-AI203 and SIO2-A1203, the flow reactor was used with carrier flow rate of 100ml min "1 and cumene feed rate of 14ml min "1 under the pressure of 30 arm. For the reaction over Pt/SO42"-ZrO2, the physical mixture of Pt/SiO2 and HZSM-5, and H-ZSM-5, the pulse reactor was used. A dose of cumene, 0.5ml (0.036mmol), was passed over the catalyst in a carrier flowing at 50ml min "1, and the products were trapped at 77K before being flash-evaporated into a gas chromatographic columns(PEG-20M and VZ-7). 3. R E S U L T S

3.1 IR study of the effect of h y d r o g e n on the acid site The acid-site types were examined by IR spectroscopy of adsorbed pyridine. For Co.Mo/SiO2-A1203, no pyridine remained on the surface after outgassing at 673K. There are no acid sites strong enough to retain pyridine against outgassing at 673K. Therefore, the catalyst was exposed to pyridine and outgassed at 423K, and then the catalyst was heated at 523K or 623K in the presence of 300 Torr of hydrogen. The variation of the percentages of protonic acid sites and Lewis acid sites as a function of the temperature at which the Co.Mo/SiO2-AI203 was heated in Fig. 1 Change of the acid sites on Co.Mo/SiO2A1203 with hydrogen t r e a t m e n t and the the presense of hydrogen is following outgassing. shown in Fig. 1.

527

Although most of the acidic sites on the surface are Lewis acid, the fraction of protonic acid sites increased as the catalyst was heated in the presence of hydrogen. The increase in the fraction of protonic acid sites was extended as the heating temperature was raised. This indicates that protonic acid sites were generated by heating in the presence of hydrogen. The protonic acid sites thus formed were eliminated by outgassing gas phase hydrogen at 673 K. The protonic sites are generated and eliminated in response to heating in the presence and absence of hydrogen. These observations are essentially the same as those observed for Pt/SO42"-ZrO2 catalyst, though the protonic acid sites formed and eliminated are much stronger for Pt/SO42--ZrO2 than for Co.Mo/SiO2-A1203. For presulfided Co.Mo/SiO2-Al203, the formation and elimination of the acid sites were almost the same as those observed for non-sulfided Co.Mo/SiO2-AI203. Presufidation with hydrogen sulfide did not affect much the conversion of the acid sitescaused by gas phase hydrogen. For Co.Mo/A1203 catalyst, no IR bands ascribed to pyridine were appreciable after outgassing at 423K. The acid sites on Co.Mo/A1203 are much weaker than those on Co.Mo/SiO2-A]203. For the physical mixture of Pt/SiO2 and H-ZSM-5, pyridine was retained on the surface after outgassing at 673K. On t~e surface of the mixed catalyst, the acid sites as strong as those on H-ZSM-5 exist as expected. The IR spectrum of adsorbed pyridine changed markedly when the sample was heated in the presence of hydrogen. On raising the temperature, the IR band ascribed to protonic acid sites developed with concomitant decrease in the intensity of the IR band ascribed to Lewis acid sites. The variations of the amounts of protonic acid sites and Lewis acid sites as a function of the heating temperature in the presence of hydrogen are shown in Fig. 2. "7

-~

H2 exposure

Evacuation

9 8 i

,~6 a

o4

i2 o

~0

o


e'-

0 6

' "" "

o

..........

O ,4-,, C

t 9

. ....

.9 ~-10 tO

o

c.c

30

60 90 Time (min)

120

150

Figure 4. Conversion-time curves on a regenerated 2NiSZ(s) sample obtained at intermitent reaction periods after which the flow of n-butane was stopped, leaving the catalyst at the reaction temperature under a 100 cm/min flow of He for a few minutes, and then resuming the flow of n-butane. The arrows indicate 15 ktl injections of l-butene.

=.5-_ -

100

200

300

400

Temperature (*C) Figure 5. n-Butane conversion to isobutane as a function of temperature in a temperature programmed reaction experiment conducted over 0.4 g of 1NiSZ(s) catalyst under an n-butane/ hydrogen mixture (n-C4 molar fraction = 0.34) at a constant heating rate of 2C/min.

Fig. 4 also shows the result of injecting 1-butene in the feed during one of the induction periods. We previously demonstrated [7] that on FMSZ catalysts, the addition of olefins significantly modified the form of the induction period. When 1-butene was added to the He stream before starting the feed of n-C,I-I~0, the induction period was very much shortened and the activity rapidly increased. By contrast, when the addition of olefins was done several minutes after contacting the catalyst to the n-CA-I~0feed, there was almost no difference in the shape of the conversion-time curve. Similar results were obtained on the NiSZ(s) catalyst reported here. When 1-butene was added before starting the feed of n-Cd-I~0 (not shown)the induction period was very much shortened. The arrows in Fig. 4 indicate the injection of 1butene, both at the beginning and almost at the end of the fourth induction period. The

560 activity jump caused by the injection was more pronounced at the beginning, when the surface started to accumulate olefins, than at the end, when it was almost at its saturation point. In a separate experiment, we varied the length of time during which the catalyst was left under He before resuming the butane flow. When this time was 10 rain or longer, we observed that upon resuming the n-Cd-I~0 flow the conversion started from zero, as shown in Fig. 4. By contrast, when the catalyst was left in He for 2 rain or less, the conversion started at about the level that was before shutting the n-Cd-I~0 flow off. This result indicates that the surface species can be removed during the 10 min. periods, but remain on the surface when the period without n-CA-It0was short. Compared to the runs with n-Cd-I~0/He mixtures, those conducted n-C4H10/H2 mixtures exhibited a much lower activity and did not show the rapid induction period exhibited by the fresh catalysts under the n-Cd-I10/He mixtures.

3.3. Temperature Programmed Studies In addition to the isothermal runs, we carried out a temperature programmed reaction (TPRx) experiment to study the evolution of n-Cd-I~0 isomerization activity of the NiSZ catalysts as a function of temperature in the presence of H2. In this experiment, we increased the temperature linearly at a heating rate of 2~ and simultaneously monitored the evolution of products, sampling the reactor outlet every 5 min. The variation of conversion of n-C4H~0 to isobutane in the linear temperature ramp is illustrated in Fig. 5. It can be observed that the conversion initially increased with temperature, reached a maximum at about 250~ and then rapidly decreased. It is interesting that at about 250~ when deactivation begun, the concentration of Ct, C3, and C5 products started to increase, but, when the temperature reached 400~ the concentration of all products dropped to zero. There are several possible causes for the rapid deactivation observed above 250~ One of them is the formation of carbonaceous deposits. This form of deactivation, however, should not be substantial in the presence of H2. Other possible causes of catalyst deactivation are the loss of sulfate groups and the conversion of Ni sulfate into Ni sulfide. In our TPR experiments, we detected the evolution of H2S, but this only occurred at temperatures close to 500~ not at the relatively low temperatures at which deactivation started. As described above, we attempt to explain the observed catalytic behavior in terms of the accumulation of reaction intermediates which involve the participation of olefins [4, 7, 10]. To study the interaction of C4 olefins with the Ni-promoted and the unpromoted sulfated zirconia catalysts, we exposed two samples (1NiSZ(s) and SZ(s) previously treated in air at 600~ to a 10 Tort of 1-butene during 10 rain at room temperature. The catalyst was then exposed to pure He and the temperature was increased with time at a constant heating rate of 8~ The desorption products, continuously monitored by MS, started to appear at about 100~ and had a maximum at about 210~ To identify the products desorbed at the begining of the desorption process and at the point of maximum resorption rate, we sent pulses to a GCMS at 150~ and 210~ The total amount of products detected in the second pulse (maximum desorption rate) was more than 20 times greater than in the first pulse. The product distribution significantly varied for the two pulses. For instance, the 150~ exhibited the presence of butenes, and C5 olefins as main products. This distribution indicates that some of the adsorbed butenes undergo oligomerization

561 and cracking, while others may have just desorbed after a simple double-bond isomerization. The 210~ exhibited a much higher alkane/olefin and the presence of isobutane as the main desorption product. The concentration of isopentane and hexane was also significant. The presence of these products would indicate that, in addition to the oligomerization and cracking steps, a H2 transfer process takes place in which the coke deposits provides the H2 for hydrogenating the olefins. The differences observed in the product distribution and desorption temperatures following the adsorption of 1-butene for the promoted and unpromoted catalysts are relatively minor. This suggests that the important effects of Ni are most evident under reaction conditions. Although olefins do play a role [7] the promoting effect of Ni cannot be ascribed to a direct interaction between the promoter and the olefin but perhaps to a concerted effect in which more than one species is participating. To probe the formation of coke we conducted TPO measurements on samples previously used in the butene TPD experiments. The TPO profiles corresponding to catalyst 1NiSZ(s) are shown in Fig. 6. Significant evolution of CO2 was detected, indicating the formation of coke during the adsorption/desorption of 1-butene. The 1NiSZ(s) catalyst exhibited almost twice as much CO2 as the unpromoted SZ sample.

to'J .

m

O')

t.j t~ CO t~ Ill

0

200 400 T e m p e r a t u r e (C)

600

Figure 6. TPO of carbonaceous deposits lett on the surface of I NiSZ(s) after the adsorption/desorption of 1-butene. Evolution of CO2 and SO2 (curves A and B, respectively). Evolution of SO2 during the butene TPD (curve C, added for comparison). The loss of sulfate during the reaction steps or during regeneration may become a critical issue when analyzing the potential of these materials as commercial catalysts. Sulfate losses during the butene TPD, made evident by the evolution of SO2 (rn/e=64), started to occur at about 500~ We have previously demonstrated the evolution of SO2 in the presence of adsorbates such as ammonia, benzene, or pyridine at temperatures much lower than those required to produce SO2 from clean sulfated zirconia [ 14]. For instance, A treatment in He at 600~ causes drastic losses which result in a significant drop in activity (see Fig. 3). It is

562 important then to compare the evolution of SO2 under different conditions. For example, during the TPO experiment, very small quantities of SO2 were evolved contrasting with relatively large amounts evolved during the TPD of the adsorbed olefins. 4. CONCLUSIONS Several conclusions can be drawn from the present work: 9 The promoting effect on the n-C~8-I~0isomerization increases with the amount of Ni, and further increases with the addition of Mn. We have presented evidence that on the Nipromoted sulfated zirconia catalysts, the nickel is, at least partially, in the form of sulfate. 9 The induction period exhibited by these catalysts is related to the accumulation of reaction intermediates on the surface. 9 When hydrogen is present in the feed, the catalytic activity of NiSZ is much lower than when it is not present. Under hydrogen, the activity increases with temperature up to 250~ but it rapidly deactivates at higher temperatures. 9 Sulfate losses are important above 500~ under reducing conditions and in the presence of adsorbed olefins. They are much less significant under oxidizing conditions such as those found during the catalyst regeneration in air. ACKNOWLEDGMENTS We thank Ms. Wei-Chee Tan for the preparation of the catalyst samples. This work was supported by the National Science Foundation (CTS-9403199), the international cooperative program NSF-CONICET (INT-9415590), and the Exxon Education Foundation. We thank the University of Mar del Plata for a fellowship (WEA), as part of the international exchange program sponsored by the University of Oklahoma and the University of Mar del Plata. REFERENCES l.

2.

.

9. 10. 11. 12. 13. 14.

US Pat. 4,918,041 (1990) C. Y. Hsu, C. R. Heimbush, C. T. Armes, and B. C. Gates J. Chem. Soc. Chem. Comm. (1992) 1645 M. Hino and K. Arata, J. Chem. Soc., Chem. Comm. (1980) 573 Adeeva V., Lei G. D., and Sachtler W. M. H., Appl. Catal. A 118, L11 (1994) A. S. Zarkalis, C. -Y. Hsu, and B. C. Gates, Catal. Lett. 29, 23 5 (1994) Guisnet M. R., Acc. Chem. Res. 23,392 (1990) M. A. Coelho, W. E. Alvarez, E. C. Sikabwe, R. L. White, and D. E. Resasco, Catal. Today (in press) A. Jatia, C. Chang, J. D. MacLeod, T. Okabe, and M. E. Davis, Catal. Lett. 35, 21 (1994) J. Tabora and R. J. Davis, J. Chem. Soc. Faraday Trans. 91, 1825 (1995) M. Coelho, D. E. Resasco, E. C. Sikabwe, and R. L. White, Catal. Lett. 32, 253 (1995) F. R. Chen, G. Coudurier, J. F. Joly, and J. C. Vedrine, J. Catal. 143, 616 (1993) R. A. Comelli, C. R. Vera, and J. M. Parera, J. Catal. 151, 96 (1994) B. A. Morrow, R. A. McFarlane, M. Lion, J. C. Lavalley, J. Catal. 107, 232 (1987) E. C. Sikabwe, M. A. Coelho, D. E. Resasco, and R. L. White, Catal. Lett. 34, 23 (1995)

J.W. Hightower, W.N. Delgass, E. Iglesia and A.T. Bell (Eds.) 11th International Congress on Catalysis - 40th Anniversary

Studies in Surface Science and Catalysis, Vol. 101 9 1996 Elsevier Science B.V. All rights reserved.

563

RARE EARTH MODIFIED SILICA-ALUMINAS AS SUPPORTS FOR BIFUNCTIONAL CATALYSIS Stuart L. Soled*, Gary McVicker, Sal Miseo, William Gates, and Joe Baumgartner, Exxon Research and Engineering Co., Rt. 22 East, Annandale, NJ 08801 USA ABSTRACT We have explored rare earth oxide-modified amorphous silica-aluminas as "permanent" intermediate strength acids used as supports for bifunctional catalysts. The addition of well dispersed weakly basic rare earth oxides "titrates" the stronger acid sites of amorphous silicaalumina and lowers the acid strength to the level shown by halided aluminas. Physical and chemical probes, as well as model olefin and paraffin isomerization reactions show that acid strength can be adjusted close to that of chlorided and fluorided aluminas. Metal activity is inhibited relative to halided alumina catalysts, which limits the direct metal-catalyzed dehydrocyclization reactions during paraffin reforming but does not interfere with hydroisomerization reactions. INTRODUCTION The worldwide interest in solid acid catalysis that has developed over the last few years has focused primarily on oxide-based strong "super" acids as possible replacements for conventional liquid or halide-containing acids (1-3). Additional opportunities exist in processes that use intermediate strength acids. In particular, hydrocarbon conversions that require the simultaneous participation of both a metal hydrogenation-dehydrogenation function and an acid function, such as catalytic reforming and paraffin hydroisomerization, have traditionally used bifunctional catalysts consisting of group VIII metals dispersed on chlorided or fluorided alumina supports. In some applications, halide is stripped off these catalysts (by water generated from the reaction, for example) and periodic halide replacement by treatment with organic halides or HCI is required. Availability and use of halogen-containing treatment gases, howver, are coming under increasing environmental pressure. In addition, variation in halide level often produces unwanted changes in acid strength that result in selectivity changes during reaction. Consequently, there is interest in developing alternative environmentally-compatible halide-free bifunctional catalysts (4). Silica-aluminas contain acid sites stronger than those found in most halide-treated aluminas. We have attempted to tailor the acidity of amorphous silica-alumina by adding varying levels of "permanent" inorganic basic titrants. Such titrants were chosen in order to meet the following four criteria: 1) be weakly basic- in order to gradually titrate the acid sites so that their strength matches that found in halided-aluminas. This precludes the use of alkali and alkaline earth oxides. 2) be easily dispersible- the weakly basic oxides should spread as a monolayer on the SiO 2A120 3 surface. If the oxides agglomerate, the acid sites would not be uniformly titrated. 3) be non-reducible- because of the metal present on bifunctional catalysts, we want to avoid any possible alloy formation on the support surface during reduction. 4) be able to disperse group VIII metals- it has been known that it is difficult to maintain high metal dispersions on pure SIO2-A120 3. To obtain an effective metal/acid balance, enhanced

564 dispersion of the group VIII metal on modified SIO2-A120 3 relative to neat SIO2-A120 3 would be advantageous. We have found that rare earth oxides are mild bases that spread uniformly on SIO2-A120 3 supports and systematically reduce the acid strength into the range of halided aluminas. Of course, solid acidity is a complex parameter that encompasses acid site number, type, strength, hardness etc. so that matching the entire spectrum of acidic properties of a particular support is not likely. In fact, subtle differences in acid and metal properties provide opportunities to exploit differences in catalytic performance. Rare earth oxides do not reduce during catalytic reactions of interest, and we show, surprisingly, that they provide advantages in improving Pt metal dispersion. Consequently, we evaluated Pt on rare-earth oxide-modified silica-aluminas as environmentally compatible bifunctional catalysts. Acidity has to be closely balanced with an active and dispersed group VIII metal, usually Pt, to minimize either metal or acid cracking side reactions (5). Both physical and chemical characterization probes as well as selected model compound reaction tests were used to characterize these new catalysts. Acid site strengths have been determined using an olefin isomerization test, metal dispersions estimated with dihydrogen chemisorption and benzene hydrogenation, and bifunctional properties monitored during heptane reforming and n-C12 isomerization. We have characterized both acid and metal functions and compared them with conventional bifunctional catalysts. EXPERIMENTAL Preparation and Characterization The SiO2-AI20 3 (Si-AI) powder samples were obtained from Davison; both MS25 (75% wt. S i t 2, 25% wt. A1203) and MS13 (87% wt. S i t 2, 13% wt. A1203) were evaluated. The rare earths were added by incipient wetness impregnation of nitrate solutions, dried overnight at 120~ and then air calcined to oxides at 500-600~ We included yttrium oxide in our studies, since it behaves similarly to rare earth oxides. Group VIII metals were added by incipient wetness impregnation. For those samples containing Pt as the group VIII metal, we used a standardized chloroplatinic acid solution as the platinum precuror. X-ray diffraction spectra, collected on a Rigaku D-Max diffractometer with CuK~

radiation, was used to check for crystalline phases. ESCA spectra, which monitored the rare earth oxide dispersion, were measured on a Perkin Elmer 5600 XPS hemispherical electron analyzer using Mg Kt~ (1256 ev) radiation with a pass energy of 35ev. Samples were mounted on double sided tape. Peak positions were shifted to place carbon at 285 ev. Surface area and pore size distributions were measured by N 2 gas desorption. Platinum dispersions were measured by hydrogen chemisorption at 25~ assuming a 1:1 H:Pt surface stoichiometry. Chemisorption isotherms were obtained by measuring dihydrogen uptakes between 25 and 200 torr and extrapolated to zero pressure to obtain the total chemisorption uptakes. The strong chemisorption uptakes, which are reported here, were obtained by subtracting the back adsorption isotherm (i.e. the hydrogen chemisorption measured after evacuating the cell following the total adsorption isotherm) from the total uptake (6). Reaction tests The 2-methylpent-2-ene (2MP2) isomerization test was carried out as described previously (7). The formation rates and rate ratios of the product hexene isomers of this test reaction reflect the relative acid site concentration and strength of the catalyst,

565 respectively. The product hexene isomers formed include 4-methylpent-2-ene (4MP2), t-3methylpent-2-ene (t-3MP2), and 2,3 dimethylbute-2-ene (2,3 DMB2). 4MP2 requires only a double bond shift, a reaction occurring on weak acid sites. 3MP2 requires a methyl group shift (i.e., a stronger acidity requirement than double bond shift), whereas the double branched 2,3DMB2 product requires even stronger acidity. For a homologous series of solid acids, differences in t-3MP2 rates normalized with respect to surface area reflect the density of acid sites possessing strengths sufficient to catalyze skeletal isomerization. Since skeletal isomerization rates generally increase with increasing acid strength, the ratio of methyl group migration rate to double bond shift rate should increase with increasing acid strength. The use of rate ratios, in lieu of individual conversion rates is preferable since differences in acid site populations are normalized. For the n-C 7 reforming and n-C 12 isomerization reactions the catalysts were run in a fixed bed micro reactor equipped with on-line GC analysis. The catalyst, together with a quartz powder diluent, was added to a 6 inch reactor bed. A thermocouple was inserted into the center of the bed. The catalysts were calcined at 350-500~ immediately prior to use and reduced in H 2 at 350-500~ for 1 hour. n-Heptane or dodecane (Fluka, puriss grade) were introduced via a liquid feed pump. The runs were made at 100-175 psi with a H2/n-heptane (or n-C12) feed ratio of 7 and a weight hourly space velocity of 6-11. Benzene hydrogenation was used to probe metal site activity. A 12/1 H2/benzene feed was passed over the catalysts at 700 kPa with a weight hourly space velocity of 25. The temperature was set to 100~ and the conversion of benzene to cyclohexane was measured after 2 hours at temperature. The temperature was then increased at 10~ increments and after two hours, the conversion remeasured. R E S U L T S AND D I S C U S S I O N Samples with different loadings of rare earth and yttrium oxides on MS25 Si-AI were prepared and characterized. Figure 1 shows the x-ray diffraction spectra of a sample of 25% Nd203/Si-AI compared with the unmodified Si-AI.

0

IO

20 two

30

40

50

~ea~m(des,~)

Figure 1. X-ray diffraction spectra of Si-A1 and 25% Nd/Si-AI. In both cases, we observe an amorphous pattern; no crystallites of rare earth oxide appear even at 25% wt. loading. This indicates that oxide particles remain less than -30,~, in diameter. The surface area, pore volume and pore size distribution of the starting Si-AI support also change on impregnation. Table 1 lists the values for yttria-modified samples of

566 MS25. The results show that as the yttria disperses over the entire surface, the smaller pores of Si-AI are blocked first making a fraction of the surface inaccessible to physisorbed N 2. Furthermore, the average pore size increases as a result of preferential blockage of small pores (Figure 2). Table 1. Change of Surface Area and Pore Volume For Yttria-modified Si-A1 (75/25% wt.) surface area (m2/g) 302 235 208 183

Si-A1 4%Y203/Si-AI 9%Y203/Si-AI 15%Y203/Si-A1.

,e~ r.,,n,~ ,~ :~,Jo-lis~vl' ~ ,,,..---.I ~ L'p*,, ,,t,=~ 0.~-0.~,,sgj

20 16

;i02-AI, Oj t "~

~12

pore volume (cc/g) 0.78 0.67 0.65 0.52

, , ~ ' ~

4.Y.O~SIO.-AI.O.

4 sio,.~o;' "f...... /_ I

o

-

0

10

-

20

30

40

SO r,

60

?t

$0

90

""

100

(~)

Figure 2. Pore size distribution for Y203-modified Si-AI (75/25% wt. SIO2-A1203) We examined a series of Nd203-1oaded silica-aluminas by ESCA to further monitor rare earth oxide dispersion. As seen in Figure 3, the ESCA Nd/(Si+A1) signal, which measures the surface concentration of Nd (relative to (Si and AI), to about a -30,& penetration depth) increases linearly with increasing Nd20 3 loading. This result strongly suggests that N d 2 0 3 continues to wet and spread on the surface (i.e. remain below monolayer coverage) up to 30% wt. loading. 0.25

o 0.2i o 0.15 + 0.1-

0.0S 10

15

20

2$

30

wt~ NdzO~ on SiO,-AI203

Fig. 3. Surface Nd ESCA Signal Increases Linearly with Loading The titration of silica-alumina acid sites with neodymium oxide is clearly seen in Figure 4. The relative acid strength, as estimated by the 2,3DMB2/4MP2 ratio during 2MP2

567 isomerization, systematically decreases with increasing N d 2 0 3 loading. The ratio does not change dramatically for the lowest neodymia loadings under these conditions (250~ 1 hr run time) since it lies near its equilibrium value. Under less severe, lower temperature isomerization

Fig. 4. SiO2-A1203 Acid Strength Decreases with Nd203 Addition conditions, larger differences occur at low rare-earth oxide loading levels. Comparative data for chlorided and fluorided aluminas is also shown in Figure 4 and indicate that the acid strength of rare earth modified silica-aluminas can be adjusted to match that of conventional halide-modified transitional aluminas. We investigated variations in thermal treatments on rare-earth modified Si-AI as well as using different rare earth oxides. We are including Y203 with the rare earth oxides, since it behaves similar to them. In Figure 5, we show that 25% wt. Nd203/Si-A1 and 18% wt. Y203/Si A1 which contain comparable similar atomic loadings have nearly identical t-3MP2/4MP2 rate ratios and both ratios are just slightly larger than that shown by 0.9% wc C1/AI20 3.

Fig. 5. Equimolar Rare Earth Oxide Concentration Display Similar Acid Strengths An interesting difference occurs with the stronger acid sites. Figure 5 shows that 0.9% wc CI/A120 3 contains more of the stronger acid sites, as measured by the 2,3DMB2/4MP2 ratio than the 25% N d 2 0 3 or 18% Y 2 0 3 catalysts. We suspect that this difference in distribution in site strengths results from preferential titration of stronger acid sites on Si-AI by the basic rare earth oxides. This subtle difference in acid strength distribution between halided aluminas and rare-earth oxide modified silica-aluminas should have an impact on their catalytic properties. We also see (Fig. 6) that raising the calcination temperature of rare earth modified silicaaluminas from 500 to 600~ may slightly sinter the rare earth oxide. Although the x-ray pattern

568 still remains amorphous, sintering would lower the effective surface coverage of the neodymium oxide titrant and consequently increases support acidity. This speculation is consistent with the data shown in Fig. 6.

Fig. 6. Apparent Acid Strength of 25% Nd203/SiO2-AI20 3 Increases with Calcination Temperature Published studies have shown that differences in basicity among rare-earth oxides is small; significantly less than differences among alkaline earth oxides (8). Furthermore, basicity does not change uniformly across the periodic table. A sequence of base strengths for the different rare earth oxides has been proposed: La > Pr = Nd > Sm > Gd -- Eu > Tb -- Ho --- Er > Dy --- Tm = Yb -- Lu > Ce (9). As seen in Figure 7, aside from CeO 2, most of the rare earth oxides we tested did not show substantial differences in basicity in the tests used here. Only impregnation with ceria did not appreciably titrate the acidity, perhaps as a result of the lower basicity of the Ce +4 ion.

Fig. 7. Different Rare Earth Oxides at Equimolar Loadings on SIO2-A120 3 We measured the dispersion of Pt (impregnated from a chloroplatinic acid precursor, calcined at 450~ and reduced at 500~ on a series of Nd203-1oaded silica-aluminas (Fig. 8). We find, unexpectedly, that dispersion increases with increasing rare earth oxide loading up to about 18% N d 2 0 3, where it plateaus at between 40 and 50%, compared to 10% with unmodified Si-AI. This compares with dispersions o f - 6 0 - 8 0 % measured on similarly Pt-loaded transitional A1203 catalysts. Transmission electron micrographs confirmed the decrease in particle size with rare earth content on Si-A1.

569

0.8

0.57

0.6

~

OA

i

0.2

5

10

15

20

25

30

% Nd203 In 0.3%Pt/Nd~O~/Si-A!

Fig. 8 Pt dispersion on Nd203-modified SIO2-A120 3, 0.3% wt. Pt, catalysts reduced 450~

As an additional probe of metal activity, we monitored benzene hydrogenation activity. As seen in Figure 9, Pt-containing rare earth catalysts have lower hydrogenation activity than chlorided alumina catalysts; this result reflects inhibition of metal activity on these supports relative to conventional transitional alumina supports. Whereas the acid strength can be adjusted close to that of chlorided and flourided aluminas, metal activity is somewhat inhibited on these catalysts relative to halided aluminas. This inhibition is not due to dispersion, and perhaps indicates a SMSI interaction between Pt and the dispersed Nd20 3 phase. 100 ., 80~0.3%Pt/0.9%CI/AI

60 ~9 40

"

)r

i

.3%Pt/25%Nd20~Si-A1

~ 20 ~-

: 7%Nd2Ob/Si.Ai

100

120

140

160 ('C)

180

200

Temperature

Fig. 9. Benzene hydrogenation on Nd203-modified SIO2-A120 3 (100~ kPa; 25 WHSV).

12/1 H2/C6H 6, 700

We compared Pt-containing rare-earth Si-AI catalysts with 0.3% Pt/0.9% C1/AI203 under n-heptane reforming conditions. The direct metal catalyzed dehydrocyclization of n-C 7 produces toluene, a desirable product. The formation of methyl hexanes proceeds primarily through a bifunctional mechanism. Both metal and acid sites can crack heptane or methyl hexanes with CH 4, C2H 6, and n-C4H10 originating from metal hydrogenolysis activity, while i-C4H10 arises principally from acid catalyzed beta-scission reactions. Small quantities of propane can form from

570 either metal or acid catalyzed reactions. As seen in Table 2, total conversion under identical conditions is about two times higher over 0.3% Pt/0.9%C1/AI20 3 than over 0.3% Pt/25%Nd203/Si-AI catalysts. Table 2. Heptane reforming at 175 psi, 6/1 H ~/n-C 7, 500~ 11 WHSV, 10 h on stream. toluene C5methyl hexane total selectivity (%) selectivity (%) selectivity (%) conversion

(%) 0.3%Pt/0.9%Ci)A1203 0.3 %Pt/25.%Nd~O3/Si-AI 0.6%Pt/Si-AI

79 44 53

27 11

4

18 28 78

.,.

51 57 15

,.

This is surprising in light of the close matching of acid strengths suggested by the 2MP2 reaction test, the comparable support surface areas, and similar platinum dispersions. Selectivites also differ dramatically in this reaction. Toluene selectivity is lower on rare earth Si-A1 than on chlorided alumina, with the rare earth catalyst preferentially forming C 7 isomers. The rare earth oxide catalyst shows slightly lower cracking selectivity than chlorided alumina and substantially lower cracking selectivity than Pt/Si-A1. These selectivity differences, indicate an inhibited metalcatalyzed dehydrocyclization activity, and when coupled with the lower observed benzene hydrogenation activity suggest that metal activity on the rare earth silica-alumina catalysts is inhibited relative to halided alumina catalysts. Apparently, platinum can readily dehydrogenate/hydrogenate n-C 7 on rare earth silica-alumina catalysts, although it does the more demanding direct dehydrocyclization poorly. The inhibition of metal activity relative to halided alumina catalysts limits direct metalcatalyzed dehydrocyclization reactions during heptane reforming but should not interfere with hydroisomerization reactions. Direct paraffin reactions require stronger acidity than olefin reactions, since paraffins are weaker bases. In classical bifunctional reactions involving paraffins, the metal initiates the reaction by dehydrogenating the paraffin to a more reactive olefin. This olefm reacts on an acid site to form a carbenium ion which in turn rearranges (isomerizes), and then decomposes (via proton loss) to an isomerized olefin. The latter is easily rehydrogenated to an isomerized paraffin over the metal. The bifunctional sequence normally has olefin isomerization on the acid site as the rate determining step, which gives an overall negative 1st order H 2 rate dependence. The acid strength and metal site activity must be "balanced" to avoid excessive metal or acid cracking. Alternative pathways found on stronger acid sites which involve intermolecular hydride transfer to heal carbocations producing isomefized paraffins (10) do not occur to any extent on any of the intermediate strength acids discussed here. We compared Pt/silica-alumina, yttria-modified silica-alumina, and fluorided alumina for n-C 12 isomefization. Not surprisingly, increasing yttria content lowers catalyst activity at a fLxed space velocity (Fig. 10). The 9% Y203/Si-AI catalyst compares closely to the l%F/AI203 catalyst in activity. Of the catalysts evaluated here, the 9%Y203-1oaded Si-Al had higher isomefization selectivity at equal conversion (Fig. 11).

571

8o

p'

0.02

-

"

o.o5

-o.~J

o.il

o.'14

I/Wl/SV

Fig. 10. n-C12 conversion over 0.3% Pt on silica-alumina, yttria-modified silica-alumina and FAI20 3 at 100 psi, 12"1 H2/n-C12, 350~

-

8s

O...,,

65

55

,

9 20

, 30

-

. 40

.

. 50

. 60

.

%n-~] Converslen

70

80

90

Fig. 11. i-C12 selectivity over 0.3% Pt on silica-alumina, yttria-modified silica-alumina and FAI20 3 at 100 psi, 12"1 H2/n-C12, 350~ CONCLUSIONS The addition of rare earth oxides to amorphous silica-alumina produces a modified support with a dispersed rare earth oxide surface phase and with acid properties similar to either chlorided or fluorided aluminas. X-ray diffraction does not show rare earth oxide crystallites (>30,~) up to 30% wt. rare earth oxide loadings. The ESCA (rare earth)/(Si + A1) signal increases linearly with increasing rare-earth oxide loading, strongly suggesting that the rare earth oxide continues to spread on the surface (i.e. remain below monolayer coverage) up to 30% wt. loading. The surface area and pore volume in the support decrease, but the average pore diameter increases as the dispersed rare-earth oxide blocks smaller pores, making a fraction of the surface inaccessible to physisorbed N 2. The olefin model compound reactions show that acid strength can be adjusted close to that of chlorided and fluorided aluminas with the addition of 520% wt. of rare earth oxides to the amorphous SiO2-A120 3 support. During benzene hydrogenation and n-C 7 reforming tests, metal activity on platinized rare earth oxide-modified silica-alumina remains inhibited relative to platinized halided-alumina catalysts, even at high rare earth oxide loadings, where the measured H 2 chemisorption and thereby the assumed dispersion equal that of halided aluminas. This reduced metal activity limits the direct metal-catalyzed dehydrocyclization reactions during heptane reforming, but facile platinum catalyzeddehydrogenation/hydrogenation reactions are not disturbed, so that hydroisomerization readily proceeds.

572 ACKNOWLEDGMENTS We thank Lenny Yacullo and John Ziemiak for their experimental assistance. REFERENCES 1. Misono, M. and Okuhara, T., Chemtech, 23-29 (1993). 2. Thomas, J.M., Scient. Amer., 266 (4) 112-118 (1992). 3. Arata, K., Trends Phys. Chem., 2 1-24 (1991). 4. Washington Bulletin, National Petroleum Refiners Association, Feb. 18, 1994. 5. Degnan, T.F. and Kennedy, C. R., AICHE Journal 39(4), 607-11 (1993). 6. Sinfelt, J. H., Carter, J.C., and Yates, D.S.C., J. Catal. 2..44,283-9 (1972). 7. Kramer, G.M. and McVicker, G.B., Acct. Chem. Res. 19, 78-86 (1986). 8. Maitra, A.M., Appl. Catal., A 85(1), 27-46, (1992). 9. Maitra, A.M., J. Therm. Anal. 36(2), 657-75, (1990). 10. Iglesia, E., Soled, S. and Kramer, G., J. Cat. 144, 238-45 (1993).

J.W. Hightower, W.N. Delgass, E. Iglesia and A.T. Bell (Eds.) 11th International Congress on Catalysis - 40th Anniversary

Studies in Surface Science and Catalysis, Vol. 101 9 1996 Elsevier Science B.V. All rights reserved.

573

W h a t N M R Has Told Us about Solid Acidity J. F. Haw a and J. B. Nicholas b aChemistry Department; Texas A&M University; College Station, TX 77843 USA bEnvironmental Molecular Sciences Laboratory; Pacific Northwest Laboratory; Richland, WA 99352 USA A wide variety of NMR methods are being applied to understand solid acids including zeolites and metal halides. Proton NMR is useful for characterizing Brensted sites in zeolites. Many nuclei are suitable for the study of probe molecules adsorbed directly or formed in situ as either intermediates or products. Adsorbates on metal halide powders display a rich carbenium ion chemistry. The interpretation of NMR experiments on solid acids has been greatly improved by the integration of theoretical chemistry and experiment.

1. OVERVIEW AND PERSPECTIVE In 1962 Olah reported the 13C NMR spectrum of the t-butyl cation in superacid solution, [ 1] and NMR was thenceforth the experimental method of choice for studies of intermediates in solution acid chemistry. The inhomogeneous nature and diversity of solid acid systems will ensure that no one experimental technique will so completely dominate as NMR has in solution studies, but the contributions and potential of NMR to solid acid studies are clearly such as to put it on an equal footing with reaction studies, infrared, TPD, diffraction methods and calorimetry. NMR of solids is a very diverse collection of methods, and the practice of applying it to chemisorption is changing dramatically as a result of advances in theoretical chemistry including reliable chemical shift calculations. The wide application of NMR to solid state chemistry grew out of the revival of magic angle spinning in the mid 1970's. This line narrowing technique is very effective for dilute or low ~,, spin 1/2 systems including most examples involving 13C, 31 p, 19F, 29Si and 15N as well as 1H at the chemical concentrations found in typical solid oxides including zeolites. Magic angle spinning is also partially to highly effective in narrowing the central transition for many quadrupolar nuclei including 27A1. More sophisticated line narrowing techniques involving more complicated mechanical averaging or elegant application of multil?le quantum evolution have extended the practice of high resolution solid state NMR to include 170, 23Na and many other nuclei that may sometimes be of interest in studies of catalytic materials. In addition to sample rotation, a particular solid state NMR experiment is further characterized by the pulse sequence used. As in solution NMR, a multitude of such sequences exist for solids; many exploit through-space dipolar couplings for either signal enhancement, spectral assignment, internuclear distance determination or full correlation of the spectra of different nuclei. The most commonly applied solid state NMR experiments are concerned with the measurement of spectra in which intensities relate to the numbers of spins in different environments and the resonance frequencies are dominated by isotropic chemical shifts, much like NMR spectra of solutions. Even so, there is considerable room for useful elaboration; the observed signal may be obtained by direct excitation, cross polarization from other nuclei or other means, and irradiation may be applied during observation or in echo periods prior to

574 observation to selectively remove or reintroduce dipolar interactions. Molecular dynamics on very diverse time scales also reveal themselves to suitable NMR techniques probing line shapes and relaxation times. Other kinds of NMR experiments applied to solid acids and other catalysts have little resemblance to the more familiar forms of spectroscopy. For example, diffusion in a magnetic field gradient interferes with refocusing of spin echoes, and this is applied in diffusion measurements. Other applications of gradients lead to a variety of imagining experiments which can be used to profile macroscopic inhomogeneities. Thus, the NMR characterization of a solid acid can range from, for example, obtaining a simple proton or silicon MAS spectrum using a routine method to an elaborate negotiation between the spectroscopist and the spi~ systems naturally contained within or introduced to probe the acid. The reader is referred the recent book by Bell and Pines [2] for a more complete overview of the various methods and objectives in NMR studies of solid acids and other heterogeneous catalysis. In the present contribution we illustrate the application of 1H, 13C and 19F MAS NMR to two archetypal solid acids, BrCnsted sites in zeolites and solid metal halides such as aluminum chloride and bromide powders which exhibit "Lewis superacidity". An important characteristic of the more recent work is the integration of quantum chemical calculations into the design and interpretation of the NMR experiments.

2.

P R O T O N NMR OF ZEOLITE BRONSTED SITES IN HZSM-5

We first illustrate the application of NMR to solid acids with some purely experimental results and then motivate theoretical methods and the integration of theory and experiment. Figure 1 shows 1H MAS spectra of a dehydrated sample of zeolite HZSM-5. Spectra obtained in the vicinity of room temperature seem to suggest no more than two isotropic signals as well as a series of spinning sidebands. The isotropic peak at 2.0 ppm is due to silanols, while that at 4.3 ppm to BrCnsted sites. The apparently singular isotropic shift for the Br~nsted sites in this material was long used to argue that all such sites are "identical", and indeed many non spectroscopic methods support the idea that the acid strengths are uniform in HZSM-5. Carefully inspecting the downfield (high shift) side of the 4.3 ppm resonance in the spectrum obtained at 296 K suggests a very broad line shape component. When spectra are run at much lower temperatures, this sharpens into an additional isotropic peak at 6.9 ppm. We used 1H{27Al } double resonance experiments to demonstrate that both the 6.9 and 4.3 ppm signals have strong dipolar coupling to aluminum; thus, it would appear that there are at least two NMR distinguishable Brcnsted sites in HZSM-5 [3]. The exact nature of the new site is under investigation in several laboratories. Figure 1. Proton MAS NMR spectra of dehydrated HZSM-5 acquired over a range of temperatures. Note in particular the broad signal at 6.9 ppm which is resolved at lower temperatures.

!''"

25

I~ : " 1 : : ' ' 1 ' : : : 1 ' " ' 1 ~

20

15

10

::':

5

0

-5

!::::1

-10 -15 ppm

575

3.

ZEOLITE PROBE MOLECULES AND REACTION INTERMEDIATES

It is often said that the property of acidity is manifest only in the presence of a base, and NMR studies of probe molecules became common following studies of amines by Ellis [4] and Maciel [5, 6] and phosphines by Lunsford [7] in the early to mid 80s. More recently, the maturation of variable temperature MAS NMR has permitted the study of reactive probe molecules which are revealing not only in themselves but also in the intermediates and products that they form on the solid acid. We carried out detailed studies of aldol reactions in zeolites beginning with the early 1993 report of the synthesis of crotonaldehyde from acetaldehyde in H Z S M - 5 [8] and continuing through investigations of acetone, cyclopentanone [9] and propanal [ 10]. The formation of mesityl oxide 1, from dimerization and dehydration of

H

1

2

acetone, was a very useful observation, because Farcasiu had recently proposed the use of the 13C shifts of this molecule as a measure of liquid acid strength [ 11 ]. In solution, the 13C shifts of I reflect the protonation equilibrium and the contributions of the resonance structures for the protonated ketone 2. In the acidic zeolites, acetone-2-13C exhibits a downfield shift (that is not nearly so large as that seen in 100 % H2SO4 solution), and 1 gives shifts that can be interpreted with Farcasiu's solution measurements to suggest that HZSM-5 has an effect on mesityl oxide that is similar to the effect of ca. 70 % H2SO4 [9]. Indeed, a large number of purely experimental observations of carbenium ions and related species underscore the fact that zeolites are strong Brcnsted acids but not superacids. Some of the more interesting "probe molecules" studied on solid acids by NMR have been carbenium ions generated in situ. These include the indanyl carbenium ion formed from the reaction of styrene dimers in HZSM-5 [ 12]. Although this result convincingly demonstrated the formation of a free cation as persistent intermediate species, the high stability (low acid strength) of this cation suggests that results like this are exceptional. 4.

THEORY, EXPERIMENT

AND H A M M E T T I N D I C A T O R S

The probe molecules of greatest historical interest in catalysis are the Hammett indicators [ 13]. The difficulty of making reliable visual or spectrophotometric observations of the state of protonation of these species on solids is well known. We have recently carried out the first NMR studies of Hammett indicators on solid acids [14]. This was also the occasion of the first detailed collaboration between the authors of this article, and theoretical methods proved to strongly compliment the NMR experiments. The Hammett story is told after first reviewing the application of theoretical chemistry to such problems. Central to the application of any physical method in chemistry is the process of modeling the relationship between the observables and molecular structure. However often one does this, it is rarely an exact process. One can rationalize almost any trend in isotropic chemical shift as a function of some variation in molecular structure - after the fact, but the quantitative prediction of such trends in advance defies intuition in most nontrivial cases. Even though the NMR spectrum is a function

576 of molecular structure, molecular structure is not a function of the NMR spectrum. Even if one correctly "assigns" spectral data to a gross chemical structure, the proposed structure rarely specifies exact bond lengths and angles, and the testable predictions that one makes from it are necessarily qualitative. Theoretical chemistry provides the means to model the relationship between structure and spectra, and this modeling process is revolutionizing the practice of interpreting NMR spectra. One strategy is as follows. 1. NMR experiments are interpreted to suggest a qualitative model of the chemisorbed state, the structure of an intermediate or some other finding. 2. The appropriate level of theory is then used to test the reasonableness of this finding as well as alternate explanations. 3. Finally one calculates spectroscopic observables (e.g., chemical shift tensors, vibrational frequencies) for the refined structure and compares to experiment. Agreement between the values of a variety of observables, determined by both theoretical and experimental means, provides powerful verification of our interpretation. Clearly, the integration of experiment and theory is much preferable to "hand-waving arguments" for the assignment of unusual chemical shifts to exotic species, as has often been done. There are many considerations that critically affect the accuracy of a quantum mechanical calculation, the most obvious of which are the degree to which the atomic coordinates of the theoretical model relate to the experimentally observed compound, the flexibility of the basis set and the extent to which electron correlation is accounted for [ 15]. For chemisorption studies in zeolites, cluster models of appropriate size have been used to elucidate a wide range of zeolite chemistry [ 16]. The perturbation theory of Mr and Plesset is the most straightforward way of improving a Hartree-Fock calculation; such treatments truncated at second order (MP2) usually provide much of the correlation effect at a reasonable computational cost. As an alternate means of including electron correlation, density functional theory (DFT) methods have been demonstrated to give results comparable to high level postHartree-Fock calculations with considerably less computational cost. Recent studies have validated the DFT approach in the computation of geometries, vibrational frequencies, energetics, and other molecular properties. More important to catalytic applications, DFT has also been shown to give good results for reaction barriers and transition state geometries in proton transfer reactions. Reliable NMR chemical shift calculations are now possible with a variety of means, including the IGLO, LORG, GIAO approaches, and with the inclusion of electron correlation by either MF2 or density functional methods. We used DFT to optimize the geometries of various Hammett bases on cluster models of zeolite Br~nsted sites. For p-fluoronitrobenzene and p-nitrotoluene, two indicators with strengths of ca. -12 for their conjugate acids, we saw no protonation in the energy minimized structures. Similar calculations using the much more strongly basic aniline analogs of these molecules demonstrated proton transfer from the zeolite cluster to the base. We camed out 19F and 15N experimental NMR studies of these same Hammett indicators adsorbed into zeolites HY and HZSM-5. Figure 2 compares the results of theory and experiment for the specific case of pfluoronitrobenzene. Inspection of the calculated structure shows that the proton is still on the zeolite, and the 19F shifts are more like chloroform solution than superacid solution. Furthermore, when the 19F chemical shift was calculated for the theoretical structure, it was found to agree with the experimental result.

577

Figure 2. The Hammett indicator p-fluoronitrobenzene on the BrCnsted site of zeolites. Left Theoretical structure calculated with DFT at the BLYP/DNP level. Right - Experimental 19F MAS NMR spectra on HY (top) and HZSM-5 (bottom). 5.

N M R S T U D I E S OF M E T A L H A L I D E SOLID ACIDS

We have discovered that a wide variety of carbenium ions and related electrophilic species can be prepared by the direct contact of suitable precursors with metal halide powders at reasonable temperatures and in the absence of solvent, permitting direct observation of reaction chemistry and measurement of the principal components of the 13 C chemical shift tensors [ 17]. The Friedel-Crafts alkylation and acylation reactions, first reported in 1877, are among the most important in chemistry. The generalized reaction, as surveyed in Olah's books [18, 19] proceeds on a wide variety of catalysts; the most familiar is AIC13, but other metal halides including ZnCI2 are valued for their lower activity and sometimes greater selectivity. Metal halide powders or molten salts are also employed as catalyst components in several industrial processes. We readily prepared the t-butyl 3 and acylium 4 cations on aluminum chloride powder from the corresponding halides. 3 and 4 are, respectively, the archetypal Friedel-Crafts alkylation and acyclation intermediates. From the 13C spectra of these and other cations, we were able to measure the principal components of the chemical shift tensors. The isotropic shift in solution is the average of the three principal components; it should be clear

i /C+

H

H3C

3

m C+

~CH 3

I 4

578 that correlations between structure and spectra should be far more meaningful when made for the principal components instead of their average alone. We have also been carrying out extensive theoretical calculations of the structures and chemical shift tensors of the cations observed in the experimental studies. As we develop greater confidence in the methodology we will be able to explore subtle effects due to ion pairing and other interactions with the media. This work is also valuable in that it assists our ongoing efforts to understand the nature of electropMlic species in zeolite solid acids. Even acetone shows appreciable shifts following contact with metal halides - the magnitude of which correlates with the expected strength of the Lewis acid-base interaction.

~.. 231 ppm

J"

I

227 ppm

Figure 3. 75.4-MHz 13C NMR spectra of a c e t o n e - 2 - 1 3 C on various substrates. The isotropic shift is strongly dependent on the strength of the Lewis acid. Note that ZnCI2 and zeolite ZnY yield the same shifts.

Zn !

231 ppm ZnY zeolite C P, 198 K . ~.4-

245 ppm

Alll 3 I '''

400

i l ' '""'I 'I

350

" ''

300

'"I ' ' "'

250

I "'""l"'""i'""i'"~""l

200

150

' ''"i'l

100

' '''

50

I

0

ppm Figure 3 shows 13C MAS spectra of acetone-2 -13C on various materials. Two isotropic peaks at 231 and 227 ppm were observed for acetone on ZnC12 powder, and appreciable chemical shift anisotropy was reflected in the sideband patterns at 193 K. The 231 ppm peak was in complete agreement with the shift observed for acetone diffused into ZnY zeolite. A much greater shift, 245 ppm, was observed on A1CI3 powder. For comparison, acetone has chemical shifts of 205 ppm in CDC13 solution, 244 ppm in concentrated H2SO4 and 249 ppm in superacid solutions. The resonance structures 5 for acetone on metal halide salts underscore the similarity of the acetone complex to carbenium ions. The relative contributions of the two canonical forms rationalizes the dependence of the observed isotropic 13C shift on the Lewis acidity of the metal halide.

579

CH 3

H

T

H3

H3 C

~O

.

.

.

n+ x

Mn+X n

m

n

5 There is good reason to believe that the potential of NMR studies of carbenium ions on solid metal halides exceeds that of corresponding studies in superacid solutions. Of course the advantages of working in solids include the possibility of very low temperatures and the mass transport restrictions of frozen media. Thus, Mehre and Yannoni were able to characterize the sec-butyl cation in frozen SbF5 by NMR [20] and Schleyer and coworkers have obtained infrared evidence of the allyl cation in the same medium [21 ]. So far, we have been successful in every case in which we have tried to duplicate known solution carbenium ion chemistry on appropriate metal halide solid acids. For example, we have recently prepared a number of acomplexes by the reaction or alkyl halides on aluminum halides. Figure 4 shows the case for the formation of the toluenium -13C7 ion. 32 ppm

O

+ CH3Br

139 ppm AIBr 3 "~

m~. 201 ppm AI2Br7"

178 pp n ~ 50ppm 32 ppm 178ppm

201ppm | 213 K I ''

250

~ ''

139ppm

Figure 4. 75.4 MHz 13 C NMR spectrum of benzene13C 6 and bromomethane13C reacting on A1Br3 powder. The spectrum shows the formation of t o l u e n i u m - 1 3 C 7 ion as indicated in the scheme. The resonance at 129 ppm is due to an excess of benzene- 13C6"

5

Q I'

200

'''

I'''

150

'I''

100

''

I'

50

'''I

0

ppm

REFERENCES .

2. .

G. A. Olah, Angew. Chem., Int. Ed. Engl., 34 (1995) 1393. A. Bell and A. Pines (eds.), NMR Techniques in Catalysis, Marcel Dekker: New York, 1994. L. W. Beck, J. L. White and J. F. Haw, J. Am. Chem. Soc., 116 (1994) 9657.

580

o

.

.

.

8. 9. 10. 11. 12. 13. 14. 15. 16. 17. 18. 19. 20. 21.

W. H. Dawson, S. W. Kaiser, P. D. Ellis, and R. R. Inners, J. Am. Chem. Soc., 103 (1981) 6780. G. E. Maciel, J. F. Haw, I.-S. Chuang, B. L. Hawkins, T. E. Early, D. R. McKay and L. Petrakis, J. Am. Chem. Soc., 105 (1983) 5529. J. F. Haw, I.-S. Chuang, B. L. Hawkins and G. E. Maciel, J. Am. Chem. Soc., 105 (1983) 7206. P. Chu, D. D. Mallmann, and J. H. Lunsford, J. Phys. Chem., 95 (1991) 7362. E. J. Munson, and J. F. Haw, Angew. Chem., 105 (1993) 643. T. Xu, E. J. Munson, and J. F. Haw, J. Am. Chem. Soc., 116 (1994) 1962. T. Xu, J. Zhang, E, J. Munson and J. F. Haw, J. Chem. Soc. Chem. Comm. (1994) 2733. D. Farcasiu and A. Ghenciu, J. Am. Chem. Soc., 115 (1993) 10901. T. Xu and J. F. Haw, J. Am. Chem. Soc., 116 (1994) 10188. L. P. Hammett and A. J. Deyrup, J. Am. Chem. Soc., 54 (1932) 2721. J. F. Haw, J. B. Nicholas, L. W. Beck, T. R. Krawietz, and D. B. Ferguson, J. Am. Chem. Soc., in press. W. J. Hehre, L. Radom, P. v. R. Schleyer, and J. A. Pople, Ab Initio Molecular Orbital Theory; Wiley & Sons: New York, 1986. J. Saurer, Chem. Rev., 89 (1989) 199. T. Xu, P. D. Torres, L. W. Beck and J. F. Haw, J. Am. Chem. Soc., 117 (1995) 8027. G. A. Olah, Friedel-Crafts Chemistry; Wiley & Sons: New York, 1973. G. A. Olah and/~. Moln~, Hydrocarbon Chemistry; Wiley & Sons: New York, 1995. P. C. Myhre and C. S. Yannoni, J. Am. Chem. Soc., 103 (1981) 230. P. Buzek, P. v. R. Schleyer, H. Vancik, Z. Mihalic, and J. Gauss, Angew. Chem., Int. Ed. Engl., 33 (1994) 448.

J.W. Hightower, W.N. Delgass, E. Iglesia and A.T. Bell (Eds.) I l th International Congress on Catalysis - 40th Anniversary

Studies in Surface Science and Catalysis, Vol. 101 9 1996 Elsevier Science B.V. All rights reserved.

581

N o v e l M i c r o p o r o u s Solid "Superacids": C s x H 3 - x P W 1 2 0 4 0 (2 < x < 3 ) t T. Okuhara a, T. Nishimurab,$, and M. Misonob aGraduate School of Environmental Earth Science, Hokkaido University, Sapporo 060, Japan bDepartment of Applied Chemistry, Graduate School of Engineering, The University of Tokyo, Bunkyo-ku, Tokyo 113, Japan

It was confirmed that pore size of the acidic Cs salts (CsxH3-xPWl2040) was controlled by the Cs content. For example, Cs2.5H0.5PW12040 had mesopores and micropores larger than 8.5 /~ in diameter, and C s 2 . 2 H 0 . s P W 1 2 0 4 0 micropores of about 7/~. Shape selective catalysis was clearly observed for the latter catalyst. The acidic Cs salts as well as H3PW12040 possess very strong acidity and were more catalytically active than zeolites and SO42-/ZRO2 for decomposition of esters and alkylation in liquid-solid reaction system. 1. I N T R O D U C T I O N Heteropolyacids are good catalysts for acid-catalyzed reactions and have been used in large-scale catalytic processes [1-5]. The formation of Cs salts brings about high surface area [6-9] and high catalytic performance for both acidcatalyzed reactions [9,10] and oxidation reactions [11,12]. We indicated previously that H 3 P W 1 2 0 4 0 and its acidic Cs salt possess "superacidity", as far as the H a m m e t t indicator test and temperature programmed desorption of NH3 are concerned [13]. Recent data of calorimetry of NH3 absorption by Lefevre et al. [14] demonstrate that H3PW12040 is a very strong acid. Excellent catalytic performances of H3PW12040 and its Cs salts, especially Cs2.5H0.5PW12040, have been shown for acid-catalyzed reactions such as direct decomposition of ester, alkylation [13], acylation [15], skeletal isomerizati0n [16], etc. tCatalysis by Heteropoly Compounds. XXVII. Part XXVI: S. Shikata, T. Okuhara, and M. Misono, J. Mol. Catal., 100 (1995) 49. SPresent address: Mitsui Toatsu Chemicals, Inc., Kasama-cho, Sakae-ku, Yokohama 247, Japan.

582 Shape selective catalysis as typically demonstrated by zeolites is of great interest from scientific as well as industrial viewpoint [17]. However, the application of zeolites to organic reactions in a liquid-solid system is very limited, because of insufficient acid strength and slow diffusion of reactant molecules in small pores. We reported preliminarily t h a t the microporous Cs salts of H 3 P W 1 2 0 4 0 exhibit shape selectivity in a liquid-solid system [18]. Here we studied in more detail the acidity, micropore structure and catalytic activity of the Cs salts and wish to report that the acidic Cs salts exhibit efficient shape selective catalysis toward decomposition of esters, dehydration of alcohol, and alkylation of aromatic compound in liquid-solid system. The results were discussed in relation to the shape selective adsorption and the acidic properties. 2. E X P E R I M E N T A L Acidic Cs salts of H3PW12040, CsxH3-xPWl2040 (abbreviated as Csx), have been prepared by titration of the aqueous solution of H 3 P W 1 2 0 4 0 with an aqueous solution of Cs2CO3 [13]. To the solution of H3PW12040 (8 x 10 -2 tool din'3), the appropriate amount of aqueous solution of Cs2CO3 (Cs: 2.0 x 10 -1 tool dm "3) was added dropwise at a rate of about 1 ml rain "1 at room temperature. The resulting white colloidal solution was allowed to stand overnight at room temperature and then was evaporated at 318 K to the solid. An isotherm of N2 adsorption and surface area were obtained by using a N2 adsorption system (Micromeritics ASAP 2000). The adsorption capacities for various molecules having different molecular size were measured by microbalance (Shimadzu TG-30) connected directly to an ultrahigh vacuum system [18,19]. The adsorption of isopropylacetate or cyclohexylacetate was measured in a liquid-solid reactor at 303 K using decane as solvent. The shape selective catalysis was examined by choosing five kinds of the reactions, Eqs. (1) - (5), that is, dehydration of 2-hexanol, decompositions of three kinds of esters and alkylation of 1,3,5-trimethylbenzene with cyclohexene. Catalytic reactions were performed in a three-neck flask (about 100 ml) after the catalyst (about 100 rag) was pretreated in a He flow at 573 K. Reactions were

/vr,

Hexenes

+

H20

(I)

--OCOCH3

,/%

+

CH3COOH

(2)

--OCOCH3

Butenes

+

CH3COOH

(3)

+

CH3COOH

(4)

OH

H3PW12040 - Cs2.5H0.5PW12040 > HZSM-5 > SIO2-A1203. This order is consistent with the results of Hammett indicator test. Table 1 Acid strength of various solid acids Solid acid HO a - AHads(NH3) b Tdes(NH3) c Ref. /kJ mol-1 /K H3PW12040 -13.16 195 -850 13, 14 C s2.5H0.5PW 12040 - 13.16 -830 13 SO42"/ZRO2 -14.52 165 - 1000 e 13, 21 (190) d HZSM-5 -12.70 150 -670 13, 20 Si02-A1203 -12.70 145 -600 13, 22 aAcidity function measured by Hammett indicator test. bInitial heat of NH3 adsorption (absorption). CDesorption temperature of NH3; the figures show the temperature of the desorption peak at the highest temperature, dAt less than 5 ~tmol g-1 of NH3 adsorbed, eDesorbed as N2 by the reaction with catalyst [13]. -

584 Table 2 demonstrates the high catalytic performance of Cs2.5. For the decomposition of cyclohexylacetate (Reaction (4)), the activity (per unit weight) of Cs2.5 is about 40 and 130 times as high as SO42"/ZRO2 and HZSM-5, respectively, and furthermore it is more active than H2SO4 by a factor of two orders of magnitude. Cs2.5 is also highly active for the alkylation of 1,3,5trimethylbenzene with cyclohexene (Reaction (5)). The high activity of Cs2.5 has been inferred to be due to the high acidity (the acid strength and amount) and the soi~ basicity of the polyanion [13]. Table 2 Catalytic activities of various acids for decomposition of cyclohexylacetate and aUTlation of 1,3,5-trimethylbenzene in liquid-solid reaction system Catalyst Surface area Reaction rates a m 2 ~-1 Reaction (4) b Reaction (5) c Cs2.5H0.sPW12040 135 130 58 H3PW12040 5 50 18 SO42"/ZRO2 93 "3 5 HZSM-5 332 1 < 0.5 SIO2-A1203 458 0.6 1 H2SO4 ammol g-1 h-1. b a t 373 K. CAt 343 K.

1

3

Figure 1 shows more detailed data of the change of the surface area as a function of the Cs content, x in CsxH3-xPW12040. The surface area of the acid form (x = 0), 5 m 2 g-l, decreased as x increased to 2 (the surface area of Cs2 was only 0.5 m 2 g-l). However, the surface area greatly increased as x exceeded 2 and

Figure 1. Surface area and surface acidity of CsxH3-xPWl2040.

585 became higher t h a n 130 m 2 g-1 when x > 2.5. This change was in contrast with the acidic Na salts; the surface area decreased monotonically as the Na content increased [23]. The acid amount on the surface, which is called surface acidity hereafter, was estimated from the surface area and the formal concentration of proton attached to polyanion. Here, the number of polyardons was calculated by assuming t h a t the hemisphere of each polyanion (11 A in diameter) contributes to the surface area. The number of protons was then estimated on the assumption t h a t each polyardon on the surface possesses (3 - x)/2 protons. Since solid state NMR revealed t h a t protons of the acidic Cs salts distribute almost uniformly t h r o u g h the whole bulk [24], this estimation of the" surface acidity m a y be reasonable. The surface acidity thus estimated is shown in Figure 1. The surface acidity decreases at first with the Cs content, but sharply increases when x exceeds 2. The maximum appeared at x =2.5. In Figure 2, the catalytic activities of Csx ( x = 0, 2.2, 2.3, 2.5, 2.7, 2.9 and 3) for the decomposition of isopropylacetate (Reaction (2)) are plotted a g a i n s t the surface acidity. A good correlation between the activity and the surface acidity was observed for these acidic Cs salts. As described above, the acid s t r e n g t h of Cs2.5 was very similar to that of H 3 P W 1 2 0 4 0 , indicating t h a t the effect of the formation of acidic Cs salts little influences the acid strength. Therefore, it is reasonable that the catalytic activity of Csx for this reaction is proportional to the surface acidity, that is, the number of protons on the surface. HZSM-5 and SO42/ZrO2 were almost inactive for this reaction under the same reaction conditions. ._

80

0

c~ loo

v

Cs2.2

~ L-

~" b~

v-

.5

.7

,-

. mo.

_

20

80 60 40

.

, 40

.

, 60

8O

Figure 2. Catalytic activities of CsxH3-xPW12040 for decomposition of isopropylacetate as a function of the surface acidity. The reaction was carried out at 373 K in liquid-solid reaction system.

6 0 x: 0

E

tt~

v

I

C

9 9

0

40 E

-

I -

o

Surface Acidity/llmol g-1

SiO2/AI20 3 > H-MontmoriUonite >> y-AI20 3 [6]. 3.2. Nature of the Active Acid Sites High resolution XPS Nls analysis of adsorbed pyridine provides a surface sensitive tool capable of distinguishing Lewis and BrOmted acid centers. Refinement of XPS intensities by application of photoelectron cross-sections and inelastic mean free path allows for quantitative evaluation of surface acid centers. Adsorption of pyridine on a Lewis acid site increases the Nls binding energy from the free pyridine value of 398.0 eV [7] to ~400 eV due to the increase of the positive charge on the N atom in proximity of cationic Lewis acid centers. Adsorption on BrOnsted acid sites results in a larger shift of the Nls binding energy to about 401.5 eV due to the formation of a pyridinium ion.

604 Reactions exemplified above were found to ~ on catalysts containing strong Brcnsted acid centers but not on weak Lewis acid centers such as those on zirconia. Only Brr acid centers were found on sulfonic acid resin catalysts, e.g. Nafion-H. Poisoni-~ the Br~nsted centers by incremental exchange of the protons with K + yielded after full exchange an inactive catalyst for alcohol coupling, as well as for alcohol dehydration, where quantitative XPS demonstrated there was one K § acid group. Sulfate-doped zirconia was found to have both Lewis and Brr acid sites on the surface, where the sulfate groups were the carder of the protonic site. The ratio of Lewis to Brcnsted sites was found to vary depending on heat treatment, but the percentage of sulfate groups associated with a proton was found to be constant at 12-17%. 3.3. Oxygen Retention/Rejection During Ether Synthesis Over Resin Catalysts *SO-labelling studies were carried out to determine if there were a common intermediate to these two ethers. Experiments were carried out with *SO-methanolp60isobutanol = 1.0/3.2 and *SO-ethanolp60-isobutanol = 1.0/5.0 reactant mixtures such that conversion levels were 98% of 1SO from the starting 1SO-ethanol, indicating that no isotope scrambling occurred. Data in Table 4 demonstrate that 1SO was retained in the mixed ether and ethanol attack of the acidactivated 2-pentanol via an axial SN2 rear-attack was the predominant synthesis pathway. Evidently, the shape selectivity induced by the ZSM-5 zeolite channel structure (Figure 2) plays an important role in achieving the remarkably higher configuration inversion

608 than either the Nafion-H or H2SO 4. The cross coupling of ethanol and 2-pentanol could not proceed inside either channel of HZSM-5 because once the 2-pentanol is adsorbed at the acid site, there is not enough space for ethanol to attack the activated 2-pentanol at the rear of the asymmetric carbon. However, the intersection of the channels can accommodate the transition state of the coupling reaction as illustrated in Figure 2 [14]. The transition state geometry was optimized by using the Spartan program at the R t ~ / S T O - 3 G level with the constraints that the O-C-O angle at the activated asymmetric carbon atom was linear and the two C-O bond distances (represented by dashed lines between the asymmetric carbon atom and the oxygen atoms from the attacking ethanol and the leaving water molecule) were 2.0 ,/~ This result indicates that acid sites that catalyze the dehydrative coupling of ethanol and 2-pentanol are located at the intersection of the two channel systems.

F'~mre 2. Transition state complex in the ethanol + 2-pentanol SN2 reaction activated by the proton at the channel intersection of HT~M-5 [14]. The zeolite pore structure is represented as a wire-frame section of the intersecting channels produced by the MAPLE V software package. The zeolite proton that activates the 2-pentanol molecule is marked with *.

In contrast, the Nafion-H catalyst bears a similarity to acid solution, very likely due to its flexible backbone carrying the sulfonic acid groups. The OH group of 2-pentanol is the preferred leaving group, after being activated by the surface H + and subjected to concerted nucleophilic attack by the light alcohol. The minor non-inversion path (23% over Nafion-H and 3% over HT_~M-5) can be accounted for by a less efficient carbenium ion (C +) or olefin (C z) intermediate mechanism. This minor path is corroborated by the observation of the 3-ethoxypentane side product, which could only be formed v/a carbenium ion or olefinic intermediates, over Nafion-H and in the liquid H2SO 4. However, 3-ethoxypemane was not observed with the H Z S M - 5 catalyst, demonstrating that its formation was suppressed by shape selectivity. The 3-ethoxypentane is more branched and would pass through the HZSM-5 channel at a slower rate than the 2-isomer even if formed by carbenium ion or olefm reaction at the channel intersections. Indeed, reaction over H7_~M-5 of ethanol with 3-pentanol resulted in 73% of the product being 2ethoxypentane (as a racemic mixture, along with 27% 3-ethoxypentane), while in HeSO 4 solution, 95.6% of the C7 ether product was 3-ethoxypentane [14]. These experiments demonstrate that the surface-catalyzed SN2 reaction is far more efficient than either the C + or C- pathway for the dehydrative coupling of alcohols over the solid acid catalysts tested. High selectivity to configurationally inverted chiral ethers ensues, especially in the case of the HZSM-5 catalyst, in which the minor C + or C: paths were further suppressed by "bottling ~ of 3-ethoxypentane by the narrow zeolite channels.

609 3.6. Kinetic Characteristics of SN2 Surface-Catalyzed Reactiom All evidence presented herein points to the generality of the SN2 path for the surface acid-catalyzed coupling of alcohols to ethers. The product composition, the ~SO flow from reactants to products, and chirality inversion demonstrate convincingly that mechanistic patterns of the SN2 reaction including oxygen retention, steric hindrances due to axial rear-attack of an alcohol activated by proton attachment by a second alcohol, and the accompanying conformational changes are in place. Additional features specific to surfaces, in contrast to solutiom, are (i) kinetics that display competitive adsorption of the two reactant alcohols on polymeric sulfonic acid catalysts [15,16] and (ii) shape selectivity in zeolites [11,14]. These features of the surface-catalyz~ S ~ paths have the following important consequences. ADO): On catalysts where dual-site competitive adsorption of the reactants occurs, the rate displays an opt/mum rather that "saturation" dependence on the concentration of either reactant alcohol. A comparison between the SN2 solution and surface kinetics is featured in Figure 3 to illustrate this point. Furthermore, selectivity is significantly influenced by relative strengths of the adsorption bond of the reacting alcohols. This is exemplified by the fact that methanol/isobutanol coupling over Nafion-H gives preferentially MIBE over DME (Table 1 and Ref. 15), while the Amberlyst catalysts give large quantities of DME (Table 1 and Ref. 16). The Langmuir-Himhelwood kinetic model is successful in casting the selectivity patterns in quantitative terms, and simply accounts for the preference to MIBE by relatively stronger isobutanol-to-methanol adsorption bond strength over Nation compared to the Amberlyst resins [15,16].

Figure 3. Schematic comparison of surface and solution kinetics for the SN2 path of ROH + R ' O H =~ R O R ' . Surface SN2 kinetics was taken from Ref. 16 and the solution SN2 kinetics was taken in the form of Rate = ktA]tA'I[H +]/(1 + Ir~,tA] + I~'[A' ]). Plots are of {Rate/ k[A][A'][H*]} v s [A] with Kb = lff 3 mol "1and assuming ILo' [A' ] < < 1, with A = ROH and A ' = R ' O H .

!

7 • 8O0 0r g

f

/,.

6o0 I

~" ~.. ,oo 7

[ .I "I"

~ 2o0 ~, ._. 0 0 tr

~o

7

I 0O0

2OOO

3OOO

4OOO

Concentration of ROH while keeping R'OH constant

Ad(ii): On catalysts with pores and cavities of molecular dimensions, exemplified by mordenite and ZSM-5, shape selectivity provides constraints of the transition state on the SN2 path in either preventing axial attack as that of methyl oxonium by isobutanol in mordenite that has to "turn the comer" when switching the direction of flight through the main channel to the perpendicular attack of methyl oxonium in the side-pocket, or singling out a selective approach from several possible ones as in the chiral inversion in ethanol/2-pemanol coupling in HZ~M-5 (14). Both of these types of spatial constraints result in superior selectivities to similar reactions in solutions.

610 4. SUMMARY Conclusive evidence has been presented that surface-catalyzed coupling of alcohols to ethers proceeds predominantly by the SN2 pathway, in which product composition, oxygen retention, and chiral inversion is controlled by "competitive double parking" of reactant alcohols or by transition state shape selectivity. These two features afforded by the use of solid acid catalysts result in selectivities that are superior to solution reactions. High resolution XPS data demonstrate that Brcnsted acid centers activate the alcohols for ether synthesis over sulfonic acid resins, and the reaction conditions in zeolites indicate that BrCnsted acids are active centers therein, too. Two different shapeselectivity effects on the alcohol coupling pathway were observed herein: transition-state constraint in HZ~M-5 and reactant approach constraint in H-mordenite. None of these effects is a molecular sieving of the reactant molecules in the main zeolite channels, as both methanol and isobutanol have dimensions smaller than the main channel diameters in ZSM-5 and mordenite. REFERENCF~ 1. 2. 3. 4. 5. 6. 7. 8. 9. 10.

11. 12. 13. 14. 15. 16.

IC Klier, R.G. Herman, and C.-W. Young, Preprints, Div. Fuel Chem., ACS, 29(5) (1984) 273. J.G. Nunan, C.E. Bogden, IC Klier, KJ. Smith, C.-W. Young, and R.G. Herman, J. Catal., 116 (1989) 195. J.G. Nunan, R.G. Herman, and K. KUer, J. Catal., 116 (1989) 222. R.G. Herman, in "New Trends in CO Activation," ed. by L Guczi, Elsevier, Amsterdam (1991) 265. J. Nunan, K. Klier, and R.G. Herman, J. Chem. Soc., Chem. Commun. (1985) 676. K. Klier, R.G. Herman, M.A. Johansson, and O.C. Feeley, Preprints, Div. Fuel Chem., ACS, 37(1) (1992) 236. R. Nordberg, R.G. Albridge, T. Bergmark, U. Ericsson, J. Hedman, C. Nordling, IC Siegbahn, and B.J. Lindberg, Arkiv Kemi, 28 (1968) 257. J.L Schlenker, J.l. Pluth, and J.V. Smith, Mat. Res. Bull., 13 (1978) 901; 14 (1979) 751. D.H. Olson, G.T. Kokotailo, S.L Lawton, and W.M. Meier, J. Phys. Chem., 85 (1981) 2243. J. Bandiera and C. Naccache, Appl. Catal., 69 (1991) 139. O.C. Feeley, M.A. Johansson, R.G. Herman, and IC Klier, to be submitted. E.D. Hughes and C.K. Ingold, J. Chem. Soc. (1935) 254; W.A. Cowdrey, E.D. Hughes, C.K. Ingold, S. Masterman, A.D. Scott, J. Chem. Soc. (1937) 1252. R.T. Morrison and R.N. Boyd, Organic Chemistry, 2nd Ed., Allyn and Bacon, Inc., Boston, (1972). Q. Sun, R.G. Herman, and K. Klier, J. Chem. Soc., Chem. Commun. (1995) 1849. J.G. Nunan, K. Klier, and R.G. Herman, J. Catal., 139 (1993) 406. L Lietti, Q. Sun, R.G. Herman, and IC Klier, Catal. Today, in press.

I I th International Congress on Catalysis - 40th Anniversary

611

Studies in Surface Science and Catalysis, Vol. 101

9 1996 Elsevier Science B.V. All rights reserved.

Characterization of two different framework titanium quantification of extra-framework species in TS-1 silicalites.

sites

and

L. Le Noc a, D. Trong On a, S. Solomykina a, B. Echchahed a, F. BUand a, C. Cartier dit Moulin b and L. Bonneviot a D6partement de Chimie a, CERPIC, Universit6 Laval, Ste-Foy, G1K 7P4, Qu6bec, Canada. Laboratoire pour l'Utilisation du Rayonnement Electromagn6tique b, CNRS-CEA-MENJS, Bfitiment 209d, 91405, Orsay Cedex, France.. The quantification of the extra-framework titanium species in titanium silicalites of MFI structure, TS-1, was performed using either XANES at the Ti K-edge or XPS Ti (2p) photolines. In addition, two different framework sites, [Ti(OH)(OSi)3 ] and [Ti(OSi)4], were characterized in dehydrated samples using Diffuse Reflectance UV-visible, multiple scattering analysis of EXAFS, ~H and 29Si NMR spectroscopies. 1. INTRODUCTION The need for heterogeneous catalysis using the environment friendly hydrogen peroxide as oxidant agent for mild oxidation reactions has grown considerably. More and more high surface area materials containing titanium (IV) ions incorporated in SiO2 matrices of various structural properties (amorphous gels, mesoporous materials of MCM or HMS types and microporous silicalites or zeolites) have been found active for those reactions such as epoxidation of olefins (Shell process) and hydroxylation of aromatics (Enichem process) [15]. Among these solids, only the microporous silicalites S-1 and S-2 of MFI and MEL structures containing Ti, i.e. TS-1 and TS-2 respectively, are active for the catalytic oxidation of alkanes into alcohols and ketones (oxyfunctionalization) [6,7]. Despite an increasing number of characterization studies on those materials, there has been no detailed description of the framework Ti site structure nor quantification of extra-framework species [8]. It is therefore not yet established if micropores are necessary to activate alkane oxyfunctionalization reactions or if a specific Ti environment is required. To tackle the problem, a panel of techniques were applied to a series of dehydrated TS-1 with a (Ti/Ti+Si) ratio varying from 0.4 to 4.6 %. 2. E X P E R I M E N T A L The S-1 and TS-1 were synthesized from gels containing a mixture of TEOS, tetraisopropoxytitanium, and TPAOH. H202 was only added for the synthesis of TS-1 containing 3.4wt% Ti [9]. Hydrothermal crystallization, filtration and calcination at 540~

612 followed standard recipes. The Ti loading was obtained by atomic absorption measurements. The X-ray diffraction patterns were recorded as reported earlier [9]. The diffraction angle was calibrated using Si as internal reference. The unit cell parameters were obtained from Rietveld refinement of the powder patterns. The XPS data were measured using a VG Scientific Escalab Mark II as reported earlier [10]; a binding energy of 103.3 eV for Si(2p) was chosen as reference. The UV reflectance spectra were collected on a Perkin-Elmer Lambda 5 spectrometer equipped with a reflectance attachment provided by Harricks, interfaced with an IBM computer. A titanium free silicalite was taken as a reference. The X-ray absorption data we're collected in transmittance mode at LURE (France) [9-11]. The XANES spectra were normalized following a standard procedure and the energy calibrated on the edge of a metal Ti foil. The EXAFS signal was extracted as reported earlier [9]. The structure of the [Ti(OH)(OSi)3] and [Ti(OSi)4] clusters were generated using Cerius 2 from Molecular Simulation TM. Their spatial coordinates were fed into the FEFF6 code for MS ab initio calculations. The N M R measurements were performed at RT using a Brucker ASX-300 MHz spectrometer with a 7 mm MAS probe. The zirconia rotors were filled with the same volume of powder, evacuated under vacuum at 480~ during 24 hours, filled with a dry gas and sealed. For 29Si NMR, pure oxygen was used to increase the spin-lattice relaxation rate [12,13]; the recycle time for full relaxation was then decreased from 60 (in dry air) to 15 s for S-1 and from 80 to 30 s for the [ 1.5ITS-1. A 5gs-rff2 onepulse sequence was applied while the rotor was spun at a rate of 5 kHz (no cross polarization). The filling procedure insured a proportionality within ca. +1% between the silicon spectrum area and the sample weight introduced in the rotor. This allowed the normalization of the Si and H spectra to the number of silicon atoms. The gas was changed to pure nitrogen for the proton measurements in order to decrease the spin-spin relaxation rate in comparison with the dead time of the probe (3040~ts). An 7r./2-x-rc pulse sequence was used with a recycling time of 4s. A series of 15 experiments at a z delay varying from 200 to 5000 gs was performed. The number of proton was calculated from the intensity of the signal extrapolated at time x = 0 s using a two exponential relaxation law to fit the data. Another series of experiments in static conditions was performed with x varying from 40 to 100 gs; in this case, a single exponential relaxation process was used to fit the data and calculate the number of protons. The intensity factor between the proton signal and the silicon signal was obtained from the pure silicalite sample for which the number of protons matches exactly the number of [Si(OH)(OSi)3] species, Q3 of the 298i spectrum. The number of Q3, n(Q3), was obtained from the integration of the half of the Q3 line between-90 and-104 ppm to avoid the overlap range between Q3 and [Si(OSi)4 ] Q4 lines (vide infra, Figure 5). The accuracy on the number of Q3 species was ca. 2% and ca. 3% on the proton signal intensity for both the TS-1 and the reference S-1 sample leading to a precision on the number of proton, n(H) of ca. 10%. 3. E X T R A - F R A M E W O R K Ti C H A R A C T E R I Z A T I O N AND Q U A N T I F I C A T I O N A limit of ca. 1.8% of Ti incorporation into the framework of the S-1 silicalite has been calculated by titration using Raman spectroscopy [14] or using voltametric measurements [ 15]. XRD and Rietveld calculation of the unit cell expansion led to higher limit of ca. 2.5 % [ 16]. The extra-framework Ti was identified as belonging to a segregated TiO2 anatase phase. Higher level of incorporation of as much as ca. 6 and 8 % were claimed, using either silicon

613 alkoxides with higher hydrolysis rate [17 ] or Ti alkoxides with lower hydrolysis rate [181 in their sol-gel synthesis route, respectively. In both cases, no unit cell expansion above the maximum value of ca. 5390 ./t 3 tbund bv Millini et ai [16] were reported and. no clear evidence of the absence of extra-framework Ti was brought by the authors. Thereti~re. the nmxirnum of Ti incorporation is still a controversial matter that prompted us to develop a new technique to quantify them. Recently, some of us showed that the whole XANES Ti K-edge of dehydrated [2.4]TS-I and [4.6]TS-1 samples can be fitted with a linear combination of the experimental edges of TiO2 anatase and a model organosiloxytitanium compound, HDPOSST, as octahedral and tetrahedral references, respectively [9,11 ]. The latter compound mimics a substitutional site with its [Ti(OSi)4] core. Since the edge profiles are characteristic of a given phase where the absorber atom is located [19], these fits confirmed anatase as the phase in which the Ti extraframework species develop and strongly support the hypothesis of isomorphous substitution of Ti to Si in TS-1. In addition, the fits provide a quantification of extm-ti'amework and framework species assuming that these species are properly represented by anatase and HDPOSST compounds respectively. The data clearly indicate that up to 1.5% Ti the samples are free of extra-framework species. Above 1.5%, the anatase phase develops and accounts for 25% of the Ti in the [2.4]TS-1 samples for a framework substitutional level of 1.7% in agreement with some of the literature data [ 14,15].

Fig. 1. Quantification of framework Ti sites and unit cell expansion versus Ti contents using XPS 2p transition lines ( O ), XANES profiles at Ti K-edge ( A ) and Rietveld refinement of the XRD powder patterns ( 1-1). Solid data points for TS-l prepared as in ref. [9].

614 XPS can also be used as a quantifying tool. It has been shown that a [1.6]TS-2, a titanium silicalite of MEL structure, exhibits a single Ti 2p3/2 line at a binding energy of ca. 459.8 eV assigned to framework species [10]. At higher loading, a second line appears at 458.3 eV and dominates the spectrum for a Ti content of ca. 4.2% which is assigned to octahedral extra-framework oxidic titanium phase. In TS-1, a Ti photoline associated to framework species also arises at the same position while the extra-framework sites are characterized by the same binding energy of ca. 457.7 eV as the binding energy found for anatase mixed with silica. The XPS spectra of TS-1 samples were simulated using the same set of curves as for anatase and the [1.5]TS-1. This gave excellent fits and provided a quantification of the extra-framework species reported in Figure 1. Though the quantification using XPS gives the same trend as XANES, there is a systematic underestimation of the framework species for the former technique. This is due to a low energy tail of the Ti 2p3/2 photolines present in all the spectra of low Ti content TS-1. This line may arise from i) a slight reduction under the XPS measurement conditions (15 h accumulation time) ii) TiOH groups according to a surface study of futile TiO2 [20] and, iii) minute amounts of surface segregated TiO2. Works are in progress concerning this matter. Anyhow, the unit cell expansion versus the titanium loading is fully consistent with the data of Millini et al.[ 16] and reinforces the idea that there is effectively no extra-framework species up to at least 1.5% Ti in our samples. 4. S T R U C T U R E OF T H E F R A M E W O R K SITES 4.1 UV reflectance spectra Up to now, no techniques were able to differentiate Ti tetrahedral species in glasses, Tisilicalites or Ti-MCM materials. Very recently, some of us showed that the electronic spectra in the UV range significantly differs for dehydrated TS-1 and dehydrated TiO2-SiO2 silica gel containing 1.5% Ti [21 ]. The titanium in a tetrahedral environment, according to XANES [8], exhibits a strong UV transition peaking at 40,900 cm -1. This transition falls in the predicted range of energy for the ligand to metal electron transfer (LMET) band according to the empirical optical electronegativities theory [22]. By contrast, the electronic transitions are blue shifted by about 10,000 cm ~ for the TS-1 samples and, are about 6 to 7 times more intense than in the glass. A proper dilution of the sample (here six times for the [1.5]TS-1, Figure 2) was necessary to obtain a reflectance spectra proportional to the true absorption using the Kubelka-Munk function (F(R,o) = (1-R~2/2R~. The electronic spectra are composed of three Gaussian components centered at 50,200, 44,060 and 40,290 cm -~ The relative intensity of the first two and more intense ones is constant for samples of increasing Ti loading up to 1.5%; above 1.5%, the UV absorption due to anatase in the 32,000-36,000 -1 cm range interferes with the electronic bands of the tetrahedral species. The third component of the fit increases with the titanium loading in and accounts for the differences on the low energy tail between the spectra (Figure 2, insert). The first two bands are assigned to Ti tetrahedral sites found only in the silicalite matrix while the third one to sites similar to those found in amorphous S i O 2 matrices. These spectral differences are related to different Ti-O-Si bond angle of the Ti sites. Indeed, an angle opening will shift the bridging oxygen hybridization from sp 3 to sp 2 and

615 eventually to sp. This mechanism will favor a n donation into the electronic empty "e" level of Ti in Tu symmetry. As a consequence, the non-bonding "e" level will split into a filled bonding "en" level and an empty "en*" anti-bonding level. The latter electronic level is the so-called LUMO implied in the LMET of the lowest energy ~-., (there is no d-d transition tbr a d ~! Ti 4* ion). A larger angle will lead to a higher ~ donation and a blue shift of the LMET. Along this reasoning, the dominant sites in TS-1 would be characterized by larger Ti-O-Si angles than in those in amorphous SiO2 matrices.

F 0.5 0.4

n~ u..

s

0.3 ,~ 9

_

0.3

~. 0.2 ~

9 /[15%]TS-1

~" 0.1~

~.~[0-5%]TS-1

o.oL,

v

0.2

,

0.1 0.0

50

45

40

35

wavenumber (103 cm -~) Figure 2. Diffuse reflectance UV spectra of dehydrated [1.5]TS-I experimental data and fit residue (dotted lines), gaussian fit components (solid lines); insert" details for two different TS-1.

4.2 EXAFS Ti-O-Si angle calculations The calculation of the Ti-O-Si angles are based on the multiple scattering analysis of the EXAFS signal at the Ti K-edge. For the first attempt published earlier, the only availaible experimental data to compare with the theorical data was the [2.4]TS-1 containing unfortunately extra-framework species [11]. The adjustment to the EXAFS signal of the [1.5]TS-I lead to a more reliable value of ca. 163 ~ wich is very close to the average angle of ca. 166 ~ of the HDPOSST compounds used as reference [23]. Using the same Debye-Waller factor tbr the oxygen and silicon shells to fit the EXAFS data of both the model compound and the [1.5ITS-1 sample, the number of Si second neighbors was found to be smaller than 3. This led us to the hypothesis of the existence of [Ti(OH)x(OSi)4.x] [11] or more likel.x~ [Ti(OH)(OSi)3] species [23] according to the possibility to substitute Si for Ti where Q~ [Si(OH)(OSi)3 ] naturally occurs in S-1 silicalite (vide infra). This type of so called open site is more likely to relax its structure and exhibits larger Ti-O-Si angles than closed sites of [Ti(OSi)4] type which fully undergo a compression stress from the lattice where it substitutes a smaller T site of [Si(OSi)4] type [23].

616

0.8

_1

'

I

'

I

'

I

' 9

0.4

~

I

0.5

'a~ -

o

0

-0.4 9

o

- y -0.8 4

12

I

6

!'

i

8

'

10

'

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i

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4

14

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.

, I,,,

--0.5

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,

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i

,

!,,,

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6

8

i

,

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,

i

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i

,

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..=..

f

8

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i

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---|

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4

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R(,g.)

R(~)

f

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-2 --t

2.2

!

,

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,

! "':

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3.8

2.2

I

,

I , 2.6

!

t

! 3

,

!

,

I , 3.4

I

,? 3.8

R(A) Figure 3. X-ray absorption at the Ti K-edge of TS-I samples, experimental data (full lines) and ab inido calculation (dotted lines); b) to f) the calculation are compared with the [ 1.5]TS-1 experimental data; left hand side (dTi-O = 1.80 ,,~, dsi-O = 1.63 A, Ti-O-Si angle of c.a. 163 ~ and 3 Si second neighbors) and right hand side, mixture of sites 30% with angle of 154 ~ and 70% with angle of 165 ~.

617 Further MS calculations were performed to test the hypothesis of a mixture of sites with different average Ti-O-Si angles. For the sake of simplicity, the Ti(O)(OSi)3 and Ti(OSi)4 clusters were set as perfect tetrahedral sites with the same Ti-O and Si-O distance in both cases and the Debye-Waller factor kept identical to those found for the fit of HDPOSST. To find out what could be the angle difference between the two sites, the angle for the open site was set to the average value of ca. 166 ~ of the HDPOSST model compound while the angle was varied to obtain a satisfactory resemblance between the ab initio and the experimental data. To obtain a better fit the angle was finally decreased down to 165 ~ for the open site and found to be at least 154 ~ for the closed site with about 30-a:l0% of the titanium engaged in the latter sites. The ab initio EXAFS calculation based on the single site model reproduces most of the experimental features but does not give satisfactory results for the two first oscillations (Figure 3a). On the other hand, the double site model drastically improves the results, particularly for the second oscillation (Figure 3b). The fourier transform shows that the intensity of the second shell peak is better fit as well (Figure 3d) while the phasing match for this shell is maintained (Figure 3e and 30. Work is in progress to further improve the model.

Q4

a

v

or) t-

, I , , , , l l , , , I

90

t-

. m

100

....

110

I ....

-120 -130

Q3 integration -90

-95

-100

-105

chemical shift (ppmFI MS) Figure 4. Q3 line of the 29Si NMR spectra: a) S-l, b) [0.5]TS-land c) [1.5]TS-1; insert: total 29Si spectrum of the S-1 silicalite.

618

4.3 Characterization of TiOH groups from quantitative IH and 29Si NMR The 29Si MAS-NMR spectrum is typical of TS-1 (Figure 4, insert) [24]. The Q3 line corresponds to the Si(OSi)3OH sites. The silicon atoms linked to the titanium atoms are not resolved and resonate in the Q4 line. The latter line does not shift when the titanium loading increases; only a slight broadening (+0.2 ppm) is observed. This indicates that there are no changes in the average Si-O distances and Si-O-Si angles and, a slightly broader distribution of these parameters [25]. On the other hand, the Q3 line clearly decreases when the Ti loading increases (Figure 4). In fact, the number of silanol groups disappearing is approximately equal to the number of incorporated titanium atoms (Table 1). This observation prompted us to verify on the IH NMR spectra if this was due to the replacement of SiOH by TiOH groups. An orthorhombic titanium free silicalite (dashed line of fig. 5) exhibits two narrow proton resonance lines a and b, at 1.8 and 2.1 ppm, and a broad line c at ca. 2.5 ppm. All these signals correspond to silanol defects (7.7% of total Si) which can be removed under a strong hydrothermal treatment leading to the monoclinic, defect-free S-1. The latter only exhibits the known structured Q4 line due to the 24 inequivalent Si sites of the monoclinic structure. We found that the monoclinic S-1, exposed during one year to ambient air, evolves back to a defective structure (6% of defects) with the same a, b and c proton resonances as the original orthorhombic structure. It is therefore unlikely that these resonances comes from silicon vacancies. They are consequently assigned to terminal silanols located in different sites of the structure. This shows that the lattice relaxes by Si-O-Si bridge breaking leading to a retraction of the unit cell from 4365 to 4330/~3.

600 a

a

b/..

500 --- 4 0 0

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.=_ 200

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.

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,

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,

~

,

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,

,

,

,

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,

,

1 3 2 chemical shift ( p p m / T M S )

0

-1

-2

Figure 5. IH NMR spectra of [1.5]TS-1 (solid line) and S-1 (dashed line); inserts a) signal intensity versus echo delays "r ([1.5]TS-1, solid line and S-l, crosses) and b) S-1 spectra acquired at various x delays.

619 The titanium silicalites exhibit very similar IH spectra to the defective S-1. Besides, a continuous increase of the a line at the expense of the b line is observed upon the Ti loading (Figure 5). It could correspond to either a redistribution of the silanol defects, or to the replacement of silanol defects by titanol defects at 1.8 ppm. The latter assumption is reasonable considering the ~H resonance of TiOH terminal groups in TiO2 anatase [26]. The total proton titration shows that the number of defects n(H) is always greater than the number of Q3 silanol groups (Table 1). That clearly demonstrates the existence of the TiOH groups in flae dehydrated TS-1. The predominance of the latter sites suggests that the structure preferentially relaxes by breaking of Ti-O-Si bonds instead of Si-O-Si bonds. Table. 1 Q3 defects and proton titration using Solid State NMR. static MAS Q3 Ti n(Q 3) n(H) n(H) per n(H) n(H) per -4 (10-4 mol) Ti atom (10-4 mol) Ti atom (%,+0.2) (%) (10 mol) S-1 7.7 0 3.28 3.28 3.28 0.4 + 1.4 [0.5]TS-1 7.1 0.51 3.00 3.16 0.7+ 1.4 3.08 0.7+0.5 [1.5ITS-1 6.4 1.52 2.68 3.19 0.8 + 0.5 3.10

5. C O N C L U S I O N Both XANES and XPS allow to titrate framework tetrahedral Ti and extra-framework TiO2 anatase species. There was no detectable extra-framework Ti up to 1.5wt% Ti. Open sites, [Ti(OH)(OSi)3], are the dominant framework sites according to UV-visible spectroscopy and ~H solid state NMR. They are characterized by a large average Ti-O-Si angle of ca. 163 ~ and two very intense ligand to metal electronic transitions. The closed sites, [Ti(OSi)4], are marginal at low loading. Their smaller Ti-O-Si angle than those of the open sites accounts for a higher structural stress undergone by these sites than the open sites. According to our late results during the referring process on amorphous TiO2-SiO2 system, the Ti-O-Si angle for closed sites is 30 ~ smaller compared to open sites. This is more than suggested here. Calculations taking into account this new information are in progress. 6. R E F E R E N C E S 1. 2. 3. 4. 5. 6.

M.G. Clerici, G. Bellussi and U. Romano, J. Catal., 129 (1991) 159. T. Sato, J. Dakka and R. A. Sheldon, Stud. Surf. Sci. Catal., 84 (1994) 1853. T. Blasco, M. A. Camblor, A. Corma and J. Perez-Pariente, J. Am. Chem. Soc., 115 (1993) 11806. P.T. Tanev, M. Chibwe, and T. J. Pinnavia, Nature, 368 (1994) 321. R. Hutter, D. C. M. Dutoit, T. Mallat, M. Schneider and A. Baiker, J. Chem. Soc., Chem. Comm. (1995) 163. D . R . C . Huybrechts, L. De Bruycker and P. A. Jacobs, Nature, 345 (1990) 240.

620 7.

8. 9. 10. 11. 12. 13. 14. 15. 16. 17. 18. 19. 20. 21. 22. 23. 24. 25. 26.

A. Bittar, D. Trong On, L. Bonneviot, S. Kaliaguine and A. Sayari, Proc. 9th Int. Zeolites Conf., R. von Ballmoos, J. B. Higgins and M. M. J. Treaty eds., ButterworthHeinemann, Boston, 1 (1993) 453. S. Bordiga, S. Coluccia, C. Lamberti, L. Marchese, A. Zecchina, F. Boscherini, F. Genoni, G. Leofanti, G. and Petrini, G. Vlaic, J. Phys. Chem., 98 (1994) 4125. D. Trong On, S. Kaliaguine and L. Bonneviot, J. Catal., 157 (1995) 235. D. Trong On, L. Bonneviot, A. Bittar, A. Sayari and S. Kaliaguine, J. Mol. Catal., 74 (1992) 233. C. Cartier dit Moulin, C. Lortie, D. Trong On, H. Dexpert and L. Bonneviot, Physica B, 208&209 (1995) 653. D. J. Cookson and B. E. Smith, J. Magn. Reson., 63 (1985) 217. J. Klinowski, T. A. Carpenter and J. M. Thomas, J. Chem. Soc., Chem. Commun., (1986) 956. G. Deo, A. M. Turek, I. E. Wachs, D. R. C. Huybrechts and P. A. Jacobs, Zeolites 13, (1993) 365. S. Castro-Martins, A. Tuel and Y. Ben TMdt, Stud. Surf. Sci. Catal., 84 (1994) 501. R. Millini, E. Previde Massara, G. Perego and G. Bellussi, J. Catal., 137 (1992) 497. A. Tuel and Y. Ben Tagtrit, Applied Catal. A, 110 (1994) 137. A. Thangaraj, M. J. Eapen, S. Sivasanker and P. Ramasamy, Zeolites, 12 (1992) 943. P. Behrens, J. Felsche, S. Vetter, G. Schulz-Ekloff, N. I. Jaeger and W. Niemama, J. Chem. Soc. Chem. Commun., (1991) 678. T. K. Sham and M. S. Lamnas, Chem. Phys. Letters, 68 (1979) 426. D. Trong On, L. Le Noc and L. Bonneviot, J. Chem. Soc., Chem. Comm., accepted. C. K. J~rgensen, Prog. Inorg. Chem., ed. S. J. Lippard, Wiley, New York, 12 (1970) 101; J. A. Duffy, Structure and Bonding, 32 (1979) 147. L. Le Noc, C. Cartier dit Moulin, S. Solomykina, D. Trong On, C. Lortie, S. Lessard and L. Bonneviot, Stud. Surf. Sci. Catal., 97 (1995) 19. A. Tuel and Y. Ben TMrit, J. Chem. Soc., Chem. Commun., (1992) 1578. G. Engelhardt and D. Michel, High-Resolution Solid-State NMR of Silicates and Zeolites, John Wiley & Sons, Chichester (1987). V. M. Mastikhin and A.V. Nosov, React. Kinet. Lea., 46 (1992) 123.

J.W. Hightower, W.N. Delgass, E. Iglesia and A.T. Bell (Eds.) 11th International Congress on Catalysis - 40th Anniversary

Studies in Surface Science and Catalysis, Vol. 101 9 1996 Elsevier Science B.V. All rights reserved.

621

I n f l u e n c e o f sulfur d i o x i d e on the selective catalytic r e d u c t i o n of N O by d e c a n e o n C u catalysts. F. Figueras 1", B. Coq 1, G. Mabilon 2, M. Prigent 2 and D. Tachon 1 1Laboratoire de Mat6riaux Catalytiques et Catalyse en Chimie Organique, URA 418 du CNRS, ENSCM, 8 rue de l'Ecole Normale, 34053 Montpellier Cedex, France. 21nstitut Franqais du P6trole, BP 311, 92506 Rueil Malmaison, France Abstract The selective catalytic removal of NO in oxygen rich atmospheres has been investigated in the presence of sulfur dioxide on a series of Cu catalysts. The reactivities correlated with the reducibility of Cu species determined by temperature programmed reduction with hydrogen. Without sulfur dioxide in the feed, the activity is related to the reducibility of Cu species. The addition of SO2 to the solid shifts the TPR peaks to higher temperatures. The magnitude of this effect is lower for acid zeolites such as MFI and B E A . Sulfation results in a small inhibition of the reactivity for deNOx in the case of Cu/AI203, no or little change in the case of Cu/zeolites, and a promotion of activity in the case of Cu/TiO2 and Cu/ZrO2. The oxidation of decane on Cu/TiO2 and Cu/ZrO2 is inhibited by SO2 at low temperatures, but remains close to 100% in presence or absence of SO2 on CuffiO2 above 600K. In the case of Cu/ZrO2 the addition of SO2 increases the rate of oxidation above 640 K. The positive effect of SO2 on deNOx is attributed to the promotion of a bifunctional mechanism in presence of strong acid sites.

1. I N T R O D U C T I O N The introduction, in most industrial countries, of very stringent regulations for nitrogen oxides emissions has forced the development of catalysts for the abatement of NOx emission from vehicles. If viable solutions now exist for stoichiometric gasoline engines with the so-called three-way catalysts, the future of diesel and lean burn engines may depend on the discovery of new catalysts active in the presence of large excess of oxygen [1]. In fact, the selective catalytic removal of NO in presence of excess oxygen remains a challenge. Most of the current studies involve C1-C4 hydrocarbons as reductants and zeolites as catalysts, among which Cu-exchanged MFI zeolites are considered as one of the most active [2]. The reductant shows a complex influence in this reaction: it has been thus reported that a Cu/ZrO2 catalysts are active with propene but show low activity with propane as reductant [3]. For a practical use reduction by higher alkanes would be attractive, since it would be easier to handle in a vehicle. It is well known also that higher alkanes suffer radical gas phase oxidation above 723 K. Therefore, their use requires catalysts active and selective for deNOx at lower temperatures. The mechanism of NOx elimination is still debated: a redox mechanism involving Cu ions is probable, and isolated Cu cations exchanged into MFI [4,5] or mordenite [6] have been found to be more active than CuO clusters. It must be emphasized, however, that acid zeolites exhibit good activity at high temperature, and acid mechanisms have been proposed [7-10]. In presence of Cu this acid mechanism disappears probably due to the decrease of the acidity of mordenite upon Cu exchange [6]. According to * present address: Institut de Recherchessur la Catalysedu CNRS,associ6 h l'Universit6 C. Bernard, 2 avenue A. Einstein, 69626 Villeurbanne Cedex, France. D.T aknowledgesa PhD grant from liP.

622 Sasaki et al. [8] this acid mechanism could proceed by nitration of a product of the oxidation of the hydrocarbon, followed by autoreduction of this intermediate to nitrogen. Poisoning of deNOx catalysts by SO2 could also be a problem since diesel fuels contain small amounts of sulfur compounds. Only a few studies deal with this subject [11-13]. It appears from the literature that for Cu catalysts the use of MFI as a support reduces the inhibition by SO2. Support effects also appear in the case of Co since Co/MFI is much less sensitive to SO2 than Co/ferrierite [13]. Since this support effect may be related to acidity, it becomes important, to investigate the influence of SO2 on the properties of Cu catalysts supported on SiO2, A1203, MFI, BEA and unpromoted or sulfate promoted TiO2 and ZrO2. These latter have been reported active for deNOx [14]. 2. E X P E R I M E N T A L

2.1. Preparation of the solids The supports used are listed in Table 1. Commercial supports were used except for zirconia. In that case a high surface area sample was prepared by precipitating a solution of zirconyl chloride at a constant pH of 10, then drying at 400 K. Cu on alumina, silica or zirconia catalysts were prepared by impregnating these solids with a solution of Cu acetate in distilled water. The amount of metal was altered by changing the concentration of the Cu solution. Zeolites were exchanged from their Na forms, by a solution of Cu acetate. After exchange, washing and drying at 400 K, the solids were calcined at 773 K. The chemical analyses were performed after dissolution of the solids by the Service Central d'Analyse (Solaize, France) using plasma atomic absorption spectroscopy.

2.2 Characterizations Surface areas were determined from the adsorption isotherms of nitrogen at 77 K, using a Micromeritics ASAP 200 instrument. Powder X-ray diffraction patterns were obtained with a CGR theta 60 instrument using CuKa monochromated radiation. Reducibility and the amount of Cu species were determined by temperature programmed reduction (TPR) with H2 (H2/Ar: 3/97, vol/vol). The experimental set up has been described previously [6]. Accessibility to Cu sites was determined by temperature programmed desorption of NO (NO TPD), using an experimental setup similar to that used for TPR, except the detector was a quadrupole mass spectrometer (Balzers QMS421) calibrated on standard mixtures. The samples were first activated in air at 673 K, cooled to room temperature in air, and saturated with NO (NO/He: 1/99, vol/vol). They were then flushed with He until no NO could be detected in the effluent, and TPD was started up to 873 K at a heating rate of 10 K/min with an helium flow of 50 cm 3 rain -1. The amount of NO held on the surface was determined from the peak area of the TPD curves.

2.3 Catalytic properties They were determined in a fixed bed reactor, at a GHSV of 63 000 h-1 with a reaction mixture composed by oxygen (9 vol%), n-decane (1200 ppm), NO (1000 ppm) and helium (qs 100%). For the experiments where SO2 was added to the feed, a special reaction mixture containing 20 ppm of SO2 was used. The solid could also be presaturated with SO2 by injection in the carrier gas, using a six port chromatographic valve equipped with a 200 ~tL loop. The products were analyzed continuously by sampling on line to a quadrupole mass spectrometer Balzers QMS 421 equipped with a Faraday detector, and nitrogen and CO were discriminated by analyzing on line on a GC equipped with on 13X molecular sieve column and a catharometer. While CO traces were detected among the products when Cu/MFI is used as catalyst, no CO was formed on Cu/ZrO2 either sulfated or not.

623 Table 1 Characteristics of the supports used. - Support

,

Origin

BEA MFI ~,Alumina Silica Titania Zirconia

'

Reference

Zeocat ~ t Procatalyse Rhone Poulenc Degussa home made

PB-1 PZ-3/30H GFS=200 Z175MP P25 ,

'''i'''1'''1'''1'''1'''1'''

100

9

9

Siirface (m2/g) 758 200 175 40 300

9

9

=,

9

9

9

oo

80 v

60 ~40 rj O

20

--"

u

9

o o

oo~ 0

k-,t

633

653

673

693

713

733

753

773

area Si/AI 9.85 15

Cationic form Na + H+

The catalytic conversion of NO was investigated first in absence of catalyst (blank). The results reported in Fig. 1 show that the homogeneous gas phase oxidation of the alkane starts at 650 K. No reduction of NO is observed in the homogeneous process, then the production of N2 can be ascril:r_xi to the catalytic reduction. The catalytic properties were determined by temperature programmed reaction (ramp: 2 K rain-l). The temperature was increased from 523 to 673 K and back.

Temperature (K) Figure 1. Conversions of the alkane (filled points) and of NO (open points) in the blank experiment. The catalytic properties were characterized in a simplified manner by two parameters: the maximal conversion CM and the corresponding temperature TM. The selectivity of NO conversion to N2 is always very high (> 98%). The formation of NO2 is marginal on these Cu catalysts. 3. R E S U L T S

3.1. Chemical composition The chemical analyses of the samples are reported in Table 2. The X-ray diffraction spectra of these solids do not show the presence of any Cu oxide phase with size larger than 3-4 nm, excepted in the cases of Cu(3)SiO2 and Cu(4)ZrO2 on one side, Cu(146)Na(6)FAU-10 and Cu(146)Na(28)MFI-15 on the other side, where the lines characteristic of CuO do appear with a line broadening corresponding to a particle size of about 4 nm.

624

3.2.

Redox

properties

The thermograms of Cu reduction in silica-, alumina-, titania- and zirconia-supported catalysts show only one peak, the maximum of which is reported in Table 3. The amount of hydrogen consumed by the reduction corresponds, within experimental error, to the theoretical amount required for the reaction: Cu 2+ + H 2 - - > Cu + 2H + In agreement with Shimokawabe e t al. [15], the desorption profiles of NO show two peaks at around 450 and 670 K, which correspond to Cull-NO and (CulI-O)-NO species respectively. The latter NO species usually represent no more than 10% of the total amount. Table 2 Chemical analyses of the different catalysts Catalyst

Support

Si/AI

wt% Cu

Ctt/Al

Na/AI

C'u exch

Cu(46)Na(22)BEA CufI28)Na(7)BEA Cu(48)Na(x)MFI Cu(146)Na(2~)MFI Cu(E)Al203 Cu(3)AI203 C'u(1)SiO2 Cu(3)SiO2 Cu(1)TiO2 Cu(E)TiO2 Cu(4)Zr02

BEA BEA MFI MFI A1203 A1203 SiO2 SiO2 TiO2 TiO2 Zr02

11.8 9.7 15 15.4

1.80 4.95 1.05 4.40 2.60 2.95 0.80 2.55 1.60 2.40 4.20

0.23 0.64

0.22 0.07

69 135 48 175

-

-

0.73

0.28

(~)

The comparison of the results of TPR with those of NO TPD on Cu on silica and alumina suggests that a better dispersion of the Cu species induces a higher reducibility of Cu, as reported elsewhere [ 16]. On the other hand, a strong support effect appears in the case of Cu supported on zirconia or titania, which shows the easiest reducibility in spite of a low dispersion as judged from XRD. Table 3 Maximum temperatures and hydrogen consumption during the temperature programmed reduction of Cu/oxides by hydrogen, and NO taken up by Cu atom. Catalyst

Reduction temp.

H2/Cu

NO/Cu

Cu(1)SiO2 Cu(3)SiO2 Cu(E)AI203 Cu(3)AI203 Ct(2)TiO2 Cu(4)Zr02

550 630 600 540 440 470

1.05 1.07 1.11 0.97

0.30 0.03 0.04 0.12 -

(K)

.......

In the case of zeolites (Table 4), a very broad reduction peak is observed on the sample partially exchanged, with a poorly defined maximum. On the other hand, two reduction peaks appear for Cu exchanges higher than 100%. In both cases, however, the consumption of hydrogen corresponds to the reduction of Cu 2+ to metallic Cu. A similar situation was previously reported for Cu-MOR [6]. By

625 analogy with the results obtained with Cu-MOR, we assign the low temperature peak to the reduction of CuO clusters to Cu 0 and isolated Cu 2+ cations to Cu +, and the high temperature peak to the reduction of Cu + to Cu 0. TPO experiments show that the re,oxidation of Cu occurs between 400 and 460 K for Cu/AI203 and Cu/SiO2 respectively. Reoxidation occurs then at a lower temperature and is thus faster than reduction of the Cu surface. Table 4 Maximum temperatures and hydrogen consumption during the temperature programmed reduction of Cuzeolites by hydrogen. Catalyst

1rst peak

2/kl peak

Cu(46)Na(22)BEA a Cu(128)Na(7)BEA Cu(48)Na(x)MFI a Cu(146)Na(28)MFI

630 530 710 600

630 720

(K)

H2/Cu ., 1.02 1.05

(K)

CuO/Cu2'40/100 35/65 0/100 50/50

aCu was postulated to be fully dispersed in the 10w exe "l'maged zeolites. Hydrogen uptake was difficult to quantify precisely.

The effect of sulfur addition on the TPR profiles for some of these catalysts is reported in Table 5. The saturation of the solid by SO2 shifts the TPR profiles to higher temperatures and a second reduction peak appears at high temperature. The maximum temperatures measured for this second high temperature peak are reported in Table 5 (column 6). The effect of sulfur addition depends on the nature of the support, and zeolites appear less sensitive to SO2. 300

250 v

-

Cu(4)ZrO-no SO2 2 Cu(4)Z.rO-1 S02

200

2

- - , o - - Cu(4)ZrO-6S02 2

a. 1 5 0 c-

2 "0 I

100

50

300

400

500

600

700

800

900

Temperature (K) Figure 2. Temperature programmed reduction profiles of Cu on zirconia and sulfated zirconia with increasing amount of SO2.

626 TPR experiments were performed at different amounts of sulfur on a Cu/ZrO 2 (Fig. 2). SO2 addition induces a modification of the TPR profile: the reduction peaks are shifted toward higher temperatures. The total hydrogen consumption is decreased by 50% for the first SO2 dose, then remains constant. 3. 3 Catalytic properties In the absence of sulfur, for two samples of comparable dispersion (NO/Cu = 0.04), the activity is related to the reducibility of the CuO phase. The highest conversion (98% at 670 K) is achieved on CuBEA, which appears as active as Cu-MFI. In order to investigate the effect of sulfur, the catalytic properties were determined: i) in absence of SO2 in the feed before and after addition of SO2 to the catalyst, up to a molar ratio SO2/Cu= 3.6, ii) in presence of sulfur dioxide in the feed (20 ppm). Table 5 Influence of the presaturation of the solid by SO2 on the redox properties of some Cu catalysts. 'Catalysts

Temperatures of the reduction peaks (K) No addition of SO2 Presaturation with SO2

(so2/cu = 2.85) Cu(3)Al203 Cu(3)SiO2 Cu(4)ZrO2 Cu(146)Na(28)MFI Cu(128)Na(7)BEA

'frst peak 540 630 470 600 520

2 ~ peak 720 630

lrst peak 620 660 480 650 560

2nd peak 800 590 660

Due to the small amount of SO2 in the feed, sulfation of the catalyst by the charge is very slow, and the effect of sulfur appears clearly only when the catalyst is sulfated separately. Dosing of SO2 onto the solids shows that sulfur is indeed adsorbed at the surface since Cu/TiO2 retains about 0.4 wt% S and Cu-BEA 0.1 wt% S. It is interesting to notice that the effect of SO2 clearly depends on the type of support. The different catalysts are compared in Table 6. The addition of sulfur dioxide has no effect on activity in the case of zeolites and slightly inhibits CR/AI203, but promotes NO conversion on Cu(1)/TiO2 and Cu(4)/ZrO2. The activity is practically doubled for Cu(4)/ZrO2 with a small shift of the optimal temperature. The promotor effect of SO2 increases with the amount added to the reaction medium (Fig.3). An effect of the addition of sulfur dioxide has also been observed on the oxidation of decane with an increase of the activation energy e ~ t e d for such a poisoning. This addition leads to a noticeable decrease of the rate of oxidation at low temperature, where Cu sulfate is stable, but the effect becomes negligible at about 600 K. At this temperature, the conversion of decane estimated by the evolution of the peak e/m = 57, characteristic of the hydrocarbon, is close to 100% with Cu/TiO2 catalysts in presence or not of SO2 (Figure 4). With Cu/ZrO2 SO2 inhibits decane oxidation below 640 K. At 640 K a conversion of about 60% is observed in both the presence or absence of additive and an acceleration of oxidation is noticed at higher temperatures.

627

100

.

.

.

.

.

.

.

.

.

.

.

.

.

I

|

!

!

i

!

Cu/-nq

~80

--0--

0

v

~

E

,--60

Cu/'l'iO2 (20 ppm $ 9

CurnO~(sojcu/= 0.35) curnq (sojcu = 2.as)

0

>.

"40

0 o

0 z

20

0 450

500 55O 600 Reaction temperature (K)

6,5O

Figure 3 . Conversion of NO on Cu(1)/TiO2 in presence of different amounts of 502 added to the catalyst. The reaction is performed without SO2 in the feed for the test, and in presence of 20 ppm of SO2 in the other cases. Table 6 Influence of sulfur dioxide on the NO maximum conversion (CM at.TM) over Cu catalysts. Catalyst Pretreatment by sulfur dioxide no SO2 in the feed 20 ppm SO2 in the feed so2/cu - 0 SO2/Cu = 0 SO2]Cu _ 3 TM (tC) CM(%) TM (K) CM (%) TM (K') CM (%) 710 42 '-Cu(Z)/AI203 590 44 640 37 Cu(3)/SiO2 600 14 600 34 600 66 Cu(1)/TiO2 530 41 640 61 Cu(4)/ZrO2 620 31 Cu(146)Na(28)MFI 710 55 650 62 Cu(46)Na(22)BEA 770 70 770 70 770 72 680 100 CuQ28)Na(7)BEA 680 98 -

-

-

-

4. D I S C U S S I O N Reasonable NO conversion can be achieved using n-decane as reductant. In the absence of sulfur dioxide, the catalytic activity is roughly related to the reducibility of the Cu phase of Cu ions in zeolites: the reaction temperature needed to reach 20% NO conversion parallels that of the TPR peak (Table 7). This relation also practically holds for Cu on simple oxides, therefore a redox mechanism in which reduction of Cu 2§ cations is the slow step could account for the results.

628 The addition of sulfur dioxide completely changes the activity patterns since sulfur has different effects depending on the support. On Cu-BEA zeolite, the addition of SO2 has little effect, neither on activity or reducibility. Indeed only trace amounts of S are retained by the solid in that case, which is consistent with a low interaction of SO2 with the solid. On C-k~iO2 and Ca/ZrO2, the addition of SO2 decreases the reducibility of Cu species, but increases the catalytic activity. Therefore the former relation between activity and reducibility no longer holds. These changes of activity are not related to the changes of alkane oxidation, since at the temperature of the maximum activities of Ctt/TiO2 and Cu/ZrO2, the concentrations of alkanes remaining in the reaction mixture are comparable in both eases. For TiO2 and ZrO2, it is well known that sulfation induces a strong increase of acidity [ 17] and the participation of an acid mechanism could then account for this promotion of activity. This mechanism can be described as a bifunctional process: oxidation of NO to NO2 on Cu sites, and nitration of a product of the oxidation of deeane on the acid function(8). The preparation of the catalyst must have a great influence on the activity. This has been shown by the comparison of three Cu/TiO2 catalysts prepared in different conditions: one in which titania is first treated with sulfuric acid, then by Cu acetate (denominated Cu/SO4/TiO2, containing 0.5 wt% Cu, 0.6 wt% S), one in which Cu is

100 80 t-" 0 .~_.

'~ k_

60

r O

~

40

c-

'-'

20 V

~

. . . I . . . I . . . I . . . I , i i i . i i i . . . I . . .

0

460

480

500

520 540 560 T e m p e r a t u r e (Ix')

580

600

620

Figure 4. Conversion of decane as a function of the reaction temperature on Cu(1)/TiO2 in absence of SO2 either in the feed or the solid (Experiment 1), on the same solid but in presence of SO2 in the feed (Experiment 2), and with the same Cu(1)/TiO2 catalyst, presaturated by SO2 (SO2/Cu _ 3, Experiment

3). introduced first, then sulfates using the same procedure (SO4/Cu/TiO2, containing 0.42 wt% Cu, 0.35 wt% S), and a Cu(2)/TiO2 sample sulfated by dosing SO2 in the gas phase (containing 2.38 wt% Cu and 0.49 wt% S). The comparison of these samples reported in Fig. 5 shows that the activity is higher at the higher Cu content. The slow step would then be the oxidation of NO at the Cu surface, as is commonly observed [18]. This is consistent with the observation of only trace amounts of NO2 in the effluents.

629

Table 7 Comparison of the temperatures needed to obtain 20% NO conversion and of the reducibility expressed by the temperature of the first peak in TPR by hydrogen. Catalyst Cu(46)Na(Z2)BEA Cu(128)Na(7)BEA Cu(48)Na(x)MFI C'u(146)Na(28)MFI Cu(1)TiO2 Cu(2)Al203 Cu(4)ZrO2

T2X)%(K') 640 530 670 560 500 540 570

I rst TPR peak (K') 630 520 720 600 530 590 620 .

The same behaviour has been found with Cu/ZrO2. A highly dispersed Cu phase was obtained at the surface of zirconia by reacting the support with Cu acetylacetonate [ 19]. This procedure yields an active catalyst. This catalyst was selective for N2 formation at low temperature (< 550 K), but produced only NO2 when the temperature becomes higher than 650 K. However, the same type of catalyst prepared from sulfated zirconia did not produce NO2 but selectively reduces NO to N2 whatever the temperature, with a yield of about 40% at 670 K, and a GHSV of 70000 h -1, using only 300 ppm of decane.

4~/'%

i~.s

i

i

i

I

-o80

i

,

i

I

so/c

i

i

i 'I

rns

i

i

i

I

i

i

i

I

i

i

'i

i

i

,

,

cu/sorn%

cu(2)rno + so

0

E ~

cO

60

40 0

0 z

20 0

540

560

580

600

620

640

660

680

Reaction temperature (K) Figure 5. NO conversion on CuffiO2 catalysts prepared according different procedures. The feed is the standard one containing 20 ppm of sulfur dioxide.

In conclusion, Cu on TiO2 or ZrO2 show a unique and interesting bebaviour since their deNOx activity is promoted and not intfibited by the presence of sulfur in the feed. This effect can hardly be attributed to a selective inhibition of the oxidation of decane, and is better explained by the promotion of a bifunctional mechanism involving the acid sites created on the support by the reaction of SO2.

630 REFERENCES 1. B. J. Cooper, Platinum Metals Rev. 38 (1994) 2. 2. M. Iwamoto ; "Future Opportunities in Catalysis and Separation Technology" ; M. Misono, Y. MoroOka, S. Kimura Eds.; Elsevier, Amsterdam, 1990, p.121. 3. tC A. Bethke, D. Alt, M. C. Kung, Catal. Lett. 25 (1994) 37. 4. S. Sato, Y. Yu-u, H. Yahiro, N. Mizuno and M. Iwamoto, Appl. Catal. 70 (1991) L1. 5. Z. Chajar, M. Primet, H. Praliaud, M. Chevrier, C. Gauthier and F. Mathis, Appl. Catal. B 4 (1994) 199. 6. B. Coq, D. Tachon, F. Figueras, G. Mabilon and M. Prigent, Appl. Catal. B 6 (1995) 271. 7. H. Hamada, Y. Kintaichi, M. Sasaki, T. Ito and M. Tabata, Appl. Catal. 64 (1990) L1; H. Hamada, Y. Kintaichi, M. Sasaki, T. Ito and M. Tabata, Appl. Catal. 64 (1990) LS. 8. M. Sasaki, H. Hamada, Y. Kintaichi and T. Ito, Catal. Lett. 15 (1992) 297. 9. J. O. Petunchi and W. IC Hall, Appl. Catal. B 2 (1993) L17; J. O. Petunchi, G. Sill and W. K. Hall, Appl. Catal. B. 2 (1993) 303. 10. IC Yogo, M. Umeno, H. Watanabe and E. Kikuchi, Catal. Lett. 19 (1993) 131. 11. M. Iwamoto, H. Yahiro, S. Shundo, Y. Yu-u and N. Mizuno, Appl. Catal. 69 (1991) L15. 12. E. Kikuchi, K.Yogo, S. Tanaka, M. Abe, Chem. Lett. (1991) 1063. 13. Y. Li and J.N. Armor, Appl. Catal. B 5 (1995) L257. 14. H. Hamada, Y. Kintaichi, M. Tabata, M. Sasaki, T. Ito, Chem. Lett. (1991) 2179. 15 M. Shimokawabe, N. Hatakeyama, K. Shimada, K. Tadokoro and N. Takezawa, Appl. Catal. A 87 (1992) 205. 16. S. D. Robertson, B. D. McNicol, J. H. de Baas, S. C. Cloet and J. W. Jenkins, J. Catal. 37 (1975) 424. 17. K. Arata, "Advances in Catalysis", Academic Press, San Diego, 1990, p.165, and references therein. 18. K. A. Bethke, C. Li, M. C. Kung, B. Yang and H. H. Kung, Catal. Lett. 31 (1995) 287. 19. G. Delahay, B. Coq, E. Ensuque and F. Figueras, Catal Lett. in press.

j.w. Hightower, W.N. Delgass, E. Iglesia and A.T. Bell (Eds.) 1 l th International Congress on Catalysis - 40th Anniversary

Studies in Surface Science and Catalysis, Vol. 101 9 1996 Elsevier Science B.V. All rights reserved.

631

CoPt C l u s t e r s in N a M o r d e n i t e s as C a t a l y s t s for S C R o f NO• L. Gtmerrez, A Ra~ootta,A. Boix and J. Petunchi.

lnstituto de Investigaciones en Cat~li~s y Petroquimica- INCAPE (FIQ,UNL- CONICET) Santiago del Estero 2829 - 3000- Santa Fe- Argentina.

A new material based on Pt and Co exchanged in NaMordenite for the selective catalytic reduction (SCR) of nitric oxide with methane in the presence of excess oxygen is studied. The incorporation of 0.5% weight of Pt and 2% weight of Co to the zeolitic matrix alter calcination and reduction on ~ flow for 1 h yields a solid converting 100% of NO to N 2 and, simultaneously, 100% o f C H 4 t o CO: with a CH4/ NO ratio = 3 and 2% of oxygen in the feed at 450~ When the oxygen concentration in the feed varies, the NO conversion goes through a maxinmm for 2% at 450~ The incorporation ofl~ also promotes Co reduc~flity; 1% is reduced to Co ~ in the monometallic sample and 13% in the bimetallic sample. XPS results reveal that in the calmed samples Co 2+is at exchange position and, after being reduced, there appear thinly dispersed Co ~particles and exchanged Co 2+ions. A greater reducibility and a shift ofthe maxima in the temperature-programmed reduction profiles suggest a Pt-Co interaction. In order to get an efficient catalyst for nitric oxide abatement it is necessary that the highly dispersed Co ~ and Pt ~ particles and the Co 2+and I-I+ ions at exchange positions be in intimate contact inside the mordenite channels. 1. I N T R O D U C T I O N Since the pioneer work of Li and Armor (1,2,3), who found that methane could be used as a selective reductant ofnirrogen oxide over Co-ZSM5, several solids have been used with the same purpose (4,5,6,7). Metals such as Co, Rh and Pt, exchanged in ZSM5 (8), were similarly studied. In the absence of O: the three metals have a high selectivity but when 25% of oxygen is added to the feed, platinum is non-selective, the NOto N 2conversion being only 3% with 25% ofNO 2production. Under these conditions Rh-ZSM5 presented 25% of NO conversion with approximately 10% of NO: at 450~ Kikuchi and Yogo (9) in their studies ofSCR of NO using different protonic zeolites found that activity and selectivity are functions of acidity. The same authors have reported that the incorporation of Cra or In to HZSM5 notably increases the NO to N: conversion. Li and Armor (10) also studied Ga-HZSM5 and they found that this solid presented a high selectivity to NO reduction, even at high temperatures. The NO to N 2 conversion kept constant over 450~ Pd-based catalysts on different acid supports including HZSM5 were studied by Loughran and Kesasco (11) and Nishizaka and lVfisono (12). These authors found a behavior ~ to that ofGa-HZSM5, atm~uting the activity and selectivity ofthose solids to a bifunctional metal-acid action All materials effective for nitric oxide abatement are sensitive to the presence of water in the feed; however, Inui et aL (13) have recently reported that a protonated Co-contained silicate was effective with CH4for NO reduction in conditions equivalent to diesel engines exhausts, the NO to N: conversion on this material not being altered by a I-I:O concentration in the feed ofup to 10%. This work studies a new catalyst based on Co and Pt exchanged in mordenite. It presents an

632 ~ l y s i s ofthe effect o f Pt on activity and selectivity in the selective NOreduction, and finally discusses an atten~t to characterize active sites. 2. EXPERIMENTAL

Catalysts Prepara#orL Catalysts were prepared by ionic exchange starting from NaMordenite Zeolon Si/AI = 6. Monomet~c solids were prepared using for the exchange ~xted solutions of Co(NO3)2 and Pt(NI-13)4('NO3)2,respectively. Bimetallic solids were obtained by successive exchanges beginning with Pt salt for PtCoMordenites, and with Co salt for CoPtMordenites. The same support was exchanged with a solution ofNH4NO 3to prepare NH4Mordenite which was then exchanged with the Pt and Co salts, respectively, to finally obtain PtCoH-Mordenite. The exchange time for each ~rnple was 24 h at room temperature and pH = 5-6, the zeolite/solution ratio employed was 2g/dm3 and all the solids were filtered, washed and finally dried at 120~ for 8 h. Monometallic ~mples with Co were pre-treated according to the standard method, i.e., heating at 2~ in oxygen flow up to 110~ with isotherm of 2 h, then up to 210 ~ with the same isotherm, and finally at 400~ for 8 h. Mono and bimetalfic solids with Pt were calcined following the method reported by Callezot et aL (14) consisting in r a j ~ g temperature at 0.5~ in oxygen flowup to 350~ keeping this temperature for 2 h. The PtCoNH, M samples was then treated in oxygen at 450~ for 8 hs. The Co/celite sax~le was prepared byinvregnating Co(NO3)2 on the low surface silica (17 m2/g). X-rayDiffraction Diffractogramswere obtained with a S h i m a ~ XD-D 1 insmmaent with monochromator using ChtKa~ radiation. It was operated in continuous scan mode at 0.5 ~(20) min". TPRExperiments. They were performed with 0.100g of catalyst with an Okura TP-2002 S insmnnent, with a heating rate of 8~ using as reducing gas a 4.8% H 2in argon flow. All samples were calcined according to the standard method, prior to each e x p ~ t . XPS Spectra. They were obtained with an Esca 750 Shimadzu instnnnent, using MgKcx radiation This spectrometer is driven by a computer system (Escapac 760)~41ich allows both the accunmlation and the processing of data. The spectra were obtained at room temperature. Platinum was analyzed in the area corresponding to the binding energy of the 4forbitals, because the signals corresponding to Pt4d overlap with the AUGER signals o f Na. The (Co/Si)s atomic ratios were calculated using the area under the Co2p, Pt4f and Si2p peaks, the Scofield photoionization cross sections the mean bee paths ofthe electrons and the i n s t a l function was given the ESCA manufacttlre:l'. The samples were subjected to reduction treatment with H 2in the reaction chamber directly attached to the ESCA spectrometer. CatalyticMeasurements. The reaction was carried out using 0.500 g of catalyst placed in a fixedbed flow reactor. This was a 12 mmi. d. tubular quartz reactor with an internal thermowelL The typical reacting mixture consisted of 1000 ppm of NO, 1000 ppm of CH, and 2-10% of 05, balanced at 1 atm with He (GHSV = 6500 h-~). The catalytic activity and the composition ofthe reacting gases were analyzed with a Varian 3700 chromatograph with 2 cokum~, one with a 5 A molecu~~ sieve, and the other containing Chromosorb 102. The NO x conversion (C~o) was calculated from the N, production. Selectivity was defined as S = Cso/Ccm where CcH4is the CH 4to CO x conversion. For all the solids analyzed, the only reaction products were N: and CO s, N:O was not detected and the carbon balance was always better than 98%. The conversion reported was determined after reaching steady state (usually after 1 h oftime-on-stream). A reaction cycle was defined as follows: The sample was kept on stream increasing temperature flora 200~ to 500~ staying 1 h at each ten~erature. After the highest temperature was reached the solid was treated in O 5 o.v. at 450~ and cooled up to 200~ in He flow. In some cases, after pre~treatment, the samples were reduced in pure ~ flow for 1

633 h at the desired temperature (350 or 500~ Table 1. Prepared Solids. Samples

'

Metallic Content

% b

Nonfiml unit cell composition

Co

Pt

Pt(o.5)M

2.51 -

0.50

Pt(o.~)Co(2.o)M

2.05

0.56

Co(2.o)Pt(o.5)M

2.96

0.57

Co,.63Pto.ogNa3.56(AdO2)7(si0 2)41

Ptr

1.73

0.50

Pt0.09Co o.9oH,.o,(AI 0 2)7(Si02 ),u

COo.o/Celite

3.00

-

Co(2.o)M

C o,.3Na4.4(AIO2)7(SiO2)4, Pto.09Na6.r~(A10,)7(SiO2)4,

Pto.o9C~

(SiO2)4,

aSubscripts represent theoretical % in weight of cation. bDeterminedby atomic absorption and CPS. 3. RESULTS In agreement with the results obtained by other authors, CoM was active for SCR o f NO xwith CH4 (3,5). In fact, Co~2.0)Mwith 42% of B.E.C. (Base Exchange Capacity) presented a ma3dmum NO to N: conversion of 31% at 450~ (Fig. 1), the selectivity at this temperature being 0.74. When the sample was reduced at 350 ~ and 500~ a decrease of NO conversion was observed (Table 2). The incorporation of 0.5% ofl~ (2.25% of B.E.C.) increased the initial activity ofthe oxidized sample. How100 ever, after various reaction cycles the NO to N 2 cow 90 version stabilized at a value close to that of Cot2.0)M but with a si~ificantly lower s e l ~ . 80 When Ptt0.~)Cot2.0)M was reduced at 350~ for =~ 70 1 h in I-[2flowthe result obtained was totally different. In effect, a si,.~nificant increase of NO conversion and selectivity was observed (Fig. 1). Besides, the sample remained stable during the successive reaction cycles and 30 also through time at constant temperature. The order in 9 ~ 20 which the exchange was performed did not have signifilo / cant importance on the N: production but increased the CH4to CO, conversion in the whole temperature range. 0 350 400 450 500 5 5 0 In order to explore the effect of acid sites on TEMPERATURE, ~ bimozllic solids a san~le was prepared fromthe annnonic Fig 1. Nitric Oxide ( I-l, II) and methane form of the mordenite. In this case, a decxease of about ((3 ,O ) conversion as a function of tem30% in the NO conversion was observed when it was perature on : CoCo>Fe corresponds with literature [4]. The N20 pressure dependency for Co-ZSM-5 is given in figure 2. Due to the integral reactor behaviour the relation between conversion and partial pressure shows a curvature, but the reaction order equals 1 for Co, and slightly lower

644

1.0 t X(N20)

1.0

Cu-ZSM-5 .'"

0.8 f

0.8 X(N20)

,/

9

,, ""

0.6,,

"

Co-ZSM-5

,r

9

/ 650

0.2 L

5, 700

800

750

T/K

773 K

0.6 o4 9l

0.4

0.2'

Co-ZSM-5

o.~u . u .......

/~I--~...=..J"

.........".......

./" ..-"" o;s' o.o~

' '

~

.4................... . ..........

' o:, 0.1 ....

748 K 723 K

698 K o.~s 0.15 ....

0.2

P(N20 ) / kPa

Figure 1 Conversion as a function of Figure 2 Partial pressure dependency of temperature 0.1 kPa N20 and space time Co-ZSM-5 at a space time of 1.52"105 1.44" 10 s g.s/mol g.s/mol

Table 2

Apparent activation energies (kJ/mol) and reaction orders only N20 3%02 CO/N20=2 NO/N20=l.5 all data *) order n *) Co 106 110 115 134 106+15 1 Cu 132 170 187 140 138+ 17 0.88_-_-~.11 Fe 173 187 78 182+31 0.79!-0.15 *) 95% confidence limits values are found for Fe and Cu. The fitting results of apparent reaction orders and activation energies 08 ~ ~ ~-''~'743 K for the different experiments are given in table 2. The presence of 02 hardly affects the reaction over Fe- and Co-ZSM-5, but it inhibits slightly for the 0.4i / F~ZSM-5 / Cu system, although this effect seems to level off at higher oxygen concentrations (figure 3). The 0.21- . ~ 7 ~ - - : _ -. . . . . ---._ ~ apparent E,, for Cu increased by nearly 40 kJ/mol. 773 K In the presence of NO the reaction is enhanced for ~176 2 4 6 8 10 Fe, and not affected for Co and Cu (figure 4). The P(02) / kPa reaction over Fe-ZSM-5 was accelerated Figure 3 The effect of oxygen on the tremendously; at temperatures where the N20 conversion at 0.1 kPa N20 and decomposition does not noticeably take place high space time W/FN2o=2.87* 105 g.s/mol, conversions were obtained by addition of NO. In the product mixture of the Fe and Co samples NO2 was observed. With Cu-ZSM-5 no NO2 was found. Figure 5 gives a product composition for Co-ZSM-5. Nearly all the converted NO and N20 results in NO2 and N2. Similar results are obtained for Fe-ZSM-5. Addition of CO also enhances the N20 conversion, by about a factor of two for Co and tremendously for Fe (figure 6). For Cu a maximum in the N20 conversion appears as a function of the CO/N20 ratio in the feed. This maximum shifts to higher values with increasing temperature (figure 7). The apparent activation energy for Co is hardly altered, for Fe it decreased nearly 100 kJ/mol, while for Cu it is increased by 50 kJ/mol. The presence of H20 inhibits the reaction and gives rise to deactivation for Cu and Fe. 1.0

xl.o,

"-

645

Concentration / ppm 70o

1.01 k

0.8

/

~9

600

l X(N20) ( 0.6~-

Fe-ZSM-5 (673 K) ,~

0.4~

~

500

~'IF~''-'-'-'-'-~

?--.L-" ....... L

30O

"

2O0

Co-ZSM-5 (723 K) ............ A

100

"

....

......

015 . . . .

z

..........

1'.0 . . . .

NO

400

.. Cu-ZSM-5 (673 K)

.

T / o~1

Co-ZSM-5

11 . . . .

2.0

N2 NO 2

, .~~~_.,_~__/?

~io

O~ ,

700

molar NO/N 2~ ratio

750

800

T/K

Figure 4 Effect of NO on the N20 Figure 5 Product composition for a conversion at 0.1 kPa N~O and space NO/N20 feed mixture over Co-ZSM-5 at time W/FN2o= 1.52"105 g.s/mol. Included 0.1 kPa N20 and space time W/FN2o= 1.52" 105 g.s/mol. is the NO conversion (open symbols). 1.0

1.01 L

Cu-ZSM-S (673 K ~ _ 0.8 ~//I i "~'--...._ X(N20)

i

k

jr"

./

!

f

>1 the rate determining step is eq. (3) and the sites are oxidised, while for kdk2 1

(9)

with: E °~' = E ~ l - a n ~

So, in the latter case the apparent activation energy is increased by the heat of adsorption of CO, amounting to about 40-60 kJ/mol as calculated from the IR experiments. Hence, for both the Co and the Cu samples Eal is slightly larger than Ea2 (table 2) while for iron Ea~ is considerably lower. All these values are compatible with values reported in the literature for Fe-zeolites [6,7,10,11 ] or dilute solid solutions of Co in MgO [31 ]. The kinetic and IR results with NO indicate that, like CO, it can remove the oxygen from the

649 surface of the Co and Fe catalyst, too, thereby forming NO2. The overall reaction is given by eq. (10). Although this occurs less efficient than with CO, it is thermodynamically quite feasible [3]. The NO does adsorb on the Co catalyst, evidenced by IR, yielding a slight increase in apparent activation energy. So, although at 723 K no enhancement is observed, but some is expected at higher temperatures. Kinetically, the blocking of sites by the observed NO adsorption on the Fe catalyst is negligible compared to the enhancement achieved by the oxygen removal effect. Over Cu-ZSM-5 steady state NO2 formation is not observed. Speculation offers two explanations; either it is not formed or it decomposes back to NO. IR measurements show by the presence of the 2134 cm -~ band that NO2 can be formed on CuZSM-5 [29,32], but from TPD experiments it appears that it decomposes rapidly in the range of 600-700 K [33]. So, in our case this could mean the occurrence of reaction (10) backwards or NO2 reacts with an oxidised site to NO and 02, acting as an oxygen carrier and hence competes with N20 in this respect. This does, however, neither result in an acceleration of the N20 decomposition nor in a changing temperature dependency. The NO2 formation offers the potential use of these catalysts in nitric acid plants off-gas treatment, where about equal amounts of NO and N20 are present. The produced NO2 can be reused in the nitric acid process [3]. SO2 also increases the N20 conversion in the case of Fe, probably according to eq. (11), although no product analysis is available. Co-ZSM-5 is inhibited by the SO2 or SO3, but not irreversibly. Copper is completely deactivated, as for other reactions over copper zeolites. NO

+

NzO

-->

SOE

+

O*

--~

NO2

+

Nz

(10)

SO3

+

*

( 11 )

l0

Water exerts both a deactivating and inhibiting influence on Cu and Fe samples, while the reaction over Co is only inhibited. The deactivation of Fe- and Cu-ZSM-5 is clearly due to migration and the sintering of the active component in H20 atmospheres [34]. The Co-ZSM-5 catalyst is much more hydrothermally stable in wet gas conditions [34,35]. The inhibition by water can be accounted for in a similar way as for CO via competitive adsorption on active sites, like in selective NO reduction studies [34]. For N20 decomposition this yields an expression like eq. (12). At 793 K Kl2 amounts to about 0.7 kPa ~ .

r=

klNrPu2o kl 1 + ~ + K12PH20 k2

(12)

Evaluating the results a clear kinetic picture of the catalysts has been obtained. In the steady state the active sites in Fe- and Cu-ZSM-5 are nearly fully oxidized, while for Co only -50% of the sites are oxidized. The former catalysts operate in an oxidation reduction cycle, Fe2+/Fe 3§ and Cu+/Cu 2§ Co a§ in zeolites is hardly oxidized or reduced, but ESR studies on diluted solid solutions of Co in MgO indicate that Co3§ formation is possible, rapidly followed by a migration of the deposited oxygen to lattice oxygen and reduction back to Co 2§ [36]. For Fe-ZSM-5 such a migration has been observed, so a similar model can be proposed for the zeolitic systems. Furthermore, it is obvious that application of these catalysts strongly depends on the composition of the gas that has to be treated.

650

ACKNOWLEDGEMENT These investigations have been supported by the European Union under contract no. JOU2CT92-0229 and J. R.-M. by a Human Capital & Mobilty grant.

REFERENCES P.L. Crutzen, J. Geophys. Res. 76 (1971) 7311. [1] G.G.d. Soete, Rev. Inst. Franc. Petr. 48 (1993) 413. [2] F. Kapteijn, J. Rodriguez-Mirasol and J.A. Moulijn, Appl. Catal. B: Env. accepted (1996) [3] Y. Li and J.N. Armor, Appl. Catal. B: Env. 1 (1992) L21. [4] J. Valyon, W.S. Millman and W.K. Hall, Catal. Lett. 24 (1994) 215. [5] V.I. Sobolev, G.I. Panov, A.S. Kharitonov, V.N. Romannikov, A.M. Volodin and K.G. Ione, [6] [7] [8] [9] [10]

[11] [12]

[13] [14]

[15] [16] [17]

[18] [19] [20] [21] [22] [23]

[24] [25]

[26] [271 [28]

[29] [30] [311 [32] [33] [341 [35] [36]

J. Catal. 139 (1993) 435. J. Leglise, J.O. Petunchi and W.K. Hall, J. Catal. 86 (1984) 392. C.M. Fu, V.N. Korchak and W.K. Hall, J. Catal. 68 (1981) 166. Y.-F.Chang, J.G.McCarty, E.D.Wachsman and V.L.Wong, Appl. Catal. B: Env. 4 (1994) 283. G.I. Panov, V.I. Sobolev and A.S. Kharitonov, J. Mol. Catal. 61 (1990) 85. Y.-F. Chang, J.G. McCarty and Y.L. Zhang, Catal. Lett. 34 (1995) 163. R.L. Garten, W.N. Delgass and M. Boudart, J. Catal. 18 (1970) 90. R.A. Dalla Betta, R.L. Garten and M. Boudart, J. Catal. 41 (1976) 40. A. Cimino, Chimica e Industria, 56 (1974) 27. F.S. Stone, J. Solid State Chem. 12 (1975) 271. J.O. Petunchi and W.K. Hall, J. Catal. 78 (1982) 327. G.D. Lei, B.J. Adeiman, J. Sarkany and W.M.H. Sachtler, Appl. Catal. B: Env. 5 (1995) 245. G. Spoto, A. Zecchina, S. Bordiga, G. Ricchardi and G. Martra, Appl. Catal. B: Env. 3 (1994) 151. M.A. Kohler, N.W. Cant, M.S. Wainwright and D.L. Trimm, J. Catal. 117 (1989) 188. E.E. Mir6 and J.O. Petunchi, J. Chem. Soc. Faraday Trans. 88 (1992) 1219. M. Anpo, M. Matsuoka, Y. Shioya and H. Yamashita, J. Phys. Chem. 98 (1994) 5744. M.W. Anderson and L. Kevan, J. Phys. Chem. 91 (1987) 4174. M.A. Makarova, E.A. Paukshtis, J.M. Thomas, C. Williams and K.I. Zamaraev, J. Catal. 149 (1994) 36. Y.-Y. Huang, J. Am. Chem. Soc. 95 (1973) 6636. Z. Chajar, M. Primet, H. Praliaud, M. Chevrier, C. Gauthier and F. Mathis, Appl. Catal. B: En v. 4 (1994) 199. Y. Li, T.L. Slager and J.N. Armor, J. Catal. 150 (1994) 388. K.-I. Segawa, Y. Chen, J.E. Kubsh, W.N. Delgass, J.A. Dumesic and W.K. Hall, J. Catal. 76 (1982) 112. M. Iwamoto, H. Yahiro, N. Mizuno, W.-X. Zhang, Y. Mine, H. Furukawa and S. Kagawa, J. Phys. Chem. 96 (1992) 9360. T.E. Hoost, K.A. Laframboise and K. Otto, Catal. Lett. 33 (1995) 105. P.A. Jacobs and H.K. Beyer, J. Phys. Chem. 83 (1979) 1174. A. Cimino and F. Pepe, J. Catal. 25 (1972) 362. J. Valyon and W.K. Hall, J. Phys. Chem. 97 (1993) 1204. Y. Li and J.N. Armor, Appl. Catal. 76 (1991) L 1. Y. Li, P.J. Battavio and J.N. Armor, J. Catal. 142 (1993) 561. J.N. Armor and T.S. Farris, Appl. Catal. B: Env. 4 (1994) L11. V. Indovina, D. Cordischi, M. Occhiuzzi and A. Arieti, J. Chem. Soc. Faraday Trans. 1, 75 (1979) 2177.

J.W. Hightower, W.N. Delgass, E. Iglesia and A.T. Bell (Eds.) 11th International Congress on Catalysis - 40th Anniversary

Studies in Surface Science and Catalysis, Vol. 101 9 1996 Elsevier Science B.V. All rights reserved.

651

O n t h e role o f free r a d i c a l s NO2 and 02 in t h e s e l e c t i v e c a t a l y t i c r e d u c t i o n (SCR) of NOx w i t h CH4 o v e r C o Z S M - 5 and H Z S M - 5 z e o l i t e s Dmitri B. Lukyanov ~ Julie L. d'ltri b, Gustave Sill c and W. Keith Hall c" "Department of Chemistry, UMIST, P.O. Box 88, Manchester M60 1QD, UK bChemical Engineering Department, and =Chemistry Department, University of Pittsburgh, Pittsburgh, PA 15260 USA The reactions of CH4 over CoZSM-5 and HZSM-5 zeolites with the mixtures of NO + O2 and NO2 + 02 and with three oxidizing components separately were studied. Differential reaction rates were determined. Comparison of the "light-off" temperatures as well as the activation energies of these reactions led to the conclusion that the SCR of NO. into N 2 and, consequently, CH 4 oxidation into COx are initiated by the reaction of CH 4 with NO 2. At low temperatures (300-400~ 02 does not compete with NOx for CH,, and its role is limited to the oxidation of NO into NO2. However, at higher temperatures a strong competition between NOx and 02 for CH4 results in a decrease in the selectivity of the SCR process. It is shown that this competition is stronger with CoZSM-5 catalyst than with HZSM-5, and this explains the higher selectivity observed with the latter catalyst. Based on these observations the formation of the CH3o free radical is postulated and possible pathways of the SCR of NOx into N 2 are discussed. 1. INTRODUCTION

Recently, CoZSM-5 and HZSM-5 zeolites were reported [1-4] to be effective catalysts for the SCR of NO with methane in the presence of oxygen. It was shown that 02 greatly enhances the NO conversion into N 2, and an important role of NO2 in the initiation of the reactions has been delineated [3-9]. In the present paper we wish to shift the focus from reduction of NO into N 2 to the competitive oxidation of the hydrocarbons [4] by NO. vis-a-vis 02. Evidence favoring NO2 as the dominant reagent has been presented [6-10], but the reaction of CH4 with 02 alone has not been considered. Hence it was of interest to better define the role of NO2 and 02 in the SCR of NO. with CH4. Consequently, the relative effectiveness of the several oxidizing agents has been determined for the reactions of CH 4 over CoZSM-5 and HZSM-5 catalysts. The same reactions were studied over the NaZSM-5 zeolite and in the homogeneous gas phase [10]. " To whom all correspondence should be addressed.

652

2. EXPERIMENTAL The reactions were carried out in the steady state flow mode as described previously [ 11 ]. Differential kinetics were determined from plots of conversion vs. W/F. Three catalysts CoZSM-5, HZSM-5 and NaZSM-5 (Si/AI = 11) were studied in this work. The catalyst preparation and the standard pretreatment used prior to reaction have been described previously [11]. It involved dehydration in flowing dried 02 as the temperature was raised slowly to 500~ The feed comprised CH 4 (0.28%), NO (0.21%) or NO2 (0.21%), and/or 02 (2.6%) in He. The flow rate was 75 ml/min and the gas hour space velocity (GHSV) was varied between 4 , 5 0 0 and 2 5 0 , 0 0 0 h 1 by changing the weight of catalyst samples.

3. RESULTS AND DISCUSSION 3.1. "Light-off" temperatures of CH 4 combustion in the presence of different oxidizing compounds Studies of CH4 reactions with NOx in the presence or absence of 02 have shown that the "light-off" temperature of methane combustion coincides in every case with the temperature at which N2 formation is initiated [10]. At a GHSV of 22,500 h 1 with CoZSM-5 catalyst this temperature was about 300~ (see Fig. 1) regardless of the nitrogen oxide used. Moreover, in the absence of 02 the "light-off" occurred with NO2 at the same temperature while with NO 450~ was required. In the absence of NO,,, oxidation of CH4 by 02 was observed at about 400~ (Fig. 1B). Taken together with the fact that over Co-containing zeolites [12] NO can be oxidized by 02 into NO2 at temperatures as low as 200~ our data suggest that the SCR of NO into N2 and, consequently, CH4 oxidation into CO2 are initiated by NO2. During NO oxidation to NO2, the catalyst is definitely involved [ 10]. Hence, O atoms held by the catalyst may interact with CH4 molecules, but as shown in Fig. 1 B, not as efficiently as when NO2 is present. Hence, it may be concluded that it is NO2 that initiates the SCR reaction. The same conclusion may be drawn on the basis of our data for the SCR reaction over HZSM-5 catalyst, although in this case the "light-off" temperature is higher. At a GHSV of 22,500 h ~ it is about 350~ for the reactions of methane with the mixtures of NO~ + 02 and with NO2 alone. With NO or 02 alone CH4 oxidation is observed only at temperatures higher than 550~ Evidently Brensted acidity is not essential.

3.2. Coupling of Combustion and SCR From the results discussed above as well as from the literature data [5-10,12-14] it follows that an important role of 02 in the SCR process is to convert NO into NO2. The latter then initiates methane oxidation into COx and is itself reduced into NO and N2. Both NOx and 02 may participate in CH4 oxidation (Fig. 1 B) and the ratio between the rates of these competitive oxidation reactions will be critical for the selectivity of the SCR process. Hence, the absolute rates of CH4 oxidation by 02 were compared with those occurring in the SCR process. The rates of these reactions were determined under different reaction conditions (using the

653

100 v

o~ 80 z o =: 6 0 m.. X

~100

A

0 CH4*NO*O 2

9CH4*NO CH4*N02*0 2 A CH4*N0 2

0

~

O z 40-

c

6O

0

40

0

O

:; 2 0 -

~

C:

C 0

0

o

0 260

860 460 Temperature (OG)

550

0

2O 0 2gO

460 T e m p e r a t u r e (OG) ,950

660

Figure 1. CH4 reactions with different oxidizing compounds over CoZSM-5 catalyst; conversion of NOx into N 2 (A) and of CH 4 into CO 2 (B) as a function of temperature. Catalyst weight was 100 mg, feed contained 0.28% CH4, 0.21% NO or NO2 (when used), and 2.6% 02 (when used) in He at a flow rate of 75 ml/min (GHSV = 22,500 h'l).

E..2o, ko.,,mo, I 4

~

Ea -14.5 kcal/mol

I

2

2. z

=.. o -

_j

E -14.2 kcal/mol 8

...I

0

1.2

-2[.

1.4

~.e

1000 IT (K -1)

.1

E=-29 kcallmol 1 .$

1.6 IO00/T (K -11

B

1.7

Figure 2. Arrhenius plots of differential rates of NO reduction (e) and of CH4 oxidation (o) during the SCR reaction, and for CH4 oxidation by 02 alone (A) over CoZSM-5 (A) and HZSM-5 (B) catalysts. Feed contained 0.28% CH4, 0.21% NO (when used) and 2.6% 02 in He.

654

linear portions of conversion vs. W/F plots) and are compared in Fig. 2. Several points n o w become clear. First, CH4 oxidation into COx and NO reduction into N 2 proceed with the same activation energy and, below 500~ at about the same rates (Fig. 2) indicating that these two reactions are coupled and have the same rate limiting step. Second, the Arrhenius dependence for NO reduction into N 2 is valid up to 500~ i.e., in the temperature range where thermodynamics favors NO2 over NO [6,13]. At higher temperatures the equilibrium constant becomes < 1 causing the observed deviation from linearity.Third, the activation energy of the SCR reaction is higher with CoZSM-5 catalysts than with HZSM-5, and reaction of CH4 with 02 has much the highest activation energy. Fig. 1B shows that at all temperatures the rate of CH4 oxidation by 02 alone is lower than the rate of CH4 oxidation during the SCR reaction, e.g., at 400~ with CoZSM-5 catalyst the difference between these rates is about 10 times. With increasing temperature this difference diminishes due to the different activation energies of these reactions (Fig. 2). At high temperatures these rates become comparable (in considering Figs. 1B and 2 recall that the rate of CH 4 oxidation during the SCR process includes a contribution from the rate of CH4 oxidation by 02 alone). These data suggest that below 5000C 02 does not compete effectively with NO,, for CH4, but that at high temperatures such a competition must exist. The data of Table I support this view. At 400~ an increase in 02 concentration results in an increase in conversions of both NO into N 2 and CH 4 into CO2. At the same time, variation of O2 concentration by a factor of 13 has practically no effect on the Table 1 Effecl; of 02 r

on the $CR reaction* over CoZSM-5""

At 400~ A B C

Conv. of NO into N 2 (%) Conv. of CH4into CO= (%) NO convCrl;ed into N~(mol) CH 4 converted into C02 (mol)

0.4

Concentration of 02 (%) 1.0 2.6 5.2

13.8 6.4 1.62

17.9 8.6 1.56

20.4 9.7 1.58

20.7 10.3 1.51

21.5 12.5 1.29

25.8 19.2 1.01

24.1 28.8 0.63

21.9 44.2 0.37

At 550 ~ D Cony. of NO into N 2 (%) E Cony. of CH4 into CO2 (%) F N O converted inl;o N~(mol) CH. converted into CD2 (mol)

*The reactant gas contained 0.28% CH4 (9.375 pmol/min), 0.21% NO (7.03 pmol/min) and x% 02 in He at a total flow rate of 75 ml/min. * *Steps A, B and C correspond to 400~ and GHSV of 45,000 hl; steps D, E and F correspond to 550~ and GHSV of 250,000 h~.

655

selectivity of the SCR process (defined as a ratio of the number of NO molecules converted into N2 to the number of CH4 molecules converted into CO2). This result can be easily Understood in terms of the higher rates of NO2 formation at the lower temperature (and, consequently, in the higher rates of the SCR process) at higher 02 concentrations. A different picture is observed at 550~ In this case an increase in 02 concentration results mainly in an increase in CH4 combustion (due to the higher activation energy for this reaction) and, consequently, in a decrease in the selectivity of the SCR process. Fig. 3 shows the effect of temperature on the selectivity of the SCR of NO with methane in the presence of oxygen. The data show that the highest selectivity of the SCR process is observed at low temperatures and low conversions of the reactants (NO and CH4) and does not differ greatly for CoZSM-5 and HZSM-5 catalysts. This fact can be easily understood since under these conditions CH4 should be oxidized, according to our data, solely by NOx. An increase in temperature or a decrease in NOx concentration should be followed by enhancement of the competition between 02 and NO2 for methane, and, consequently, by a decrease of the SCR selectivity as observed. Moreover, it follows from Fig. 2 that at any given temperature the difference between the rates of CH4 oxidation by the mixture of NOx + 02 and by 02 alone is larger for HZSM-5 catalyst than for CoZSM-5 . Thus, the competition between NOx and 02 for CH4 is stronger with CoZSM-5 than with HZSM-5. This explains the fact that the SCR selectivity (Fig. 3) is higher with HZSM-5 than with CoZSM-5.

r

8

"~" 8 [ 9400~ J O 450~

9400~ O 450~

.=..

E --~ 6

L/

O

E

o z

"

65ooc

4

q..

0

O

:; 2 r O

o

o

0

2

4

6

8

10

C o n y . of CH 4 ( ~ m o l l m l n )

g'o

0

2

4

6

8

10

C o n y . of CH 4 (~mollmln)

Figure 3. Effect of temperature on the selectivity of the SCR reaction over CoZSM-5 (A) and HZSM-5 (B) catalysts. Feed contained 0 . 2 8 % CH4, 0 . 2 1 % NO and 2.6% 02 in He at a flow rate of 75 ml/min ( flow rates of CH4 and NO were 9.375 and 7.03 pmol/min).

656 100

100

80

8O

v r

O z

6O

O 60 z

O

40

o 40

O

20

o 2O

q=-

o

O

0 260

360

0

450

660

Temperature (oc)

660

250

380

460

650

660

Temperature (~

Figure 4. Conversion of NO2 into different products by reaction w i t h CH4 over C o Z S M - 5 (A) and HZSM-5 (B) at various temperatures. Feed contained 0 . 2 8 % CH 4 and 0 . 2 1 % NO2 in He at a f l o w rate of 75 ml/min. GHSV was 2 2 , 5 0 0 h "1.

6p

4

A

ol

,

9NO 2 into N2 O NO 2 into NO A CH4 into COx

B 4

A

~ 3

_.52

kcal/mol

=,21 -

'

0

1.3

i

1.4

1.5

1.6

IO001T (K "1)

1.7

1.8

0

,

1.3

1.4

EI I -13.6 kcal/mol i

1.6

I

1.6

I

1.7

1.8

I O 0 0 / T (K "1)

Figure 5. Arrhenius plots of differential rates of NO2 reduction into NO and N 2, and for CH4 oxidation into COx over CoZSM-5 (A) and HZSM-5 (B). Feed contained 0 . 2 8 % CH4 and 0 . 2 1 % NO2 in He.

657

3.3. Reaction of CH4 with NOz

This reaction was studied in detail and some typical results are shown in Fig. 4. Previously, it was demonstrated [10] that in the absence of CH4, NO2 achieves equilibrium with NO + 89 02 over CoZSM-5 and HZSM-5 catalysts at temperatures higher than 350~ As shown in Fig. 4, in the presence of CH4 t w o additional reactions occur: (a) reduction into NO, and (b) reduction into N 2. The oxygen released during these two reactions is used to oxidize CH4 into CO2over CoZSM-5 catalyst or into CO2 + CO over HZSM-5. In the presence of CH4the conversion of NO 2 into NO + 89 decreases significantly and at 400~ and a GHSV of 22,500 h1 this decrease is about 9 times with CoZSM-5 (Fig. 4A) and 3 times with HZSM-5 (Fig. 4B). Invariably more of the NO 2 is reduced into NO than into N 2. At low temperatures (300-400~ the rate of this conversion, under differential conditions, was about 2 times higher than the rate of N2 formation. With increasing temperature the concentration of 02 in the tail gas decreases (probably, due to the reaction with CH4) and at temperatures higher than 500~ it is not observed in the reaction products. Simultaneously, the total conversion of NO 2 (into NO and N 2) reaches 100% at about 500~ Temperatures above 500~ resulted in a decrease in NO2 conversion into NO and in a corresponding increase in the conversion into N 2. Over HZSM-5 (Fig. 4B) this was not observed. These data clearly show that NO2 reduction with CH4 into NO and into N2 proceeds over CoZSM-5 and HZSM-5 catalysts via processes that have at least one common reaction step, viz., the activation of CH4. It is of interest to compare the activation energy data for the reactions of NO2 (Fig. 5) with those for the SCR reaction (Fig. 2). Within experimental error, they are the same; 22 kcal/mol and 21 kcal/mol, respectively for CoZSM-5 and 14 kcal/mol for HZSM-5 processes. The conclusion that combustion of CH4 and reduction of NOx processes are coupled and have the same rate limiting step is strongly supported by these facts. We therefore conclude that the SCR of NO with CH 4 in the presence of 02 is initiated by the reaction of NO2 with CH 4. Additional information on the reaction system (NO2 + CH,) follows from Fig. 4A. These data were obtained over CoZSM-5 catalyst at a GHSV of 22,500 h 1 . At this space velocity and at temperatures higher than 450~ the direct reaction between NO and CH 4 occurs resulting in N 2formation. Obviously, this reaction must also occur in the system of NO2 + CH4 and explains a decrease in NO2 conversion into NO and a corresponding increase in conversion into N 2 at temperatures higher than 550~ However, in the SCR of NOx in the presence of 02 the reaction between NO and CH, does not play an appreciable role, since it proceeds only at high temperatures and with a rate much lower than the rate of CH4 oxidation by 02. Interestingly, Shelef et al. [7] fed approximately equilibrium mixtures of NO and NO2 with and without 02 over their catalyst (CuZSM-5) together with C3H8 and found most of the NO 2 fed had been converted into NO or N2; the system moved away from equilibrium by the selective reaction of NO2 with hydrocarbon faster than it could be produced by NO reacting with 02.

658

3.4. Pathways of the SCR of NO with CH4 in the presence of O z We have already established that oxidation of CH 4 and consequently the catalytic reduction of NOx are initiated by interaction of the former with NO2. At a GHSV of 2 2 , 5 0 0 h 1 "light-off" occurs at 300 ~ and 350~ with CoZSM-5 or HZSM-5 catalysts, respectively. In the empty tube the "light-off" of CH4 combustion with NO 2 occurs about 450~ [10]. Under these conditions the homogeneous oxidation of CH4 and the conversion of NO2 into NO (formation of N 2 was not observed) increase sharply with temperature and at 600~ reaches 100%. With NaZSM-5 catalysts this picture does not change much, the sole difference from the homogeneous reaction being the difference in the products of CH4 oxidation. In the presence of the catalyst, CH 4 is oxidized into CO and CO2 (CO/CO2 ratio is about 2) and the carbon balance closed at all temperatures. In the empty tube the carbon balance did not close at 500 ~ and 550~ possibly due to the formation of formaldehyde [ 10]. Thus, we suggest that over NaZSM-5 catalyst a rapid oxidation of intermediate oxidation products occurs. Yokoyama and Misono [8] reported that N 2 was formed in the reaction of NO2 with C3He over NaZSM-5. Many years ago Wojciechowski and Laidler [15] studied the homogeneous decomposition of CH4 and C2He in the presence of NO and concluded that the stable free radical NO could abstract an H atom from a paraffin molecule. Since NO2 is also a stable free radical and a much stronger oxidizing agent than NO, the reaction CH4 + NO2 --" C H 3 9+ "OH + NO (or HONO)

(1)

seems very likely. This finds support in the work of Cant and co-workers [16,17] who found a first order isotope effect on the rate of N 2 formation when CD4 was substituted for CH4. Coupling of CH3 9radicals would not be expected because of much higher concentration of NO,, playing the role of a radical trap, and was not observed. The formation of CH3NO2 (and/or CH3NO) seems much more likely. Yokoyama and Misono [8] reported that the former, when introduced as a reactant, reacts with NO2 forming CO2 and H20. Similarly, formaldehyde may be produced by (2),

CH3 9 + ONO --- CH20 + HNO

and may react further with NO2 producing CO that, in turn, may be oxidized to CO2. i.e. CH20

+

NO 2 ~ CO +

CO + NO2 ~

CO2 + NO

H20 +

NO

(3), (4).

Equations (2) - (4), considered together with reactions like eOH + HNO ~ H20 + NO and CH3 9+ HNO -~ CH4 + NO, explain the homogeneous reduction of NO2 into NO and the corresponding oxidation of CH4 into CH20, H20, CO and CO2. The same

659

reactions may occur over NaZSM-5, but in this case, the rates of the reactions (3) and (4) are much higher than in the empty reactor. Moreover, it seems likely that the reactions (2)-(4) occur also over CoZSM-5 and HZSM-5 catalysts forming the pathway for NO2 reduction into NO. In contrast to NaZSM-5 zeolite, introduction of CoZSM-5 or HZSM-5 zeolite in the reaction system shifts the "light-off" temperature and modifies the chemistry; n o w not only NO but N 2 is formed. Hence, some intermediate species required for N 2 formation must be stabilized on the catalyst surface. The "light-off"temperature shifts observed with CoZSM-5 and HZSM-5 catalysts may result from the enhanced redox capacity provided by these catalysts or from the NOJNO equilibrium achieved more readily than with NaZSM-5. Moreover, equilibrium is approached at a somewhat lower temperature over CoZSM-5 than HZSM-5, and much lower than with the empty reactor (see Fig. 1 of Ref. lO).The decomposition reaction of NO2 into NO + 89 occurs readily on these catalysts and the "light-off" temperature of both combustion and SCR is lower in comparison with that of the homogeneous reaction. Summarizing, our findings permit us to speculate on possible mechanisms occurring in the SCR reaction. Over these catalysts NO 2 is reduced into NO and N 2. These t w o reactions are coupled and have the same rate limiting step, i.e., activation of CH4 by NO2. The results suggest that this reaction results in formation of the CH3e radical and this idea finds support in the isotope effects for SCR [17] vis-a-vis CH4 coupling [ 16] and in the fact that CH4 coupling catalysts also catalyze SCR[18,19]. Finally, recently reported [20] calculations on CH4 activation with an O atom adsorbed on a surface demonstrate that the abstraction of an H atom from a CH4 molecule with formation of CH3o free radical should be energetically preferable compared with the other possible mechanisms of CH 4 activation. Thus, the chemistry of Eq. (1) seems plausible. It is not unreasonable to think that free radical chemistry can occur within the zeolite pore system, just as it can in solution. However, the lack of molecular sieving effects with larger hydrocarbons [4] leads us to believe that part of the reaction occurs homogeneously. The NO formed as product may desorb from the catalyst, react with 02 to produce NO2, or with the CH3 9radical to produce CH3NO molecule. The adsorbed CH3NO molecule may be transformed into HCN via dehydration, i.e., Z-CH3NO---Z-HCN + H20

(5).

The formation of HCN and adsorbed CN species during the SCR process over CeZSM-5 and CuZSM-5 zeolites have been reported previously by several research groups [8,21,22]. HCN may undergo various transformations in the presence of NO2 and 02 [23,24]. Some of these produce eCN and eNCO radicals. The latter radical, as has been shown by Cooper et al. [25] reacts with NO producing the ON-NCO molecule which, in turn, decomposes thermally into N2 and CO2. The transient experiments performed by Yogo and Kikuchi [9] over GaZSM-5 and InZSM-5 together with the results of Hayes et al. [22] suggest that eCN may play a role in forming the N-N bond.

660 Another possible path for the formation of eNCO radical may include formation of CH3NO2 which may undergo transformation on the active sites of the catalyst into H20 and HC(N)O species. The latter may be converted, according to the literature data [24], into oNCO radical. In conclusion, although at present time we cannot unambiguously define the reaction steps leading to N2 formation, we have attempted to suggest steps that seem to be reasonable and hope that our proposals will be useful for further investigation of the detailed mechanism of the SCR process. Acknowledgement Thanks are due to the Department of Energy, Division of Basic Energy Sciences, under Grant No. DE-FGO2-95ER14539. REFERENCES o

2. 3. 4. 5. .

7. 8. 9. 10. 11. 12. 13. 14. 15. 16. 17. 18. 19. 20. 21. 22. 23. 24. 25.

Y. Li and J.N. Armor, Appl. Catal. B, 1 (1992) L31. Y. Li and J.N. Armor, Appl. Catal. B, 2 (1993) 239. K. Yogo, M. Umeno, H. Watanabe and E. Kikuchi, Catal. Lett., 19 (1993) F. Witzel, G.A. Sill and W.K. Hall, J. Catal., 149 (1994) 229. J.N. Armor and Y. Li, Preprints of Symposium on NOx Reduction. ACS, Division of Petroleum Chemistry, Inc., 39 (1994) 141. J.O. Petunchi and W.K. Hall, Appl. Catal. B, 2 (1993) L17. M. Shelef, C.N. Montreuil and H.W. Jen, Catal. Lett., 26 (1994) 277. C. Yokoyama and M. Misono, J. Catal., 150 (1994) 9. K. Yogo and E. Kikuchi, Stud. Surf. Sci. Catal., 84 (1994) 1547. D.B. Lukyanov, G. Sill, J.L. d'ltri and W.K. Hall, J. Catal., 153 (1995) 265. J.O. Petunchi, G. Sill and W.K. Hall, Appl. Catal. B, 2 (1993) 303. Y. Li, T.L. Slager and J.N. Armor, J. Catal., 150 (1994) 388. K.A. Betkhe, C. Li, M.C. Kung and H.H. Kung, Catal. Lett. 31 (1995) 287. H. Yasuda, T. Miyamoto and M. Misono, in "ACS Symposium Series 587", Chapter 9 (1994). B.W. Wojciechowski and K.J. Laidler, Can. J. Chem., 38 (1960) 1027. N.W. Cant, E.M. Kennedy and P.F. Nelson, J. Phys. Chem., 97 (1993) 1445. A.D. Cowan, R. Dumpelmann and N.W. Cant, J. Catal., 151, (1995) 356. X. Zhang, A.B. Waiters and M.A. Vannice, J. Catal., 146 (1994) 568. X. Zhang, A.B. Waiters and M.A. Vannice, Appl. Catal. B, 4 (1994) 237. M.Yu. Sinev, L.Ya. Margolis and V.N. Korchak, Russian Chem. Rev., 64 (1995) 349 (English translation). F. Radtke, R.A. Koeppel and A. Baiker, Appl. Catal. A, 107 (1994) L125. N.W. Hayes, W. Grunert, G.J .Hutchings, R.W. Joyner and E.S. Shpiro, J. Chem. Soc., Chem. Commun., (1994) 531. J.A. Miller and C.T. Bowman, Prog. Energy Combust. Sci., 15 (1989) 287. M.C. Lin, u He and C.F. Melius, Intern. J. Chem. Kinetics, 24 (1992) 1103. W.F. Cooper, J. Park and J.F. Hershberger, J. Phys. Chem., 97 (1993) 3283.

J.W. Hightower, W.N. Delgass, E. Iglesia and A.T. Bell (Eds.) 11th International Congress on Catalysis - 40th Anniversary

Studies in Surface Science and Catalysis, Vol. 101 9 1996 Elsevier Science B.V. All rights reserved.

661

An infrared study of N O reduction by CH4 o v e r C o - Z S M - 5 A. W. Aylor, L. J. Lobree, J. A. Reimer, and A. T. Bell Center for Advanced Materials, Lawrence Berkeley National Laboratory and Department of Chemical Engineering, University of California, Berkeley, CA 94720, USA An in situ infrared investigation has been conducted of the reduction of NO by CH4 over Co-ZSM-5. In the presence of 02, NO2 is formed via the oxidation of NO. Adsorbed NO2 then reacts with CH4. Nitrile species are observed and found to react very rapidly with NO2, and at a somewhat slower rate with NO and 02. The dynamics of the disappearance of CN species suggests that they are reactive intermediates, and that N2 and CO2 are produced by the reaction of CN species with NO2. While isocyanate species are also observed, these species are associated with Al atoms in the zeolite lattice and do not act as reaction intermediates. A mechanism for NO reduction is proposed that explains why O 2 facilitates the reduction of NO by CH4, and why NO facilitates the oxidation of CH4 by 02. 1. INTRODUCTION There has been considerable interest recently in the use of metal-exchanged zeolites for the selective reduction (SCR) of NO by methane [ 1-20]. Amongst the various catalysts tested, CoZSM-5 and Co-ferrierite have shown particularly high activity. Attempts to explain the mechanism by which CH4 reduces NO over these catalysts have largely been based on studies of reaction kinetics. Several studies have shown that the rate of NO reduction is considerably more rapid when 02 is added to the feed [1, 2, 10, 13, 14, 18], and that under such circumstances NO2 is formed rapidly by the reaction of NO with 02. It has also been found that the reaction of NO2 with CH4 occurs at a comparable rate to that observed for the reaction of NO with CH4 in the presence of 02 [ 13, 14, 18]. Based on this evidence it has been proposed that in the presence of 02 the reduction of NO is initiated by the oxidation of NO to NO2, then subsequent reaction of CH4 with either gas-phase or adsorbed NO2 [11, 13-16, 18]. Isotopic labeling studies have demonstrated that the reduction of NO in the presence of 02 occurs more rapidly with CH4 than with CD4, suggesting that C-H bond scission is very likely the rate-limiting step in SCR [17]. Infrared studies of NO adsorption on Co-ZSM-5 and Coferrierite have identified the presence of both mono- and di-nitrosyls at room temperature [ 15]. Over Co-ferrierite both of these species can be converted to adsorbed NO2 in the presence of oxygen at elevated temperatures. The adsorbed NO2 is thought to react with CH4 thereby producing CH3 radicals and adsorbed HNO2. More recently, studies involving infrared spectrsocopy and mass spectrometry have shown that CH4 reacts with adsorbed NO 2 to form N2 and CO2 [20]. The goal of this work was to determine surface species present under reaction conditions, and to investigate the interactions of adsorbed NO and NO2 with CH4. Infrared spectra were collected under reaction conditions, and in various mixtures of NO, Oz, NO2, and CH4. It was of particular interest to make direct observations of the factors affecting the formation of NO2 and its reaction with CH4, since NO2 has been suggested as an intermediate in the reduction of NO by CH4. A further objective was to elucidate the pathway by which N2 and CO2 are formed.

662 2. EXPER/MENTAL Na-ZSM-5 was obtained from UOP. About 15 g of the zeolite was added to a 3.5 L solution of 0.01 M cobalt acetate. This mixture was stirred at 25 ~ for 21 hr, then at 60 ~ for 33 hr, and finally at 70 ~ for 19 h [1]. The zeolite was faltered, washed, and dried overnight in a vacuum oven at 120 ~ Elemental analysis of the catalyst determined the Si/A1 ratio to be 18.4, and the Co/A1 ratio to be 0.28. For infrared spectroscopy, 20-50 mg of the cobalt-exchanged zeolite was pressed into a self-supporting wafer and placed into an infrared cell similar to that described by Joly et al. [21]. Spectra were recorded on a Digilab FTS-50 Fourier-transform infrared spectrometer at a resolution of 4 cm-l. Typically, 64 or 256 scans were coadded to obtain a good signal-to-noise ratio. A reference spectrum of Co-ZSM-5 in He taken at the same temperature was subtracted from each spectrum. Gases were supplied to the infrared cell from a gas manifold. 4.99% NO in He and 2.14% CH4 in He were obtained from Matheson. Oxygen and helium were obtained on-site. The He, NO, and CH4 cylinders were passed through an oxysorb trap, an ascarite trap, and a molecular sieve trap, in that order, for additional purification. The 02 was passed through an ascarite and a molecular sieve trap. Prior to each experiment the catalyst was: (1) heated at 500 ~ in 02 for about 30 min., (2) heated at 500 ~ in He for about 3 h, and (3) cooled to the desired temperature in helium. The temperature ramp experiments were run at about 0.6 ~ To determine its activity, the catalyst was placed in a quartz microreactor. Reactants were supplied through mass flow controllers and the product composition was determined by mass spectrometry. A typical reaction mixture contained 3,600 ppm NO, 1.06% CH4, and 6.0% 02, with the balance He. A 0.05 g sample of the catalyst was used with a total flow rate of 100 cm3/min, resulting in a GHSV = 60,000 (based on an apparent bulk density of the zeolite of 0.5 g/cm3). The conversion of NO was based on the amount of N 2 formed and the conversion of CH4 was based on the amount of CO2 formed. Carbon and nitrogen mass balances were closed to within 10%. 3. RESULTS Figure 1 shows the effects of temperature on the conversion of NO to N2. In the absence of O2 significant conversion of NO occurs above 500 ~ When 02 is present the light off temperature decreases to 350 ~ Similar light off temperatures are observed for the conversion of CH4 to CO2. The ratio of NO to CH4 consumption increases from 0.7 to 1.04 as the temperature rises from 550~ to 700~ for the reaction of NO with CH4 in the absence of O2. This ratio is significantly lower than 4.0, the ratio for the reaction CH4 + 4 NO -~ 2 N2 + CO2 + H20, suggesting that additional CO2 may be produced by steam reforming of CH4. In the presence of 02 in the feed, the ratio of NO to CH4 consumption decreases from 0.84 to 0.18 as the temperature rises from 400~ to 600~ Similar ratios of NO to CH4 consumption have been reported previously [2, 13], and the decline in this ratio with increasing temperature has been attributed to the production of CO2 as a consequence of CH4 combustion by O2. Figure 2 shows a series of infrared spectra taken during NO desorption from the catalyst in He after it had been exposed to 6,600 ppm of NO for 42 min at room temperature. At 25 ~ sharp bands are observed at 2132, 1941, 1894, and 1815 cm-1, together with shoulders at 1874 and 1799 cm-l, and a doublet centered at about 1400 cm-~. Weak bands are also seen at 1633, 1599 and 1528 cm-~. Purging the sample in He for 4 min at 25 ~ causes the disappearance of the band at 1633 cm-l, and an increase in intensity of the features at 1599 and 1528 cm-l.

663

N O + I:::I't4 + O 2

- - . - - N O + C~-I,, + 02 ---,.-NO + C~FI,,

40

~ 4o

o -~ 30

~9 30

O Z 2o

~ 2o

et

~NO

50

+

o

i

IQ

C

L

200

300

4OO

500

60O

700

200

300

400

500

600

700

Temperature (~

Temperature (~ Figure I a. N O conversion versus temperature. [NO] = 3600 p p m . [CH 4] = 1.06%, [02.] = 6.0%.

Figure I b. CH~ oonvea~ion versus temperature. [NO] = 3600 ppm, [CH,] = 1.06%, [02] = 6.0%.

Total flow rate = 100.0 cm3/min, GHSV = 60,000.

Total flow rate = 100.0 cm3/min, GHSV = 60,000.

4

-

-

-

i

-

-

.

i

.

-

-

v

-

.

-

i

-

-

-

3.5

3.5

3

3

2.5

2.

0.5 !~5o'c~ ~400~ 0

l

2350

.

.

.

-2

O.5

. .

.

2150

~'~2

.

4.00~ .

.

.

.

1950

.

.

.

.

.

1750

.

.

' . . . i

1550

1350

wavenumber (cm-!) Figure 2. 6600 ppm NO was preadsorbed for 40 min. at room t e m p e r a r ~ and desorbed in He, during temperatu~ ramp.

0 2350

_ 2150

1950

__ 1750

1550

1350

wavenumber (cm-1) Figure 3. 6400 ppm NO was pxeadsorbed for 40 min at room t e m p e ~ u r e , and desorbed in 2.14% CH,/He during temperature ramp.

664 Based on previous studies [15, 22-25], the band at 1941 cm-1 is assigned to Co2*(NO), and the pair of bands at 1894 and 1815 cm-~, to Co2+(NO)2. The shoulders at 1874 and 1799 cm-~ may be due to a second dinitrosyl species. While little is known about the location and coordination of the Co2§ in ZSM-5, it is likely that cobalt ions are associated with both [Si-OA1]- and [A1-O-Si-O-A1]2- structures in the zeolite. In the former case, the cobalt cations are assumed to be present as Co2+(OH-) cations and in the latter case as Co2§ cations. The presence of cobalt cations in different environments could account for the appearance of two sets of dinitrosyl bands. The band at 2132 cm-1 is present not only on Co-ZSM-5 but also on H-ZSM-5 and Na-ZSM-5, and has been observed by several authors on Cu-ZSM-5 [26-28]. A careful investigation of this feature suggests that it is attributable to NO2 ~ associated with both Bronsted acid and M§ cations [29]. The doublet at 1400 cm-~ is identical in position and appearance to that observed in a sample of Na-Y doped with NaNO3 [30]. We therefore assign this peak to a NO3- species affiliated with the residual sodium in the catalyst. The position of the band at 1528 cm-l is very similar to that for nitrito species in Co-A and Co-Y [22, 23] and is, therefore, assigned to Co-ONO. The features at 1599, and 1574 cm-I are best assigned to Co-O2NO [30]. The band at 1633 cm-1 is similar to that observed on H-, Na-, and Cu-ZSM-5. We believe that this feature is best assigned to nitrito (NO2) or nitrate (NO3-) species. Temperatures in excess of 150 ~ are required to desorb NO from Co-ZSM-5 (see Fig. 2). The ratio of the intensity of either of the dinitrosyl peaks to the mononitrosyl peak decreases with increasing temperature suggesting that NO in dinitrosyls is less strongly bound than that in mononitrosyls. The NO2 a§ band decreases in intensity slowly and disappears completely at temperatures above 100 ~ NaNO3 is not very stable, and decomposes completely by the time the temperature reaches 100 ~ Switching from NO to He at room temperature leads to a sudden increase in the intensity of the bands associated with NO2 and NO3. These features increase in intensity up to 150 ~ whereafter they decrease, disappearing completely above 350 ~ Analysis of the effluent gas during TPD reveals that NO desorbs from Co-ZSM-5 without decomposition. A series of spectra taken during the TPD of NO into a stream containing 2.14% CH4 in He are shown in Figure 3. At temperatures up to 300 ~ these spectra are identical to those presented in Figure 2. However, at 300 ~ the nitrito peak at 1515 cm-I [shifted from its position at 1528 cm-1 at room temperature] is notably absent, presumably due to the reaction of Co-ONO with CH4. Spectra obtained during TPR of NO are shown in Figure 4. Nitrosyls are the principal species observed as the temperature is elevated, and in the presence of gas phase NO there is no adsorbed NO2 at temperatures above 150 ~ The absence of adsorbed NO2 is very likely due to its displacement by NO, since it was found that NO readily displaces adsorbed NO2 at 300 ~ During exposure of the catalyst to NO at 300 ~ the intensity of the band at 1528 cm-1 goes through a maximum with time. This suggests that a small amount of NO2 is probably being formed throughout the temperature ramp and is simultaneously being displaced by NO. As the temperature increases the ratio of mononitrosyl to dinitrosyl species increases. Figure 5 shows a series of infrared spectra taken during the TPR of NO with CH4. At temperatures less than 350 ~ the spectra in Figure 5 are virtually identical to those for NO TPR seen in Figure 4. Above 350 ~ the nitrosyl bands are more intense in the presence of CH4. A new peak appears at 2270 cm-I when the temperature is raised to 400 ~ and above, and another one appears at 2173 cm-~ at 450 ~ Neither of these bands was observed when the reaction mixture was passed over Na-ZSM-5. When 15NO was substituted for 14NO, the two bands appeared at 2256 cm-1 and 2144 cm-l, and when 13CH4 was substituted for 12CH4, the bands shifted to 2237 cm-1 and 2132 cm-l. Based on previous studies [32-34], the band at 2270 cm-1 is best attributed to NCO species adsorbed at AI3§ sites. The observed shift in the

665

3.5

3.5

3

1

2.5

2 25oc

1.5

.

lsooc

___...-~/~ ~ _

1 25o*c

0.5

0.5 ~

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2350

2150

1950

1750

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1350

.

,

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,

.

,

I

_

,

2150

,

.

i

.

1950

.

1750

1550

1350

wavenumber (cm" 1)

wavenumber (cm-t)

Figure 5. 5100 ppm NO + 9800 ppm CH 4 was passed over the catalyst for 20 min at room temperature, before beginmng temperatur~ ramp.

Figure 4. 6000 ppm NO was pxeadsorbed for 20 min at room temperature, and maintained over the catalyst during Icmperatmr ramp.

3.5 - , - , . , . , . , . , . , - , - , - , - , . , . , . , . , . , - , . , .

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'400~

0.5

0.8

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. l . l . l . l

. l l l . l l i . l . l l t , l . l . i . l . t . i . l .

2300 2200 2100 2000 1900 1800 1700 1600 1500 1400

wavenumber (cm" l)

Figure 6.

3800 ppm NO + 5.4% 02 was passed over the c.atalyst for 20 min at room temperature, before beginmng temperature ramp.

2300 2200 211111 211131)1900 1800 1711(I 1600 151113141713 wavenumber (cm-t) Figure 7. 3600 ppm NO + 1.1% CH,t + 5.6% 0 2 was passed over the catalyst for 20 min at room temperature, before beginning temperature ramp.

666 frequency of the of this feature upon isotopic substitution is consistent with the assignment to the band to NCO species. The band at 2173 cm-1 is more difficult to assign. Based on its position, this feature might be attributed to NCO species adsorbed on Co2+ [31]. This would also be consistent with the observation of a band at 2185 cm-I for NCO species adsorbed on Cu2+ cations in Cu-ZSM-5 [34]. However, the observed frequency shifts upon isotopic substitution are inconsistent with this assignment, and are in much better agreement with the assignment of the 2173 cm-1 band to CN species on Co2§ While the position of this band is about 30 cm-1 lower than that reported for Co complexes containing CN ligands, the presence of electronegative ligands such as NO and NO2 within the coordination sphere of cyano complexes is known to shift the frequency of the C-N vibration upscale by as much as 30 cm-l [31]. Since the Co2§ cations in Co-ZSM-5 are assumed to occur as Co2+(OH-), the presence of OH- may be responsible for the upscale shift in the vibrational frequency of the adsorbed CN species. A series of spectra taken during TPR of a mixture of NO and 02 are presented in Figure 6. Bands are observed for both mon- and dinitrosyls, together with bands characteristic of NO2 and NO3- species. As the temperature rises, the ratio of nitrosyl to NO2/NO3 bands increases, consistent with what is expected on the basis of equilibrium considerations for the reaction NO + 1/2 02 = NO2 [35]. Figure 7 shows spectra recorded during a TPR experiment in which a mixture of NO, 02, and CH4 are passed over the catalyst. At room temperature several new bands are present. These are located at 2189, 1878, and 1747 cm-~. The peak at 2189 cm-~ is most likely due to NO2r~ [36, 37], since this band is observed upon adsorption of NO2 at room temperature (see Figure 8). The band at 1747 cm-I is assigned to N204 [38], and the feature at 1878 cm-I is probably due to N203 [30, 39]. Elevating the temperature removes all three of these bands. The NO2/NO3 bands are quite intense at room temperature relative to the mono- and dinitrosyl nitrosyl bands. As the temperature rises, the ratio of nitrosyl to NO2/NO3 band intensities increases in a manner similar to that seen in Figure 6. Above 350 ~ the intensities of the NO2 and NO3 bands are smaller than those observed in the absence of CH4, a pattern identical to that already noted in the comparison of Figures 2 and 3. When the temperature is raised to 450 ~ the only features remaining are weak bands located at 2264, 1934, and 1635 cm-1. The first two bands are attributed to A13+-NCO and Co2§ respectively, and the third is due to adsorbed H20. Since the formation of NO2 can occur homogeneously, it was of interest to establish whether adsorbed NO could be oxidized. NO was adsorbed at 225 ~ after which the infrared cell was purged with He and subsequently a stream of 10.1% 02 in He was allowed to flow over the catalyst. Prior to the introduction of the O2-containing stream, the only features evident were those for mono- and dinitrosyls. In the presence of 02 at 225 ~ the intensities of the bands for both mono- and dinitrosyl species attenuated and new features appeared at 1628 and 1518 cm-1, corresponding to nitrate and nitrito species, respectively. A similar experiment carried out in the absence of 02, showed only a small decrease in the intensity of the nitrosyl bands due to NO desorption and the absence of bands for nitrate and nitrito species during a 30 min purge in He at 225 ~ To investigate the behavior of adsorbed NO2 further, TPR experiments with both NO2 and an NO2/CH4 mixture were performed. These results are shown in Figs. 8 and 9. Figure 8 shows that upon exposure of the catalyst to NO2, bands appear at 2189, 2133, 1747, 1672, and 1528 cm-l. Of these features, only the band located at 1672 cm-I represents a feature not seen previously in the spectra of adsorbed NO. While this band can be attributed to some form of adsorbed NO2, the exact structure is not known. Upon raising the temperature to 50 ~ all of the bands below 1700 cm-1 disappear, the band at 1528 cm-l decreases in intensity, and the

667

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2300 2200 2100 2000 1900 1800 1700 1600 1500 1400

wavenumlxa" (cm-t)

wavenumber (cm" t)

Figure 8. 5100 ppm NO2 was passed over the catalyst for 20 rain at room tempetatmc, before bcginmng t c m l x a ~ m ramp.

Figure 9. 4900 ppm NO2 + 1.1% CH 4 was passed over the catalyst for 20 rain at room tcmpcmun=, before bcginning temperature ramp.

I

Co ~

+ NO

~

Co 2+ (NO)

2

Co 2+('NO) + NO

~

Co2§

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5

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6

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$

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9

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02

~

C~

+

2 )

+ NO2

+ OH.

Co2+(CH2NO) + HO2- (H20) Co2+(CN) + H 2 0

+ NO + CO

Figure 10. Proposed reaction mechanism for NO reduction by CH 4 in the presence of O 2 .

668 band at 1672 cm-I increases in intensity. A new feature also appears at 1633 cm-1. At temperatures above 50 ~ the band at 1672 cm-1 disappears, the band at 1633 cm-1 decreases in intensity, and the band at 1528 cm-1 first increases and then decreases in intensity. Comparison of Figures 8 and 9 shows that the spectra in Figure 9, taken in the presence of CH4, are identical to those shown in Figure 8 below 400 ~ To determine whether the isocyanate and/or nitrile species are reaction intermediates, experiments were performed to compare the reactivity of these species in He, NO, 02, and NO2. Isocyanate and nitrile species were first produced at 450 ~ by reaction of NO with CH4. At t = 0 the reactant flow was replaced by either He or a mixture containing 1.0% NO, 0 2, o r NO2 in He. Upon cessation of the flow of NO and CH4, there was an immediate rise in the intensity of the NCO band (2270 cm-1) and a decrease in the intensity of the CN band (2173 cm-1) caused by partial oxidation of CN to NCO species by residual NO or NO2 in the reactor. The intensity of the CN band then decreased with an apparent first-order rate coefficient of 2.1x10-3 s-1 in He, 1.0x10-2 s-1 in NO, 9x10-3 s-l in 02, and > 10-1 s-1 in NO2. The intensity of the NCO band decayed with an apparent fin'st-order rate coefficient of 3.9x 10 .3 s-l, in He, NO, and 02, and lx10-2 s-1 in NO2. Based on the initial intensity of the CN band, it is estimated that about 1% of the C02* cations are occupied CN species at the start of a transient experiment. Using this figure, the turnover frequency for the disappearance of CN species is estimated to be about 10-4 s-1 in NO and > lx10-3 s-1 in NO2. The latter figure is close to the turnover frequency measured during steady-state reaction of NO with CH4 at 450~ (8.4x10-4 s-l), lending support to the suggestion that CN species are reaction intermediates. NCO species adsorbed on A13§ react with NO2 at a rate that is an order of magnitude lower than the rate of reaction of Co2+CN species, and at an even lower rate when the reaction of NCO is with NO or 02. For this reason, it seems unlikely that NCO species play a primary role in the formation of N2 and CO2, the final products of NO reduction. 4. DISCUSSION The infrared observations presented in Figures 3 and 9 clearly demonstrate that CH4 reacts with adsorbed NO2 rather than adsorbed NO. Infrared spectroscopy also reveals that NO2 is readily formed when NO and 02 are present in the feed stream and that NO2 is more strongly adsorbed than NO. The temperature at which adsorbed NO2 begins to react with CH4, 350~ is identical to that at which significant NO conversion is observed during the reaction of NO with CH4 in the presence of 02. The higher lightoff temperature for NO reduction by CH4 in the absence of 02 (see Figure 1) is, thus, attributable to very low concentrations of adsorbed NO2, rather than to a higher barrier for the activation of CH4. Figure 10 illustrates a possible mechanism for the reduction of NO by CH4 in the presence of 02 based upon an amalgamation of elementary steps previously suggested in the literature [11, 13-16, 18, 20] and those deduced from the experiments presented here. The sequence begins with the adsorption of NO adsorption to form both mono- and dinitrosyl species (reactions 1 and 2). The later species undergo oxidation in 02 (reaction 3) to form adsorbed and gas phase NO2. The reaction of CH4 with adsorbed NO2 (reaction 4) is assumed to form an hydroxyl radical and adsorbed CH3NO. Weiner and Bergman [39] have reported the formation of CH 3NO and other nitrosoalkanes by migratory insertion of coordinated NO into Co-C bonds of cobalt alkyl species. Subsequent reaction of adsorbed nitrosomethane with either OH radicals or 02 (reactions 5) followed by the elimination of water (reaction 6) leads to the formation of adsorbed CN species. The nitrile species are hypothesized to react with NO2 to form N2 and CO2 (reaction 7), or with NO to form N2 and CO (reaction 8). The reaction of nitrile species with O2 (reaction 9) could result in the formation of NO and CO. While not

669 indicated, the CO released in reactions 8 and 9 is envisioned to undergo further oxidation to

CO2. The reaction sequence presented in Figure 10 is consistent with the mechanistic arguments given previously by Li et al. [ 15], who proposed that the first step in the reduction of NO by CH4 over Co-ferrierite is the reaction of gas-phase CH4 with adsorbed NO2, but differs in regard to the sequence of reactions leading to N2 and CO2. It is noted that the possibility of CN serving as a precursor to N2 and CO2 was suggested recently by Li et al. [32], based on studies conducted with Cu/ZrO2 and by Hayes et al. [41] based on studies conducted with CuZSM-5. The proposed mechanism is attractive in that it explains not only the manner in which NO2 initiates the reaction of CH4, but also the pathway to CO2 and N2. This mechanism would also explain why NO facilitates the combustion of CH4 by 02 [13, 18]. TPD experiments conducted in our laboratory have shown that Co-ZSM-5 will not adsorb 02, whereas it will adsorb NO2. If the product of the reaction of CH4 with NO2 is retained as an adsorbed species, then it is easy to see how NO2 (derived from the oxidation of NO) could facilitate the oxidation of CH4 by 02. 5. CONCLUSIONS In situ infrared observations show that the primary species present during the reduction of NO by CH4 over Co-ZSM-5 are adsorbed NO2 and CN. When 02 is present in the feed NO2 is formed by the homogeneous and catalyzed oxidation of NO. In the absence of 02, NO2 is presumed to be formed via the reaction 3 NO = NO2 + NzO. The CN species observed are produced via the reaction of methane with adsorbed NO2, and transient response studies suggest that CN species are precursors to N2 and CO2. A mechanism for the SCR of NO is proposed (see Figure 10). This mechanism explains the means by which NO2 is formed from adsorbed NO and the subsequent reaction sequence by which adsorbed NO2 reacts with CH4 and O2 to form CN species. N2 and CO or CO2 are believed to form via the reaction of CN with NO or NO2. CH3NO is presumed to be formed as a product of the reaction of CH4 with adsorbed NO2. The proposed mechanism explains the role of O 2 in facilitating the reduction of NO by CH4 and the role of NO in facilitating the oxidation of CH4 by 02. 6. ACKNOWLEDGMENT This work was supported by a grant from the Gas Research Institute. REFERENCES 1. Y. Li and J. N. Armor, Catalytic Reduction of NOx Using Methane in the Presence of Oxygen, U. S. Patent No. 5 149 512 (1992). 2. Y. Li, and J. N. Armor, Appl. Catal. B, 1 (1992) L31. 3. Y. Nishizaka and M. Misono, Chem. Lett., (1993) 1295. 4. Y. Li, J. Battavio and J. N. Armor, J. Catal., 142 (1993) 561. 5. R. Burch and S. Scire, Appl. Catal. B, 3 (1994) 295. 6. T. Tabata, M. Kokitsu and O. Osamu, Catal. Letc, 25 (1994) 393. 7. J.N. Armor and T. S. Farris, Appl. Catal. B, 4 (1994) L11. 8. R. Gopalakrishnan, P. R. Stafford, J. E. Davidson, W. C. Hecker and C. H. Bartholomew, Appl. Catal. B, 2 (1993) 165. 9. J.L. d'Itri and W. M. H. Sachtler, Appl. Catal. B, 2 (1993) L7. 10. Y. Li and J. N. Armor, Appl. Catal. B, 2 (1993) 239. 11.J.O. Petunchi and W. K. Hall, Appl Catal. B, 2 (1993) L17. 12. R. Burch and S. Scire, Appl. Catal. B, 3 (1994) 295.

670 13. F. Witzcll,G. A. Silland W. K. Hail, J. CataI., 149 (1994) 229. 14. Y. Li and J. N. Armor, J. Catai.,150 (1994) 376. 15. Y. Li, T. L. Slager and J. N. Armor, J. CataI., 150 (1994) 388. 16. E. Kikuchi and K. Yogo, Catal.Today, 22 (1994) 73. 17.A.D. Cowan, R. Dumplemann and N. W. Cant, J. Catal., 151 (1995) 356. 18. D. B. Lukyanov, G. Sill,J. L. d'Itriand W. K. Hall, J. Cat~., 153 (1995) 265. 19. Y. Li and J. N. Armor, Appl. Catal. B, 5 (1995) L257. 20. B. J. Adelman, T. Beutcl,G.-D. Lci, and W. H. M. Sachtler,J. Catal., 158 (1995) 327. 21.J.F. Joly, N Zanier-Szyldowski, S. Colin, F. Raatz, J. Suaussy and J. C. LavaIIey, Catal. Today, 9 (1991) 31. 22. K. A. Windhorst and J. H. Lunsford, J. Am. Chem. Soc. Farad. Trans.,97 (1975) 1407. 23.J.H. Lunsford, P. J. Hutta, M. J. Lin and K. A. Windhorst, Inorg. Chem.,17 (1978) 606. 24.M.C. Kung and H. H. Kung, Catal. Rcv. Sci. Eng., 27 (1985) 425. 25. W. Zhang, H. Yahiro, M. Iwamoto and J. Izumi, J. Chem. Soc. Farad. Trans., 91 (1995) 797. 26.M. Iwamoto, Y. Hidenori, N. Mizuno, W.-X. Zhang, Y. Mine, H. Furukawa and S. Kagawa, J. Phys. Chem., 96 (1992) 9360. 27. A. W. Aylor, S. C. Larsen, J. A. Reimer and A. T. Bell, J. Catai., 157 (1995) 592. 28. J. Valyon and W. K. Hall, J. Phys. Chem., 97 (1993) 1204. 29. T. E. Hoost, K. A. Laframboise and K. Otto, Catal. Lett., 33 (1995) 105. 30. C. -C. Chao and J. H. Lunsford, J. Am. Chem. Soc., 93 (1971) 71. 31. K. Nakamoto, Infrared and Raman Spectra of Inorganic and Coordination Compounds, John Wiley and Sons, New York, 1986. 32. V. A. Bell, J. S. Feeley, M. Deeba and R. J. Farrauto, Catal. Lett., 29 (1994) 15. 33. C. Li, K. Bethke, H. H. Kung and M. C. Kung, J. Chem. Soc. Chem. Commun., (1995) 813. 34. F. Solymosi and T. Bansagi, J. Catal., 156 (1995) 75. 35. Y. Li and W. K. Hall, J. Phys. Chem., 94 (1990) 6145. 36. J. C. Evans, H. W. Rinn, S. J. Kuhn and G. A. Olah, Inorg. Chem., 3 (1964) 857. 37.J.W. Nebgen, A. D. McElroy and H. F. Klodowsky, Inorg. Chem., 4 (1965) 1796. 38. G. M. Begun and W. H. Fletcher, J. Molec. Spec., 4 (1960) 388. 39. I. C. Hisatsune and J. P. Devlin, Spectrochimica Acta, 16 (1960) 40 I. 40. W. P. Weiner and R. G. Bergman, J. Am. Chem. Soc., 105 (1983) 3922. 41. N. W. Hayes, W. Grunert, G. J. Hutchings, R. W. Joyner, and E. S. Shpiro, J. Chem. Soc. Commun. (1994) 531.

J.W. Hightower, W.N. Delgass, E. Iglesia and A.T. Bell (Eds.) 11th International Congress on Catalysis - 40th Anniversary

671

Studies in Surface Science and Catalysis, Vol. 101 9 1996 Elsevier Science B.V. All rights reserved.

Precious metal loaded In/H-ZSM-5 for reduction of nitric oxide with methane in the presence of water vapor M. Ogura and E. Kikuchi* Department of Applied Chemistry, School of Science and Engineering, Waseda University, 3-4-10kubo, Shinjuku-ku, Tokyo, Japan The catalytic activity of In/H-ZSM-5 for the selective reduction of nitric oxide (NO) with methane was improved by the addition of Pt and Ir which catalyzed NO oxidation, even in the presence of water vapor. It was also found that the precious metal, particularly Ir loaded In/H-ZSM-5 gave a low reaction order with respect to NO, and then showed a high catalytic activity for the reduction of NO at low concentrations, if compared with In/H-ZSM-5. The latter effect of the precious metal is attributed to the enhancement of the chemisorption of NO and also to the increase in the amount of NO2 adsorbed on In sites.

1. Introduction Emission of nitrogen oxides (NOx) in combustion exhausts is a serious problem that should be solved in order to protect the earth from acid rain, or other environmental concerns [1]. Recent progress in catalytic removal of NOx is the emergence of an impressive SCR process, which takes an advantage of using unburned hydrocarbons in exhausts as the reductant. Since the discoveries by Iwamoto et al. [2] and Held et al. [3] that the reduction of NO with hydrocarbons could be catalyzed by Cu-ZSM-5 zeolite, interest in this field has been increasing and various catalysts have been proposed for this reaction [4-10]. However, no catalysts having enough activity and selectivity in a practical use have yet been developed, particularly the catalyst having enough durability against the inhibition by water vapor, which is inevitably contained in practical combustion exhausts. In our previous work [11], it has been shown that the reduction of NO with CH, on GaJ and In/H-ZSM-5 catalysts selectively proceeds in the following two stages: NO + 1/2 02 --> NO2 NO2 + CH, + NOx --> N2 + COx + H20

[ on zeolite acid sites ] [ on Ga or In sites ]

(1) (2)

The catalytic activity for NO oxidation [reaction(1 )] was strongly inhibited by water vapor, because this reaction occurs on Lewis acid sites of zeolite as

672

proposed [12]. Although the reduction of NO2 with CH4 [reaction (2)] was also inhibited by water vapor, the catalytic activity of In/H-ZSM-5 for this reaction was more durable against H20 than that of Ga/H-ZSM-5 [11]. Furthermore, the catalytic activity of In/H-ZSM-5 for the reaction of NO-CH4-O2 even in the presence of water vapor was found to be improved by the addition of precious metals such as Pt, Rh, and Ir, because these metals catalyzed NO oxidation instead of the Lewis acid sites of the zeolite [13]. The effects of precious metals on In/H-ZSM-5 was found not only to simply catalyze NO oxidation but also to enhance NOx chemisorption. It is noted that NO conversion on the Ir/In/H-ZSM-5 exceeded NO2 conversion in NO2-CH402 reaction on In/H-ZSM-5, when the concentration of NOx was decreased [14]. This study shows the catalytic activities of In/H-ZSM-5 promoted by precious metals for the removal of low concentration NOx and the promotive effects of the precious metal will be discussed.

2. Experimental Na-ZSM-5(a molar SIO2/AI203 ratio=23.8) provided by Tosoh Corp. was used. In(4wt%)/H-ZSM-5 and Ir(lwt%)/H-ZSM-5 catalysts were prepared by the ion exchange method using NH4-ZSM-5 derived from the Na-ZSM-5 with aqueous solutions of In(NO3)3 at 368 K for 8 h and IrCI(NH3)5CI2 at room temperature for 24 h, respectively. Addition of precious metals, lwt% platinum and iridium to In/H-ZSM-5 was carried out by impregnating the In/NH4-ZSM-5 in aqueous solutions of Pt(NH3)4CI2 and IrCI(NH3)5CI2, respectively. The catalysts were calcined at 813 K for 3 h. Reaction was carried out in a fixed-bed flow reactor mainly by passing a reactant gas mixture of 100-1000 ppm NOx (NO or NO2), 1000 ppm CH4, 10% 02 and 0 or 5% H20 in He at a rate of 100 cm3(STP)omin -1 over 0.1 g of catalysts (GHSV = 36000 h-l). Kinetics studies employed higher GHSV to obtain low levels of NOx conversion below 30%. Water vapor was admitted by passing He through a saturator heated in water bath controlled at 313 K, and the wet He was mixed with other reactants. The reactant stream-line was heated at a higher temperature than the temperature of the saturator to avoid H20 condensation. Reaction products were analyzed by means of on-line gas chromatography with a TCD detector and chemiluminescence NOx analysis. The catalytic activity was evaluated by the conversion of NOx into N2. Chemisorption of NOx was measured by the pressure swing adsorption method reported by Zhang et al. [15]. A gas mixture containing 100 - 1000 ppm NOx was admitted on to 0.1 or 0.2 g catalyst at 673 K, and the concentration of NOx passing through the catalyst bed was detected using a NOx analyzer. Typical response curve and breakthrough point are shown in Fig. 1. The amounts of adsorbed NOx were calculated from these breakthrough curves: the areas shown by a and b correspond to the amounts of totally adsorbed and reversibly adsorbed NOx, respectively.

673

He

I_ F .

(1) u) E 0 Q.

iI

ii

_1 -I

NOx+O2 m

i

ii

ii

al

ii

i

i

m

ii

ii

ii

m

m

am

lille

I

He

I

i

.

e e

e e e e ! I |

n,-

e o | ! o e e | | i | | |

Time Figure 1. Measurement of NOx adsorption" a, corresponds to the total amount of adsorbed NOx; b, the amount of reversibly adsorbed NOx.

3. Results Figure 2 shows the effect of NOx concentration on the conversion of NOx reduced by CH4 in the presence of 5% H20. In the NO-CH4-O2 system, In/HZSM-5 showed low catalytic activity in the whole range of NO concentration. On the other hand, this catalyst was active for the NO2-CH4-O2 reaction, while the conversion of NO2 decreased with decreasing concentration of NO2. The catalytic activity of In/H-ZSM-5 for the reduction of 1000 ppm NO was enhanced by the addition of Ir and Pt almost to the level of NO2 reduction on In/H-ZSM-5, indicating that these precious metals worked as the catalytic sites for NO oxidation, which is a necessary step for NO reduction with CH4. With decreasing NO concentration to 100 ppm, however, the increase in NO conversion was observed on Ir/In/H-ZSM-5 and the conversion of NO exceeded that of NO2 on In/H-ZSM-5. This can not simply be explained by the catalytic activity of Ir for NO oxidation. Kinetic parameters for NOx reduction are summarized in Table 1. It is obvious that the addition of Ir to In/H-ZSM-5 led to the decrease in reaction orders with respect to NO, CH4, and 02 in the NO-CH4-O2 reaction. The decrease in the order for NO can explain that Ir/In/H-ZSM-5 was effective for the reduction of NO at low concentrations. On the contrary, the reaction orders with respect to NO2, OH4, and 02 in the NO2-CH4-O2 reaction were not significantly changed by the addition of Ir. The retarding effect of CH4

674

100 o~

80

OJ

z O *-

60

._o L_

>

40

0 z

20

co o x

f 0

A

A

I

I

I

i

I

200

400

600

800

1000

NOx concentration / ppm Figure 2. Catalytic activities of In/H-ZSM-5 (e), Pt (4) and Ir ( 0 ) l o a d e d In/H-ZSM-5 for NO reduction and that of In/H-ZSM-5 (n) for NO2 reduction in the presence of 5% H20 as a function of NOx concentration. Catalyst weight, 0.1 g. Reaction temperature, 773 K. Table 1. Summary of kinetic data for NOx-CH4-O2 reaction on In/H-ZSM-5 and Ir/In/H-ZSM-5 catalysts. Reaction

NO-CH4-O2

NO2-CH4-O2

Catalyst

/ m~

ro

19

Reaction orders ,2 with respect to 1

NO

NO2

02

OH4

In/H-ZSM-5

1.5 x 10-6

0.80

-

0.80

0

Ir/In/H-ZSM-5

6.1 x 10-6

0.65

-

0.13

-0.18

In/H-ZSM-5

1.0 x 10-5

0.44

0

0.47

Ir/In/H-ZSM-5

1.5 x 105

0.47

0

0.44

I

19 r0, the initial rate reaction under the standard conditions: NOx, 100 ppm; CH4, 1000ppm; and 02, 10%. 29 Concentrations: NOx, 100 - 1000 ppm; CH4, 500 - 2000 ppm; 02, 5 - 12 %. Reaction temperature, 673 K. Catalyst weight: 5.0 - 40 mg for In/H-ZSM-5; 3.0 - 7.0 mg for Ir/In/H-ZSM-5.

675

observed as the negative order of reaction in the reduction of NO on Ir/In/HZSM-5 might be due to the competitive oxidation of NO and CH4 on Ir sites. This effect was not found in the reduction of NO2. Chemisorption of NOx on In/H-ZSM-5 and Ir/In/H-ZSM-5 was measured as a function of NOx concentration. The isotherms of reversible adsorption are shown in Fig. 3. A larger amount of NO2 was adsorbed on In/H-ZSM-5 than NO. Chemisorption of NO was remarkably enhanced by the addition of Ir leading to a larger amount of NO being adsorbed on Ir/In/H-ZSM-5 than NO2 on In/H-ZSM-5. Figure 4 shows the amounts of adsorbed NO from the feed of NO, NO-O2, NO2, or NO2-O2 on In/H-ZSM-5, Ir/H-ZSM-5, and Ir/In/H-ZSM-5. It is apparent from these results that Ir/H-ZSM-5 adsorbed little NOx. NO could hardly be adsorbed on every catalyst in the absence of 02. It is interesting to note that Ir/In/H-ZSM-5 adsorbed larger amounts of NOx from the mixture of NO and 02 than from NO2, and it was also larger than the amount of NOx adsorbed on In/H-ZSM-5 from NO2.

40 ! "7 O)

o

E

30

0 x

~ x 0 z

20

.~

10

L_

0

0

=F,'-

0

I,

I

A

I

I

200 400 600 800 10001200 NOx concentration / ppm

Figure 3. Isotherms of reversible adsorption at 673 K for NO on In/H-ZSM-5 (e) and Ir/In/H-ZSM-5 (0) and for NO2 on In/H-ZSM-5 (m) and Ir/In/H-ZSM-5

(n).

Catalyst weight: 0.2 g for In/H-ZSM-5; 0.1 g for Ir/In/H-ZSM-5.

676

Catalysts" Ir/H-ZSM-5

NO N O + 02 NO2 NO2 + 02 NO

In/H-ZSM-5

Ir/In/H-ZSM-5

~['-]NO + 02

"I////A rl///i//~

IN 02 IN 02 + 02

7] NO Z/////I////IA ~//I IN02 ~////////,/A

0

5

NO + 02

INO~2+ 02 10

15

Amount of adsorbed NOx / 10"6 mol.(g-cat)1 Figure 4. Comparison of chemisorption of NOx at 673 K from various kinds of NOx mixtures admitted on to In/H-ZSM-5, Ir/H-ZSM-5, and Ir/In/H-ZSM-5 catalysts: E] parts correspond to the reversible adsorption (b in Fig. 1). NOx, 100 ppm; 02, 0 or 10%. Catalyst weight: Ir/In/H-ZSM-5, 0.1 g; In/H-ZSM-5, 0.2 g; Ir/H-ZSM-5, 0.1 g.

4. Discussion As shown in Table 1, the reaction order with respect to NO2 on In/H-ZSM-5 was smaller than that of NO. This is in accordance with the proposed reaction sequence that NO is firstly oxidized to NO2 and the NO2 reacts with CH4. Coincidence in the order of reaction for NO2 between In/H-ZSM-5 and Ir/In/HZSM-5 catalysts means that NO2 react on a common active site which should be In species. Chemisorption data shown in Fig. 4 show that either NO or NO2 was hardly adsorbed on Ir site. Furthermore, chemisorption of pure NO was negligibly small on In/H-ZSM-5 and Ir/In/H-ZSM-5 at 673 K, while NO in the presence of 02, as well as NO2, could significantly be adsorbed. These results also support our supposition that chemisorption of NO2 is important and that of NO is less important on these catalysts. In the presence of 02, chemisorption of NOx, both NO and NO2, was enhanced by the addition of Ir on to In/H-ZSM-5.

677

It has been reported that GaJ and In/H-ZSM-5 had low activities for the dissociative adsorption of oxygen, and this result can explain why these catalysts show high selectivities for NO reduction with hydrocarbons [16]. Reaction order with respect to 02 was lowered by the addition of Ir to In/HZSM-5 in the NO system. It is noteworthy that the added precious metals promote the adsorption of oxygen, which is an indispensable component for HC-SCR. 02 is activated on Ir to adsorb dissociatively. On the other hand, the similar reaction orders were obtained on In/H-ZSM-5 and Ir/In/H-ZSM-5 in the NO2 system. NO2 adsorption on In is not inhibited by the adsorbed oxygen on Ir. It is previously reported that the CH4 selectivity to NO reduction was lowered by added precious metals, Ir and Rh, which were active for CH4 oxidation with oxygen [13]. The selectivity of Ir/In/H-ZSM-5 decreased when the concentration of CH4 increased, while it increased with increasing NO concentration. Therefore, 02 can easily be adsorbed on Ir and the adsorbed oxygen can activate not only NO but also CH4. These results suggest that NO

"7,

5

4

x

3

o;2 -s o

1

0

0

2

4

6

8

10

12

1/(NOx concentration) x 10 3 / p p m 1 Figure 5. 1/P - 1/V plot for NO adsorption on In/H-ZSM-5 (o) and Ir/In/H-ZSM-5 (0) and for NO2 adsorption on In/H-ZSM-5 (ll) and Ir/In/H-ZSM-5 (El). Reaction conditions and symbols are the same with those in Fig. 3.

678

and CH4 competitively react with adsorbed oxygen on Ir, resulting in the inhibition of the NO oxidation and in the negative order with respect to CH4 in the NO system on Ir/In/H-ZSM-5. It is noted that the amount of NO adsorption on Ir/In/H-ZSM-5 in the presence of 02 was remarkably larger than that of NO2 on In/H-ZSM-5. To understand the chemisorptive property of Ir/In/H-ZSM-5 in a comparison with that of In/H-ZSM-5, the chemisorption data were analyzed in detail. The data shown in Fig. 3 are relatively well fitted to the Langmuir isotherm, as shown in Fig. 5. From these relations, the equilibrium constant (K) and the amount of NO2 adsorbed at saturation (Vo) were determined according to the following equations: V = KPVo 1 +KP

(3)

1 = 1+1 x 1 V Vo KVo P

(4)

Here, 0 and V represent the surface coverage and the amount of adsorbed NOx, respectively. Calculated K and V0 are summarized in Table 2. If the important adsorption site on these catalysts consists of In species, and if NO2

Table 2. Equilibrium constant of adsorption (K) and adsorbed amount at saturation (Vo) for NOx at 673 K on In/H-ZSM-5 and Ir/In/H-ZSM-5. []

I

NOx in the reactant feed

V0 / 10"6 mol (g-cat)-1

K / 103

NO

9.3

1.4

NO2

21

2.0

i

In/H-ZSM-5

i

Ir/In/H-ZSM-5

NO

57

1.6

NO2

63

1.4

Reactant feed: NOx, 100 - 1000 ppm; 02, 10%. Catalyst weight: 0.2 g for In/H-ZSM-5; 0.1 g for Ir/In/H-ZSM-5.

679

is a sole common adsorbed species, then the adsorption equilibrium constant would be unique. The observed values lie in the range of (1.6 4- 0.4)x10 -3. Taking into consideration the scatter in the observed data probably due to interconversion between NO and NO2 during chemisorption measurements, these values could be regarded almost within experimental errors. The value of Vo sounds more meaningful. The adsorbed amount of NOx at saturation should correspond to the number of adsorption sites. These are in the range of 1 x 10 -5 to 2 x 10 -5 mol-g-catalyst -1 for In/H-ZSM-5. They are extremely small compared with the amount of In in the catalyst (4.8 x 10 -2 mol-g -1). By the addition of Ir, Vo became 3 to 6 times greater. This seems to explain why Ir/In/H-ZSM-5 gave a greater rate of NOx reduction than In/HZSM-5, as shown by ro in Table 1. Although the exact reason of the larger Vo for Ir/In/H-ZSM-5 is not certain at present, a possible explanation is that NO can diffuse into zeolite pores more easily than NO2 and diffused NO can be oxidized on Ir in the pore to NO2 which is then reduced on In sites also existing in the pore. In our previous work [11], it has been shown that the reduction of NO with CH4 on Ga and In/H-ZSM-5 catalysts proceeds through the reactions (1) and (2), and that CH4 was hardly activated by NO in the absence of oxygen on these catalysts. Therefore, NO2 plays an important role and the formation of NO2 is a necessary step for the reduction of NO with CH4. In the works of Li and Armor [17] and Cowan et al. [18], the rate-determining step in NO reduction with CH4 on Co-ferrierite and Co-ZSM-5 catalysts is involved in the dissociative adsorption of CH4, and the adsorbed NO2 facilitates the step to break the carbon-hydrogen bond in CH4. It is suggested that NO reduction by use of CH4 needs the formation of the adsorbed NO2, which can activate CH4. From the result that the conversion of NO at a low concentration (100 ppm) on Ir/ln/H-ZSM-5 exceeded that of NO2 on In/H-ZSM-5, it can be concluded that there are some correlations between the increase in the quantity of adsorbed NO2 and the enhancement of the catalytic activity for the reduction of NO with CH4.

4. Conclusion The catalytic activity for the selective reduction of NO with CH4 was significantly enhanced by the addition of precious metals, particularly Ir, to In/H-ZSM-5. The role of added Ir was not only to promote NO oxidation which is a necessary step for NO reduction, but also to enhance the abilities of the catalyst to adsorb NO and to increase the amount of NO2 adsorbed on in sites.

680

5. Acknowledgment This work was supported by the Grant-in-Aid for Scientific Research on Priority Areas from the Ministry of Education, Science, and Culture of Japan.

References [1] H. Bosch and F. Janssen, CataL Today, 2, 369 (1988). [2] M. Iwamoto, H. Yahiro, Y. Yuu, S. Shundo, and N. Mizuno, Shokubai(Catalyst), 32, 430 (1990). [3] W. Held, A. KSnig, T. Richter, and L. Puppe, SAE Paper, 1990, 900496. [4] H. Hamada, Y. Kintaichi, M. Sasaki, and T. Itoh, App/. Cata/., 64, L1 (1990). [5] M. Misono and K. Kondo, Chem. Lett., 1991, 1001. [6] K. Yogo, M. Ihara, I. Terasaki, and E. Kikuchi, Catal. Lett., 17, 303 (1993). [7] Y. Li, P. J. Battavio, and J. N. Armor, J. Catal., 142, 561 (1993). [8] J.O. Petunchi, G. Sill, and W. K. Hall, App/. Catal. B2, 303 (1993). [9] Y. Nishizaka and M. Misono, Chem. Lett., 1993, 1295 [10] K. Yogo, M. Umeno, H. Watanabe, and E. Kikuchi, Cata/. Lett., 19, 131 (1993). [11] E. Kikuchi and K. Yogo, CataL Today, 22, 73 (1994). [12] J. G. M. Brandin, L. A. H. Anderson, and C. U. I. Odenbrand, CataL Today, 4, 187 (1989). [13] E. Kikuchi, M. Ogura, N. Aratani, Y. Sugiura, S. Hiromoto, and K. Yogo, Catal. Today, in press. [14] M. Ogura, S. Hiromoto, and E. Kikuchi, Chem. Lett., 1995, 1135. [15] W.X. Zhang, H. Yahiro, N. Mizuno, J. Izumi, and M. Iwamoto, Langmuir, 9, 2337 (1993). [16] T. Tabata, M. Kokitsu, and O. Okada, Catal. Lett., 25, 393 (1994). [17] Y. Li and J. N. Armor, J. Catal., 150, 376 (1994). [18] A.D. Cowan, R. D0mpelmann, and N. W. Cant, J. CataL, 151,356 (1995).

J.W. Hightower, W.N. Delgass, E. Iglesia and A.T. Bell (Eds.) 11th International Congress on Catalysis - 40th Atmiversary

Studies in Surface Science and Catalysis, Vol. 101 9 1996 Elsevier Science B.V. All rights reserved.

681

Interfacial RhOx/CeO 2 sites as locations for low temperature N20 dissociation

J. Cunningham%J.N. Hickey%R. Cataluna~'%J-C Conesab,J. Sofia b and A.Martinez-Arias a Physical Chemistry Laboratories, University College Cork Ireland b CSIC Institute for Catalysis, Universidad Autonoma, Cantoblanco, Madrid, Spain.

SUMMARY Temperatures required for extensive N20 dissociation to N 2, or to N 2 plus 02, over 0.5% RhOx/CeO 2 materials, and over polycrystalline Rh203 or CeO2, are compared for preoxidised and for prereduced samples on the basis of conversions achieved in pulsed-reactant, continuous-flow and recirculatory microcatalytic reactors. Influences of sample prereduction or preoxidation upon those measurements and upon results from parallel ESR and FTIR studies of N20 interactions with such materials are presented and compared. Over partially reduced 0.5% RhOx/CeO 2 materials complete dissociation of N20 pulses to N 2 plus 02 is obtained at temperatures 50-100 ~ lower than those required for extensive dissociation over prereduced Rh203. Furthermore, N: was the sole product from the latter. Higher ongoing N20 conversions to N 2 plus 02 at 623 K over 0.5% Rh/CeO2 in pulsedreactant than in continuous-flow mode point to regeneration of active sites under helium flushing between pulses. The TPD profile for dioxygen release from Rhodia containing samples at temperatures 350-550 K is presented. ESR measurements reveal complementary effects of outgassings at temperatures, Tv, > 573 K upon the availability at RhOx/CeO 2 surfaces of electron-excess sites reactive towards N20. Differences from observations over Rh203 and CeO2 can be understood by attributing the low-temperature activity of RhOx/CeO 2 to electron excess sites at microinterfaces between the dispersed Rhodia component and the Ceria support.

1. INTRODUCTION Since nitrous oxide, N20, is a designated "greenhouse" gas, and may contribute to depletion of the ozone layer, its removal from emissions to atmosphere is desirable [1]. However, there are several reports that N20 can be formed at low selectivity as an undesirable by-product of NO+CO conversions during the initial warm-up-from-cold periods in three-way-catalytic (TWC) converters or components thereof [1-3]. TWC's commonly contain Rhodium and Ceria and although N20 dissociation over Rh203 has been extensively studied [4], the following are among mechanistic possibilities as yet

682 incompletely resolved concerning N20 formation and its eventual catalytic removal over binary Rhodimn+Ceria materials: (a) relative importances of the reduced and oxidised forms of Rhodium and Ceria components and of microinterfaces between them in providing active sites [5-7]; (b) relationship of rates for N20 adsorption and dissociation on clean Rh/CeO2 surfaces to corresponding rates on surfaces having some active sites blocked by adsorbed O2 [4]; and (c) direct or indirect influences of anion-vacancies/lattice-oxygens from the ceria support in promoting/inhibiting catalytic dissociation [8]. The present work seeks insights into those possibilities, mainly through comparative studies of N20 interactions with 0.5% Rh/CeO 2, CeO2 and Rh203 samples in both preoxidised and prereduced forms. Comparison between their "initial" and "steady-state" activities have been obtained with a microcatalytic reactor system designed to switch N20 reactant from pulsed to continuous flow. Complementary information concerning the nature of species formed on surfaces of the materials upon contacting N20 and/or 02 with samples activated by prior vacuum-outgassing at increasing temperatures (Tv) has been obtained from measurements by electron-spin-resonance, (ESR) and FTIR, and compared with previous observations of electron localization by 02 at such surfaces. [9,10].

2. EXPERIMENTAL Materials and Characterizations: Rh203 powders commercially available from Aldrich, Rh203(A), and Johnson Matthey, Rh203 (JM) with BET=I6 m2g-~, were used after calcinations/reductions indicated in the text. Wet impregnation of Rhodium acetylacetonate onto CeO 2 powder (BET---IIO m2g~ from Rhone Poulenc) from solution in tetrahydrofuran or methanol yielded 0.5% RhOx/~CeO2 and 4% RhOx/wCeO2 after calcination in 02 at 823 K. [8]. Identical procedures were followed to achieve dispersions of oxidised Rhodium species upon 27% CeO2-AI203. Analagous procedures using Rhodium nitrate as precursor [9] were employed in preparing samples used in ESR and FTIR studies. Coprecipitation by NH4OH from solutions containing nitrate(s) of Rhodium and/or Cerium was the first step of an alternative preparation yielding powders having rhodium ions initially dispersed throughout Ceria e.g. 0.5% RhOxPCeO2. TPR profiles of aliquots precalcined at 823 K were obtained under 3% H2/Argon, with particular attention to position and magnitude of any rhodia-related features at 373 ~ 473 K. Use of 0.5% RhOx/sCeO2 material prepared by wet impregnation of Rhodium acetylacetonate onto CeO2 (rp) which had been sintered at 1273 K overnight, was found necessary in studies of oxygen temperature programmed desorption from 0.5% RhOx/CeO2, otherwise 02 desorption from surface oxygens of h.s.a. CeO2 in the range 473-773 K obscured a rhodia-related O2-desorption feature in that range. X-band ESR spectra were taken at 77 K on a Bruker system (ER 200D), g-values being calibrated with DPPH (g = 2.0036). Ca. 20 mg of sample were placed in a quartz cell with double greaseless stopcocks, and were subjected to outgassing or gas adsorption; prior to taking the spectra, any excess 02 was pumped out at 77 K where necessary to avoid magnetic dipolar broadening of ESR lines. FT-IR spectra were taken with a Nicolet 5ZDX spectrometer, accumulating 200 scans at 4 cm "~ resolution. Self-supported wafers of ca. 25 mg/cm 2 were placed in a high vacuuna cell, provided with NaCI windows, for outgasing and adsorption treatments. All IR data were obtained with the cell at room temperature, and the sample at the temperature induced by the IR beam (estimated to be .~ 320 K). Outgassing and gas treatments of samples for spectroscopic experiments were made in a

683 conventional high vacuum line capable of keeping a dynamic vacuum of I O -3 Nm 2. Research grade 02 and N20 (SEO) were purified by freeze-thaw cycles before use.

Microcatalytic Reactor Systems: Most catalytic results were obtained with a flow reactor system which operated at I atm total pressure, comprised mainly of helium or argon carrier gas with only low content (usually 3%) of N20 therein. With the aid of a 10-port valve the reactant flow could rapidly be switched between: Configuration A in which individual N20 pulses passed over the catalyst and thence straight to a pair of GC columns for on-line analysis, upon completion of which a second pulse could be delivered etc. During each 1015 min. interval between pulses the sample remained at reaction temperature under a flow of pure helium; Configuration B in which a continuous flow of 3% N20/He became established over the same sample, but with gases exitting continuously therefrom being sampled only periodically to the GC columns for analysis. A separate, evacuable, all-glass recirculating reactor system operating at low pressures 1-5 mbarr and equipped for mass spectrometric analysis was used to investigate possibilities of oxygen isotope exchange between lattice- ~60 and N2~80, in parallel with temperature programmed catalytic dissociation (TPCD) of the nitrous oxide.

3. RESULTS AND INTERPRETATION

3.1 Microcatalytic Reactor Studies Rh203: Parts (A) and (B) of figure 1 illustrate results obtained from delivery of the first I0 pulses (A), and then a very large number of N20 pulses (B), at 673 K to a sample of Rh203 powder prereduced in H2 at 473 K. The latter treatment was chosen on the basis of a TPR profile indicating reduction of Rh203 to be complete at 473 K and with a view to facilitating later comparison of N20 pulse interactions over prereduced Rh203 with those over prereduced 0.5% Rh/PCeO2 and PCeO:. In Fig. 1A, titration-like trends are evident in the level of N20 conversion achieved upon delivery of the first ten N20 pulses over prereduced Rh203 at 673 K: conversion to N 2 being complete in the first pulse, but then declining rapidly towards zero for pulses "4 --> ~10, and conversely for survival of nondissociated N20. This behaviour could be understood in terms of the development of an inert Rhodium plus Oxygen overlayer on the metallic Rhodium, thereby passivating it. Such inert overlayers have been reported in other systems [I1], sometimes as precursors of identifiable Rh203 overlayers [12]. Data in Fig. 1B for the N:O conversions achieved in later N20 pulses of the long pulse sequence at 673 K over prereduced Rh203 demonstrate how, following decline of N20dissociation activity almost to zero for pulses ~8 to ~12, such activity increased progressively until reaching 80% after 85 pulses. Peak area comparisons indicated that the product was predominantly N 2 and certainly not N2+l&O2, thereby pointing to an autocatalytic oxidation of Rhodium by N20 rather than catalytic N20 dissociation. Bearing in mind that the sample was exposed to flows of Helium at 673 K during each 10 min interval between pulses, and during two overnight interruptions in delivery of the long train of pulses, ample opportunity existed for progression of structural reorganization of any initial inert oxygen overlayer a Rhodia surface layer [12] at 673 K and for thickening of such rhodia overlayer in an autocatalytic reaction with N20. Results similar to Fig. 1B did not develop in N,O pulse sequences at 423 or 523 K over prereduced Rh203. Furthermore, periodic analyses of exit

684

Figure 1. Relative GC peak areas for N2-only product and any residual N20 reactant from passage of N20 pulses at 673 K over prereduced Rh203: (1A) declining conversion in pulses 1 -~ 10; (1 B) reversal of initial decline for pulses 12 - 85.

gases from a continuous flow of N20 over Rh203 prereduced at 473 K showed no measurable conversion to N2 or O 2 during 2 h on-stream at 623 or 673 K. The following results were obtained from passage of N20 pulse sequences over Rh203 aliquots which had not been prereduced: (i) full ongoing conversion to N2+1/~O2 was achieved at 623 K over Rh203 previously flushed for 1 hr with helium at that temperature, whereas ca. 30% ongoing conversion resulted at 573 K; (ii) ongoing conversion at 28% level to N2+1/~O2 was observed from N20 pulses at 573 K over an aliquot which had been preoxidised for 1 hr at 773 K under a flow of IO% OJHelium. 0.5% RhOx/CeO2: At reaction temperatures ca. 50 ~ lower than those indicated above as necessary for Rh203, the following features were observed over aliquots of coprecipitated 0.5% RhOx/CeO 2 materials pretreated and tested in conditions closely similar to those used for Rh203 aliquots: (i) analyses of exit gases from N20

pulses delivered at 623 K over prereduced aliquots showed complete dissociation to N2 as sole product from the first ten pulses, whereupon a rapid switch-over at pulses ~10-12 was observed in composition of gases emitting from complete N20-pulse dissociation to yield N2+~AO2 from pulses ~13-~30 (cf. Fig. 2A). This switch-over could be understood in terms of a predominance of surface-site reoxidation process, N20+S%ed ~ N2+O"/Sox, durmg the first ten pulses [13] after which came an onset of N20-dissociation to N2+~A02 proceeding via redox cycling at such sites via N20+0n/Sox ~ N2+02+S"~d etc; (ii) Figure 2B illustrates that analyses of exit gases after passage of N20 pulses at 623 K over preoxidised aliquots provided evidence for presence of N 2 plus O2 product already after the second pulse and continuing to be present in subsequent pulses thereafter with undissociated N20 at a level indicating only 50% ongoing conversion. (iii) Analyses of exit gases from N20pulses delivered at T~x=523 or 423 K over prereduced aliquots yielded only an initial dissociation to N2 as sole product for the first 7 or 4 pulses respectively i.e. similar to Fig. 1A. Analyses of exit gases after higher pulse numbers at these relatively low temperatures over the prereduced aliquots showed only N20 with zero N2 or O 2 product, thereby

685 establishing an absence of ongoing dissociation which contrasted with results from N20pulses over the same material at 623 K. Tests made in the pulsedN20 mode at 423, 523 and 623 K upon a "reference" CeO2 material prepared by NH4OH precipitation from cerium nitrate and pretreated in similar fashion to the 0.5% RhOx/CeO 2 material did not discover any activity for ongoing N20 dissociation, but only a small initial dissociation to N 2 similar to Fig. IA. Following determination of the % conversions of N20 pulses to N 2 plus 02 at 623 K just noted over prereduced or preoxidised aliquots of 0.5% RhO]CeO 2, the reaction system was instantly reconfigured to flow N20 reactant continuously over the samples at that temperature. In each case the result was an approximate halving of the N20 conversion relative to that achieved in pulsedN20 configuration, despite similar flow rates in both configurations and continuation of flow for up to 2 h. It seemed probable from this halving, and from an abovementioned enhancing effect of prior helium flush at 623 K upon activity of Rh203 for N20 dissociation to N2 plus 02, that the periodic 10 min flushed with pure helium at 623 K experienced by 0.5% RhOx/CeO ~ between Figure 2. Relative GC peak areas of exit gases from N20 pulses over 0.5% RhOx/CeO2 at 623 K: (2A) Complete conversion to N2-only from pulses 1 ~ 10 over prereduced coprecipitated material followed by a switch to N 2 plus 02 products for pulses 12 ~ 27; (2B) partial conversion at 623 K over the same material when preoxidised; (2C) complete conversion to N 2 plus 02 products for pulses 1 ~ 21 over material prepared by wet impregnation.

686 N20 pulses, contributed to formation/regeneration of Rhodium-containing sites having activity for dissociation of each incoming N20 pulse to N 2 plus 02 at 623 K. Possibilities for such formation/regeneration of active sites include the diffusion/desorption of blocking species away from Rhodium-containing sites, which are further considered below in respect of (N202) "n and 02. Part c of Fig. 2 demonstrates that ongoing complete dissociation of N20 pulses at 623 K to N2+1/202 occurred over 0.5% RhOx/CeO 2 prepared by wet impregnation rather than the coprecipitation method used for the material used in part a of the figure. The extensive dissociation over material prepared by wet impregnation is consistent with greater localization of the Rhodium component in surface and near-surface regions. Further support for this emerged from differences in extent of dissociation at pulse numbers greater than ~15 in N20 pulse sequences at 523 K: no ongoing dissociation then being detectable over the coprecipitated material, whereas dissociation to N2 plus 02 continued at pulse numbers "16 -~ '27 over the material originated by wet impregnation. Evidence was obtained for significant Oxygen isotope exchange between N2~80 and ~60" species on the surface of a preoxidised sample of 4% Rh/wCeO2-Al203. Data were obtained by introducing N2~80 at 3 mbarr pressure into a recirculatory reactor system containing the preoxidised catalyst and applying a temperature ramp at 10 ~ min ~ to the reactor segment whilst analysing by mass spectra at 30 sec intervals the composition of gases recirculating in the system. Results showed onsets of a decrease in PCN2~80) @ m/e=46, together with an increase in P(N2) @ m/e=28 at ca. 520 K, confirming dissociation of N2~80 at this and higher temperatures. An interesting effect, not revealed by other studies, was indicated by parallel onset at 520 K of another process producing a gaseous species, having rn/e=44 and identified as N2~60. (The possibility that this might be due to CO2 could be rejected on the basis that no equivalent rise in signal at m/e=44 was observed to onset at this temperature from a CeO2-AI203 sample subjected to identical in-situ preoxidation at 823 K prior to testing isotopically with N2J80 in conditions similar to those for the 4% Rh/CeO2-A1203 material). Comparison of the rates of rise @ m/e=28 and m/e=44 with temperature in the range 520-773 K showed that the process yielding N2160 proceeded with ca. 50% of the efficiency of that yielding N2~g~ product formation, possibly via (N2~sO~60)" intermediates at Rhodia-containing sites. The observed increase in P(N2160) up tO ca. 725 K implies presence of some such surface precursor up to those temperatures and makes it a candidate for blockage of Rhodia-containing active sites. 3.2. Oxygen Desorption Profiles for temperature programmed desorption of O2(TPD) were compared for 0.5% RhOx/CeO2, 4% RhO~/sCeO2, Rh203 and sCeO 2 materials pretreated in-situ to a flow of 02 to 823 K, cooled down in O2 to 300 K and flushed overnight in helium at 300 K to remove physically adsorbed oxygen. Results are summarised in Fig. 3, which shows that all the Rhodia-containing materials yielded a detectable TPD feature in the range 350 -~ 550 K. Facile detection of that Oxygen TPD feature from Rh203 (plot c) contrasted with greater difficulty for its detection from 4% RhO• (plot d) because of a large background 02 desorption from that high surface area Ceria support. Attainment of the indicated clear differentiation between zero desorption from (plot a) and a small,but observable feature from 0.5~ RhO• (plot b) only became possible by dispersing 0.5% Rhodium upon CeO 2 presintered overnight at 1250 K and by using the sintered CeO2 as reference (plot a).

687

3

t

J

Aldrich RhOx

Release of ~602 from preoxidised 4% RhOx/CeO 2 at temperatures > 473 K into a low pressure of Oxygen-18 enriched 1802 was reported previously on the basis of measurements made with the recirculatory reactor system [8]. An analagous experiment in the present study, featuring a low pressure of N2160 in contact with preoxidised 4% RhOx/CeO2 while temperature was raised 300-500 K, likewise showed a small release of 02 (together with N2). However, 02 release was not observed when that sample was mildly prereduced nor when CeO2 was similarly tested. Preoxidised rhodia dispersed upon Ceria is thus seen to enhance the ease of release of 02 at temperatures < 550 K.

1

)

4.50E-09

OOOO- 0.5% Rh on i CeO2 calcined [ at 1273K

O-O

4.00E-09

4 0 RhCeO2

Ia

!

1 I

CeO2 calcined i at 1273K. I

]

3.50E-09

J ~3.00E-09 t..

~2.50E-09 ,~

~2.ooF~-o9

3.3. Spectroscopic Studies of NzO --Surface Interactions: Previous work has

1.50E-09

1.00E-09

5.00E-lO

0.00E+00 [ 300

i 500

: 700

\ Temperature/-K. Figure 3. Oxygen TPD profiles from the indicated preoxidised materials.

900

demonstrated the value of adsorbing 02 at 77 or 300 K upon Ce02 surfaces activated by prior vacuum outgassing at increasing temperatures (Tv) to serve as an electron-accepting surface probe yielding information, via ESR measurements on the number and type of O~ radicals detected at 77 K concerning relative densities of reduced surface centres, such as electron-vacancy complexes (Ce3"-Vo), [9,10]. However, the application of similar procedures, using N_,O rather than 02 as adsorbed probe, to samples of CeO2 outgassed up to Tv=773

K did not produce any new paramagnetic species, despite FTIR observations confirming appearance of IR features attributable to adsorbed N20 (2234 and 1256 cm" in Fig. 4a) upon contact with N20 at 300 K. Stepwise decreases in magnitude of those IR features were, however, observed in each of a sequence of FTIR spectra taken after separate N20 adsorptions at increasing adsorption temperatures (TO up to 573 K (Fig. 4b-d). From these FTIR observations it could be inferred that increased T a for contact between N20 and vacuum-outgassed CeO 2 resulted in increased fractional decomposition of the N20 introduced. FTIR spectra did not show bands due to peroxide species after N20 adsorption.

688

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-

~

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

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.

.

.

.

.

.

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-

.

_

-

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iiii

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I _

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1500

Wavenumbers (cm- 1)

Figure 4: FTIR spectra after N20 adsorption on CeO 2 preoutgassed at T~ = 773 K; T~ = RT (a) and Ta = 473 K (b). Adsorption of N20 at RT on RhO• 2 preoutgassed at T,, = 473 K (c) and T v = 773 K (d). Parallel ESR measurements did not reveal the accompanying formation of O or 0 2 which was to be expected if Ce3*-Vo or other one-electron reducing sites initiated sequences known over other oxides [13] viz N20 --~ N 2 0 --> N 2 + O N__2,O N2 + 02") This leaves N20 --> N 2 + Os2" as the likely alternative pathway for the FTIR-observed increase in decomposition of adsorbed N20 at the higher Ta and also helps to account for dissociation of the first few pulses in profiles similar to Fig. 1A. Information on the fraction of reduced surface sites not oxidised by contacts with N20 was sought using a procedure involving the following steps: (i) vacuum outgassings at Tv = 373-773 K to produce (Ce 3*Vo) centres, such outgassing pretreatment being performed before each N20 adsorption in order to regenerate a desired initial surface condition; (it) N2O adsorption at a selected Ta, followed by desorption at 300 K (verified by FTIR) in order to avoid coverage of adsorption sites by residual N20; (iii) oxygen adsorption at 300 K. Application of this strategy to CeO 2 activated by Tv < 473 K showed that availability of (Ce3*-Vo) or other oneelectron reducing sites for reaction in step iii to yield O 2 was no___]significantly affected by N20 adsorptions even at T~ = 473 K. For CeO 2 more severely outgassed at T~ > 573 K, however, intensity of the 0 2 ESR signal achieved by step iii did progressively decrease with increasing T a for N20 until only a small residual 0 2 signal was detected after T a = 473 K (Fig. 5A,c). Much greater capability to react with N20 at T~ = 473 K and destroy the sites having potential to produce 02 - from 02 was thus established for sites generated on the CeO2 surfaces by T~ > 573, than was the case for reducing sites generated by T, _< 473 K. No new ESR signals were observed upon adsorption of N20 onto RhOx/CeO 2 outgassed

689 at Tv --- 473 K at 300 K, although a decrease was observed in a broad signal with < g > =2.16, previously assigned to Rh 2§ cations in Rhodium oxide clusters [9,10]. Unlike the CeO2 results, N20 adsorption at 300 K onto RhOx/CeO2 after T~ _> 573 K did produce new ESR signals termed 01 (g2 = 2.035, gx 2.016 and g y = 2.011) and RO (g~ = 2.19, g2 = 2.00, g3 = 1.986), accompanied by decreases in broad signals assigned to Rh 2§ cations in rhodium oxide clusters - Figure 5B. ESR signals very similar to these were previously observed to form upon contacting 02 with RhOx/CeO2 > 573 K, and were assigned respectively to Ce 4§ 02-centers stabilised at isolated vacancies and to weak [Rh-OO] 2§ adducts. The contributions by signals 01 and RO changed not only with T~ but also with N20 contact time. For example, during 3 h contact of N20 at 300 K with RhOx/CeO2 after Tv = 573 K intensity of signal RO experienced a large increase although signal 01 remained almost constant (cf. Fig. 5B,plots b and c). This indicated that formation of 02 and related surface complexes is measurable (with ESR), but slow at 300 K. The slower build-up of signal RO may be due to weaker character of its adsorption bond, as evidenced by its disappearance upon outgassing at 300 K while signal 01 remains with substantial amplitude (Figure 5B,d). Neither of these ESR signals was detected in the present study whenever Rh203 powder, after vacuum outgassing at T~ > 573 K, contacted N20 at room temperature. Upon such interaction with N20 the spectra did, however, show a decrease of the broad Rh 2§ signals present in such outgassed rhodia specimens, indicating that electron transfer to the adsorbed molecule did occur, probably with formation of O2-type diamagnetic species as explained above for CeO2.

z.d,, m

_

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1.5

Z,ltq C

I", o

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I

e,l

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3275

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|

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5A

3375

MAGNETIC FIELD (G) Figure 5A. ESR spectra after 02 adsorption at RT on CeO2 preoutgassed at Tv = 573 K. Without N20 preadsorbed (a). With N20 preadsorbed at T N = 373 K (b) and T~ = 473 K.

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2.oH 27CX)

32(X~

MAGNETIC FIELD (G) Figure 5B. ESR spectra of RhOx/CeO2. After outgassing at Tv = 573 K (a). Subsequent N20 adsorption at RT during 1/2 h (b) and 3 h (c). The latter outgassed at RT (d).

690 4. CONCLUSIONS The failures of CeO2 alone or Rh203 alone to yield RO or 01-type paramagnetic centers, when subjected to the same experimental strategy as that which did yield such centers over RhOz/Ce02 after Tv >- 573 K, pointed to synergism between the RhOx and CeO2 components of the latter in producing at 300 K the dioxygen species required for 01 and for RO-type signals, possibly through the intermediacy of Rh-O-n-Ce and Rh-O-O-Ce interfacial sites.

Evidence emerges from both the microcatalytic and spectroscopic studies for enhanced dissociative interaction of N20 with preoxidised RhO]CeO2 materials exposed to heliumflush or vacuum-outgassing at T > 573 K. This is consistent with generation of reduced active-sites, such as electron-vacancy complexes, by such treatments.

Prereduced 0.5% RhOx/CeO2 was not passivated under N20 pulses at 623 K in analagous fashion to the passivation of rhodia powder prereduced to metal, thereby again implying synergy between CeO2.x and Rhodia dispersed thereon.

Acknowledgements. Financial support for research at both laboratories under EC contract SCI-CT91-0704 is gratefully acknowledged by all co-authors. The coordinator (JC) also thanks M. Jauch, F. Farrell, J.A. Sullivan and M. O'Neill for valuable assistance.

REFERENCES

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

10. 11. 12. 13.

K.C. Taylor, Catal.Rev.Sci.Eng., 35(1993)457. C.H.F. Deden, D.N. Belton and S.J. Schmeig, J.Catal.,155(1995)204. D.N. Belton and S.J. Schmeig, J.Catal.144(1993)9. E.R.S. Winter, J.Catal.34(1974)431; 19(1970)32; 15(1969)144. C. Lamonier, G. Wrobel and J.P. Bennell, J.Mater.Chem.,4(1994)1927. S.E. Golunski, H.A. Hatcher, R.J. Rajaram and T.J. Truex, Appl.Catal.B.Envir.,5(1995) 367. J. El-Fellah, S. Boujara, H. Dexpert, A. Kiennemanov, J. Magerus, O. Tourtet, F. Villain and F. Le Normand, J.Phys.Chem.,98(1994)5522. J. Cunningham, D. Cullinane, F. Farrell, J.P. O'Driscoll and M.A. Morris, J.Mater.Chem.,5(1995)1027. J. Soria, A. Martinez-Arias and J.C. Conesa, J.Chem.Soc.Farad.Yrans.,91(1995). J. Soria, A. Martinez-Arias and J.C. Conesa, Vacuum(1992)437. A.D. Logan and A.K. Datye, Surf.Sci.,245(1991)280. D.G. Castner and G.A. Somorjai, Applic.Surf.Sci.,6(1980)29. J. Cunningham, J.J. Kelly and A.L. Penny, J.Phys.Chem.,74(1970)1992.

J.W. Hightower, W.N. Delgass, E. Iglesia and A.T. Bell (Eds.) 11th International Congress on Catalysis - 40th Anniversary

Studies in Surface Science and Catalysis, Vol. 101 9 1996 Elsevier Science B.V. All rights reserved.

691

The activity of VOx/ZrO2 for the selective catalytic reduction of NO V. Indovinaa, M. Occhiuzzi a, P. Ciambelli b, D. Sannino b, G. Ghiotti c and F. Prinettoc, aChem. Dept., "La Sapienza" University, Roma, P.le A. Moro 5, 00185, Roma, Italy bChem and Food Engineering Dept., Salemo University cChem. Inorg., Chem. Phys. and Materials Dept., Torino University Abstract

Samples VOx/ZrO2, prepared by i) adsorption from aqueous NH4VO3 solutions at pH=l-4, ii) dry impregnation with the same solution, or iii) adsorption from vanadyl acetylacetonate solutions in toluene, were characterized by means of ESR, XPS and IR spectroscopies. In the selective catalytic reduction of NO with NH3 in the presence of 02 (SCR), VOx/ZrO2 catalysts were active and stable. In the NO+NH3 reaction, they had much lower catalytic activity. Their activity depended only on the vanadium content, not on the method used for preparing the catalysts. Catalytic activity (molecules nm -2 s -1) markedly increased with the vanadium concentration up to 3 atoms nm -2 and changed little thereafter, paralleling the increased concentration of specific polyoxovanadates, detected by IR. The surface concentration of NH~ also paralleled the SCR activity. The results suggest a possible role in SCR for NH; ions and adjacent chelating nitrates, also identified by IR.

1. INTRODUCTION The selective catalytic reduction of NO with NH3 in the presence of 02 (SCR) has been extensively studied mainly on VOx supported on TiO2 [1-4]. The commercial catalysts for the SCR of flue gases from stationary sources are V205-TiO2 and V205(-WO3)/TiO2. Many studies have investigated the dispersion, the nuclearity and the oxidation state of vanadium supported on TiO2 [5-14]. All these properties might depend on the support and it was therefore of interest to extend the study to other supports and particularly ZrO2. Szakacs et al. [15] have studied the SCR activity of VOx/ZrO2 catalysts prepared by adsorption on ZrO2 of VO(acetylacetonate)2 from toluene solutions. The SCR mechanism has been investigated by isotopic tracers [16-18], and the surface species on VOx/TiO2 by spectroscopy [19-22]. Takag! et al. [23] proposed a mechanism involving the reaction of adsorbed NO2 with NH~. Although Tops~e et al. [22] evidenced the participation of NH~ species, they are in favour of a mechanism involving reaction with gas phase NO or weakly adsorbed NO. Several investigators have proposed a redox mechanism involving vIv/v v species [1, 3, 6-8, 24] and have pointed out the need for two adjacent V sites [2, 4, 9, 15]. In this paper we report (i) the catalytic activity for SCR of VOx/ZrO2 samples prepared by various methods (adsorption from aqueous metavanadate solutions at different pH values, dry impregnation, and adsorption from VO(acetylacetonate)2 in toluene), (ii) sample characterization (nuclearity, dispersion and oxidation state) by means of XPS, ESR and FTIR and (iii) the nature and reactivity of the surface species observed in the presence of the reactant mixture. Catal .ytic results are here reported in full. Characterization data relevant to the discussion of the catalytic activity will be given, whereas details on the catalysts preparation and

692

characterization will be reported elsewhere.

2. E X P E R I M E N T A L 2.1. Sample preparation

The zirconia support was prepared by hydrolysis of zirconium oxychloride with ammonia, as already described [25]. Before its use as support, the material was calcined in air at 823 K. VOx/ZrO2 samples were prepared by three methods: (i) adsorption from a solution of ammonium metavanadate (AV) at pH values from 1 to 4, adjusted by nitric acid, (ii) dry impregnation with AV solutions and (iii) adsorption from a solution of VO(acetylacetonate)2 in toluene. VOx/ZrO2 catalysts were designated as ZVx(y)pHz, where x gives the analytical vanadium content (weight percent), y specifies the preparation method (a, adsorption, i, impregnation or acac, acetylacetonate) and z the AV solution pH. The V-content was determined by atomic absorption (Varian Spectra AA-30) after the sample had been dissolved in a concentrated (40%) HF solution. BET surface areas (SNm2g -1) were measured by N2 adsorption at 77 K. The SA of ZrO2 was 49 m2g -1. The SA of some ZV samples was determined after the various treatments. All these samples had SA values ranging from 45 to 49 m2g -1, slightly smaller than those of zirconia.

2.2. Procedure and characterization techniques

Specimens were placed in a silica reactor that was equipped with two side tubes for XPS and ESR measurements and connected to a circulation apparatus, described elsewhere [25, 26]. The catalysts, dried at 383 K, were characterized as prepared (a.p.), after heating in dry oxygen at 773 K (s.o.), or after reduction with CO. In some experiments, as specified, samples were exposed to NO, NH3, or various mixtures NO-O2-NH3. Electrons per V atom (eN) were determined from the CO consumed. The average oxidation number of vanadium was calculated as 5 - eN. FT-IR spectra were recorded at RT on a Perkin-Elmer 1760-X spectrophotometer equipped with a cryodetector, at a resolution of 2 cm -1 (number of scans -100). In the 1070-960 cm -1 region, band integration and c u r v e fitting were carried out by "Curve fit, in Spectra Calc." (Galactic Industries Co.). Powdered materials were pelleted in self-supporting discs of 25-50 mg cm -z and 0.1-0.2 mm thick, placed in an IR cell allowing thermal treatments in vacuo or in a controlled atmosphere. The ESR measurements were made at RT or 77 K on a Varian E-9 spectrometer (X-band), equipped with an on-line computer for data analysis. Spin-Hamiltonian parameters (g and A values) were obtained from calculated spectra using the program SIM14 A [26]. The absolute concentration of the paramagnetic species was determined from the integrated area of the spectra. Values of g were determined using as reference the sharp peak at g = 2.0008 of the E'I center (marked with an asterisk in Fig. 3); the center was formed by UV irradiation of the silica dewar used as sample holder. XPS measurements were obtained with a Leybold Heraeus LHS 10 spectrometer operating in FAT mode and interfaced to a 2113 HP computer. Mgk(~ (1253.6 eV) radiation (12 kV and 20 mA) was used. The a.p. sample was pressed onto a golddecorated tantalum plate attached to the sample holder. After the various treatments (s.o., or reduction with CO), the specimen was transferred into the above mentioned XPS tube without exposure to the atmosphere. The spectra were collected by the computer in a sequential manner (figure in parenthesis gives the kinetic energy): Ols (719.0 eV), V2pl/2 (724.5 eV), V2P3/2 (732.0 eV), Zr3d3/2 (1062.5 eV) and Zr3d5/2

693

(1066.0 eV). The binding energy (BE) of Ols (530.0 eV) was taken as reference. The spectrum analysis involved (i) satellite subtraction of Mgk(z components; (ii) inelastic background removal by a linear integral profile; (iii) curve-fitting by a least-squares method, using a mixed Gaussian-Lorentzian function and (iv) determination of the peak area by integration. Satellite subtraction of ec3 and (~4 oxygen components allowed V2pl/2 to be partially resolved. 2.3. Catalytic experiments Catalytic experiments were done in an apparatus consisting of a flow measuring and control system (mass flow controllers, Hitech), fixed-bed flow microreactor, electrically heated and equipped with a temperature programmer-controller (Ascon), two on-line IR analyzers, one for NO (Radas 1G, Hartmann & Braun) and the other for NH3 (Siemens, Ultramat 5E), and an on-line gas chromatograph (Dani 86.10 HT), equipped with a 2 m length column (AIItech CTR) for the analysis of 02, N2 and N20. Typical experiments were conducted in the temperature range 473-723 K, feeding a gas mixture containing 700 ppm NO, 700 ppm NH3 and 3.6 % 02 in helium. The effect of 02 partial pressure was also tested. The flow rate of the reactant gas was 60 L/h (W/F = 5x10 -6 g h cm-3). NH3 was oxidyzed by feeding a gas mixture containing 700 ppm NH3 and 3.6 % 02 in helium. The nitrogen mass balance was better than 90%. Catalytic data were expressed as NO or NH3 conversion percent, or calculated as apparent kinetic constants (k/NO molecules nm -2 s-l), assuming the occurrence of a single reaction (4NO + 4NH3 +02 = 4N2 + 6H20), first order with respect to NO and zero order with respect to NH3.

3. RESULTS AND DISCUSSION 3.1. Vanadium uptake For samples ZV(a)pH1, up to about 2.5 mmol L -1, the vanadium uptake (atoms nrn -2) was proportional to the AV concentration, but as the concentration increased further, the uptake tended to level off (Fig. 1, a). o,I

"E

a

'E t-

r

E o

15

>

10

t"

E o

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,,

2

5

;

10

20

30

NH4VO 3 concentration/mmol L1

40

0

5

10

15

V available/atoms n m "2

Figure 1. Section a V-uptake (atoms nm -2) vs. AV concentration (retool L-l) at (z~) pH=l, (El) pH=2, (O) pH=3 and (o) pH=4. Section b: V-uptake (atoms nm -2) vs. the vanadium available in the AV solution (referred to the ZrO2 surface area) at (z~) pH=l, (El) pH=2, (O) pH=3 and (o) pH=4. For samples ZV(a)pH2-4, up to about 2.5 mmol L-1 the uptake increased linearly

604 and thereafter remained about constant up to 6.5 mmol L-1 . In contrast to ZV(a)pH1, in ZV(a)pH2-4 samples, at AV concentration > 6.5 mmol L -1, the vanadium uptake increased sharply. The increase depended on the precipitation of a vanadium phase on the zirconia support. Accordingly, in these samples, the X-ray analysis showed the presence of a segregated phase, which transformed into V205 after calcination in air at 773 K or s.o. treatment. Because V-uptake is a surface phenomenon, we plotted (Fig. 1, b) the V adsorbed per unit area of zirconia_(atoms nm -z) as a function of V atoms available in the AV solution (V atoms nm-z, referred to the ZrO2 surface area). For ZV(a)pH2+4 samples, up to 2.5 atoms nm -z, all the available V was adsorbed. As the available V increased further, uptake reached an extended plateau, corresponding to about 3 V atoms nm -2. By contrast, for ZV(a)pH1 samples, V-uptake progressively increased throughout the region of the available V. The maximum V-uptake was about 2.4 atoms nm -2. 3.2. XP$ characterization For all samples, both a.p. and s.o., irrespective of the preparation method, the experimental intensity ratios, V2p/Zr3d, increased proportionally to the V-content up to 3 atoms nm -2 (Fig. 2). The ratio approaches those calculated with the "spherical model" proposed recently by Cimino et al. [27] (full line in Fig. 2). For ZV samples with V-content _< 3 atoms nm -2, this finding shows that vanadium species are uniformly spread on the ZrO2 surface. On ZV catalysts with a larger V content (not shown in Fig. 2), the intensity ratios were markedly larger than the corresponding values yielded by the spherical model. The results obtained on samples with V-content > 3 atoms nm -z point therefore to a V surface enrichment.

Figure 2. Intensity ratio, V2p/Zr3d, vs. V-content. Samples: ZV(a)pH1 (o) a.p. and (l) s.o.; ZV(a)pH2 (z~) a.p. and (&) s.o.; ZV(a)pH3 (El) a.p. and (11) s.o.; ZV(a)pH4 () a.p. and (O) s.o.; ZV(i) ( v ) a.p. and ( v ) s.o.; ZV(acac) (+) a.p. and (x) s.o.

Figure 3. ESR spectra at RT of reduced VOx/ZrO2 samples (CO at 623 K). Samples: (a)ZV0.05(a)pH1; (b) ZV0.33(a)pH4; (c) ZV0.58(a)pH1; (d) ZV1.09(a)pH4.

Because of the intensity of V2ppeaks, which is much lower than that of the nearby Ols peak, the vanadium oxidation state could be reliably ascertained only for ZV(a) and ZV(i) specimens with a V loading > 1% (2.5 atom nm-2). The binding energy value of the V2p3/2 component, obtained by curve fitting of the region Ols-V2p,

695

showed Vv only (517.1 eV), in a.p. and s.o. samples, and complete reduction to V Iv (516.6 eV), after reduction with CO at 500 K. On the same sample the redox cycle showed eN=l, corresponding to an average vanadium oxidation state of 4. 3.3. E SR characterization In a. p. and s.o. ZV(a) and ZV(i) samples, no ESR signals were detected. In a.p. ZV(acac), a weak ESR signal of vanadyl species was detected (5% of total V), absent after the s.o. treatment. The spectra of samples reduced with CO at 400 to 623 K consisted of a signal showing a resolved hyperfine structure (Vh), overlapping a broad (,~Hpp = 300 Gauss) and nearly-isotropic band (Vb, giso = 1.97) (Fig. 3). When recorded at 77 K, both Vh and Vb maintained the same shape as at RT, and their intensity as a function of temperature followed the Curie law. The spectroscopic features of Vh (gll=1.913, gL=1.983, and All=204 Gauss, Aj.= 76 Gauss, line-width, dependence on recording temperature, and number of lines), allow the signal to be assigned to mononuclear isolated V TM in a square pyramidal configuration (vanadyl species). The absence of a hyperfine structure and the large value of the line-width of Vb, both features arising from dipolar and exchange interactions among paramagnetic species, suggest its assignment to magnetical!y interacting V IV, formed by the reduction of polyoxoanions anchored to the zlrcon=a surface. The sequence of spectra, referring to ZV samples with increasing V-content, shows that in the low-loading ZV samples up to 0.2 atoms nm -2 isolated mononuclear V Iv species prevailed, whereas with increasing V-loading interacting VIV became prevalent (Fig. 3). Exposure of s.o. samples to NH3-NO at 623 K caused the formation of a weak Vh signal. A subsequent treatment with NO-NH3-O2 mixtures, containing increasing amounts of 02, caused a progressive decrease in Vh and its disappearance with 1-2% 02. Exposure to NO-NH3 of reduced samples (CO at 623 K), therefore containing Vh and Vb, caused a decrease in the ESR detected V IV by 65%. After this treatment Vb species were nearly absent. Exposure of reduced samples to NO-NH3-O2 caused the complete oxidation of V Iv species. 3.4. FTIR characterization in a.p. ZV(a) and ZV(i) samples, broad bands arising from hydrated vanadates were detected in the 800-1100 cm -1 region. Metavanadate-like species (band at 920 cm -1) prevailed on ZV samples with V-content < 1.5 atoms nm -z and decavanadates (bands at 850-880 cm -1 and 960-990 cm -1) in the range 1.5-3 atoms nm -2. A.p. ZV(acac) samples showed bands from CH3 and C=O, suggesting the adsorption of VO(acac)2 as such (spectra not reported). Spectra of s.o. samples differed markedly from those of a.p. samples and were unaffected by a subsequent evacuation up to 673 K (Fig. 4, a). Spectra consisted of a composite envelope of heavily overlapping bands at 980-1070 cm -~, with two weak bands at 874 and 894 cm -1. Irrespective of the preparation method, the integrated area (cm -1) of the composite band at 980-1070 cm "1 was proportional to the V-content up to 3 atoms nm -z. An analysis of spectra by the curve-fitting procedure showed the presence of several V=O modes. The relative intensity of the various peaks contributing to the composite band depended only on the V-content and did not depend on the method used for preparing the catalysts. Samples with V > 3 atoms nm -~" had IR-spectra features similar to those of pure V205 (spectrum 8 in Fig. 4, a). According to the dependence of the intensity of the various peaks on the V-content, we distinguished vanadates with different nuclearities (roughly three types). The first, corresponding to peak 3 and prevailing in most dilute samples (spectra 2 and 3 Fig. 4, b), is a low nuclearity species possibly mononuclear (type-I). Type-II vanadates had an increasing concentration in the vanadium range 0.4-0.8

696

atoms nm "2, peaks 1 2 and 4 (Fig. 4, b and r Type-Ill vanadates had a markedly increasing concentration in the vanadium range 1.5-3 atoms nm -2 and increased little thereafter, peaks 5, 6 and 7 (Fig. 4, b and r

Figure 4. IR spectra of s.o. samples. Section a: ZrO2, curve 1; Z V0.18(a)pH4, curve 2; ZV0.30(acac), curve 3; ZV0.58(a)pH4, curve 4; ZV0.83(a)pH4, curve 5; ZVl.05(i), curve 6; ZV1.21 (a)pH4, curve 7; ZV4.65(a)pH4, curve 8. Section b: curve fitting of the band at 980-1070 cm-1; 990-1000 cm -1, peak 1, 1007-1008 cm -1, peak 2, 1017-1020 cm -1, peak 3, 1025-1029 cm -1, peak 4, 1034-1038 cm -1, peak 5, 1042-1045 cm -1, peak 6 and 1050-1052 cm -1, peak 7. Sample s.o. ZV0.58(a)pH4. Section c: as in section b. Sample s.o. ZV1.21(a)pH4.

Figure 5. IR spectra of s.o. samples after various treatments. Section a: after adsorption of NH3 (1 mbar) at RT: ZrO2 (curve 1), ZV0.58(a)pH4 (curve 2), ZV1.21(a)pH4 (curve 3). Section b: ZV0.58(a)pH4 sample after adsorption of NO+O2 at 623 K (curve 1), after subsequent adsorption of NH3 (1 mbar) at RT (curve 2) and after subsequent heating at 623 K (curve 3). Bands assigned to bridged bidentate nitrates (*) and to chelating nitrates (**). Section c: the same treatments as in section b on s.o. ZV 1.21 (a)pH4. At RT, NH3 adsorbed on Lewis acid sites, Zr Iv and V v. Accordingly, the intensity of bands from NH3 decreased little with the V-content, by 15% at most, as expected on account of the similar Lewis acid strengths of Zr Iv and Vv. The symmetric bending

697

of NH3 was 1157 cm -1 on pure ZrO2, and shifted to higher frequency on ZV. In particular, on most dilute ZV the frequency band was 1195 cm -1 and with increasing V-content it progressively increased up to 1208 cm -1 (Fig. 5, a). NH~ did not form on ZrO2 and ZV samples with V-content < 1.5 atoms nm -2, whereas it did form on more concentrated samples and markedly increased with V-content up to 3 V/atoms nm -2 (Fig. 5, a). At RT, NO adsorption on s.o. ZV samples gave weak bands from N20 and nitrites. The same species formed on pure ZrO2. Adsorption of NO+O2 gave strong bands from bridged bidentate-nitrates (spectra 1 in Fig. 5, b and r On both ZV0.58(a)pH4 and ZV1.21(a)pH4 after evacuation at RT, the subsequent addition of NH3 at RT gave more intense NH~ bands than those on s.o. samples and caused the concomitant transformation of bridged nitrates into chelating nitrates (spectra 2 in Fig. 5, b and r NH~ species were much more intense in ZV1.21(a)pH4 than in ZV0.58(a)pH4. A subsequent heating at 623 K caused the disappearance of chelating nitrates in ZV1.21(a)pH4 (spectrum 3 in Fig. 5, r and their decrease (to 50%) in ZV0.58(a)pH4 (spectrum 3 in Fig. 5, b), while surface-OH formed and H20 and N2 were detected by analysis of the gas phase. Exposure of s.o. samples to NH3-NO at 623 K, caused a slight reduction of V v to V iv, whereas exposure to NO-NH3-O2, did not affect the vanadium oxidation state. Exposure of reduced samples (CO at 623 K) to NH3-NO caused slight oxidation, whereas exposure to NO-NH3-O2 oxidized all V IV. Catalytic activity On all catalysts, the activity for SCR was stable as a function of the time on stream and the ratio NO/NH3 remained very close to unity. 3.5.

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Figure 6. NO conversion (%) vs. temperature. Section a: (n) ZV0.18(a)pH4, () ZV0.34(a)pH4, (m) ZV0.58(a)pH1, ( v ) ZV0.60(a)pH4, (~) ZV0.83(a)pH4, (o) ZV1.17(a)pH4, (zx) ZV4.79(a)pH4. Section b: ()ZV0.17(i), (v)ZV0.32(i), (~)ZV0.64(i), (m)ZVl.05(i), (z~)ZV0.30(acac), (o)ZV0.96(acac), (D)ZV1.36(acac)and (x)ZrO2. In samples ZV(a) (Fig. 6, a), ZV(i) and ZV(acac) (Fig. 6, b), NO conversion increased with V-loading at all temperatures. In the whole temperature range, the selectivity to N2 was ve~ high. A small amounts of N20 (< 3%) were detected only above 573 K. Pure zirconla showed some SCR activity at T > 573 K, comparable with

698

that of ZV0.18(a)pH4. The best performance was obtained with the catalyst ZV1.17(a)pH4, containing 2.8 V atoms nm -2, namely a V-content close to that of the adsorption plateau in ZV(a) samples (3 atoms nm-Z). On the sample ZV4.79(a)pH4 (12.6 V atoms nm-2), containing segregated V205, NO conversion reached a maximum at 623 K and decreased thereafter. Higher reaction temperatures resulted in the formation of very large amounts of N20 (>10 %) arising from the oxidation of NH3; NH3 conversion still monotonically increased with the temperature. In all ZV catalysts, the apparent activation energy (Ea/kJ mol "1) was nearly independent of the V-content (42 + 4 kJ mol-1). Therefore, the dependence of catalytic activity on the V-content can be conveniently inspected through k ~ values, the pre-exponential factor of the Arrhenius equation. The finding that k ~ values, irrespective of the method used for preparing the catalysts, stay on the same curve, shows that the SCR activity is mainly controlled by the vanadium content (Fig. 7, a). The marked and non-linear increase of k ~ with the V-content clarifies that the concentration of the active vanadium is not proportional to the V-loading. Namely, only specific configurations are active. To identify the active vanadium configuration, we divided k ~ values by the intensity of i) the composite band at 980-1070 cm -1, ii) peak corresponding to type-I vanadates, iii) peaks of type-II polyvanadates, and iv) peaks of type-Ill polyvanadates. Normalized k ~ values monotonically and markedly increase by a factor 8 to 11, but normalized k ~ values for type-Ill polyoxovanadates remain about constant, well within a factor of two (Fig. 7, b).

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Figure 7. The dependence of catalytic activity on the V-content (a) and its correlation with type-Ill polyoxovanadates (b). Section a: 10-3 k ~ vs. V-content on ZrO2 (0), ZV(a) (El), ZV(i) (z~), and ZV(acac) (O). Section b: 10.2 k ~ divided by the total integrated area (cm -'1) of bands in the region 980-1070 cm -1 (El), and 10-~k ~ divided by the sum of components 5, 6 and 7 (0) areas (cm-1). In the absence of 02, NO reduction continued, however at a rate about ten times lower than that in the presence of 02. During 20 h experiments NO conversion remained constant. On 02 addition, the catalytic activity increased with 02 content in the mixture up to about 1000 ppm, and changed little thereafter. We noticed that increasing the 02 concentration caused NO conversion to become lower than that of NH3, probably due to changes in the stoichiometry of the overall reaction (the NO/NH3 ratio passed from 1.5 to 1). Catalytic tests of NH3 oxidation with 02 yielded high selectivity to N2 (66-90%), which decreased with the higher loading catalysts. in special experiments we tested the activity of ZV samples prereduced in an NH3

699

flow (700 ppm in He) for 1 h at 623 K. In the SCR and NO + NH3 reaction the prereduced ZV1.04(a)pH4 sample showed the same activity as the s.o. sample. 4. C O N C L U S I O N S The heating in 02 at 773 K of VOx/ZrO2 stabilizes surface vanadates of various nuclearities. Provided that the V-content in the samples is < 3 atoms nm -2, XPS shows effective spreading of vanadates on the ZrO2 surface. As the V-content increases, the concentration of monomers increases little, whereas that of polyoxovanadates increases markedly, particularly that of type-Ill polyoxovanadates. The relative abundance of the various vanadates depends only on the V-content, not on the method used for catalyst preparation (impregnation, or adsorption from both aqueous and toluene solutions) and pH of the solution used for adsorption (pH =1-4 for AV solutions). This finding strengthens our earlier proposal [28, 29], based on the results obtained on the related MoOx/ZrO2 system, that the ZrO2 surface has a buffer effect. When samples are heated, the buffer effect causes the condensationdecondensation of vanadium species, therefore making the surface composition independent of the nature of the precursor-adsorbate. In agreement with the results from the characterization, the SCR activity of VOx/ZrO2 also depends only on the V-content, not on the method used for catalyst preparation. The marked increase in SCR activity with the V-content shows that only specific vanadium configurations are active. Although we assess the V=O modes associated with these active configurations, IR analysis did not specify the structure of active polyoxoanions. The presence of V Iv on the surface before catalysis is unessential for ca.tai~ic activity. We cannot however rule out an SCR redox mechanism involving Vv-v Iv. ESR and IR results show that the oxidation state of surface vanadium at the reaction temperature is controlled mainly by the composition of the reactant mixture. "7,

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Figure 8. The catalyt!c activity for SCR (10 -3 k~ vs. concentration of NH~ (from the IR area of 700 *C; whereas, the conventional catalytic studies of Vannice and co-workers have been carried out at temperatures < 700 ~ Therefore, for comparison with our low pressure results, the effect of temperature on the rate of N 2 formation at a total pressure of 760 Torr has been investigated up to 875 ~ for three gas compositions: (i) 1% NO, (ii) 1% NO + 0.25 % CH4 and (iii) 1% NO + 0.25 % CI-I~ + 0.5 % 02. The results are summarized in Figure 7. At 700 ~ the trends with respect to the effect of gas composition are qualitatively the same as those reported by Vannice; i.e., the addition of CH4 enhanced the rate of N2 formation, and the addition of 02 enhanced the rate even more. At 700 ~ the effect of adding CH4 was to increase the rate two-fold, while the effect of adding CH4 + 02 was to increase the rate five-fold. Unexpectedly, the rate of N 2 formation in the presence of CI-I4 and 02 went through a maximum at 700 *C. Furthermore, in the temperature range 700 *C - 800 *C, the conversion of CH4 approached nearly 100%. Although it is not possible to measure directly the amount of O2 under reaction conditions, one may presume that some remains even at temperatures of ca. 700 ~C. Nevertheless, at higher temperatures as 02 is consumed its positive effect diminishes. Moreover, as shown in Figure 3, the production of CH3" radicals reached a maximum at 750~ The decrease in oxygen concentration and the maximum in CH3" radicals formation may cause the maximum in the rate of N2 formation that is described in figure 7. In addition to the maximum at 700 ~ the rate of N2 production also went through

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Teml~mture,~ Figure 7. Rate of N 2 formation over Sr/La203 in a conventional catalytic reactor with the following reagent gases:, O 1% NO;A, 1% NO + 0.25% CH4; ll, 1% NO + 0.25% CH 4 + 0.5% 02. The total pressure was 760 Torr.

718 a minimum at 800 ~ Such a minimum is expected since, at sufficiently high temperatures, the direct decomposition of NO would become significant. In fact, by 850 ~C, the direct decomposition of NO was about twice as fast as when CH4 and 02 were added as reagents. This apparent negative effect of CH4 and 02 probably is a result of CO2, which is known to be a strong poison for these basic oxide catalysts during oxidative coupling reactions [16,17]. It is relevant to the low pressure studies, particularly those of Figure 5, that at 735~ CI-L + 02 had a positive effect on the rate of N2 formation and on the rate of CH3" radical production, relative to the case when only CH4 + NO were the reagents. These results are consistent with the view that CH3" radicals are intermediates in the reaction of CH 4 with NO when 02 is present, but in the absence of 02, the reduction of NO by CH4 over Sr/La203 may occur via another pathway that does not involve CH3" radicals.

ACKNOWLEDGMENT This research was supported by a grant from the U. S. Department of Energy, Office of Basic Energy Sciences. The authors wish to thank Professor M. A. Vannice for providing us with preprints of his unpublished papers and for helpful discussions.

REFERENCES 1. 2. 3. 4. 5. 6. 7. 8. 9. 10. 11. 12. 13. 14.

X. Zhang, A. B. Waiters and M. A. Vannice, J. Catal., 146 (1994) 568. X. Zhang, A. B. Waiters and M. A. Vannice, Appl. Catal. B, 4 (1994) 237. X. Zhang, A. B. Waiters and M. A. Vannice, J. Catal., 155 (1995) 290. X. Zhang, A. B. Waiters and M. A. Vannice, Appl. Catal. B, in press. K. D. Campbell, E. Morales and J. H. Lunsford, J. Am. Chem. Soc., 109 (1987) 7900. K . D . Campbell and J. H. Lunsford, J. Phys. Chem., 92 (1988) 5792. Y. Feng and D. Gutman, J. Phys. Chem., 95 (1991) 6558; Y. Feng, J. Niiranen and D. Gutman, J. Phys. Chem., 95 (1991) 6564. T. Johnston and J. Heicklen, J. Phys. Chem., 70 (1966) 3089. G. Hermig and H. Gg. Wagner, Berm. Bunsenges, Phys. Chem., 98 (1994) 749. M. Xu and J. H. Lunsford, Catal. Lett., 11 (1991) 295. G. H. Smudde, X. D. Peng, R. Viswanathan and P. C. Stair, J. Vac. Sci. Technol., 9 (1991) 1985. M. Xu, T. H. Ballinger and J. H. Lunsford, J. Phys. Chem., 99 (1995) 14494. H. Yamamoto, H. Y. Chu, M. Xu, C. Shi and J. H. Lunsford, J. Catal., 142 (1993) 325. A. Maschke, B. S. Shapiro and F. W. Lampe, J. Am. Chem. Soc., 86 (1964) 1929.

719 15. M. T. Macpherson, M. J. Pilling and M. J. C. Smith, J. Phys. Chem., 89 (1985) 2268; I. R. Slagle, D. Gutman, J. W. Davies and M. J. Pilling, J. Phys. Chem., 92 (1988)2455. 16. M. Xu, C. Shi, X. Yang, M. P. Rosynek and J. H. Lunsford, J. Phys. Chem., 96 (1992) 6395. 17. M. Xu, Ph.D. Dissertation, Texas A&M University, 1994.

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