Livro Importante

Studies in Surface Science and Catalysis 45

TRANSITION METAL OXIDES: Surface Chemistry and Catalysis

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Studies in Surface Science and Catalysis

Advisory Editors: B. Delmon and J.T. Yates

Vol. 45

TRANSITION METAL OXIDES: Surface Chemistry and Catalysis

Harold H. Kung

Chemical Engineering Department, The Technological Institute, Northwestern University, 2 145 Sheridan Road, Evanston, IL 60208, U. S. A.

E LS EVI E R Amsterdam - Oxford - New York - Tokyo 1989

ELSEVIER SCIENCE PUBLISHERS B.V. Sara Burgerhartstraat 25 P.O. Box 2 1 1, 1000 AE Amsterdam, The Netherlands

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V

To my mentor

Professor Robert L. Burwell, Jr

and my wife

Mayfair

vii

PREFACE

In the past twenty years there have been rapid advances in the understanding of surface phenomena as a result of the availability of many new and powerful techniques to study solids and their surfaces. Clever experimentation coupled with techniques such as laser Raman spectroscopy, solid state NMR, extended X-ray absorption fine structure spectroscopy, as well as a host of surface sensitive electron spectroscopies have provided information on the atomic structure and composition of surfaces and near surface regions. To date, the bonding geometries, the positions of adatoms. the orientations of adsorbed molecules, the structures of overlayers, the extent of relaxation of surface atoms, and even the atomic positions of surface defects have become accessible. Indeed, many of these data have been obtained for metallic surfaces, especially transition metal surfaces where the attention of most surface scientists has been focused. These data have resulted in significant advances in the understanding of the surface chemistry and catalytic properties of these materials.

It is anticipated that a similar rapid increase in our level of understanding of surface phenomena will be forthcoming for transition metal oxides. Indeed, a growing number of investigators have turned their attention to oxidic materials. Simultaneously, new catalytic properties of transition metal oxides are being discovered. New reactions are being reported, such as metathesis, selective oxidation of butane to maleic anhydride and various photo-enhanced processes, and new catalytic materials are being synthesized, especially various mixed oxides. It appears that the level of understanding in the area of surface chemistry and catalysis of transition metal oxidcs is poised for a quantum leap.

In this book I have attempted to summarize existing information on the structure, electronic properties, chcrnistry, and catalytic properties of transition metal oxide surfaces so that it can serve as a uscrul source of information for investigators in this field and as a comprchensivc overview of the subject for graduate students. The book is inlended for surface physicists, chcmists, and catalytic engineers. By presenting physical, chemical and catalytic properties in one volume, it is hopcd that the interrclationship among them will become more apparent.

... VIU

The subjects in this book can be divided into three sections. The first section (chapters 1 to 3) deals with the structural, physical, magnetic, and electronic properties of transition metal oxides. Although the emphasis is on surface properties, relevant bulk properties are also discussed. The second section (chapters 4 to 7) covers surface chemical properties. It includes topics that describe the importance of surface coordinative unsaturation in adsorption, the formation of surface acidity and the role of acidity in determining surface chemical properties, the nature and reactivities of adsorbed oxygen, and the surface chemistry in the reduction of oxides. The third section is on the catalytic properties (chapters 8 to 14). Various catalytic reactions including decomposition, hydrogenation, isomerintion, metathesis. selective oxidation, and reactions involving carbon oxides are discussed. Emphasis is placed more heavily on reaction mechanisms and the role of catalysts than on kinetics and processes. A chapter on the preparation of oxide catalysts and one on photo-assisted processes are also included. Whenever appropriate, relationships among various topics are indicated.

It would have been impossible to complete this book without the encouragement and help of a number of people. The most significant of them is Professor Robert L. Burwell, Jr. who provided numerous suggestions and comments on its content. The constant encouragement and helpful discussions with Dr. Mayfair C. Kung are also greatly appreciated. The enthusiasm and dedication of my students have made the study of this subject enjoyable. Much of our research in this area has been supported by the U. S. Department of Energy without which I would not have the necessary background to undertake the task. Finally, permission by various publishers to reproduce the figures and tables in the text is also gratefully acknowledged.

Harold H. Kung 1988

ix

CONTENTS

PREFACE vii

Chapter 1. INTRODUCTION References 5

1

Chapter 2. BULK AND SURFACE STRUCTURE OF 6 TRANSITION METAL OXIDE 2.1 Bulk Structure 6 2.2 Shear Structure in Intermediate Oxides of Ti, V, and Mo 12 2.3 Structure, Stability and Reconstruction of Oxide Surfaces 15 2.4 Structure of Supported Oxides 20 References 25

Chapter 3. PHYSICAL AND ELECTRONIC PROPERTIES 27 3.1 Surface Composition 27 3.2 Ionicity of Oxides 28 3.3 Magnetic Properties of Small Oxide Particles 36 3.4 Quadrupole Splittings of Surface Ions 43 3.5 Surface Electronic Structure 46 3.6 Surface Vibration 48 References 50

Chapter 4. SURFACE COORDINATIVE UNSATURATION 53 4.1 Formation of Surface Coordinative Unsaturation 53 4.2 Chemical Properties of Surface Coordinatively

4.3 Adsorption of Small Molecules 61 References 69

Unsaturated Sites 57

X

Chapter 5. SURFACE ACIDITY 5.1 Surface Acid Sites 72 5.2 Formation of Surface Acid Sites 74 5.3 Determination of Acidity 80 5.4 Role of Acid Sites in Catalytic Reactions 88 References 89

Chapter 6. REDUCTION OF OXIDES 6.1 Introduction 91 6.2 Thermodynamics of Reduction 92 6.3 Kinetic Models 93 6.4 Mechanism of Reduction 96 6.5 Effect of Support 98 6.6 Effect of Other Components 100 6.7 Reactivity of Reduced Surfaces 100 6.8 Influence of Reduced Oxides on the Properties of

References 107 Transition Metals 102

Chapter 7. OXYGEN ON OXIDES 7.1 Nature of Adsorbed Oxygen 110 7.2 Detection of Adsorbed Oxygen 11 1 7.3 Reactivity of Adsorbed Oxygen 116 References 119

72

91

110

Chapter 8. PREPARATION OF OXIDES 121 8.1 General Considerations 12 1 8.2 Preparation of Unsupported Single Component Oxides 123 8.3 Preparation of Supported Oxides 129 8.4 Preparation of Multicomponent Oxides 13 1 References 134

Chapter 9. METATHESIS AND ISOMERIZATION 136 9.1 Metathesis 136 9.2 Isomenzation of Alkenes 140 References 144

Chapter 10. DECOMPOSITION, HYDROGENATION AND 146 RELATED REACTIONS 10.1 Decomposition of Alcohols 146 10.2 Decomposition of Nitrous Oxide 153 10.3 Hydrogenation, H-D Exchange of Hydrocarbons, and H2-D2

Scrambling 155 10.4 Reduction of NO 162 References 166

Chapter 11. SELECTIVE OXIDATION REACTIONS I 1 1.1 Introduction 169 11.2 Types of Selective Oxidation Reactions 170 11.3 Features of Catalytic Selective Oxidation 170 11.4 Chemical Factors Affecting Selectivity 175 1 1.5 Oxidation of Propene to Acrolein and Ammoxidation

11.6 Effect of Water on Propene Oxidation 194 References 195

to Acrylonitrile 181

Chapter 12. SELECTIVE OXIDATION REACTIONS II 12.1 Selective Oxidation of Butenes 200 12.2 Selective Oxidation of Butane 210 12.3 Selective Oxidation of Methane 212 12.4 Selective Oxidation of Methanol 217 References 223

Chapter 13. CATALTYIC REACTION BETWEEN HYDROGEN AND CARBON OXIDES 13.1 Introduction 227 13.2 Alcohol Synthesis on Copper-Zinc Oxide and Other

13.3 Group VIII Metalmetal Oxide Catalysts 236 13.4 Isosynthesis Reaction 240 13.5 Water-Gas Shift Reaction 244 References 248

Cu-based Catalysts 228

xi

169

200

227

Chapter 14. PHOTO-ASSISTED SURFACE PROCESSES 252 14.1 Introduction 252 14.2 Photo-Assisted Adsorption and Desorption 256 14.3 Photocatalysis: Gas Phase Reactions 258 14.4 Photocatalysis: Liquid Phase Reactions 266 14.5 Photocatalysis by Metalmetal Oxide and Oxide/Oxide

References 273 Composites 271

INDEX 277

Chapter 1

INTRODUCTION

Transition metal oxides are technologically important materials that have found many applications. For example, in the chemical industry, these oxides are the functional components in the catalysts used in a large number of processes to convert hydrocarbons to other chemicals. They are also used as electrode materials in electrochemical processes. In the electronics industry, they are used to make conductors in films. The recently discovered high temperature superconductors are multicomponent transitional metal oxides.

Among these applications, perhaps the use of transition metal oxides as catalysts is the most technologically advanced and economically important. It is also an area in which much progress has been made in recent years in terms of the understanding of the fundamental processes that occur, primarily because advances in instrumentation and experimental techniques have made it possible to study the chemistry of the interface between the transition metal oxide and the fluid phase in greater detail than ever before. In particular, developments in surface science techniques have provided very detailed pictures about the surface structures, chemical compositions, and electronic properties of the surfaces.

Some of the chemical processes that make use of transition metal oxides are listed in Table 1-1. As can be seen from the table, many of the processes require high selectivity for a particular product, and many involve oxidation of the reactant molecules. In fact, selective oxidation, ammoxidation, and selective dehydrogenation probably constitute the most important catalytic uses of transition metal oxides. The different oxidation states available in these oxides make it possible to control the selectivity in oxidation with the properties of the oxides. Some transition metal oxides can also catalyze selective hydrogenation and are used in some commercial processes. As the demand for specialty chemicals (that is, specific chemicals for spccific processes) increases in the future, demand for high selectivity will increase in a wide variety of reactions including amination, alkylation, aldol condensation, and carbonylation, in addition to those in Table 1 - 1 . It is quite possible that transition metal oxides will occupy an increasingly

1

2

Table 1-1 Examples of Chemical Processes in which Transition Metal Oxides are Catalysts

Process

Oxidation

Dehydrogenation (nonoxidative)

Dehydrogenation (oxidative)

Selective oxidation

Selective

Selective reduction ammoxidation

Metathesis Water-gas shift

Example

Production of SO3 from SO2 CO oxidation in emission control Production of styrene from ethylbenzene

Production of formaldehyde from methanol

Production of acrolein from propene, and butadiene from butenes

and maleic anhydride from benzene or butane

Production of acrylonitrile from propene

Reduction of NO, selective hydrogenation of unsaturated ketones

Production of long chain alkenes Production of hydrogen

important position as catalysts in these chemical processes. Not listed in the table is the production of methanol by the hydrogenation of CO or C 0 2 . The earlier generation of catalysts for this process are based on zinc-chromium oxide. However, there is controversy over the current copper-zinc oxide catalyst as to whether the active component is the oxide or the copper metal.

In addition to being used as catalysts, transition metal oxides are also precursors for other important catalysts. The cobalt-molybdenum sulfide catalyst for hydrodesulfurization is an example. This catalyst is prepared by sulfiding cobalt-molybdenum oxide (often supported on alumina). Another example is the chromium-based catalyst for ethylene polymerization. The catalyst can be made from supported chromium oxide as a precursor. Finally, many noble mctal catalysts are prepared by reduction of the corresponding oxides. In addition, these metal catalysts are often stored in air and are converted to their oxides during storage. It is quite conceivable that in these cases, the detailed structures, morphologies, or other properties of the transition metal oxide precursors could affect the properties of the final catalysts.

Understanding catalysis requires an understanding of surface chemistry, which deals with the bonding and reaction of an adsorbate with the surface and the influence of the surface on the bonding and reaction between adsorbates. It is apparent that an important part of any effort toward obtaining such an undcrstanding is the ability to characterize thc physical and chemical properties of a surface. In recent years, much progress has been made in the understanding of

INTRODUCTION 3

Table 1-2 Properties that are Important in the Surface Chemistry of Transition Metal Oxides

Presence of cations and anions in stoichiometric ratios and in

Possibility of covalent and ionic bonding between cations

Presence of a strong electric field normal to the surface

Presence of charged adsorbed species Presence of surface acidity and basicity Presence of cationic and anionic vacancies Ability of cations to undergo oxidation and reduction High mobility of lattice oxygen and the possibility that

Interaction of the solid with incident photons that leads

well-defined spatial (structural) relationships

and anions

due to the coulombic nature of the ionic lattice

the lattice oxygen are reactants in a reaction

to photo-assisted surface chemical processes

metallic surfaces [l-41. Progress has also been made, but at a slower pace, for the transition metal oxides because of the higher level of complexity in the experimental techniques involved. There are significant differences between the chemistry of transition metal oxides and the corresponding metals. Table 1-2 provides a list of properties that are important in the surface chemistry of transition metal oxides. Many of them either do not apply or apply only to a limited extent to the metals.

Common to many of the properties listed in Table 1-2 is the fact that transition metal oxides are made up of metallic cations and oxygen anions. The ionicity of the lattice, which is often less than that predicted by the formal oxidation states, results in the presence of charged adsorbate species, and the common heterolytic dissociative adsorption of molecules (that is, a molecule AB is adsorbed as A+ and B3. Surface exposed cations and anions form acidic and basic sites as well as acid-base pair sites. The fact that the cations often have a number of commonly obtainable oxidation states has resulted in the ability of the oxides to undergo oxidation and reduction and the possibility of the presence of rather high densities of cationic and anionic vacancies. As can be seen throughout the discussions in this book, these properties have determining effects on the interaction of molecules with the oxide surfaces.

This book starts with discussions of the structural (Chapter 2) and physical and electronic properties of transition metal oxides (Chapter 3). Knowledge of surface structure is as important as knowledge of molecular structure in understanding surface chemistry. Physical properties of these oxides, especially those for small crystallites, are often used for identification and characterization purposes. Sometimes they are used as means to monitor chemical interactions.

4

Surface electronic properties determine the mode of bonding of the adsorbates, and both surface and bulk electronic properties determine the photo-assisted surface chemical processes.

Discussions of the surface chemical properties then follow. The importance of surface coordinative unsaturation in governing the adsorptive properties of many oxides will be discussed in Chapter 4. The discussion of surface acidity will be found in Chapter 5 , and the nature and reactivities of adsorbed oxygen will be discussed in Chapter 7. Reduction of transition metal oxides, which is initiated at oxide surfaces, will be discussed in Chapter 6.

After surface chemistry, the catalytic properties of transition metal oxides will be discussed. The discussions will begin with the methods of preparation of oxides and the dependence of the final properties of the oxides on the preparation methods (Chapter 8). Then various catalytic reactions will be discussed. These are metathesis and isomerization (Chapter 9), decomposition and hydrogenation (Chapter lo), selective oxidation (Chapters 11 and 12), reactions of carbon oxides (Chapter 13), and finally photo-assisted surface processes (Chapter 14). In the last chapter, photo-assisted surface chemical reactions will be described together with photo-assisted catalytic reactions.

The emphases of the discussions will be on transition metal oxides. However, whenever appropriate, nontransition metal oxides will also be discussed, especially when they are used to illustrate certain concepts or for comparisons. In particular, properties of ZnO will be discussed rather extensively because ZnO is among the best understood oxides whose surface chemistry, catalytic, electronic and structural properties have been studied extensively, and because its behavior is in many ways similar to many transition metal oxides.

Throughout this book, reference to information obtained from various experimental techniques will be made. The readers are referred to some excellent treatises that describe these techniques 11.4-91. The following is a list of standard acronyms that are often used and their meanings:

AES Auger electron spectroscopy ELS Electron energy loss spectroscopy EPR Electron paramagnetic resonance spectroscopy EXAFS Extended X-ray absorption fine structure FTIR Fourier transform infrared spectroscopy HREELS High resolution electron energy loss spectroscopy IR Infrared spectroscopy ISS Ion scattering spectroscopy LEED Low energy electron diffraction NMR Nuclear magnetic resonance spectroscopy SEM Scanning electron microscopy STEM Scanning transmission electron microscopy TEM Transmission electron microscopy TPD TPR UPS Ultraviolet photoelectron spectroscopy

Temperature programmed desorplion or decomposition Temperature programmed reduction or reaction

INTRODUCTION 5

UV-vis Ultraviolet-visible absorption spectroscopy XANES XPS X-ray photoelectron spectroscopy

X-ray absorption near edge spectroscopy

Other techniques will also be mentioned. They include:

Cyclic voltammetry Electric conductivity measurement Magnetization measurement Mossbauer spectroscopy Raman or laser Raman spectroscopy

It will be greatly beneficial for the readers to have some general knowledge about all of these techniques, especially regarding the types of information that are obtainable, so as to better understand the discussions of the data presented.

REFERENCES

1. G. A. Somorjai, "Chemistry in Two Dimensions: Surfaces," Cornell University

2. 'The Nature of Surface Chemical Bonds." T. N. Rhodin, and G. Ertl, ed., North-

3. S. R. Morrison, 'The Chemical Physics of Surfaces." Plenum Press. NY. 1977. 4. G. Ertl, and I. Kuppers, "Low Energy Electrons and Surface Chemistry,"

5 . "Experimental Methods in Catalytic Research," Vol. 1-3, ed. R. B. Anderson,

6. W. N. Delgass, G. Haller, R. Kellerman. and I. Lunsford, "Spectroscopy in

7. D. P. Woodruff, and T. A. Delchar. "Modem Techniques of Surface Science,"

8. "Electron Spectroscopy for Surface Analysis," Topics in Current Physics,

9. 'The Chemical Physics of Solid Surfaces and Heterogencous Catalysis,

Press, Ithaca, NY, 1981.

Holland Publ. Co., NY, 1979.

Weinheim, Germany, 1985.

et. al., Academic Press, NY, 1968-76.

Heterogeneous Catalysis," Academic Press, N.Y., 1984.

Cambridge University Press, Cambridge, 1986.

vol. 4, ed. H. Ebach, Springer, Berlin, 1977.

D. A. King and D. P. Woodruff ed., Elsevier Scientific Publ., Amsterdam, 1?8l, 1983.

Chapter 2

BULK AND SURFACE STRUCTURE OF

TRANSITION METAL OXIDE

2.1 BULK STRUCTURE

With the exception of some complex oxides of unusual stoichiometries, multicomponent compounds, and oxides that are only stable at high tcmpcrature and high pressure, the bulk structures of most transition metal oxides are known. An examination of the known structures shows that transition metal oxides exist in many different crystallographic forms. There does not appear to be a simple generalization that relates the structure to the stoichiometry or the position in the periodic table. In fact, it is not uncommon to find a certain oxide in more than one crystal structure at ordinary temperatures as a result of the high activation energy in the process of transforming from a less thermodynamically stable to a more stable structure. There is perhaps one generalization, which is the fact that the ionic radii of transition metals are smaller than that of 02-. Thus the oxygen ions are usually close-packed with the smaller metal ions situated in the octahedral and tetrahedral holes among the oxygen ions.

There are many excellent treatise on the subject of crystal structures, such as the one by Wells [l]. In this section, structures of some of the oxides that are commonly used in studies of surface chemistry and catalysis are described. The structures of the other oxides are necessarily left out.

Oxides commonly studied as catalytic materials belong to the structural classes of corundum, rocksalt, wurtzite, spinel, perovskite, rutile, and layer structure. Table 2-1 lists the stable structures of some binary oxides. These arc the structures often reported for the oxides prepared by common methods under mild conditions. In some cases, other structures exist. Furthermore, the structures indicated represent the general type. The positions of the ions may not be at the ideal positions of the highest symmetry. For example, distortions are found for FeO, NiO, MnO, and COO from the cubic lattice, and V02, Nb02, Mo02, W02 from the perfect rutile structure.

The rocksalt structure is made up of a three-dimensional array of alternating

6

U R

cl 5cr C c1

C

Table 2-1 Crystal Structural Classes of Some Common Transition Metal Oxides v)

Sc2Qcs Ti0 r* VO r Cr2Q cr MnO r FeO r COO r NiO r CUO s zno w TizQ cr VzO, cr Cr02 t Mn304sp* Fe2Q n,sp Co304 sp cuzo c cl

Ti02 t,a,b VOz t* C I Q or MnzQcs* Fe@4 sp w MnOz t*,and M

vZo5 a others

AgzO c CdO r Y 2 4 cs ZrO r NbOz t* Mooz m,(t*) TcOz m.(t) RuOz t R h z Q cr* PdO s

Z r O 2 m,tet NbZO5 mt MOO, l,(or) T C Z ~ a Rho2 t

LazQmt H Q m Ta02 t WOz m,(t*) Re02 m,(t) OsO2 t m t Pt304 cub HgO or,and others

Ta205 or WO, m Re@ cub OsO, m PtOz t (perovskite)

Re207 a

r =rocksalt s =PtSscructure, w=wurtzite, t =rutile, a = anatase, b = brookite, c = interpenetrating cristobalite, f = fluoride, cr = corundum, sp = spinel, cs = C structure, 1 = layer, or = orthorhombic, tet = tetragonal, m = monoclinic. cub = cubic,

mt = multiple modifications, * = distorted or defective

8

cations and anions (Fig. 2-la). Each ion is in the center of an octahedron whose vertices are ions of the opposite type. The structure can be viewed as being made up of comer-sharing octahedra (Fig. 2-lb). A wurtzite structure is made up of a three-dimensional net of corner-sharing tetrahedra (Fig. 2-lc,d). Each ion is in the center of a tetrahedron in which the opposite ions are at the vertices.

The corundum, the rutile, and the spinel structures are made up of layers of close-packed oxygen ions. If the oxygen ions are modeled by hard spheres, (neglecting the cations for the moment,) each ion in a close-packed layer is in contact with six others (see Fig. 2-2). When one such layer is stacked on top of another such that an ion in this layer is in contact with the maximum possible number of ions in the other layer, this ion will be sitting above a triangular hole of the other layer (point B or C), and it will be in contact with three ions in the other layer. Now consider a case of two close-packed layers stacked in this manner as in Fig. 2-2. When a third layer is put on top of these two, its ions can take the positions vertically on top of the ions in the first layer such as above point B, or the positions above point C. In the former case, the spatial positions of the layers follow the sequence ababab ... The resulting structure is called hexagonal close- packing (h.c.p.), and it forms the basis for the corundum and the rutile structure. In the latter case, the spatial sequence of the layers is abcabca bc... The resulling s.tructure is called cubic close-packing (c.c.P.), and it forms the basis for the spinel structure.

Between adjacent layers of oxygen ions in both h.c.p. and c.c.P., the interstices (holes) are bound by either four or six oxygen ions (Fig. 2-2). They are commonly referred to as tetrahedral and octahedral holes, respectively. There are as many octahedral holes as the number of oxygen ions, and half as many tetrahedral holes as octahedral holes.

In the ideal rutile structure, half of the octahedral holes are filled with cations, while the tetrahedral holes are empty. Thus the compound has a formula M02 (e.g. TiOz). In the corundum structure, two-thuds of the octahedral holes are filled. The tetrahedral holes are empty, and the compound has a formula M203 (e.g. a-FezO3).

An ideal spinel structure would have one-half of the tetrahedral holes and one-half of the octahedral holes filled, and the formula of the compound is M304

(e.g. Fe304). It is easily seen that for charge neutrality, the cations must be of two different oxidation states. The most common oxidation states are +2 and +3, and the formula can be rewritten as MIMJII04. Both M r I and MI11 can occupy the tetrahedral or octahedral holes. The equilibrium distribution depends on the nature of the cation and the temperature. For entropic reasons, increasing the temperature tends to randomize the distribution. A normal spinel is one in which all MI ions are in the tetrahedral holes, and all MI1 ions in the octahedral holes. An example is ZnFez04. An inverse spinel has all MI ions in the octahedral holes. The MI1 ions are distributed equally between the octahedral and the tetrahedral holes. Examples are Fe304 and MgFe200. Mixed spinels have intermediate distributions, The fact that some of the octahedral or tetrahedral holes arc unoccupied makes it possible that other ions may occupy these holes when thcse are exposed on the surface.

OXIDE STRUCTURE 9

Figure 2-1 a. A rocksalt structure; b. Two comer-sharing octahedra, one centered around ion A. Octahedra centered around ions A and B in FIg. 2-la are edge-sharing; c. A wurtzite structure; d. Two comer-sharing tetrahedra, one centered around ion C.

Figure 2-2 Close-packed layers of oxide ions. If ions in the third layer are above B, it is a hexagonal close-packed structure. If these ions are above C, it is a cubic close- packed stucture. Ions that define octahedral and tetrahedral holes are also shown.

y-Fe2O3 is also a spinel. As suggested by the chemical formula, the compound has fewer cations than needed to complete an ideal spinel structure. Indeed, the formula is sometimes written as Fe3(Fe5@)OI2 to represent the fact that for every twelve oxygen ions (which provide twelve octahedral holes and six tetrahedral holes), there are three FeIII ions in the tetrahedral holes, and five FeIII ions in the octahedral holes. Compared to the idcal spinel structurc, the occupancy of the tetrahedral holes is the same, but the occupancy of the octahedral holes is 1/6 less. This sixth position is denoted by @ to represent a cation vacancy. Thus there is one cation vacancy for every three spinel units. When this vacancy is ordered, the repeating unit is a trispinel.

There is another way to view the spinel and the corundum structures. Instead of constructing the solid with close-packed layers of oxygen ions, the same structure can be formed using the octahedral units as building blocks. At the comers of the octahedral units are the oxygen ions, and at the centers are the cations. Of course, these cation positions do not have to be completely filled. The corundum structure is then made up of a three dimensional net of such octahedra in which some octahedra share comers, edges, or faces, whereas the spinel structure is made up of octahedra that share comers and edges. Although not apparent in these cases, it is sometimes more convenient to discuss structures with networks of octahedra.

The rutile structure is one such case. In this structure, the sheets of close- packed oxygen ions are rather distorted. The cations are in the center of octahedra of oxygen ions, as shown in Fig. 2-3a. Along the c-direction (the vertical direction), the octahedra are linked by sharing edges (Fig. 2-3b) to form a chain. Adjacent chains are connected by sharing vertices (Fig. 2-3c). For Ti02, the octahedra are distorted such that four metal-oxygen distances are of one value, and the other two are of a different value. For some others like the dioxides of V, Nb, Mo, and W, the metal-metal distances along the octahedra chain are not regular, but alternate between a longer and a shorter distance.

Molybdenum trioxide M a 3 has a unique layer structure made up of chains of octahedra that share comers. Two such chains are connected by sharing two edges of the octahedra to form a double chain (Fig. 2-4a). These double chains are then connected together in the third dimension (perpendicular to the plane of the double chain) by sharing comers to form a sheet-like structure. Thus for each octahedron, three 0 atoms are shared by three octahedra of the same double chain, two are shared by two octahedra of adjacent double chains, and one is unshared. This unshared unit is commonly referred to as a Mo=O unit. Finally, these two- dimensional sheets are stacked on top of each other with rather weak interaction between layers (Fig. 2-4b). Vanadium pentoxide V 2 0 5 also has a layer structure. The basic units are chains of tetrahedra linked through two comers. Two chains are then linked by placing a fifth oxygen ion from one chain to each V ion in the other chain to form a double chain and the basic structure of metavanadates. In this manner, each vanadium ion is regarded to have five oxygen neighbors. There are two V-0 bonds that are shorter where the oxygen is not shared. When one of these oxygen is shared bctween double chains so that they are linked, the layer structure of V 2 0 5 is formed. The oxygen ions in this structure are

OXIDE STRUCTURE 11

0 0 o T i

Figure 2-3 a. A unit cell of rutile TiO,; b. Two edge-sharing octahedra of Ti06 units from adjacent unit cells; c. A network of octahedra that makes up TiOz.

Figure 2-4 Moo3 structure. a. A double-chain unit; b. Cross-section of the layer structure of sheets of double-chains. The Octahedra are shown as squares with diagonals.

12

considered to be attached to one, two, or three vanadium ions. The one attached to one V is the shortest, and is referred to as V=O. The vanadium ion is considered to be in a distorted octahedron in which the sixth oxygen is from another layer, and is very far from the vanadium ion.

The last simple structure to be discussed is the perovskite structure. Compounds of this structure are usually of the formula M1MIIO3. An ideal perovskite structure is made up of a cubic net of corner-sharing octahedra (Fig. 2-5). The smaller and more highly charged cation, MII sits in the center of an octahedron, and the larger cation MI sits in the center of the cavity defined by a cube of eight octahedra. Thus this latter cation is coordinated to twelve oxygen ions. Usually, the MI1 ions are the transition metal ions, and the MI ions are the alkali, alkali earth or lanthanide ions. Some examples of perovskites are KTa03, SrTi03, and LaCoO3. Some compounds have ideal perovskite structures, but many others, especially those that have large cations, are often distorted.

In addition to those described, there exist many other structures, such as scheelite, pyrochlore, and wolframite. The variety increases with the number of components in the compound. The readers are referred to the text by Wells [ l] for discussions of these structures.

2.2 SHEAR STRUCTURE IN INTERMEDIATE OXIDES OF Ti, V, AND Mo

Many transition metal ions possess multiple stable oxidation states. This is evident from the compounds listed in Table 2-1 where a number of oxides of different stoichiometries are found for many elements. Among these elements, four have unusually large number of stoichiometries. They are Ti, V, Mo and W. Table 2-2 lists the known oxides of Ti, V, and Mo that have well-defined three- dimensional crystal structures. It is evident that many of these oxides differ only slightly from each other. In fact, their structures are very similar and can be constructed from the same building blocks. This contributes to the easy conversion of one oxide to another of adjacent stoichiometry by oxidation or reduction. The easy oxidation and reduction, and the existence of cations of different oxidation states in the intermediate oxides have been thought to be important factors for these oxides to possess desirable properties in selective oxidation catalysis.

As an example, the structures of M o ~ ~ ~ ~ and M 0 ~ 0 ~ ~ h a v e the common building blocks of a three-dimensional cubic network of comer sharing octahedra (the Re03 structure, which is also the building block for perovskite). The structures of these compounds are formed when the positions of some octahedra are shifted that pairs of octahedra share edges instead of comers (Fig. 2-6). The shift in the positions are regular along certain directions called shear planes. The farther is the stoichiometry from Moo3, the higher is the density of shear planes.

OXIDE STRUCTURE

Table 2-2 Intermediate Oxides of Ti, V, and Mo

13

Ti30 Ti20 Ti0 Ti203 Ti305 Ti407 Ti509 Ti02

Figure 2-5 A unit cell of perovskite MIMIIOJ. The MII ion is in the center of a comer- sharing octahedron.

14

Figure 2-6 Shear structure in molybdenum oxides. a. Re03 structure; b. MogOZ6; c. MoSOZ3. Each Moo6 octahedron is shown as a half-filled square block.

OXIDE STRUCTURE 15

2.3 STRUCTURE, STABILITY AND RECONSTRUCTION OF OXIDE SURFACES

The positions of the surface ions may or may not be the same as those defined by simple extension of the bulk structure, depending on whether the free surface reconstructs or not. The driving force for reconstruction is to lower the surface Gibbs energy per unit surface area to attain a thermodynamically more stable system. However, metastable surface structure can exist if the energy barrier for reconstruction is too high.

At present, there are few reliable experimental values for surface Gibbs energies of oxides. Reliable calculated surface energies are also difficult to find. The uncertainties in the true ionic charge, the degree of covalency, and other surface properties such as compressibilities have made it difficult to perform accurate calculations. Therefore, information on the thermal stability of a certain surface plane comes entirely from experimental data.

Table 2-3 lists the surfaces of transition metal oxides that have been studied. A surface is regarded stable (in some cases metastable), if it has been reported to yield a 1x1 LEED pattern at low temperature. A thermodynamically stable surface structure will not reconstruct on heating, provided that the temperature is not too high, while a metastable surface would usually undergo reconstruction. In the absence of knowledge of surface Gibbs energy, it is useful to develop qualitative guidelines as a first approximation to compare the stabilities of different surface structures.

Correlations with two effects appear reasonable. One is the polarity of the surface, and the other is the degree of coordinative unsaturation of a surface cation.

When a crystal of a binary oxide is cleaved to generate two new surfaces, the ions in the cleavage plane are partitioned into the two separating solids in such a manner that charge neutrality is maintained in each solid. The structure of the two newly created surfaces, however, may or may not be identical. If they are identical, or if the surface plane contains a stoichiomeuic ratio of cations and anions, the surface will be dipoleless and it is called a nonpolar surface. If they are not, the surface will probably (but not necessarily) possess a strong dipole and the surface is a polar surface. A schematic representation of the two types of surfaces is shown in Scheme 1.

Examples of nonpolar surfaces are rocksalt (100) surface, rutile (110), (loo), and (001) surfaces, pervoskite (100) surface, and corundum (047) surface. Examples of polar surfaces are wurtzite (0001) and (0001) and rocksalt (111) surfaces. With other factors being identical, a polar surface is less stable than a nonpolar surface. The presence of a dipole moment increases the surface Gibbs energy. Comparing a metal-polar and an oxygen- polar surface, the latter is usually more stable because oxygen ions are more polarizable than metal ions. Polarization lowers the surface electric field and the surface energy.

The degree of coordinative unsaturation of a surface cation measures the number of bonds involving the cation that have to be broken to form a surface.

16

Table 2-3 Stability of Oxide Surfaces

Stabil- Cau 'on coordination Polar or Surface itya bulk surface nonpolar Ref. Comment

NiO (100) + 6 5 c o o (100) + 6 5 MnO (100) + 6 5 TiO, (110) + 6 6 and 5

(100) +/- 6 5 (001) +/- 6 4

SrTiO, (100) + 6 5 (111) +/- 6 3

T$O,(047) + 6 5 vp , (047) + 6 5 v,o, (010) + 6 6 and 5 a-Fe,O, (001) - 6 3 BaTiO (001) - 6 5

(1120) + 4 3 (4041,5031) + 4 3 (Oool) +/- 4 3 (mi) +I- 4 4

MOO, (010) + 6 6 wo, (100) + 6 6

ZnO (13oio) + 4 3

n n n n n n n P

P n n n n P P n n

1-3 3 3

4-6 5

7 3 22 9 10 10 11 12 13

14-17 16

17,18 18,19 19 20 21

Footnotes: a ) -, +/-, and + represent structures of increasing stability to thermal treatment. b)reconstruct to (2x2) and (6x6) structures at high temp. (ref. 23). C)reconstruct to (1x3). (1x5), and (1x7) structures on heating. d, facet to (011) and then to (114) which is stepped (001) on heating. e, reconstruct to (2x2). (6 x &)R30" and splitting of spots. f , reconstruct to (6 x 6. g , 6-fold "1x1" at low temp., spot-splitting at 300-400°C. (2x2).

h, 6-fold "1x1". ' ) distorted "1x1".

(5x1), and (6 x a R 3 0 " above 700°C.

References: 1. F.P. Netzer and M. Prutton, J. Phys. C., 8, 240 (1975). 2. M.R. Welton-Cook, and M. Prutton, J. Phys. C., 13, 3993 (1980). 3. M. Prutton, J.A. Walker, M.R. Welton-Cook, R.C. Felton, and T.A. Ramsey,

Surface Sci., 89, 95 (1979).

OXIDE STRUCTURE 17

Table 2-3 continued 4. R.H. Tait and R.V. Kasowski, Phys. Rev. B. 20, 5178 (1979). 5. Y.W. Chung, W.J. Lo, and G.A. Somorjai, Surface Sci., 64, 588 (1977). 6. V.E. Henrich, G. Dresselhaus, and H.J. Zeiger, Phys. Rev. Lett., 36, 1335 (1976). 7. V.E. Henrich. and R.L. Kurtz, Phys. Rev. B, 23, 6280 (1981). 8. L.E. Firment, Surface Sci., 116. 205 (1982). 9. W.J. Lo, and G.A. Sornorjai, Phys. Rev. B, 17,4942 (1978). 10. R.L. Kurtz, and V.E. Henrich, Phys. Rev. B. 25. 3563 (1982); 26, 6682 (1982). 11. L. Fiermans, and J. Vennik. Surface Sci., 18, 317 (1969); 9, 187 (1968). 12. R.L. Kurtz, and V.E. Henrich, Surface Sci., 129, 345 (1983). 13. A. Aberdam, and C. Gaub, Surface Sci.. 27, 571 (1971); 27, 559 (1971). 14. M.F. Chung, and H.E. Famsworth, Surface Sci., 22. 93 (1970). 15. J.D. Levine, A. Willis. W.R. Bottoms, and P. Mark, Surface Sci., 29, 944 (1972). 16. S.C. Chang. and P. Mark, Surface Sci.. 45, 721 (1976). 17. W.H. Cheng, and H.H. Kung, Surface Sci., 122, 21 (1982). 18. W.H. Cheng. and H.H. Kung, Surface Sci., 102. L21 (1981). 19. S.C. Chang. and P. Mark, Surface Sci.. 46, 293 (1974). 20. L.E. Firment, and A. Ferretti. Surface Sci., 129. 155 (1983). 21. M.A. Langell, and S.L. Bernasek, J. Vac. Sci. Technol., 17, 1287, 1296 (1980). 22. V.E. Henrich, G. Dresselhaus, and H.J. Zeiger, Phys. Rev. B, 17,4908 (1978). 23. M. Langell. N. Cameron, Surface Sci., 185, 105 (1987).

I l l l l - M - O - M - 0 - M -

1 " l l I I - 0 - M - O - M - O - \o/ M M \o' \o'

t 0-polar surface M-polar surface

f

L 1

nonpolar surfaces t cleavage +

Nonpolar surfaces Mewl-polar (lower) and Ox ygen-polar (upper) Surfaces

Scheme 1

18

The degree of coordinative unsaturation of a surface anion can be similarly defined. Qualitatively the smaller the number of bonds broken, the more stable is the surface. Thus with all other factors being identical, a surface is more stable if its surface cations and anions are more coordinatively saturated. This factor, together with the surface dipolar factor, qualitatively explains the trends observed in Table 2-3.

When a metastable surface acquires enough thermal energy to overcome the barriers for atomic migration, reconstruction by rearrangement of surface atoms takes place to lower the areal surface Gibbs energy. The structure of a reconstructed surface is no longer a simple extension of the bulk. At present, only a few structures of reconstructed surfaces have been analyzed. One example is the TiOz (001) surface. Upon heating, this surface facets. Analysis of the LEED pattern shows that the faceted surface is a stepped surface which is composed of (011) facets at low temperatures, and (114) faccts at high temperatures [2]. Other examples are the ZnO polar surfaces. A LEED pattern showing a "1x1" pattern of six-fold symmetry is often obtained for the ZnO (0001) and (0001) surfaces, although the truncated bulk structure possesses only a three-fold symmetry. Two possible explanations for the six-fold symmetry have been proposed. One explanation comes from the LEED intensity analysis of the Zn-polar surface [3]. It has been found that the outermost Zn layer relaxes inward by about 0.02 nm (see Table 2-4). This brings the layer of Zn closer to the layer of 0 atoms, and the surface behaves like a hexagonal close-packed layer of atoms in the LEED analysis. The other explanation is that double layer steps are present on these surfaces. Since adjacent double layers are rotated by 60°, a random occurrence of these steps would result in a six-fold LEED pattern [4].

It should be emphasized that the stability of these surfaces has been studied under ultra high vacuum conditions such that the surfaces are presumably clean. The environment during catalytic reactions or during oxide powder preparation is either that of an aqueous solution or of gases of atmospheric pressure or higher. Adsorption of molccules can substantially change the stability of a certain surface structure. Adsorption often increases the stability of an otherwise unstable surface.

In vacuo, some surfaces lose oxygen atoms to lower the Gibbs energy of the system. If the oxygen atoms are lost such that the anion vacancies form a periodic structure, a superstructure is observed in the LEED pattern, This phenomenon has been used to explain the data for a TiOZ (100) surface [5]. Upon heating this surface to increasingly higher temperature, (1x3), (1x5) and (1x7) patterns arc formed consecutively. Simultaneous AES data indicate increasing loss of surface oxygen.

In addition to reconstruction, small relaxations of the surface atoms from their truncated bulk positions also occur. The bonds of covalent compounds arc not very compressible, and it would be expected that the anion-cation bond lengths would remain approximately constant during relaxation [6,7], as is observed for GaAs (110) [8]. The extent of relaxation also depends on the presence of electron

This is accompanied by reduction of the cations.

OXIDE STRUCTURE 19

Table 2-4 Relaxation of Surface Atoms of Oxides Deduced from LEED Data

Relaxationa Rumpling Ref. Surface 6M 60 %s,dJ s, -60

NiO (100) 0 to -3% - 5 t o 5 % a 0 to -3% Oto3% b

coo (100) small b ZnO (lOi0) 4 . 3 f 0.11 4 . 1 +_ 0.051 C ZnO (0001) 4 . 2 1 d

0 0.1 to 0.21 e coo (111) -15% f

Footnote: a) negative relaxation is towards the solid.

References: a) C.G. Kinniburgh, and J.A. Walker, Surface Sci., 63, 274 (1977). b) M. Prutton, J.A. Walker, M.R. Welton-Cook, R.C. Felfon, J.A. Ramsey.

c) C.B. Duke, A.R. Lubinsky, S.C. Chang, B.W. Lee, and P. Mark, Phys.

d) C.B. Duke, A.R. Lubinsky, Surface Sci.. 50, 605 (1975). e) R.E. Watson. M.L. Perlman, and J.W. Davenport, Surface Sci.. 115, 117

f ) A. Ignatiev, B. Lee, and M. Van Hove, Proc. 7th Intern. Vacuum

Surface Sci., 89, 95 (1979).

Rev. B, 15, 4865 (1977).

(1982).

Congr.. Vienna. R. Dobrozemsky, et al. ed., 1977, p. 1733.

lone pairs on the surface [9,10]. For example, the oxygen atoms at the (OOOT) oxygen-polar surface of ZnO have two electrons in the lone-pair orbital pointing outward from the surface, while the Zn atoms in the (0001) Zn polar surface do not have such filled orbitals. Surface atom relaxation is expected to be different for the oxygen and for the Zn atoms, assuming that ZnO is a covalent compound.

For ionic compounds, relaxation of a surface ion is determined by three factors [ l l ] : changes in h e ionic size because of reduced surface charge transfer resulting from a reduced Madelung potential at the surface, imbalance of the ionic forces because of the termination of the lattice, and the residual influence of covalent bonding. The extent and the direction of relaxation depend on the relative contributions of these factors. Since the restriction on bond directions and bond lengths in an ionic compound is much less severe than in a covalent coompound, it is possible that both the metal ion and the oxygen ion relax inward. This difference between covalent and ionic

20

surface

bulk

ionic covalent

SIDE VIEW

Figure 2-7 Side view of a solid showing relaxation of surface atoms in an ionic and a covalent compound. The dotted circles represent truncated surface lattice positions.

relaxation is schematically illustrated in Fig. 2-7 for the surface atoms of the zinc oxide (1010) plane [l 11.

Available data on the relaxation of surface atoms of transition metal oxides are summarized in Table 2-4. For these oxides, the relaxation is of the ionic type, and both the metal and the oxgyen atoms relax inward. These results are not too surprising as these oxides are of the first transition period. They are relatively ionic. Since the ionicity decreases on going to the lower right hand comer of the periodic table, one would expect the mode of relaxation of the surface atoms of oxides of F't, Pd, Cu, etc. to be different.

It should be emphasized that substantial amounts of defects may exist in surfaces of oxide crystallite powders prepared by conventional techniques. It has been reported that steps of different heights exist on surfaces of Moo3 crystallites. These surfaces also reconstruct readily upon thermal heating or irradiation by electron beams [12]. In fact, single crystal surfaces may also have well-ordered domains coexisting with disordered regions [ 131.

2.4 STRUCTURE OF SUPPORTED OXIDES

Supported oxides that do not interact strongly with the support form three- dimensional crystallites whose properties are similar to large bulk crystals. However, some oxides interact strongly with oxidic supports such that a monolayer of an oxide of properties different from the bulk oxide is formed. Growth of three- dimensional crystallites occurs only after a substantial fraction of the surface is covered by the monolayer. This is often the case for oxides of Cr, Mo, W, V, and

OXIDE STRUCTURE 21

Re supported on alumina, and for vanadium oxide supported on anatase titania. The absence of three-dimensional crystallites and the presence of monolayer

structure are readily observed by Raman spectroscopy. Figure 2-8 shows the Raman spectra of vanadia. The spectrum for V205 (curve a) shows intense sharp peaks at around 996, 703, 530,483, and 406 cm-' [14]. The 996 cm-' band is assigned to terminal V=O stretch, and the 703 and 530 cm-' bands are assigned to bridging VC&V symmetric and antisymmetric stretch, respectively. When 2.1 wt.% of vanadia is supported on y-AI2O3 (curve b), the spectrum does not show any of the prominent V2O5 peaks. Instead, two featureless broad peaks appear; one centers around 970 cm-' accompanied by a shoulder around 995-1000 cm-', and the other in the range between 800-830 cm-' [15]. When a larger amount of vanadia is supported on the alumina (curve c), the intense peaks characteristic of crystalline V2O5 appear. Interestingly, these characteristic peaks begin to appear at lower vanadia loadings on a silica support than on alumina [16]. This suggests that vanadia interacts more strongly with alumina than with silica so that it forms a monolayer structure with alumina but not with silica. Vanadia also interacts strongly with anatase Ti02 to form a layer structure [17].

Vanadia supported on anatase Ti02 is a much studied system. It is generally believed that the vanadia layer structure exposes the (010) plane and the V=O groups preferentially [ 18-20] because of the excellent match of the vanadate unit and the anatase structure. It has been proposed that on the (001) plane of anatase, vanadium exists as monovanadate (V04)"- groups bonded to the surface as [21]:

\ l O l O l O l I \ / \ / \ / Ti Ti Ti Ti

In this form, the V ion occupies roughly a position which another Ti ion would occupy if the bulk structure continues. The V-0 bridging bond length of 0.190 nrn is very close to the Ti-0 bond length of 0.193 nm. Indeed, EXAFS studies show that on anatase, the basic structural unit of vanadium contains two terminal V=O bonds of a bond length of 0.165 nm, and two bridging V - 0 bonds of 0.190 nm [221. The Dicture is also consistent with the observation that an average of one oxygen ion per vanadium can be readily removed on reduction [21]. However, this model has not been supported by IK data which do not show the expcctcd band- splitting due to coupling of the two V=O groups [23]. On other crystallographic planes of anatase, however, there are no sites of such a good match with the monovanadate groups, and other structures may be present.

On y-AI2O3, EXAFS and XANES studies show that vanadium ions are present in a more regular tetrahedral coordination than on anatase. Each vanadium ion is associated with two terminal V=O bonds of 0.167 nm, and two bridging bonds of 0.182 nrn bond length [21,22]. Upon reduction, each vanadium

22

1500 1000 5bO

Wavenumber cm-l

Figure 2-8 Raman spectra of supported and unsupported vanadia. a. V205; b. 2.1 wt.% V/y-A1203; c. 4.0 wt.% Vfi- A1203; d. 2.3 wt.% V/SiO2. Curves a-c are from Z i t . Phys. Chern. Neue Folge, 111, 215 (1978), curve d is from J. Phys. Chem., 84, 2783 (1980).

ion loses 0.6 oxygen atoms readily on the average. A picture that is consistent with these observations is that the vanadium ions exist as dimcric units on y-AI2O3, such as pyrovanadate units [21]:

Like vanadia supported on titania or A1203, molybdena on thcse supports also cxists as well-disperscd units or monolaycr until high loadings. This is illustratcd by Raman spectroscopy of MoO3/AI203 [24] and M003/ri02 1251. As shown in Fig. 2-9, the intense and sharp characteristic bands of crystalline Moo3 at 998

OXIDE STRUCTURE 23

I I I I 1

800 1000 1200

Wavenumber cm-’

Figure 2-9 Raman spectra of supported molybdena. a. 12 wt.% MoO3/q-AI2O3; b. 4 wt.% Mo03/A1203; c. 13.5 wt.% Mo03/Ti02; d. 1.8 wt.% MoO-,/riOZ. Curves a and b are from J. Phys. Chern., 82, 2002 (1978), and curves c and d are from J. Catal., 94, 108 (1985).

and 821 cm-’ are not present in samples of low loadings of Moo3. They appcar only in samples of high loadings. A monolayer structure is formed when the Mo loading is below about 5 Mo atoms per nm2. Below this value, the ratio of the intensity of the Mo signal to the A1 signal in XPS and ISS increases linearly with the Mo loading, consistent with the picture of a well-dispersed phase [26,27]. At higher loadings, this ratio reaches a plateau. Electron microscopy also shows a highly dispersed phase of molybdena at low loadings [28,29]. Interestingly, crystallites of MoS2 on alumina can be redispersed by reoxidation to form a dispersed phase of molybdena.

One picture that has been proposed to explain these and other evidence is that the molybdenum ions exist as monomeric MOO^^- units at loadings up to about 1 Mo per nm2. Betwecn 1 to about 4.5 Mo per nm2, heptameric Mq0246p units and octahedrally coordinated polymeric surface species are also present [30-341. The heptamer is proposed to be adsorbed as a bilayer with four Mo ions lying next to the AI2O3 and the remaining three on top [32]. At still higher loadings, crystalline Moo3 appears.

24

0 500 1000 -1

Wavenumber cm

Figure 2-10 Raman spectra of tungsten oxides. a. WO,; b. W03/A1203. (From J. Catal.. 90. 150 (1984). copyright Academic Press.)

A12(W04),; c. 10 wt.%

The existence of these well-dispersed or monolayer phases is actually quite common among supported transition metal oxides. Fig. 2-10 shows another example of WO3 supported on alumina [35,361. Commonly found in these well- dispersed phases are the terminal metal-oxygen double bonds. These M=O bonds usually have characteristic Raman stretching frequencies in the 950-1000 cm-' region. These peaks are often at somewhat different frequencies and broader than the corresponding peaks for the crystalline oxides [371. Indeed such M=O groups have been detected for W6+, Mo6+, V5+, Re7+, and Cr6+ oxides on y-AI2O3 [37,38], and Mo6+ and V5+ on titania [39,401. The frequencies of these terminal double bonds vary somewhat depending on the surface coverage of the dispersed phase and on whether the double bond is interacting with adsorbed water or surface hydroxyl groups [41 and references thercin].

Nickel oxide also forms a well-dispersed phase on y-A1203 instead of crystallites of NiO. This well-dispersed two-dimensional phax may actually be Ni occupying the octahedral and/or the tetrahedral holes at the surface of the alumina to form a layer of spinel-like NiA1204 [26,27,42].

When Ni is supported together with Mo on y-AI2O3, there is interaction between Ni and Mo which affects their locations on the support. It has been

OXIDE STRUCTURE 25

suggested that at low calcination temperatures, the Ni ions interact strongly with the molybdenum species to partially shield the oxomolybdenum ion. At high calcination temperatures, the Ni ion migrates to the tetrahedral or octahedral holes of the alumina, exposing the molybdenum ions [27,43]. Similar incorporation into the alumina has also been observed in Ni-W/AI2O3 [44] and for Co in the Co- Mo/A1203 system [451.

REFERENCES

1. A. F. Wells, "Structural Inorganic Chemistry", 4th ed., Clarendon Press,

2. L. E. Firment, Surface Sci., 116, 205 (1982). 3. C. B. Duke, and A. R. Lubinsky, Surface Sci., 50, 605 (1975). 4. V. E. Henrich, H.J.Ziegler. E. I. Solomon, and R. R. Ray, Surface Sci..

5. Y. W. Chung. W. Lo, and G. A. Somorjai. Surface Sci., 65, 419 (1977). 6. J. D. Levine, and S. Freeman, Phys. Rev. B., 2, 3255 (1970). 7. J. E. Rowe, S . B. Chrisman, and G. Margaritondo, Phys. Rev. k f t . 35, 1471 (1975). 8. A. R. Lubinsky, C. B. Duke, B. W. Lee, and P. Mark, Phys. Rev. Lett.

9. H. C. Gatos, and M. C. Levine, J . Electrochem. Soc., 107, 427 (1960).

London, 1975.

74, 682 (1978).

36. 1058 (1976).

10. H. C. Gatos, J . Appl. Phys., 32, 1232 (1961). 11. C. B. Duke, A. R. Lubinsky, B. W. Lee, and P. Mark, J . Vac. Sci. Technol.,

12. J. M. Dominguez-Esquivel, 0. Guzman-Mandujano, and A. Garcia-Borquez,

13. L. E. Firment, and A. Ferretti. Surface Sci., 129, 155 (1983). 14. I. Beattie, and T. Gibson, J . Chem. SOC. A, 2322 (1969). 15. F. Roozebaum, J. Medema, and P. Gellings, &it. Physik. Chem. Neue Folge,

16. F. Roozebaum, M. Mittelmeijer-Hazeleger, J. Moulijn, J. Medema, V. de Beer,

17. I. Wachs, R. Saleh, S . Chan, and C. Chersich, Appl. Catal., 15, 339 (1985). 18. A. Vijux, and P. Courtine, J . Solid Stare Chem., 23, 93 (1978). 19. G. C. Bond and P. Konig, J . Catal., 77, 309 (1982). 20. Y. Murakami, M. Inomata, K. Mori, T. Ui, K. Suzuki. A. Miyamoto, and T.

13, 761 (1976).

J . Catal., 103, 200 (1987).

111, 215 (1978).

and P. Gellings, J . Phys. Chem., 84, 2783 (1980).

Hattori, in Preparation of Catalysts III. G. Poncelet, P. Grange, and P. Jacobs ed., Elsevier Science Publ., Amsterdam, 1983, p. 531.

21. J. Haber, A. Kozlowska, and R. Kozlowski, J . Catal., 102, 52 (1986). 22. R. Kozlowski, R. Pettifer. and J. M. Thomas, J . Phys. Chem., 87, 5176 (1983). 23. G. Busca, G. Centi, L. Marchetti, and F. Trifio, Langmuir. 2. 568 (1986). 24. H. Knozinger. and H. Jeziorowski, J . Phys. Chem., 82, 2002 (1978). 25. Y. Liu, G. Griffin, S . Chan, and I. Wachs, J. Card., 94, 108 (1985). 26. P. Dufresne, E. Payen, J. Grimblot, and J. P. Bonnelle, J. Phys. Chem.,

85, 2344 (1981).

26

27. S. Kasztelan, J. Grimblot, and J. P. Bonnelle, J . Phys. Chem., 91, 1503 (1987). 28. T. Hayden, and J. Dumesic, J . Catal.. 103. 366 (1987). 29. T. Hayden, J. Dumesic. R. Sherwood, and R. Baker, J . Catal., 105, 299 (1987). 30. S. Kasztelan, J. Grimblot. J. P. Bonnelle, E. Payen, H. Toulhoat, and Y. Jacquin,

31. L. Wang, and W. K. Hall, J . Cafal.. 77. 232 (1982). 32. W. K. Hall, Proc. 4th Conf. Chemistry and Uses of Mo, H. F. Barry, and

33. H. Weigold, J. Catal., 83, 85 (1983). 34. N. Giordano, J. C. J. Bart, A. Castellan, and G. Martinotti, J . Cafal., 36. 81 (1975). 35. S. Chan, I. Wachs. and L. Murrell, J . Catal., 90, 150 (1984). 36. S. Chan, I. Wachs. L. Murrell. L. Wang, and W. Hall, J . Catal., 88, 5831 (1984). 37. I. E. Wachs. F. D. Hardcastle, and S . S . Chan, Spectroscopy, 1. 30 (1986). 38. S. S. Chan, and I. E. Wachs, J . Catal., 103, 224 (1987). 39. C. P. Cheng and G. L. Schrader, J . Cafal., 60, 276 (1979). 40. K. Y. S. Ng, and E. Gulari, J . Catal., 92, 340 (1985). 41. R. Quincy. M. Houalla, and D. Hercules, J . Catal., 106, 85 (1987). 42. S . Kasztelan, J. Grimblot, and J. P. Bonnelle. J . Chim. Phys., 80, 793 (1983). 43. H. Jeziorowski, H. Knozinger, E. Taglauer, and C. Vogdt, J . Cafal., 80, 286 (1983). 44. B. Horrell. D. L. Cocke, G. Sparrow, and J. Murray, J . Cafal., 95, 309 (1985). 45. F. Delannay, E. Haeussler. and B. Delmon, J . Cafal., 66, 469 (1980).

Appl. Cafal., 7 , 91 (1983).

P. C. H. Mitchell, ed., 1982, p.224.

Chapter 3

PHYSICAL AND ELECTRONIC

PROPERTIES

3.1 SURFACE COMPOSITION

The surface composition of a single component oxide is determined by the surface anion to cation ratio, which, for an ideal surface, depends on the stoichiometry of the oxide and the orientation of the exposed crystal plane. It is often important to determine whether or not a surface is stoichiometric. Nonstoichiometry often arises from preferential removal of surface oxygen leading to a slight reduction of the surface. The extent of nonstoichiometry depends on the pretreatment of the sample. Stoichiometric surfaces can often be obtained for surfaces that have low surface areal Gibbs energies (the stable surfaces) by low temperature annealing. For example, stoichiometric SrTi03 (loo), TiOz (1 lo), TiOz (100) [l-31, ZnO (1070) [4], and Moo3 (010) [5] surfaces have been prepared. High temperature annealing in vacuo or ion-sputtering preferentially removes oxygen, and the cations near the surface are reduced to lower oxidation states. For oxides that have empty d bands, reduction of the surface to generate nonstoichiornetry rcsults in the appearance of band-gap states which are readily detectable by ultraviolet photoelectron spectroscopy and electron energy loss spectroscopy .

For multicomponent oxides, in addition to the surface anion to cation ratio, the ratio of the different cations is also of interest. At least two questions are asked. If the two component oxides form a solid solution, are the surface and the bulk cation ratios the same? If the component oxides form a bulk compound, does the surface have the same chemical stoichiometry as the bulk?

Whether the surfacc cation ratio is the same as the bulk ratio depends on a number of factors. Thesc factors include: (1) the surface tension (or surface Gibbs energy) of the component oxides. The lower energy component tends to be segregated to the surface; (2) thc bulk strain of the solid solution due to mismatch of the ionic sizes or the coordination symmetry. The larger ion tends to be segregated to the surface, as docs the ion whose preferred coordination symmetry

27

28

differs from that provided by the matrix; (3) the nature of the adsorbate. Chemisorption lowers the surface energy of the solid. Thus chemisorption tends to induce surface segregation of the component that binds more strongly with the adsorbate; (4) formation of a surface compound. Even if the bulk oxide is a true solid solution, a surface compound of a certain stoichiometry may be formed. The surface composition is then determined by the surface compound. One driving force for the formation of a surface compound is the presence of surface adsorbates that causes the surface cations to have an oxidation state different from the bulk oxidation state. For example, in the presence of oxygen, the surface chromium ions in Cr203 have an oxidation state of +6, whereas the bulk ions are +3. In this case, the adsorbed oxygen are being incorporated as surface lattice oxygen ions.

There are only a few experimental results on solid solutions. In the dilute solution of ZnO in MgO, the surface is enriched with Zn [6]. This can be explained by point (2) in that Zn prefers tetrahedral coordination, while the coordination symmetry in MgO is octahedral. In the solid solutions containing Cr, which includes CoO-Cr203, NiO-Cr203 [7], and Fq03-Cr203 [8], Cr is enriched on the surface in the region of low Cr concentrations. The surface cation ratios seem to remain constant at intermediate Cr concentrations. This phenomenon has been interpreted as due to the formation of a compound that is segregated onto the surface, such as CoCrZO4 spinel [7]. The surface enrichment of Sb in the Sn02- Sb02 solution [9] can be explained by the larger size of Sb than Sn [lo].

In an oxidic multicomponent compound, the surface composition is constrained by the periodicity of the compound, unless the driving force for surface segregation is sufficiently large to disrupt the periodicity. In the limited number of cases reported, the surface has the same cation ratio as the bulk. For example, the surface Bi/Mo ratios for Bi2Mo06, Bi2M0209 and Bi2Mo-,OI2 are within 10% of the bulk values [ll].

3.2 IONICITY OF OXIDES

The formal oxidation states of transition metal oxides range from one in Cu20 to eight in Os04, for example. However, there is little doubt that the true ionic charges in many transition metal oxides are less than those predicted from formal oxidation states. Since ionicity affects oxide properties including the surface electric field gradient, the surface electrostatic potential, and the mode of surface relaxation, knowledge of its magnitude is important.

Recently, experimental measurements of the degree of charge transfer between the cation and the anion become available from X-ray diffraction data [ 12- 151, and X-ray photoelectron spectroscopy data [ 16-18]. Some theoretical calculations are also available that provide charge lransfcr information. These data are summarized in Table 3-1.

In the X-ray diffraction technique, electron dcnsity distribctions are deduced directly from X-ray data. From the distribution maps, boundaries of ions are assumed to correspond to the point of minimum (local minimum) electron density. The number of electrons associated with each ion, and thus the true cation charge,

PHYSICAL AND ELECTRONIC PROPERTIES 29

q, is then obtained by summing the electron density within the boundary. Reliabliliy of the values of q obtained depends on the validity of this method of partitioning of electrons to the various ions. When properly analyzed, the method appears to be adequate for electrostatic potential calculations using a point charge model or a shell model which disregards the nature and spatial orientation of the electron orbitals.

The method using X P S binding energy shift data was originally proposed by Siegbhan et al. [17]. The values of q reported here are obtained using the method of Bagus and Braughton [16]. In this method, it is assumed that the experimental binding energy of an ion in a solid differs from the binding energy of the ion in the gas phase by the lattice self-potential at the ion site and the final state extra-atomic relaxation effect. (The final state extra-atomic relaxation effect is an effect that affects the kinetic energy of the escaping photoelectron due to the relaxation of the electrons in the lattice surrounding the ion in response to the removal of the photoelectron from the ion). Assuming that the final state effects are the same for the cation and the oxygen ion, the following relationship is obtained [16,18]:

BE(0,XPS) - BE(M,XPS) = BE(0,gas ion) - BE(M,gas ion) + O(0) - O(M) (3-1)

The left hand side is the difference in the values of experimental binding energy of the observed electrons from the oxygen and the metal ion. The first two terms on the right hand side represent the difference in the binding energy of the same electrons from the gas phase oxygen and metal ions. The last two terms represent the difference in lattice self-potential of the oxygen and the cation in the solid. Since the electron binding energies of the gas phase ions and the term O(0) - @(M) are unique functions of the true ionic charge, it is in principle possible to calculate q using this equation. In calculating the values of q listed in Table 3-1, the electron binding energy calculated for the gas phase ions [19] have been used. The lattice self-potentials used are listed in Table 3-2.

Assuming that the experimental values of binding energy are accurate, the accuracy of the ion charge thus obtained still depends on the accuracy of the calculated binding encrgy for the gas phase ions, the validity of the assumption that the final state extra-atomic relaxation effects of the oxygen ion and the metal ion cancel out, and the error introduced by using lattice self-potentials of bulk ions in calculations involving surface and near surface ions. With these limitations, the ionic charges thus obtained are at best qualitative and should be used only for qualitative comparison of ionicity among similar solids.

The seriousness in neglecting the final state effect can best be illustrated using Si02 and CdO for which the extra-atomic relaxation effect has been included in the calculation. The charge for Si in Si02 using equation (3-1) is +4. If a relaxation of about 5 eV is included [20], the charge becomes +2. Similarly for CdO, a Cd charge of 0.4-0.8 would be obtained without relaxation compared to a charge of 1.0 with relaxation [21].

A variation of the above method has been used for a number of chromium compounds [22]. In this study, the authors make use of the fact that because of

30

Table 3-1 True Cation Charges of Oxides (From J. Solid State Chem., 52, 191 (1984), copyright Academic Press with update).

Method a Compound X P S b X-ray Calculation

MgO 1.25- 1.75 1.9 (a), 1.85 (b), 1.5 (c) CaO 1.8 (c) SrO 2 (c) BaO 2 (c) Ti0 1.52 (o), 1.35" (0) vo 0-1 (d) 0.66 (1) MnO 1.5 1.4 (a), 1.51 (b) 1.24 (o), 1.03" (0) FeO 1.29 (o), 1.03" (0) c o o 1-1.5 1.2 (a), 1.4 (b) 1.16 (o), 0.78" (0) NiO 1-1.5 0.7 (a), 0.9 (b) 1.2-1.8 (m), 1.9 (n),

1.00 (o), 0.74"(0) CUO 2 ZnO 1.2-1.7, 1.11" (h) CdO 1.05" (h) Ago Ag(1) 0.53" (h)

Ag(II1) 1.87' (h) a-Fe203 2.7 a-Cr2O3 0.15d (g), 2.6

Ti02 2.8 2.19" (p)

cro2 0.38d (g)

crO3 0.46d (g)

wo3 1.5 &2O 0.48" (h)

a-A1203 2-2.5 (j)

S i02 4, 2" (i) 1.0 (k)

zrO2 1.2

Moo3 1.5-2

Mg2Si206 (Mg iron) (MI) 1.84 (b) (M2) 1.79 (b)

LiA1Si206 (Li ion) 0.7 (b)

CaMgSi206 (Mg ion) 1.44 (b)

Mn2Si04 (Mn ion) (Ml) 1.21 (b) (M2) 1.49 (b)

(A1 ion) 2.4 (b)

(Ca ion) 1.39 (b)


Table 3-1 continued

Footnotes: a) Letters in the brackets denote references. b, Unless noted, the charges are calculated using equation 3-1 in text.

The XPS data are from references e and f except when indicated. ') The final state extra-atomic relaxation of the cation has been included

in calculating these values. See text for detail. d, These are taken directly from reference g using a modified form of

equation 3-1. See text for detail. ') These are for clusters that resemble a surface.

References: a) S. Sasaki, K. Fujino. and Y. Takeuchi, Proc. Japan Acad., 55, Ser R, 43 (1979). b) S. Sasaki, K. Fujino, Y. Takeuchi, and R. Sadanaga, Acta Cryst. A36, 904 (1980). c) G. Vidal-Valef J.P. Vidal, and K. Kurki-Suonio, Acta Cryst., A34, 594 (1978). d) M. Morinaga, and J.B. Cohen, Acta Cryst., A32, 387 (1976). e) C.N.R. Rao, D.D. Sarma, S . Vasudevan, and M.S. Hegde, Proc. Royal SOC.

f) D.D. Sarma, and C.N.R. Rao, J. Electron. Spectra. Related Phenom., 20, 25 (1980). g) T. Dickinson, A.F. Povey, and P.M.A. Sherwood, J. Chem. SOC. Faraday Trans.

h) S.W. Gaarenstroom, and N. Winograd, J. Chem. Phys., 67, 3500 (1977). i ) F. Bechstedt, Phys. Stat. Sol. B, 91, 167 (1979). j ) V.I. Nefedov, D. Gati, B.F. Dzhurinskii, W.P. Sergushin. and Ya. V. Salyn,

k) R.F. Steward, M.A. Whitehead, and G . Donnay, h e r . Mineralogist, 65, 324 (1980). 1) V.A. Gubanov, B.G. Kasimov, and E.Z. Kuxmaev, J. Phys. Chem. Solid, 36, 861 (1975). m)A.B. Anderson, Chern. Phys. Lett., 72, 514 (1980). n) P.S. Bagus, and U. Wahlgren, Mol. Phys., 33. 641 (1977). o) C. Satoko, and M. Tsukada, IMS Ann. Review, 1979, p. 33; M. Tsukada,

p) M. Tsukada, C. Satoko. and H. Adachi, J. Phys. SOC. Japan, 48. 200 (1980).

London, A367, 239 (1979).

I, 72, 686 (1976).

Russ. I. Inorg. Chem., 20, 1279 (1975).

H. Adachi, and C. Satoko, Prog. Surface Sci., 14, 113 (1982).

reduction by X-ray, XPS data could be recorded for Cr ion and reduced Cr atom from the same sample. Presumably, since both the Cr ion and the Cr atom are embedded in the same solid matrix, their extra-atomic relaxation should be the same. Then using an equation similar to equation (3-1) except that the difference between two Cr peaks are used instead of between oxygen and chromium, the Cr ion charge is calculated. It can be scen (Table 3-1) that the ionic charge thus obtained is smaller than that obtained using equation (3-1).

Finally, the ionic charge provided by some theoretical calculations are included. These calculations are cluster calculations and the ionic charges are calculated from the Mulliken populations.

The It is apparent that different methods yield different ionic charges.

32

Table 3-2 Lattice Self-potential a t the Cation Sites of Oxides a

Compound Potential, eV Compound Potential, eV

MO MgO CaO S r O BaO MnO FeO c o o NiO

CdO CUO ZnO PbO

M02 Ti02 Sn02 Si02 ca2 zrO2

-23.9 -20.9 -19.5 -18.2 -22.7 -23.3 -23.6 -24.2

-21.4 -24.3 -24.0 -20.5

M2O3 a-Cr2O3 -34.9

v2°3 -33.4

a-Fe203 -34.8 a-A1203 -36.6

Ti203 -33.6 Ga203 -35.0 R h 2 0 3 -34.1 Pb203 Pb(1) -32.8

Pb(2) -28.1 Nd203 -29.0

-29.7

4 4 . 7 M03 Cr03 -58.5 4 2 . 9 M a 3 -58.5 4 6 . 4 to -50.6b wo3 -64.5 -46.2 4 2 . 3

Footnotes: a) Taken from J. Solid State Chem., 52, 191 (1984); original data from J.Q.

Broughtori and P.S. Bagus, J. Electron Spectros. Relat. Phenom., 20, 261 (1980). and W. Van Cool, and A.G. Piken. J. Mater. Sci., 4. 95 (1969).

b, Value depends on the crystal structure.

differences, in addition to experimental variations, may be due to the fact that different methads measure ionic charges that are defined differently. Only in the limit that eleclrons associated with each ion are clearly defined and there is no overlap of electron density from different ions that the different methods could yicld the same charge. In spite of this, however, it is worth noting that if the data are analyzed as exactly as possible with minimal simplification, he ionic charge does not differ significantly among the different methods.

The values shown in Table 3-1 also confirm the fact that, in general, the trend of ionicity parallels that derived from electronegativity arguments. Thus the alkali metal oxides are rather ionic with ionicity increasing with increasing atomic masses as one goes down a given group. Transition metal oxides near the middle of the pcriod (groups V1 to IIB) have true ionic charges that arc about half of the formal


charges. Finally, Si02 which is normally thought of as covalent does possess a cation charge of one.

When the true ionic charge is not available from experimental or theoretical results, relative ionicity can still be estimated using the electronegativity scale of Pauling [23], Sanderson [24] or Phillips [251.

The basic assumption in Pauling's ionicity is that if the heat of formation of a A-B bond, DAB exceeds the arithmetic average of the heats of formation of the homopolar A-A and B-B bonds, DAA and DBB, the extra energy is due to the transfer of electron from the less electronegative to the more electronegative atom in the bond. A scale of elemental electronegativities XA and XB can be defined from the relations:

It is, therefore, ionic in origin.

The proportionality factor has the dimension of energy. XA and XD are dimensionless and increase by 0.5 with valence charge of unity for the first row of the periodic table. Fractional ionic character f,(A,B) is defined by:

to satisfy the conditions that 0 5 f, 5 1, and f,(A,B) = f,(B,A). Phillips [25] has proposed that the energy of a bond contains a homopolar, or

covalent, and a heteropolar, or ionic part, the values of which can be defined accurately in terms of spectroscopic transition energy between bonding and antibonding states. This proposal stems from the observation that the bonding states have lower energy, are centered predominantly on the more electronegative atom, and point towards the nearest-neighbor atoms. The antibonding states are centered predominantly on the more electropositive atom, and point away from the nearest neighbors.

The homopolar energy Eh is assumed to depend only on the bond length or nearest neighbor distance r, and the position in terms of the rows in the Period Table under concern. The average energy gap between bonding and antibonding state, E,, is then the geometric sum of Eh and the average ionic energy gap C:

Eg2 = Eh2 + c' (3-4)

Phillip's ionicity fi is defined as:

The values of E, can be evaluated from the optical properties of the solids such as the valence electron plasma frequencies or the low frequency electronic dielectric constants, and Eh can be extrapolated from values of E, of elemental compounds such as diamond, silicon, germanium and tin using an empirical relation that Eh K r 0 , where r is the interatomic distance, and S an empirical constant with a value of 2.48.

34

Table 3-3 Phillips' Ionicity, f,

MnO 0.887 FeO 0.873 COO 0.858 NiO 0.841

Ge20 0.730

A1203 0.796 Cr2O3 0.777 F%03 0.677 LiNb03 0.825

LiTa03 0.850 (Ta-0)

C U ~ O 0.56

Ti02 0.686

(Nb-0)

Be0 ZnO MgO CdO CaO SrO

0.65.67 0.65-.66 BaO 0.53 0.53 Ge02 0.49 0.51 Sn02

Li20 si02

0.620 0.653 0.839 0.778 0.916 0.928 0.93 1 0.730 0.784 0.57 0.57 0.766 0.570 0.57-.59 0.57-.59

Ge02 0.511 0.53 0.54

A1203 0.796 LiGa03 (Li-0) 0.815 (Ga-0) 0.653

TeOz 0.67 0.63

Footnotes: a ) From B.F. Levine, Phys. Rev. B. 7. 2591 (1973). b , From B.F. Levine, J. Chem. Phys., 59, 1463 (1973). ' ) From F. Gervais. Solid State Commun., 18, 191 (1976); I I and L refer

to the polarization directions.

The ionicities defined in this method are obtained from measuremcnts of the interaction of electrons in the solid with incident photons. Thus they are based on an effective dynamic charge e that differs from the static charge q by [26]:

e = q + R dq/dr (3-6)

whcre R is the equilibrium interatomic distance, and dq/dr is the sensitivity of the static charge transfer to small shifts in the atomic positions. The difference bctwecn e and q arises from the fact that in the optical measurement of E,, the solid interacts with an electromagnetic force which draws atoms from their equilibrium positions. From the viewpoint of electrostatic interaction, the effect of this movement is equivalent to that of placing on an unperturbed lattice an electric dipole of magnitude c and a lenglh equaling the change in the interatomic distance.


Since atoms or ions in a solid strongly interact with each other, their movements are strongly coordinated, giving rise to group vibrations (phonons). Therefore, it is possible to have different values of dynamic charges (Ze) for a given solid depending on the type, direction and frequency of the vibration. The relationship between this dynamic charge and the dielectric constants is given by Gervais [271:

WLo and WTO are frequencies for the longitudinal and transverse mode, p is the reduced mass of the ions, q is the dielectric constant of vacuum, and V is the molar volume. It can be shown that Ze is the actual transverse charge of the ions within the rigid ion model description.

Phillip's original concept has been developed for simple AB crystals. It is later extended by Levine to compounds of other structures and stoichiometries, including oxides, ternary compounds, and compounds containing transition metal cations [28,29]. The ionicity reported for a number of oxides are listed in Table 3-3. By examining a large number of compounds including halides, halogenides (sulfides, tellurides, etc.), binary oxides and ternary oxides, it is found that values of Z/Zo, where Zo is the formal charge, is roughly linearly proportional to exp(f,) [ 271 :

Z/Zo = 0.24 (efi - e) (3-8)

An examination of the static and dynamic ionic charges in Tables 3-1 and 3-3 suggests that the ionic charges of transition metal oxides seldom equal their formal charges. In fact, in most cases, the charge is about half the formal charge. Furthcrmore, the static and dynamic charges do not differ significantly.

What are the consequences to surface chemistry and catalysis that transition metal oxides are ionic? Qualitatively, the ionic character results in the presence of a strong electric field that points outward from the oxide surface. The separation of charges into cations and anions results in a strongly modulated electronic potential on the oxide surface. These effects lead to the common phenomenon of heterolytic dissociative adsorption of molecules. An example is shown for HZ:

(3-9)

Contrary to H adsorbed on metals, the two hydrogen atoms that are dissociatively adsorbed are not equivalent and carry opposite charges. It is likely that they have different reactivities. It should be noted that the charges shown in the equation do not necessarily represent the real charges which are usually unknown.

The ionic character may also increase the sticking probability of polar molecules, such as ammonia, water, alcohols, acids, ethers and amines. When these molecules approach the surface, their dipole moments interact with the electric field at the surface to orient the molecule, thus enhancing the probability of an attractive bonding interaction.

Since coulombic interaction (that is, charge-charge interaction) is a long- range interaction, the surface chemistry of an oxide depends not only on the nature of the cation or anion at the immediate vicinity of a surface adsorbate, but also on the ionicity of the rest of the oxide matrix. The consequence of this effect in dilute solid solutions has been explored. It is concluded that the ionicity of the matrix oxide in a solid solution affects the activation energy of catalytic reactions in which the rate limiting step involves charge transfer between the oxide and the surface intermediate [301. This effect is also important in determining whether new acid sites are formed in oxide solid solutions [31]. This will be discussed in Chapter 5.

3.3 MAGNETIC PROPERTIES OF SMALL OXIDE PARTICLES

In some supported transition metal oxide catalysts, the oxide is present as submicron-size crystallites. They may be so small that detection by X-ray diffraction is difficult. Recent studies on these particles indicate that some of their properties differ from those of the bulk oxides. One such property is the magnetic property which has been used to identify the presence of small crystallites of oxides and to determine their sizes.

The magnetic properties of an ion or atom are determined by the orientation and the number of its electron spins. For transition metal oxides, the individual electrons are so strongly correlated in their motion that the spin of an ion is better characterized by one total spin (atomic spin) than the individual electron spins. The atomic spins of neighboring ions may also be strongly correlated with each other to form a spin sublattice. Depending on the magnitude, the orientation, and the number of spin sublattices, transition metal oxides possess different internal magnetic fields, as well as different responses to applied magnetic field.

The readers are referred to textbooks on the subject, such as that of Cullity [32] or Morrish [33], for a more extensive discussion. Here we shall give a very brief introduction to magnetism, and then proceed to discuss magnetic properties of small oxide particles.

Fig. 3-1 shows the types of magnetism commonly found in transition metal oxides, and the associated spin orientations. The temperature dependence of the magnetic susceptibility x , which is the magnetization per unit applied field, is shown in Fig. 3-2.

Diamagnetism is related to changes in the orbital motion of electrons that occur when atomic systems are placed in a magnetic field. This induced motion of electrons (or currents) is set up in such a direction as to oppose the change in the magnetic flux, and persists as long as the magnetic field is present. The magnetic field produced by the induced current is opposite to the applied field, and the

There are different types of magnetism.

PHYSICAL AND ELECTRONIC PROPERTIES

0

0

0

0 0

0

Diamagnetism

0 P Antiferromagnetism

Figure 3-1 Types of magnetism in oxides. direction.

Diamagnetic

p” “s

Q-

ST Paramagnetism

Ferrimagnetism

37

Direction of arrows indicates the spin

Paramagnetic

Anti ferromagnetic Ferrimagnetic a

Figure 3-2 Temperature dependence of magnetic susceptibility for different types of magnetism, For antifcrromagnetism. the Nee1 temperature is shown. For femmagnetisrn, saturation magnetization 0, and the Curie temperature T, are shown.

,

38

Table 3-4 Room Temperature Magnetic Properties of Some Common Transition Metal Oxides a

Paramagnetic Antiferromagnetic Ferrimagnetic Diamagnetic

Ti203 cr203

vo2 Mn203

vo MnO

Nd02 Mn02 FeO a-Fe203 COO ( 3 3 0 4 NiO v 2 0 3

Footnote: a) Omitted are the ferromagnetic oxides, such as CrOz, in which there is only one

spin sublattice where the spins are all oriented in the same direction.

magnetic moment associated with the current is a diamagnetic moment. Paramagnetism is related to the tendency of a permanent magnetic dipole to

align itself parallel to a magnetic field. In transition metal ions, the permanent magnetic moment is associated with coupled electron spin and orbital motion of partially filled shells. In a paramagnetic material, these permanent moments are normally random and uncorrelated, but are aligned in an applied field. The magnetic susceptibility decreases with increasing temperature because thermal motion tends to randomize any spins that are aligned by the external field.

In a ferromagnetic material, the interaction among spins is so strong that the magnetic moments are aligned parallel to each other. Such an internal interaction is called exchange field. However, few oxides are ferromagnetic. In a simple antiferromagnetic material, there are two sets of strongly correlated spins that form two spin sublattices of equal magnitude but opposite orientations. These two sublattices cancel each other, resulting in a solid that is like a paramagnet and has weak magnetization. Below the ordering or Ned temperature, the behavior of an antiferromagnetic material differs from that of paramagnetic material because at such low temperatures, the spin sublattice is so rigid that the effect of small imperfect cancellation of spins becomes evident.

In a ferrimagnetic material, the magnitudes of the two opposing spin sublattices are different, which result in a net sizable spin. In an applied magnetic field, these spins align with the external field to yield a saturation magnetization M,, which can be attained at a relatively low field of about 100 Oe. On increasing


100

(0 0

X F

8 nm

8 nm

9 nm

- ---

-.-.- ........... 12 nm

-0.-- 14 nm - 22 nm

-x-x-400 nm

I 1 1

100 200 300

T e m p e r a t u r e K

Figure 3-3 Temperature dependence of magnetic susceptibility of a-Fe203 of different crystallite sizes. (From J. Phys. SOC. Japan, 17, supplement B-1, 690 (1962). copyright Physical Society of Japan).

temperature, thermal fluctuation randomizes the spin until the Curie temperature is reached when the spins are totally random that the response of the material is like that of a paramagnet. Table 3-4 lists the magnetic behavior of some common transition metal oxides.

The behavior of small oxide crystallites may differ from the bulk behavior described above because of the large contribution from ions at the surface. The different magnetic and electric fields experienced by ions in the surface may result in canting and pinning of surface spins. The surface ions may have different atomic spins than the bulk ions if they are in a different oxidation state. Adsorbates may also affect the magnetic behavior.

The magnetic susceptibilities of a number of antiferromagnetic oxides, 01-

Fe2O3, NiO and COO have been measured at different temperatures for different crystallite sizes [34-391. Fig. 3-3 shows the data for a-Fe203. The large crystallites (> 14 nm) behave like typical antiferromagnetic materials. The magnitude of the magnetic susceptibility is very small, and it decreases slowly with increasing temperature up to about 260K. For the small crystallites, the magnetic Susceptibility first increases and then decreases with increasing temperature. This is a typical behavior for superparamagnetic particles. Furthermore, the magnitude of thc low temperature susceptibility increases with decreasing crystallite size. This is also demonstrated in Fig. 3-4 which shows the magnetization as a function of applied field for small crystallites of a-F@03 supported on Si02 [401. The

40

o” 15 N

Y

c W

5 5 a I

0 10 2 0 30 40 50

MAGNETIC FIELD KCe

Figure 3-4 Magnetization at 1.7K as a function of applied magnetic field strength: a. 2.5 nm a-FezO,/SiOz; b. 7.5 m a-FezQ/SiOz; c. 9.5 nm a-Fe,Q/Si02; d. 14.5 m a- Fq@/SiOz; e. 25 m a-Fe2Q. (From J. Phys. Chem., 88, 2525 (1984). copyright Amer- ican Chemical Society).

magnetization increases with decreasing crystallite size. The shape of the magnetization curves follows the behavior for superparamagnetic particles, and can be fitted satisfactorily with the classical Langevin equation:

(3-10)

where M is the instantaneous magnetization, which is the value at the applied field, Ms is the saturation magnetization, which is the value at an infinitely large applied field, p is the magnetic moment per particle, H is the applied field, k is the Boltzmann constant, and T is the temperature.

A quantitative theory to explain the increase in magnetic susceptibility with decreasing particle size is not available. It has been suggested by Nekl [41] that spins which lie near the surface tend to orient parallel to the surface. For small particles, the number of surface spins is a large fraction of the total number of spins. Since the surface spins have different orientations from those of the bulk spins, cancellation of the antiparallel spins is no longer perfect. As the particle size decreases, the contribution of the surface relative to the bulk increases. Thus the magnetic susceptibility increases.

Let us return to the temperature variation of the magnetic susceptibility for large particles of a-Fe203 in Fig. 3-3. These particles behave like typical antiferromagnetic material until a temperature of about 250 K when the susceptibility suddenly increases before it decreases again on further increase in temperature. The temperature at which this phenomenon occurs is called the


Table 3-5 Morin Transition Temperature and Lattice Dilation of a-Fe203 (From J. Phys. SOC. Japan, 24, 23 (1968), copyright Physical Society of Japan).

cr-Fe203 Transition Particle Temp Lattice Constants Size (nm) TM (K) a (1) c (1)

80 245 5.0345 13.749 26 227 5.0365fD.0007 13.7901tO.OOO7 20 120 5.0370fD.0007 13.79020.0007 17 5.03739.0007 13.805+0.0007 14 5.04049.0 1 13.83920.01

Morin transition temperature. The increase in susceptibility for a-Fe203 just above this transition temperature is due to the flipping of the magnetically ordered spin from an orientation parallel to the c-axis to parallel to the c-plane (basal plane) of the crystal. The crystal then changes from antiferromagnetic to weakly ferrimagnetic. For small a-Fe203 crystallites, however, this transition appears to be absent. This effect has been confirmed by magnetization measurements [35,36,391 such as those of Fig. 3-3, and by Mossbauer spectroscopy [42,43] which makes use of the fact that on going from the weakly ferrimagnetic to the antiferromagnetic state, there is a decrease in the magnetic dipolar field and a decrease in the magnitude of the magnetic hyperfine splitting in the Mossbauer spectrum. From these measurements, it has been found that the Morin transition temperature decreases with decreasing crystallite size until it disappears for very small sizes (see Table 3-5). The critical size for the disappearance of the transition temperature has not been firmly established owing to the difficulties in accurate size determination. One explanation for this behavior is that the magnetically coupled spins in small crystallites fluctuate in the basal plane strongly enough that they are no longer oriented to the c-axis even at low temperatures. Another explanation is that the spins are pinned at the surface. For small crystallites in which the surface spins dominate, the magnetization vector does not change direction when the temperature is changed.

Concurrent with the lowering and disappearance of the Morin transition temperature, the lattice of small a-Fe203 crystallites is found to dilate [39,43]. The extent of dilation increases with decreasing crystallite size as is shown in Table 3-5. The lattice dilation is probably homogeneous [39,43-451, and results in clearly observable shifts in the x-ray diffraction peaks. This dilation along the a-axis may contribute significantly to the lowering of the Morin transition temperature through the change in the dipolar magnetic field. It should be noted,

For a-F%03, it is estimated to be about 20 nm.

42

Table 3-6 Crystallite Size Dependence of Magnetic Hyperfine Field

Compound Temp (K) Crystallite Hyperfine field (KOe) size (nm) A site B site

y-Fe203a 77 9.3 sphere 483 508 17.5 sphere 489 513 30.0 sphere 497 517 800 length acicular 508 525

6/1 shape ratio

a-Fe2O3b 296 bulk 83 bulk

296 18 80 18

518 542 503 527

Footnotes: a) From K. Haneda. and A. Momsh, Phys. Lett.. 64A, 259 (1977). b, From W. Kiindig, H. Bommel, G. Constabaris. and R. Lindquist, Phys.

Rev., 142, 327 (1966).

however, that this phenomenon of lattice dilation may not be general, and is not observed for y-Fe203 over the crystallite size range from 9.3 nm to over a few hundred nm [46].

Another effect observed in small magnetic crystallites is the decrease of magnetic hyperfine field with decreasing crystallite size. The magnetic hyperfine field is the magnetic field experienced by the nuclei in an oxide particle. The nuclear magnetic moment interacts with this field, and the interaction can be detected by various spectroscopic techniques, especially Mossbauer spectroscopy. Table 3-6 lists some illustrative data for a- and y-Fe203. There have been three proposals as to the origin of this decrease. The first one, based on a decrease in the Curie or Ned temperature with decreasing crystallite size can now be discounted. Magnetization measurements on small crystallites do not detect changes in these temperatures. The second interpretation assumes that surface ions have a smaller hyperfine field than the ions in the bulk. This effect has now been discounted also by Haneda and Monish [471, at least for y-FezO3. Haneda and Momsh have prepared y-Fe03 samples with a surface coating enriched in s7Fe. The magnetic hyperfine field of this enriched sample is found to compare well with the samples without enrichment. This suggests that the magnetic hyperfine field of the ions at the surface is the same as the bulk ions.

The third interpretation is that of Mdrup and Topsde [481 who suggest that at temperatures below the superparamagnetic blocking temperature, thermally excited oscillations of the magnetization about an energy minimum reduce the


average magnetization and thus the magnetic hyperfine field. The amplitude of the thermal oscillation depends on the magnitude of the energy barrier for flipping of the magnetization vector from one easy direction to another, which is expressed as a product of the anisotropy constant K and the volume of the crystallite. For small oscillation amplitudes, the hyperline field Hhr(V,T) for a crystallite of volume V at temperature T is:

where k is the Boltzmann constant. The theory predicts that Hhf decreases linearly with increasing temperature, and the rate of decrcase depends inversely on the crystallite size. Furthermore, Hhf is independent of the crystallite size at 0 K. These predictions have been separately confirmed for Fe304 [48] and 'y-Fe203 crystallites [47].

Although the magnetic hyperfine field of the surface ions is the same as the bulk ions, there is evidence that under an applied magnetic field, the canting angle of the surface spins is substantially different from that of the bulk ions. This is the origin of the much enhanced 2,5 peak in the Mossbauer spectrum of y-Fez03 [49].

Together with the variation in magnetic hyperfine field, the anisotropy constant also depends on the crystallite size. Table 3-7 shows that the anisotropy constant K decreases sharply from bulk material to 10 nm size crystallites. As yet the origin of this decrease is not known. The effect of such a decrease is that the rate of decrease of the energy barrier for the flipping of the magnetization vector in a particle and the temperature at which the particle becomes superparamagnetic (the superparamagnetic blocking temperature) decrease faster than the volume V of the crystallite, since both of these are proportional to KV.

3.4 QUADRUPOLE SPLITTINGS OF SURFACE IONS

It has been observed in Mossbauer spectra of small iron oxide crystallites that the quadrupole splittings of surface ions are different from the bulk ions. This is not surprising because the quadrupole splittings are determined by the electric field gradient at the nucleus, and it is expected that such a gradient at the surface is different from in the bulk. Fig. 3-5 shows the behavior of small iron oxide crystallites supported on SiOz [50]. In the fully oxidized form, the spectrum shows a superparamagnetic doublet due to Fe3+ ions (curve a). After reduction (curve b), the spectrum consists of two sets of doublets due to Fe2+ ions. The inner doublet shows a quadrupole splitting of 0.91 mm s-l, and the outer doublet, a quadrupole splitting of 1.74 mm spl. It is common to assign the outer doublet to Fez+ ions of higher coordination such as those octahedrally coordinated by oxygen ions in the bulk, and the inner doublet to Fez+ of lower coordination, such as those in the surface [50- 521. The assignment is supported by the fact that the inner doublet can be converted to the outer doublet by adsorption of NH3, CH30H [51], pyridine

44

Table 3-7 Dependence of Anisotropy Constant K on Crystallite Size

Compound Crystallite K erg/cm3 Ref. size (nm)

Fe3 0 4 6 1.2-1.4 x lo6 10 0.85-0.95 x lo6 12 0.85-0.90 x lo6

large 1.1 x 103

?I-Fe203 6 8 x lo5

95.5 1.2 x 106

epitaxially grown 4.6 x 104

12 3 x 105

single domain powder 3 104

cx-F%03 10 3.5 x 105 12 5-6 x 105

13-18' 4-7 x 104

f g

h.i

* These values may not be accurate, see ref. f.

References: a) L. Bickford, J. Brownlow, and R. Penoyer. Proc. Inst. Elect. Eng..

b) M. Boudart, A. Delbouille, J. Dumesic, S. Khammouma, and H. Topsde,

c) J. Coey, and D. Khalafalla, Phys. Stat. Sol. (A), 11, 229 (1972). d) A. Morrish, and E. Valtyn, I. Phys. SOC. Jpn.. 17, suppl. B1, 392. (1962). e) H. Takei, and S , Chiba. J. Phys. SOC. Jpn. 21, 1255 (1966). f) J. Amelse, K. Arcuri, J. Butt, R. Matyi, L. Schwartz, A. Shaxpiro, J.

g) S . Mdrup, and H. Tops&, Appl. Phys., 11, 63 (1976). h) W. Kiindig, H. Bommel, G. Constabaris, and R. Lindquist, Phys. Rev.,

i ) W. Kiindig, K. Ando, R. Lindquist, and G. Constabaris, Czech. J. Phys.

B104, 238 (1957).

J. Catal., 37, 486 (1975).

Phys. Chem., 85, 708 (1981).

142, 327 (1966).

B17, 467 (1967).

[50], NO [52], CO [53], and H 2 0 [54]. It is apparent from these examples that detecting the changes in quadrupole splittings of surface ions on adsorption of molecules is one method to investigate the nature of the adsorption sites.


-4 - 2 0 2 4 VELOCITY ( r n r n l s )

Figure 3-5 Room temperature Mossbauer spectra for 0.5 wt.% Fe/SiOz. a. After reduction in H, at 673K followed by oxidation in 0 2 at 423K; b. After reduction in H2 at 498K; c. After reduction in Hz at 673K. (From J. Catal., 101, 103 (1986), copyright Academic Press).

It might be expected that the electric field gradient due to the iorricity of the oxide matrix is larger at the surface than in the bulk, and would lead to a larger quadrupole splitting for the surface Fez+ ions than the bulk ions, which is contrary to what has been observed. This is due to the fact that the couloumbic field from the ions in the oxide matrix is only one contribution to the electric field at the nucleus. The smaller quadrupole splitting for surface Fez+ ions than bulk Fez+ ions has been interpreted as due to the opposing effect of larger crystal field gradicnt and smaller d electron field gradient at the surface than thc bulk [52,56,57]. Inner and outer doublcts of Fez+ ions have been observed on many samples of silica-supported small crystallites of reduced iron oxide, although thc magnitude of the quadrupole splitting differs somewhat from sample to sample. The prcsencc of inner and outer doublcts has also been observed using MgO, y- A1203, and Ti02 as supports [%I.

Clearly distinguished inncr and outer doublets of Fe3+ ions have not bccn rcportcd. However, it is well established that the average quadrupole splitting of

46

Table 3-8 Crystallite Size Dependence of Quadrupole Splitting in a-Fe203 at 298K. (From Phys. Rev., 142, 327 (1966), copyright American Physical Society).

Size (nm) Q.S. mm s-1

< 10 13.5 15.0 18.0

0.98 0.57 0.55 0.44

Fe3+ in a crystallite increases with decreasing crystallite size, as is shown in Table 3-8. This increase has been explained by the larger asymmetry in the environment of surface ions than bulk ions. Some researchers have attempted to fit their Mossbauer spectra of small crystallites containing Fe3+ ions using two sets of peaks of different quadrupole splittings. The set with a larger quadrupole splitting is assigned to the surface ions [56,58], which is supported by the observed decrease in the quadrupole splitting on adsorption of water, methanol, and ammonia [59].

3.5 SURFACE ELECTRONIC STRUCTURE

Knowledge of the detailed surface electronic structure is critical in the understanding of the chemisorptive and catalytic properties of oxides. The bulk electronic structures have been studied for a number of oxides [60]. Studies on surfaces have been concentrated on the detection of electronic states not found in the bulk that are generated by the presence of the surface. These states are called surface states. The energy of the surface states of interest in catalysis commonly lies near the Fermi energy of the solid (or valence states). These surface states are either partially filled with electrons so that they can both donate and accept electrons from the molecules interacting with the surface, or close enough to the filled valence band (or empty conduction band) such that together, they provide a pair of states to accept and donate electrons to the interacting molecules simultaneously. Consideration of electron bands in the solid is also important in photo-enhanced adsorption, desorption, and catalysis. These photo-assisted processes as well as a short description of electronic bands of oxides will be discussed in Chapter 14. The readers are referred to standard textbooks in solid state physics for more in-depth discussions.

The nature and energy of any surface state depend on, among other factors, the ionicity of the oxides and the position of the ions. The major contributions to these factors are from ions in the surface region. Therefore, it is expected that surface structural rearrangements could lead to substantial changes in the surfxe electronic structure [61,62].


Another effect expected for ionic surfaces is the reduction of the ionic charges of the surface ions. This results from the downward shift in energy of the surface orbitals compared to the corresponding bulk orbitals and the enhanced covalency of the cation-anion bond, as well as the polarization of the surface orbitals by the surface electric field. The surface-induced changes in the electronic properties are strongly localized in the outer few layers of the surface regions.

For oxides that are predominantly covalent, the surface coordinative unsaturation results in the formation of valence orbitals projecting from the surface that are from the outermost ions (sometimes being referred to as dangling bonds) [63]. There are two sets of such orbitals. One set is empty and is located mostly at the cation, and the other is filled and is located at the anion. The separation in energy of these two sets increases with increasing ionicity of the solid.

The electronic structures of some oxides have been studied by electron spectroscopies (UPS and ELS). For surfaces of oxides with empty d bands, such as stoichiometric Ti02 (loo), (001) and (110), SrTiO3 (100) [l-31 and ZnO [@I, no intrinsic surface states are identified within the bandgap. Similar results have been observed for W 0 3 [65] and Moo3 (010) [66] surfaces. The surface states of these oxides are close to the conduction or valence bands. For example, the surface states of ZnO are at -2 eV and -5 eV below the top of the valence band 185,861.

Loss of surface oxygen ions in these transition metal oxides with empty d- bands causes development of filled electronic states in the band gap. These states are detected by U P S as emission above the valence band emission (see Fig. 3-6) [67,68]. It is interesting to note that the intensity of this bandgap emission decreases when the oxide is annealed in vacuum. The decrease is due to reoxidation of the surface by oxygen diffusion from the bulk which is taking place faster than reduction by thermal desorption of oxygen. In a similar manner, this band gap state can be depopulated by adsorption of O2 [65,67,69,70]. For oxides with empty d bands, the bandgap surface state has been attributed to the cations at a lower oxidation state than the bulk, such as Ti3+ for Ti02 and SrTi03.

Oxides with partially filled d-bands, such as Ti2O3, V2O3. a-Fe203 [70], T i0 [711, MnO [72], FeO [72,73], COO [72,74], NiO [72,75], and CuO [76] have been investigated. The d states of the cations overlap significantly with the 0 2p bands. This makes it difficult to detect the presence of any surface states. Deconvolution of the experimental spectra shows that the bands are derived primarily from the metal 3d and oxygen 2p levels. The density of states of the 3d band shows two maxima near the Fermi level, and becomes more pronounced as the number of d electrons is increased. The absence of large differences between the U P S spectra of the surfaces of these compounds and the bulk band structures implies that their surface electronic structures do not differ significantly from the bulk.

The effect of surface roughness has been investigated on Ti02 [77,78]. The surface states on the atomically rough surface are found to lie deeper in the bandgap than on the smooth surface. This could result in enhanced reactivity of the rough surface as compared to the smooth surface.

Most of the studies todate concentrate on studying the surface electronic structures of the oxides. Only a few have been reported to study the interaction of molecules with these surfaces. Therefore, there is very limited knowledge of the

48

-1 2 -8 -4 Ef=O

Electron Binding Energy e V

Figure 3-6 UPS spectra of the clean ordered SrTiQ (1 11) surfaces taken at: a. 27; b. 300; c. 6000C. (From Phys. Rev. B.. 17. 4942 (1978). copyright American Physical Society).

electronic structure of the surface-molecule bonds. The formation of such bonds is essential in surface chemistry and catalysis, and it requires that the molecule can physically approach the adsorption site (steric requirement), and surface orbitals of the appropriate energies and symmetry are available. Since information on the surface orbitals are not readily available, researchers in this field have made significant use of the steric requirement. This has led to the concept of coordinative unsaturation of surface ions, which will be discussed in the next chapter.

3.6 SURFACE VIBRATION

The lattice vibration of surface atoms which is called the surface optical phonon vibration modes of some transition metal oxides have been studied using high resolution electron energy loss spectroscopy (HREELS). These oxides include NiO, TiOz, ZnO, and SrTi03, all of which exhibit phonon modes that strongly couple with the incident electron beam. The strong coupling makes it readily possible to detect not only the fundamental phonon modes but also the overtones. Fig. 3-7 shows a HREEL spectrum of a ZnO (10iO)surface which possesses a fundamental phonon mode at 68.8 meV [79]. Fig. 3-8 shows that for a NiO (001) surface whose fundamental phonon mode is at 69.5 meV [80]. The overtones in both spectra arc clearly visible. For the large ordered surfaces studied, the surface phonon vibrations generally lie in the range of 40-100 meV (about 320-800 cm-’), and the strong


286 K

h 127 K

Energy Loss meV

Figure 3-7 Energy-loss sepctrum of 7.5-eV electlons after specular reflection from the

(1070) surface of ZnO. (From Phys. Rev. Lett., 24, 1416 (1970), copyright American Physical Society).

overtones are in the region where vibrations involving surface adsorbates appear. This makes it difficult to identify adsorbates.

The frequencies of the surface phonon modes can be described by classical analyses that apply well to cases where the vibrational wavelength is much larger than lattice parameters @I]. It has been shown that the surface phonons that are excited most strongly in HREELS have small wavevectors parallel to the surface and long characteristic penetration depths [82]. Thus the energy loss spectra are rather insensitive to details of the surface structure.

Small differences are observed in the phonon frequencies of different ZnO surfaces. The observed values are 68.8 meV for the (lOi0) surface, and 67.3 meV for Lhe (0001) and (000i)surfaces [79]. These values are higher in energy than the infrared-active optical bulk modes for which the displacement of ions is pcrpcndicular to the surface. They agree well with the frequencies calculated from the classical treatment of ionic vibrations in the continuum approximation. The fundamental modcs of the NiO (100) and (111) surfaces are at about 69.5 mcV [80]. For the SrTi03 (100) surface, surface optical phonon modes at 57 meV (460 cm-') and 92 meV (740 cm-') have been observed, which are just below the bulk longitudinal modcs at 474 and 788 cm-' [831.

Recently, a Fourier transform technique has been developed to remove the combination and overtone structures in HREEL spectra and retain only the loss slructure of the fundamental modes [84]. Using this technique, the energy loss spcclrum of H 2 0 adsorbed on a SrTi03 (100) surface has been observed.

50

-200 0 200 400

Energy Loss meV

Figure 3-8 Specular electron energy loss spectrum of a NiO (001) surface. Sci.. 152/153, 784 (1985). copyright Elsevier Scientific Publ.).

(From Surf.

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Chapter 4

SURFACE COORDINATIVE

UNSATURATION

4.1 FORMATION OF SURFACE COORDINATIVE UNSATURATION

It is generally accepted that surface coordinative unsaturation is important in surface chemistry. This concept is analogous to that in coordination chemistry and arises from the fact that because of steric and electronic reasons, only a limited number of ligands or nearest neighbors can be within bonding distance of a metal atom or ion. In most transition metal oxides, the oxygen anions in the bulk form close-packed laycrs and the metal cations occupy holes among the anions as described in Chapter 2. In this picture, since the bulk oxide ions are as densely packed as possible, the oxide ion ligands around a metal cation are thought to have saturated the coordination sphere of the bulk cation, that is, the bulk cation is coordinatively saturated.

We have seen in Chapter 2 that in the formation of a surface by cleaving an oxide crystal, metal-oxygen bonds have to be broken. Therefore the surface anions and cations have fewer numbcrs of nearest neighbors than the corresponding ions in the bulk. These surface anions and cations are coordinatively unsaturated (cus). In most instances, coordinative unsaturation results in ions that are active in bonding with adsorbates. However, as will be discussed later, not all surface coordinativcly unsaturated ions are necessarily active and have a high tendency to form chemical bonds with adsorbates.

There are two approachcs to picture the formation of surface coordinative unsaturation depending on the way the surface is prepared. The fist approach applies to surfaces of microcrystalline samples prepared from aqueous solutions. The second applies to surfaces formcd from cleaving large single crystal samples.

An oxide or hydrous oxidc sample prepared by precipitation from an aqueous solution is formed by condensation and polymerization of hydroxylated mctal ions. For example, a M(II1) metal ion existing as a monomeric unit possesses a saturated coordination of three hydroxyl and three watcr ligands. Condensation of two

53

54

hydroxyl groups from two different monomeric units links the monomers with a bridging oxygen to form a dimer. Condensation of more than one pair of hydroxyl is possible and the dimeric unit has more than one bridging oxygen ion. Bridging by hydroxyl ions is also possible. These units are schematically shown below:

Monomer

Dimers

Coordinatively unsaturated sites

Hydroxylated Surface Partially Dehydroxylated Surface

When condensation occurs between many diffcrcnt monomeric units, which is orten promoted by drying at elevated temperatures, a three-dimensional network is formed. Depending on the metal and the drying condition, the network may be amorphous, semicrystalline, or crystalline. A possible configuration of a hydroxylatcd surface of a network is shown above. In this schematic drawing, thc number of ligands around a metal ion is set to remain at six, but the number of ligands for the oxygen ions increases to four in the bulk and three in the surface. Other situations are possible, and the coordination number of the metal ions may change as the solid is formed. Dchydroxylation of this surface may take place between two adjacent hydroxyl groups, and the surface mctal ions involved would be coordinated to only five oxygen ions. They bccome coordinatively unsaturated. The surface oxide ions generated are also coordinativcly unsaturatcd and have lower coordination numbers than bulk oxide ions or oxide ions of the

COORDINATIVE UNSATURATION 55

surface hydroxyls. Often the M"+(cus) site bchavcs like a Lewis acid, and the 02-(cus) ion is more basic than the bulk ions. Such an acid-base pair site participates in hcterolytic dissociative adsorption. More extensive coordinative unsaturation is possible if more dchydroxylation takes place, or if surface lattice oxygen is rcmoved by reduction. Such a process of dchydroxylation has been used by Burwell, et al. [2] to describe the formation of a partially dchydroxylated chromia surface which has surface ions that are coordinatively unsaturated by one or two ligands.

If dehydroxylation occurs randomly bctwccn pairs of hydroxyl groups, it would be difficult to achieve complete removal of surface hydroxyl groups. Evcn after high temperature evacuations, isolatcd surface hydroxyl groups are often present that show a sharp IR absorption band at high frequencies (?3700 cm-I). As an example, IR bands of surface hydroxyl groups have bcen identified on anatase at 3636, 3654, 3672, and 3707 cm-'. The intcnsitics of thcse bands decrease with increasing evacuation tcrnperature. The rate of decrease is the fastcst for the lowest frcquency band, whilc thc highest frcqucncy band is almost unchanged [ 11.

The second approach to picture the formation of surface coordinative unsaturation is by clcaving a single crystal. Commonly one s w t s with an clcctrically ncutral singlc crystal whose faces are bound by nonpolar surfaces. If thc crystal is bound by polar faces, thcsc faces nccd to be populated with ion vacancies, hydroxyl groups, or other species so that the crystal does not posscss any pcriodic dipole momcnt.

Once the crystal is dcfined, it is cleaved (conccptually) into two halves along the dcsircd direction to expose the planes of intcrcst. The common guidclincs for the partition of ions along the clcavage planc arc: (1) it rcsults in two half- crystals that are electrically ncutral; (2) it minimizes the total amount of coordinativc unsaturation as much as possiblc; (3) it avoids the formation of ions with an unusually large dcgrce of coordinative unsaturation; and (4) it involves the brcaking of as small a numbcr of bonds as possible.

Figurc 4-1 shows modcls of three ideal and ordcrcd Ti02 and Ti203 surfaces. The Ti02 surfaces would be fully oxidizcd surfaces cxccpt for the anion vacancics shown in Fig. 4-la and b. If thc anion vacancics arc fillcd by oxidc ions, thcsc surfaces can be generated theoretically by fracturing a Ti02 singlc crystal along thc dcsircd dircction, and partitioning the atoms along thc clcavagc planc equally bctwcen thc two parting faces. It should be cmphasizcd that thc modcls rcprcscnt highly ordcrcd surfaccs, and such ideal surfaccs arc not cxpcctcd to cxist ovcr large macroscopic surfaces in rcality. For cxamplc, the Ti02 (110) surfacc is pictured as one in which pcrfcct rows of oxygcn ions sit on altcrnatc rows of Ti ions in thc surface plane. Although such a configuration minimizes the total numbcr of coordinative unsaturation and thc number of bonds broken, it has low entropy. On a rcal surface, thc desire to maximize cntropy rcsults in lcss wcll- ordcrcd arrangcmcnts.

When a crystal is fractured in a rcal cxpcriinent, it is oftcn observed that clcclrons arc emitted from thc surface for somc time. Thcsc arc callcd cxoelcctrons. In addition, oxygcn ions arc obscrvcd to be evolvcd. I t is possiblc

56

-LJ

Y Pe 1

Figure 4-la Model of a Ti02 (110) surface. Solid circles are Ti ions, open circles are 0 ions. One type 1 surface lattice oxygen is removed to show an anion vacancy.

Figure 4- lb Model of a Ti02 (100) surface. Shaded circles are surface lattice oxygen ions under the plane of Ti ions. Other symbols are the same as in Fig. 4-la. One type 1 oxide ion is removed to show an anion vacancy.

Figure 4-lc Model of a Ti203 (047) surface. Symbols arc the same as in Fig. 4-lb.


that neutral atoms are also evolved without being detected. This further points out the fact that an ideal perfect surface of macroscopic dimensions may be hard to obtain.

An inspection of these models shows that different exposed surfaces possess ions of different degrees of coordinative unsaturation. On the Ti02 (1 10) surface, two types of Ti ions are present. Half of the Ti ions have six oxygen ion nearest neighbors and are coordinatively saturated. The other half have five oxygen ion neighbors. Similarly, there are two different types of lattice oxygen ions. Type one is bonded to two Ti cations that are six-fold coordinated, sits above the plane of the other ions, and, on Sn02 (1 10) [73], has been shown to be more easily removed than other anions. The other type is bonded to three Ti ions of six- and five-fold coordination. They are in the surface plane. The Ti02 (100) face has five-fold coordinated Ti ions, and the (047) face has four-fold coordinated Ti ions.

In addition to these, ions of other types of coordinative unsaturation can be created by introducing defects on the surface. An example is shown in the Ti02 (100) surface in Fig. 4-la. An oxygen ion in the row sitting above the surface is removed, perhaps by reduction. This action exposes two Ti ions of four-fold coordination. In general, the presence of anion vacancies results in cations of a lower coordination and in a lower oxidation state. The resulting surface is commonly chemically more active.

4.2 CHEMICAL PROPERTIES OF SURFACE COORDINATIVELY UNSATURATED SITES

Coordinatively unsaturated sites are responsible for chemisorption and binding of molecules to a surface in most instances. Indeed this explains very well a variety of phenomena including poisoning of a surface, competitive adsorption, and the common requirement of heating of an oxide to activate it for chemisorption and catalysis.

Activation of an oxide by heating is needed because of the strong adsorption of water on oxides. Thus an oxide surface becomes fully covered with adsorbcd water and hydroxyl groups once it is exposed to the moisture in the atmosphere. This process is shown by eq. (4-1) and (4-2):

In eq. (4-l), water is adsorbed molecularly to satisfy the coordinative unsaturation of the metal ion. In eq. (4-2), the coordinatively unsaturated cation and anion pair adsorbs a water molecule dissociatively as OH- and H+. These

58

a, c.'

2" 0 N r

Evacuation Temperature, C

Figure 4-2 Surface hydroxyl content of a ZnO sample evacuated at various temperatures. (From J. Phys Chem., 84, 2054 (1984), copyright American Chemical Society).

surface ions become coordinatively saturated and unable to adsorb other molecules. Heating causes removal of water and formation of surface coordinative unsaturation. The extent of dehydroxylation depends on the temperature, as can be seen in Fig. 4-2 for ZnO [3], and Fig. 4-3a for Cr203 [2]. These figures are representative of most oxides.

Referring to Fig. 4-2, it can be seen that water is lost continuously from ZnO when the temperature is raised. The water that is lost at low temperature are weakly adsorbed water. It is probably nondissociatively adsorbed water that is lost by a process similar to the reverse of eq. (4-l), as well as water held to the surface by hydrogen bonding. A rapid loss of water occurs around 200 to 350°C which corresponds to dehydroxylation of the surface by the reverse of eq. (4-2). Further dehydroxylation of the remaining hydroxyl groups bcyond 400°C is slow. These hydroxyl groups are probably isolated and difficult to be removed.

Removal of the weakly adsorbed molecular water docs not necessarily activate the surface, especially if the water molecules are held to the surface by hydrogen-bonding. This is illustratcd in Fig. 4-3b. The amorphous Cr203 surface only becomes active in chemisorption of CO or O2 and in hydrogenation of 1-hexcne after the oxide is activated above 200"C, although some adsorbed water is already removed by this temperature (Fig. 4-3a) 121. The chemisorptivc capacity and the catalytic activity increase rapidly as the degree of dchydrox ylation increases. The manner in which these quantities incrcasc dcpcnds on the molcculc and the reaction bccause different requirements of surface sites may bc involved. For example, the dependence of NH3 adsorption on Cr203 on the activation temperature is much less than for CO or 0 2 .

COORDINATIVE UNSATURATION

1.2

- 0 0)

e .8 0 v,N O E Q : c cN2 0 2 L u.4 0 0 ,

0 0 o E

J

-

0

59

-

/

-

-

0 I N

2

3 0 200 300 40(

Activation Temp. C

300

v, 5 -h P,

200 D ;

3

P, .. 100 .!

rn

Figure 4-3a Amount of water lost per CPf upon heating from 25OC and the surface area of amorphous Cr2O3 as a function of temperature of activation with Hz. Data from ref. 2.

Activation Temp C

300

Figure 4-3b Capacity for 0 2 or CO adsorption and catalytic activity for 1-hexene hydrogenation as a function of the tcmpcrature of activation of amorphous CrzO3 in H2. Data from ref. 2.

60

T Surface normal

Figure 4-4 Adsorption of CO on a: ZnO (0001); b: ZnO (1070) surface.

There is evidence that on a surface where the coordinativc unsaturation is along a certain direction, (that is, directional valcnce orbital), a molecule covalently bonded to the surface lies roughly in the direction of thc valence orbital. The evidence is provided by CO adsorption on ZnO surfaces shown in Fig. 4-4. Using angle-resolved U P S to monitor the orientation of the adsorbate, it has been found that CO is adsorbed linearly along the direction of the surface normal of a ZnO (0001) surface and colinear with the surface valence orbital [4]. On the ZnO (1010) surface, the axis of the adsorbed CO molcculc is along the direction near the expected direction of a tetrahedral bond of a surface Zn ion [4,5]. On these two surfaces, CO are bonded to the surface Zn ions which have thrce oxygcn ion nearest neighbors instead of four as in the bulk. The situation for the 0-polar (0001) surface is less clear. The Zn ions in this surfacc have four oxygcn ion nearest neighbors. They arc coordinatively saturatcd. In one report, it is said that adsorption of CO on this surface takcs place only on cations in the surface stcp defects where coordinatively unsaturated Zn ions are exposed [4,6]. In another report, evidence is presented that CO may adsorb to form surface carbonate [721. It is further mentioned that the amount of CO adsorbcd on this surface is rcduccd if the surface is disordercd by ion-bombardment, although one might cxpcct that such treatment would gencrate more surfacc coordinativcly unsaturated cations. Furthermore, it is reported that this bchavior is oppositc to that of C02 [72].

It follows that surface ions that are coordinatively unsaturatcd by more than


one ligand may adsorb more than one molecule. This has been observed on chromia [7,8]. When a sample of reduced chromia is exposed to a mixture I2CO and 13C0, three IR bands corresponding to Cr(12C0)2, Cr(12CO13CO) and Cr(13C0)2 are observed. If each Cr ion adsorbs only one CO, only two bands for Cr12C0 and Cr13C0 are expected.

Direct observation of the interaction between an adsorbate and a surface cation has been made. For example, UV-vis spcctroscopy shows that on a reduced Cr2O3/AI2O3 sample, Cr ions of square pyramidal coordination are converted to a distorted octahedral coordination upon adsorption of ammonia or methanol [9]. When 13C0 is adsorbed on a Coo-MgO solid solution, the EPR signal of Co2+ shows hyperfine interaction with 13C [lo]. The Mossbauer spectrum of reduced supported iron oxide shows an inner doublet and an outer doublet of Fe2+ ions. Upon adsorption of H 2 0 [ll], pyridine [12], NO [13] or other mblecules, the inner doublet is converted to the outer doublet.

Only in a few cases where surface coordinatively unsaturited oxide ions are presesnt without any associated coordinatively unsaturated metal ions. The 0- polar surface of ZnO is one such case. It has been found that without accompanying M"+(cus) ions, the 02-(cus) ions are not very reactive, particularly for species that prefer acid-base pair sites. For example, binding of water on the 0- polar surface is weak. Temperature programmed desorption of water shows a peak at 190 K from this surface, and a peak at 340 K from the Zn-polar surface [14]. The 0-polar surface is also less active than the Zn-polar surface in the catalytic decomposition of 2-propanol [15].

4.3 ADSORPTION OF SMALL MOLECULES

Substantial information is available on the interaction of small molecules with transition metal oxides and other oxides. At present, most of the information on the molecular nature of the adsorbate has been obtained by infrared spectroscopy. Because of the strong infrared absorption by oxides in the region below about 1000 cm-', there is little information on the surface-adsorbate bond. Such data may become available as new techniques such as EXAFS and HREELS become more sophisticated, especially the latter when deconvolution techniques to rcmove surface phonon spectra are developed.

In general, molecules may interact with an oxide surface in different ways:

(1) Molecular (nondissociative) adsorption in which the interaction is mainly by o-donation and/or x-bonding interaction. single surface coordinatively unsaturated ion.

(2) Dissociative adsorption in which a molecule dissociates upon adsorption. Dissociation of H 2 0 into H+ and OH- upon adsorption is an example of hctcrolytic dissociative adsorption in which the molccule is dissociated into charged species. This type of adsorption usually requires an anion-cation coordinativcly unsaturatcd pair site. Homolytic dissociative adsorption in which ncutral species are formed may also occur but less

This can take place on a

62

frequently. (3) Abstractive adsorption in which the adsorbate abstracts a species from the

surface (often a proton). This commonly occurs on acidic oxides. If a proton is abstracted from the surface and the adsorbed species becomes cationic, the adsorbate could be held to the surface by electrostatic forces, and coordinatively unsaturated ions might not be involved.

(4) Reductive adsorption in which an adsorbed molecule is oxidized while the surface is reduced. It may also be abslractive as in the case when a hydrocarbon molecule is oxidized on adsorption to a carboxylate utilizing the lattice oxygen while reducing the cation.

In addition to adsorption, a surface may catalyze reactions of the adsorbate. Examples of various catalytic reactions are discussed in later chapters.

Deprotonation and protonation of an adsorbate are examples of lssociative and abstractive adsorption, respectively. These are BrQnsted acid-base reactions between an adsorbate and a surface, and commonly occurs with molecules such as NH3, pyridine, and alcohols. They will be discussed in greater detail in Chapter 5.

Reductive adsorption is commonly observed on transitional metal oxidcs which have cations that have readily accessible multiple oxidation statcs. This type of adsorption must be considcrcd whcn interpreting adsorption data. For example, adsorption of nitrogen-containing compounds may result in their oxidation to nitrates or nitrites, and oxidation of adsorbed hydrocarbons to carbonates or carboxylates also occurs frequently in the absence of gaseous oxygen.

l a t i c e RN02(ad) + 2 OH(ad)

RC02(ad) + 3 OH(ad) RCH3 l a t t i c e 0,

(4-3)

(4-4)

Busca and Lorenzelli have summarized the infrared band frcquencies of carbonate ions, carboxylates, bicarbonates and formates of various compounds [16]. Because of the many different possiblc ways that these species can bond to a surface, there are no characteristic bands unique to each species to help thcir identification. Free carbonate ions possess two dcgcnerale IR-active asymmelric CO stretching (v3) bands at 1415 cm-’. The splilting between the two bands on adsorption may be used to identify the type of coordination of the carbonatc. If bonding is to a metal ion that has a high polarizability, the classificauon proposed by Nakamoto et a]. shown on the ncxt pagc may be uscd [171.

The assignment of v3 splitting largcr than 300 cm-’ is not certain. On oxidcs of metals having low polarizing powcr, the splitting of v3 is smaller, and the distinction between a monodcntate and a bidcntatc coordination bccomcs less clear. Then consideration of thermal stability can help. A moncdcntatc structure should normally be less stable on a surface than a bidcntatc spccics [16].

Carboxylates show absorption in thc rcgion of 1550-1760 and 1150-1200 cm-’. Sometimes the observation of two bands in thcsc rcgions is confused with the v3 splitting of a carbonate species. Adsorbcd formate ions show four bands


Coordintation of Carbonates

v3 splitting (cm-') Type of coordination zero symmetric 100 monodentate 300 bidentate or bridged

symmetric monodentate bidentate bridged

under optimal conditions: C-H stretching at about 2900 cm-', asymmetric C 4 stretching at 155CL1600 cm-', C-H bending at about 1400 cm-', and symmetric CO stretching at 1345-1385 cm-'.

It is commonly observed that the form of adsorption of a molecule depends upon the pretreatment of the oxide. This results from different modes of adsorption requiring different surface sites. Protonation requires surface hydroxyl groups and thus depends on the extent of surface dehydroxylation. Abstractive reductive adsorption removes surface oxide ions and thus depends on the oxidation state of the surface. The presence of adsorbed oxygen could also oxidize adsorbates. As illustrated earlier, whether a molecule adsorbs nondissociatively (eq. 4-1) or dissociatively (eq. 4-2) depends on whether cation-anion coordinatively unsaturated pair sites are available. In what follows, illustrative examples of the adsorption of hydrogen, hydrocarbons, alcohols, CO and NO are presented to show the different types of adsorption.

H2 Hydrogen is adsorbed on ZnO both molecularly and dissociatively.

Molecular adsorption has been observed to occur at -195"C, and the adsorbed Hz shows an IR band at 4019 cm-'. This band is shifted to 3507 and 2887 cm-' for HD and DZ, respectively [18]. Dissociative adsorption is most likely heterolytic although there is no direct confirmation of it, and occurs at room temperature:

For Hz, a band assigned to Zn-H at 1712 cm-' and C t H at 3490 cm-' are observed [19]. It is interesting that the H atoms bound to Zn-H and 0-H exchange with D2 present in the gas phase to form HD in the gas phase, but not

64

-1 1700 1650 1600 Wavenumber cm

Figure 4-5 Effect of increasing CO pressure on the Zn-H strctching band frcquencics (Pll2 = 100 Ton). (From J. Catal., 51, 160 (1978), copyright Academic I’ress).

with Zn-D or O-D [20]. This IR-active dissociative form of adsorbed H2 is involved in the catalytic hydrogenation of alkenes on ZnO.

The frequency of the Zn-H band is affected by coadsorbates. When CO is adsorbed onto a hydrogen-coverd surface, two new Zn-H bands appcar whosc intensities vary syslematically with the CO coverage. This is illuslratcd in Fig. 4-5 [21]. The two new Zn-H bands represent Zn-H cornplcxes wilh one or both of its neighboring Zn ions coordinated to a CO molcculc.

Alkenes Alkenes are adsorbed both nondissociativcly and dissociatively on ZnO.

Nondissociative adsorption is in the form a x-complex. This has k e n observed for C2H4 [22], propcnc [23 ] , and butenes [24]. Thesc x-complcxes exhibit a C=C stretching band that is shifted to a lower frequency from the corrcsponding band in the gas phase by less than 50 cm-’ . These spccics are only weakly adsorbed.

Propcne is also adsorbed on ZnO dissociatively to form x-allyl. The dissociation is probably heterolytic, and a x-ally1 anion is formcd [23,25]:

C ~ H ~ + -Zn+ 4 CH~-CH-CH; H+ (4-6) - 1 - - I Zn


This process is similar to H2 adsorption on ZnO. Upon adsorption of propene in this form, an 0-H infrared band is obscrvcd but no Zn-H band. Thus the x-ally1 species is adsorbed on the Zn ion. Another band at 1545 cm-' is also obscrvcd which is assigned to the antisymmctric stretch of thc C-C-C skeleton. This assignment is substantiated by observing shifts in this band position upon deutcrating the propene molecule at various positions and noting that the magnitudes of the shifts agree with those prcdictcd by group frequency analyses. It is also supportcd by thc fact that the hydrogen of the G H group formed arises from thc dlylic H of the propene molecule. This x-ally1 species can be reversibly removed from the surface on evacuation at slightly elevated temperatures. This fact strongly suggests that it is not an oxidized surface species. The relative amount of propene adsorbcd as x-ally1 and as n-complex depends on thc extcnt of dchydration of the surface [26].

Conclusivc evidence for the formation of n-ally1 from butene when it is cxposcd to ZnO has not been established. Although an IR band at about 1570- 1580 cm-l has been observed, it cannot be readily removed by evacuation [24]. This, together with the fact that the kinetics of its formation does not appcar to correlate with the kinetics of butene isomerization, raises the possibility that this band might be due to a surface oxidized species.

By applying the criterion that x-complex and x-ally1 can be readily dcsorbcd from the surface, it has also bcen claimcd that a 1600 cm-' band is due to the C=C strctch of a x-complcx of propcne on Ga-Mo oxide [27], and a 1580 cm-l band to that on cuprous oxide [28]. Bands at 1400 and 1440 cm-I have bcen assigncd to the n-ally1 species on thcsc two oxides, respectively. However, definitive assignmcnt of IR bands cannot be made readily without supplementary information from other observations, espccially in the absence of close analogs to the adsorbed spccics. Indccd, thc investigators caution that the 1440 cm-' band mentioned above may also be assigncd to deformation vibrations of the CH3 group in physically adsorbcd propene or a x-complcx. It is also intcrcsting to note that adsorption of allyl bromide on ZnO rcsults in a surface allyl species with a band at 1470 ern--' instead of 1545 cm-' which is observed on adsorption of propcnc [29]. Only x-complexes have bccn idcntificd on adsorption of ethene and propene on Ni- Mo-MgO, Co-Mo-MgO 1301, and Ti02 [31]. They havc also been identified on adsorption of butcncs on a V-P-0 [32], where it is observed that the red-shift in the C=C stretch on adsorption depends on whether the oxide is oxidized or reduced. Thc shift is 33 cm-' on an oxidized sample, and 23 cm-' on a reduced sample.

Like alkcncs, terminal alkynes may bc adsorbcd eithcr molecularly or dissociativcly [33,34]. It has bcen obscrvcd that acetylene and phcnylacctylcne are adsorbcd at room tcmpcraturc dissociativcly as acetylide (HCS-) and phcnylacctylidc, rcspectivcly, on ZnO [34]. Propyne is adsorbcd on ZnO to form both mcthylacetylide and propagyl spc5cs [33,35-391:

[ CH3 X - C ]

mclhylacctylidc ProPagYl

66

The propagyl species is the intermediate in the isomerization of propyne to allene. The formation of propagyl species on ZnO has been confirmed using XPS which shows only one carbon peak for the adsorbed species, instead of two peaks for the gas phase propyne molecule or the methylacetylide species on a silver surface [34]. Some of the dissociatively adsorbed species is further oxidized upon heating, resulting eventually in carbon oxides and water. The extent of dissociative adsorption depends on the crystallographic orientation of the surface. On ZnO, dissociative adsorption occurs on the Zn-polar surface, but only molecular adsorption occurs on the 0-polar surface [341.

Alcohob and Acetone With few exceptions, alcohols are adsorbed by heterolytic dissociation at

room temperature on a dehydroxylatcd surface with a proton going to a surface lattice oxygen and an alkoxide to a surface cation. This is again similar to the other heterolytic dissociative adsorption shown in eq. (4-2), (4-3, and (4-6). The process on ZnO is shown below:

ROH + Z n 4 -+ RO H I I

Z n - 0 (4-7)

For example, methoxide has been detected when methanol is adsorbed on ZnO [40] and Moo3 [41,42]. Similarly, ethoxide is formed from ethanol on Moo3 [43] and ZnO [44], 1-propoxide is formed from 1-propanol [44] and 2- propoxide is formed from 2-propanol on ZnO [45]. Methanol has been shown to be adsorbed dissociatively at as low as 105 K [46].

Oxidation of surface alkoxide commonly occurs. The oxidation of methoxide to a surface formate is well documented on ZnO [47,48]. The zinc ion is reduced to zinc metal which can be desorbed from the surface. It also occurs on Cr2O3 [49]. 2-propoxide on ZnO is oxidized to a surface enolate which eventually is dcsorbed as acetone [45].

Adsorbed acetone dissociates heterolytically into a surface enolate species with simultaneous formation of a surface OH group. This reaction:

is observed to occur on both ZnO [45] and NiO [50]. The reversibility of this reaction has been demonstrated. Complete exchange

of the hydrogen atoms of the methyl groups with deuterium has been observed when a sample of ZnO with adsorbed acetone is exposed to deuterium gas [451. Thus deuterium is adsorbed dissociatively on the oxide, and the OD group participates in the reverse reaction.


co Carbon monoxide is a rather common probe molecule in the study of the

surface chemistry of transition metal oxides. Its behavior on these oxides is quite different from its behavior on the metals. Adsorption of CO on noble metals is much stronger than on oxides. CO is adsorbed dissociatively on transition metals at the left of the Periodic Table. Dissociative adsorption of CO on oxides has not been reported. Instead, oxidation of CO to C02 or carbonate occurs.

Nondissociative adsorption of CO usually occurs on a surface coordinatively unsaturated cation. A characteristic feature of this CO is that the species has a higher infrared absorption frequency than the gas phase frequency of 2143 cm-'. A value between 2170 amd 2200 cm-' is observed [51]. In constrast, a shift to lower frequencies is observed on metals. The phenomenon is first noticed on NiO [52]. Explanations of this shift to higher frequencies on oxides have been based on the assumption that the adsorption involves an interaction between a surface cation and the carbon end of the CO molecule. This assumption has now been substantiated by single crystal studies on ZnO [4-61. In this picture, the C=O stretching frequency depends on the relative extent of a-donation from CO to the cation, which would increase the C-0 bond strength and its frequency, versus back x- bonding from the cation to the CO molecule, which has the opposite effect. It is argued that for cations in a high oxidation state, the cation size is small and its electron affinity is high. Thus the bonding with CO involves mainly a-donation and little x-bonding, which results in an increase in the CO stretching frequency. The importance of x-bonding increases when the cation is in a lower oxidation state. It has been proposed that the frequency for W'CO is above 2170 cm-', M'CO in the region of 2120-2160 cm-', and M"C0 below 2100 cm-' [53].

In addition to the molecular orbital picture of the bonding, the electric field near an oxide surface could also affect the CO stretching frequency. This effect has bccn analyzed [54,55]. It is agreed that the interaction of the surface electric field with the dipole moment of the CO molecule changes the interatomic distance in the molecule, and thus its force constant and the stretching frequency. However, quantitative predictions are not yet available.

The mode of adsorption of CO depends on the pretreatment of an oxidc. On a fully oxidized surface, CO may adsorb reductively to form carbonates [57,58]. CO may also react with surface hydroxyl groups to form formate [59].

The frequency of the CO band can also be affected by coadsorbates. When CO and COz are coadsorbed on ZnO [56], the CO frequency is shifted from 2183 to 2212 cm-'. The stretching frequency also depends on the pretreatment. ZnO evacuated at 400°C followed by an oxygen treatment yields a CO band at 2212 cm-*. If the sample is evacuated at 400°C, the band is shifted to 2187 cm-' [60]. The reasons for these shifts are not well understood.

NO NO is another probe molecule commonly used to study transition metal

oxides. Except for one additional unpaired electron, it has an electronic structure similar to CO. Thus in many ways its adsorption is similar to that of CO, and occurs on surface M"'(cus) ions, although its adsorption is often stronger.

68

Table 4-1 Infrared Absorption Peak Positions of Dinitrosyl or Dimeric Species of Adsorbed NO

Cation Peak Positions, cm-' Peak Separation, cm-1

W+, Cr3+ 1745- 1775, 1865- 1895 120 - 130 reduced MOO, 1695-1713, 1800-1817 100

Fe2+ 1810, 1910 100

reduced WO, 1685, 1795 110 co2+ 1765- 1795, 1840- 1875 80 - 90

NO may be adsorbed in one of the four different forms shown below [51]:

Adsorption as neutral species either as a single molecule or as a pair is the predominant mode observed. In a few cases, charged adsorbed species are formed. The infrared stretching frequency of adsorbed NO- is lower than the others, being in the region below 1735 cm-'. The frequency for the othcrs are mostly in the region 1740-1950 cm-'. In addition to the above species, oxidation of NO occurs occasionally to form surface nitrate or nitrite species [61-631.

Adsorption in pairs on one M"+(cus) ion is characteristic of NO adsorption, and it shows in infrared spectroscopy as a pair of absorption bands due to the symmetric and antisymmetric stretches of the adsorbed NO pair. The absorption frcqucncies of the pair dcpend on the oxide as well as on the particular sample of a given oxide. However, the separation between the two peaks is constant for a given cation and varies less from sample to sample. These separations are listed in Table 4-1 which is derived from the data summarized in ref. 51. The high tendcncy for NO to be adsorbed in pairs is due to the stabilization derived from the mutual interaction of the unpaired electron on each NO [64]. Whether the pair exists in the form of a dimcr or dinitrosyl is still unclear.

Dinitrosyl Dimcr


The arguments favoring a dinitrosyl species include the absence of absorption in the 1250-1350 cm? region, which excludes the presence of a hyponitrite species (N202-) [&I], the fact that the adsorbed NO pair is stable to quite a high temperature (on Mo, W, and Cr cations), and the two nitrosyl bands have approximately the same intensity [65].

Other evidence supports a dimeric form for the adsorbed NO pair. It has been observed that the ratio of the symmetric and asymmetric stretching from a N202 dimer is very similar to the coupled nitrosyl bands on a silica-supported chromia [66]. When a reduced molybdena is first half-saturated with 15N0 and then exposed to 14N0, no infrared bands attributable to a mixed l 5 N W 4 N O complex is observed [67,68]. Since isotopic mixing is expected for dinitrosyl complexes, this observation supports a dimeric species. However, it may also be explained (though less likely) with a dinitrosyl complex if the adsorption sites differ greatly in their binding energy for NO or in their accessibility. It should be noted that the ratio of the intensities of the asymmetric and symmetric stretch of adsorbed NO pairs is related to the angle between the two M-0 bonds [69].

It has been suggested that linear M-NO groups possess infrared absorption frequencies higher than 1850 cmpl [70]. Whether a linear group is formed depends on the degree of coordinative unsaturation of the metal cation. It appears that a F e 2 + 4 0 complex is more linear if the Fe2+ ion is less coordinatively unsaturated than more unsaturated [71].

REFERENCES

1. N. D. Parkyns, in "Chemisorption and Catalysis," ed. by P. Hepple, Elsevier

2. R. L. Burwell, Jr., G . L. Haller, K. C. Taylor, and J. F. Read, Adv. Cafaf. . 29,

3. M. Nagao. and T. Morimoto, J. Phys. Chem., 84, 2054 (1984). 4. R. R. Gay, M. H. Nadine, V. E. Henrich, H. J. Zeiger. and E. I. Solomon,

5 . K. L. DAmico, F. R. McFeely, and E. I. Solomon, J. Amer. Chem. Soc.,

6. K. L. DAmico, M. Trenary, N. D. Shim, E. I. Solomon, and F. R. McFeely.

7. A. Zecchina, E. Garrone, and E. Gulielminotti, Cdalysis (London), 6, 90 (1983). 8. B. Rebenstorf' and R. Larsson, Z. Anorg. Allg. Chem.. 453, 127 (1979). 9. V.A. Shvets, Russ. Chem. Rev., 55, 200 (1986).

Publ. Co., Amsterdam, Netherland, 1971, p. 150.

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12. G. Connell, and J. A. Dumesic, J. Cafal. , 101, 103 (1986). 13. S. Yuen. Y. Chen, J. E. Kubsh. 1. A. Dumcsic. N. Tops&. and H. Tops&.

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18. C. C. Chang, and R. J. Kokes, J. A m r . Chem. Soc., 93, 7107 (1971). 19. R. P. Eischens, W. A. Plisken, and M. J. D. Low, J. Catal., 1. 180 (1962). 20. S. Naito, H. Shimizu, E. Hagiwara, T. Onishi. and K. Tamaru, Trans. Faraday Soc.,

21. F. Boccuzzi, E. Garrone, A. Zecchina, A. Bossi. and M. Camia. J. Catal.,

22. A. L. Dent, and R. J. Kokes, J. Phys. Chem., 73, 3772, 3781 (1969). 23. A. L. Dent, and R. J. Kokes, J. Amer. Chem. Soc., 92. 6709 (1970). 24. C. C. Chang, W. C. Connor, and R. J. Kokes. J. Phys. Chem., 77, 1957 (1973). 25. T. T. Nguyen, and N. Sheppard, J. Chem. Soc., Chem. Commun., 868 (1978). 26. A. A. Efremov, and A. A. Davydov, Kinet. Catal., 21. 383 (1980). 27. A. A. Davydov, J. Tichy, and A. A. Efremov. React. Kinet. Catal. Lett., 5 ,

28. V. G. Mikhal’chenko, V. D. Sokolovskii. A. A. Filippova, and A. A. Davydov,

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31. A. A. Efremov, and A. A. Davydov, React. Kinet. Catal. Lett., 15, 327 (1980). 32. E. V. Rozhkova, S. V. Gorej, and Ya. B. Gorokhovatskii, Kinet. Katal., 15,

33. R. J. Kokes. Intra-Science Chem. Rep., 6, 77 (1972). 34. J. M. Vohs, and M. A. Barteau, J. Phys. Chem., 91, 4766 (1987). 35. C. C. Chang, and R. J. Kokes. J. Catal., 28, 92 (1973). 36. C. C. Chang, and R. J. Kokes, J. A m r . Chem. Soc., 92, 7517 (1970). 37. J. Saussey, and J. C. Lavalley, J. Chim. Phys., 75, 506 (1978). 38. T. T. Nguyen. J. C. Lavalley, J. Saussey, and N. Sheppard, J. Catal., 61, 503 (1980). 39. T. T. Nguyen, J. Catal.. 61. 515 (1980). 40. A. Ueno, T. Onishi, and K. Tamaru, Trans. Faraday Soc., 67. 3585 (1971). 41. R. P. Groff, J. Catal., 86, 215 (1984). 42. M. Ito, Vib. Surf. (Proc. Intern. Cod.), 2nd, 1980 (published 1982), p. 71. 43. K. Aika. and J. H. Lunsford, J . Phys. Chem., 81, 1393 (1977). 44. M. Nagao, and T. Morimoto. J. Phys. Chem., 84. 2054 (1984). 45. 0. Koga, T. Onishi, and K. Tamaru, J. Chem. Soc. Faraday Trans. I , 76, 19 (1980). 46. W. Hirschwald. and D. Hoffmann, Surface Sci., 140, 415 (1984). 47. S. Akhter, W.H. Cheng, K. Lui, and H. H. Kung, J. Catal.. 85, 437 (1984);

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56. J. C. Lavalley, J. Saussey, and T. Rais, J . Molec. Catal., 17, 289 (1982). 57. P. G. Harrison, and E. W. White, J . Chem. SOC. Faraday Tram. I , 74, 2703 (1978). 58. E. Guglielminotti, L. Cermti, and E. Borello, Gazz. Chim. Ital., 107, 503 (1977). 59. M. He, and J. G. Eckerdt, J . Cafal., 72, 303 (1981). 60. C. H. Amberg, and D. A. Seanor, Proc. 3rd Intern. Cong. Catal., Amsterdam,

61. P. G. Harrison, and E. W. White, J . Chem. SOC. Faraday Tram I , 74, 2703 (1978). 62. G. Busca, and V. Lorenzelli, J . Cafal., 72, 303 (1981). 63. J. W. London. and A. T. Bell, J . Catal., 31, 32 (1973). 64. A. Zecchina, E. Garrone, C. Morterra, and S. Coluccia, J . Phys. Chem.,

65. A. Kazusaka, and R.F. Howe, J . Cafal., 63, 447 (1980). 66. E. L. Kugler. R. J. Kokes, and J. W. Gryder. J . Catal., 36, 142 (1975). 67. W. S. Millman, and W. K. Hall, J . Phys. Chem., 83, 427, (1979). 68. J. B. Peri, J . Phys. Chem., 86. 1615 (1982). 69. F. A. Cotton, and G. Wilkinson, "Advanced Inorganic Chemistry". 3rd ed.,

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Chapter 5

SURFACE ACIDITY

5.1 SURFACE ACID SITES

A partially hydroxylated oxide surface has hydroxyl groups and coordinatively unsaturated metal cations and oxygen anions. Each of these species can participate in an acid-base reaction. Exposed coordinatively unsaturated cations may act as acceptors for free electron pairs of adsorbed molecules.

M"+(cus) + :B(g) + M"+:B (5-1)

Such cations are Lewis acid sites. The strength of these acid sites depends on the charge and size of the cations, both of which may vary with the oxidation number of the cation. In general, according to the concept of hard and soft acids [ l ] , cations of a higher oxidation state are harder. For cations in the same group in the Periodic Table and of the same oxidation state, those in a later period are softer. Harder cations are smaller and less polarizable. They will adsorb or bind hard bases stronger than soft or polarizable bases. Complicating these considerations, however, is the fact that cations in oxides are usually surrounded by larger and more polarizable oxygen anions. The harder, smaller and less polarizable cations are sometimes partially shielded by the oxygen anions so that binding of molecules to the cations is sterically hindered. As a result, binding is weaker than expcctcd.

Surface hydroxyl groups may act as BrBnsted acid sites. They may dissociate to protonate adsorbed bases:

The resulting conjugate acids and bases are stabilized on the surface by electrostatic interaction with each other and with the oxide. As will be discussed later, the oxide exerts an effect similar to that of an aqueous solution in slabilizing charges on the surface.

72

SURFACE ACIDITY 73

Table 5-1 Relative Acidities on Oxide Surfaces a

relative acidity order on M acidb D@(B-H)' ZnO MgO PKil kcal/mol kcal/mol

1. C~HSSH 2. HCOOH 3.7 3. CH3COOH 1. CH3COOH 4.8 4. HCN 9.3 5. C ~ H S O H 2. C6H50H 9.9 6. CH3SH 3. CH3SH 12.0

7 . f H 3 0 H 4. CH30H 15.5 C ~ H S O H 5. C ~ H S O H 17.0

9. C~HSCCH 6. C~HSCCH 18.5 7. HCCH 26

33 1.8 345.2 348.5 353.1 351.4 359.0 379.2 376.1 370.3 375.4

83.3 106 105.8 123.8 86.5 90.7 104.4 104.2

132

Footnotes: a) From R. Spitz, J. Barton, M. Barteau, R. Staley, and A. sleight, J. Phys. Chem.

90, 4067 (1986), copyright American Chemical Society. b, Homolytic band dissociation energy. From J.E. Bartmess. R.T. McIver, Jr., in "Gas

Phase Ion Chemistry", M.T. Bowers, ed., Academic Press, NY 1979, p. 87. ') Hcat of heterolytic dissociation. From D.F. McMillen, D.M. Golden, Ann. Rev.

Phys. Chem. 33, 493 (1982).

An exposed coordinatively unsaturated oxygen ion participates in an acid- basc reaction as a conjugate Brbnsted base. It is one element of the pair sites for hctcrolytic dissociative adsorption (see eq. 4-2). The heterolytic dissociative adsorption of water, hydrogen, alcohols and alkynes described in Chapter 4 are cxamplcs where the surface oxygen ions act as conjugate bases. The dissociative adsorption of bcnzaldchydc and chloroform are other examples. In general, transition metal oxides do not have strongly basic sites.

When considering the BrBnstcd acidity of surface hydroxyl groups or basicity of surface oxygen ions, it is important to know whether it is more appropriate to use the aqueous solution acidity scale or the gas phase acidity scale. The charged species are primarily stabilized by the dipoles of water molecules in an aqueous solution, and by other polarizable spccies in other solutions. In the gas phase, they arc stabilized by induction within the molecule. Thus molecules may show different strengths of acidity and basicity in different media.

While it has long bccn conjectured that the aqueous solution scale is more appropriate for oxidc surfaces, this has only been demonstrated recently [2]. In this study, the adsorption of a series of organic acids that include thiols, alcohols, acids and alkyncs was studicd (see Table 5-1). The relative strength of adsorption

74

of these molecules was monitored by displacement/ titration experiments. In these experiments, one molecule was first adsorbed. After evacuation, the oxide was exposed to a second molecule. The displacement of the first molecule by the second molecule was followed with FTIR and the extent of displacement was evaluated from the IR peak intensities. Sometimes deuterated molecules were used to avoid overlapping peaks. Although this method is not an equilibrium method and quantitative values of relative acidity could not be obtained, qualitative ordering could be derived. It was shown that the titration of acids of similar strengths is completely reversible, but is irreversible for acids of very different strengths. The results of this study are shown in Table 5-1. The orders of the strength of adsorption on both ZnO and MgO are the same and follows the same order as aqueous pK, values. The order does not follow the acid strengths in the gas phase as measured by the enthalpies of heterolytic dissociation of the probe molecules, or the bond dissociation energies of these molecules. It may be concluded that the oxide must function llke an aqueous solution to stabilize the electric charge on the dissociated molecules.

Although Table 5-1 shows the results of an experiment where the oxide ion of the solid serves as the conjugate base, the same phenomenon probably applies to the situation in which the surface hydroxyl groups act as Brdnsted acids, since the same charge stabilization mechanism is required. Therefore, it is probable that the order of acid strengths of surface hydroxyl groups on different oxides determined either by protonation of adsorbed molecules in a nonaqueous or gaseous environment or by the isoelectric points in aqueous solutions will be the same. However, quantitative comparisons between the two methods will probably be impossible until accurate calculations to evaluate charge stabiliiation become available. Complicating the matter is the fact that the acid strength of a hydroxyl group may depend on its environment including the presence of ion vacancies, impurities, and surface dislocation defects. The importance of surface impurities has been shown to be a critical factor in the variation of the isoclectric point of a particular oxide reported in the literature.

5.2 FORMATION OF ACID SITES

Exposed coordinatively unsaturated metal cations and oxygen anions on the surface are Lewis acid and Brdnsted conjugate base sites, as described above. Brdnsted acid sites are present only when hydroxyl groups are present. Therefore the number of such sites depends on the extent of hydroxylation of the surface.

As mentioned in chapter 4, hydroxylation and dehydroxylation involve the interaction between a water molecule and surface coordinatively unsaturated metal cations and oxygen anions:

SURFACE ACIDITY 75

It may seem appropriate to relate the appearance of Lewis acidity with dehydroxylation and disappearance of BrQnsted acidity, and vice versa. In some cases this can be readily demonstrated. On a ZnO sample that is fully hydroxylated, significant amounts of BrQnsted acid sites are present, as indicated by the formation of IW,+ ions on adsorption of ammonia. But when the hydroxyl density is less than half of the fully hydroxylated surfaces, hardly any w+ is observed [3]. On alumina-supported molybdena, chromia, rhenium oxide, or tungsten oxide, adsorption of pyridine results in the formation of pyridinium ions the amounts of which increases after exposing the oxides to water [4]. A similar observation has been made using the adsorption of ammonia [ 5 ] .

In some other cases this interconversion may not be demonstrated so readily. This may be due to the fact that the surface oxygen anion is very strongly basic such that its protonated form is only a very weak BrQnsted acid - so weak that it protonates only strongly basic molecules. It is also possible that in the dehydroxylated form, the surface cation behaves like a very weak Lewis acid center because of its nature or because of steric shielding by the neighboring oxygen ions. Then dehydroxylation would not result in the development of readily detectable Lewis acidity. For example, anatase TiOz prepared from titanium tetraisopropoxide possesses Lewis acid sites which are not converted by the addition of water to Brdnsted acid sites strong enough to protonate pyridine. However, Lewis acid sites on anatase prepared by the hydrolysis of titanium oxide sulfate can be converted to Brdnsted acid sites on exposure to water vapor [6] .

Since surface coordinatively unsaturated ions are important, it follows that the presence of cation or anion vacancies and other defects that results in greater exposure of the ions also affects acidity. However, very little understanding of this matter is available.

One way to generate new and perhaps stronger acid sites on an oxide is by incorporation of a second oxide. The mixed oxide of silica-alumina is a classic example where a high density of new strong acid sites is generated. This situation is found in many other mixed oxides, as is shown in Table 5-2. Also shown in the table is the fact that not all mixed oxides develop new acid sites.

Various models have been proposed to predict the formation of new acid sites in mixed oxides. Tanabe's model applies to dilute mixed oxides where a small amount of a second oxide is incorporated into the first oxide by cation substitution [7-91. This model assumes that the generation of new acid sites is caused by an excess of negative or positive charge in a model structure of a binary oxide. The model structure is constructed as follows:

i) The coordination number of a cation in the component oxide is maintained in the binary oxide.

ii ) The coordination number of the oxygen ion in the binary oxide is the same as in the major component oxide.

As an example, the model structures of a dilute mixed oxide of TiOz and Si02 are shown in Figure 5-1. In TiOz, each Ti ion is coordinated to six oxygen ions and each oxygen ion to three Ti ions. In Si02, each Si ion is four-coordinated and each oxygen ion is two-coordinated. In the model structure, the coordination numbers of Ti and Si are maintained, and the coordination number of oxygen is that of the

76

Table 5-2 Formation of New Acid Sites in Some Mixed Metal Oxides. (From J. Solid State Chem., 52, 191 (1984), copyright Academic Press).

New Acid Sitesb Matrix Substituting Oxide Oxide Experimenta Tanabe's Kung's

Model Model

za2 CdO wo3

+ -

+ +

+ +

+ -

-

+

+ +

+

? ?

+ ?

Footnotes: a) Experimental data extracted from ref. 9 except for those involving WO,

b, + means affirmative, - means negative, ? means dependent on the conditions. which arc taken from Yamaguchi, et al., J. Catal., 65, 442 (1980).

SURFACE ACIDITY 77

- I 1 7 - 1

- 0 - I I

-0 -0 'I / T-

(a) Si in T i02

I 0 I

I O 0' I /

(b) Ti in Si02

Figure 5-1 Structures used in Tanabe's model for the prediction of formation of new acid sites in mixed oxides. a. Si in a matrix of Ti02; b. Ti in a matrix of SiOz. (From Bull. Chem. SOC. Jpn., 47, 1064 (1974). copyright Chemical Society of Japan).

major component. The formal charge of each ion is then assumed to be evenly distributed over the coordinating bonds. In figure 5-la, the +4 charge of the substituting Si ion is distributed over four bonds, while the -2 charge of an oxygen ion is distributed over three bonds. Thus each of the bonds surrounding Si bears a net charge of 4/4 ~ 2/3 (=+1/3). The excess charge at Si is then 4 x 1/3 =+4/3. In this case, a Lewis acid site is assumed to appear because of the presence of an excess positive charge. In Fig 5-lb, a similar calculation results in an excess charge of -2 at the substituting Ti ion site. In this case, a Brbnsted acid site is assumed to appear, because two protons are assumed to be associated with the site to maintain charge neutrality. If there is no excess charge at the site of the substituting ion by such a calculation, as in the case of AI2O3-Bi2O3, there will be no new acid sites.

This model has been applied to many binary oxides, some of which are shown in Table 5-2. It has been found that the prediction is accurate over 90% of the time. The high success rate makes the model very useful, although it is limited by thc assumptions uscd. One limitation is the need to have a unique coordination number, which may be difficult to decide in systems of low symmetry. Another limitation is the use of formal oxidation states which may be quite different from thc real chargc (see Chapter 3). Since electron deficiency at a site refers to real charge at a site, the use of formal oxidation states may not be accurate. Finally the

78

-8 0 r

5 p -4 E z -2-

-6-

U .- 2 0 -

tl 2 -

= 4 -

w

c P, .-

6 -

0. r, 0

-

0.0

- .)a

0 0 . .

0 . 0

0

0

I I I I I

Averaged Electronegativi t y

Figure 5-2 Highest acid strengths and average electronegativities of metal ions of binary oxides (molar ratio = 1) . (From Bull. Chem. SOC. Jpn., 46, 2985 (1973). copyright Chemical Society of Japan).

model cannot predict acid strength. However, there appars to be a rough correlation between electronegativity of the cations and the strength of acid sites in many mixed oxides, as is shown in Fig. 5-2 [lo].

Another model proposed by Kung takes a rather different approach. In Tanabe's model, charge compensation at the substituting ion site by neighboring oxygen ions is important. In Kung's model, changes in electrostatic potential experienced by the substituting cation due to all the ions in the matrix oxide is important. Thus Tanabe's model is a localized model, and Kung's model i s a delocalized model.

In Kung's model [ l 11, the difference AV between the electrostatic potentials experienced by a cation A in a matrix BO, and in AO, is given by:

(5-3)

where qi is the charge of the ion at a distance ri from the A cation. The subscripts BO and A 0 denote the matrices BO, and AO,, respectively. When AV is negative, cation A in matrix BO, experiences a more negative potential than in

SURFACE ACIDITY 79

AO,. It will be electrostatically more stable. Therefore the electron energy levels of cation A arc lower in energy in matrix BO, and the cation can accept electrons inorc rcadily. It will act as a new Lewis acid site. When AV is positive, A is less rcadily in accepting electrons in matrix BO, than in matrix AO,. No new Lewis acid site is generated at the substituting A cation.

In addition to the changes in the electrostatic potential at site A, when the oxidation states of cations A and B are different, the overall charge neutrality of the solid will be maintaincd by a change in the matrix. Two possibilities exist: a substituting cation A is of a lower formal oxidation state than the matrix cation B, y < z, and the reverse case of A being of a higher formal oxidation state than B, y >

When y < z, a simple substitution of B ion by A ion would result in a solid with excess oxygen. This excess can be balanced by: (1) development of anion vacancies; (2) adsorption of protons on the surface; or (3) development of intcrstitial cation defects. (1) and (2) are intimately related if the solid is prepared by aqueous precipitation such that the surface is hydroxylated. If the surface stays hydroxylated, the protons present to balance the excess oxygen will act as new Brbnstcd acid sites not present in the component oxides. On heating, some of these protons may be removed as water which is formed with the concurrent formation of an anion vacancy. The anion vacancy site is a Lewis acid site, and could act as a BrQnsted acid site after it had been hydrated. The effect of (3) is less easily prcdictcd. It could bc important for solids that have open structures. In most common binary oxides, however, the concentration of interstitial defects is limited. Whether these defects lead to the formation of acid sites will likely depend on the nature of the cation.

When y > z, a simple substitution of B ion by A ion would result in a deficiency of oxygen. The deficiency can be removed by: (1) adsorption of ncgativcly charged oxygen species onto the A ion; (2) adsorption of OH- onto the A ion; or (3) formation of cation vacancies. When (1) or (2) operates, the consequence of electrostatic potential change at the A ion is removed because the coordinative unsaturation of A is removed by the adsorbed oxygen or OH-. Since it is not likely that adsorbed oxygen acts as an acid site, new acidity is predicted not to dcvclop. Adsorbed OH- could provide Brdnsted acidity, but the acidity would be weak. When (3) operates, new Lewis acid sites could appear because cation vacancies are clcctron deficient.

A comparison of experimental data and the predicted formation of new acid sitcs in binary oxides by Kung's model is summarized in Table 5-2.

It is interesting to note that although Tanabe's and Kung's models employ very diffcrent approaches, either method gives the same prediction in many cases. This is bccause Tanabc's model can be mathematically expressed as examining the diffcrence between the qF/c value of the substituting ion and the surrounding oxygcn ion, where qp is thc formal oxidation state, and c is the coordination nunibcr. Although the stoichiometry of an oxide and the qF/c ratio do not have a fixcd rclationship, this ratio is the same for most oxides of the same stoichiometry. Sincc the latticc potentials ( i s . , C q,/r,) are about the same for oxides of the same stoichiomctry and arc diffcrcnt from other stoichiomctries, there is a correlation

L.

80

between the difference in the qF/c values and the value of AV of eq. (5-3). Thus the two models yield similar conclusions in most cases [111. Although they may yield similar results, they involve important differences. Tanabe's model is a localized model such that any new acid sites formed are at the substituting cation site. Kung's model is a delocalized model. New acid sites can be formed on the matrix surface far away from the substitution site as well as at the site.

The two models described above apply to dilute oxide solid solutions. That is, isolated "impurity" cations occupy sites of the host oxide substitutionally. Cations deposited on an oxide surface but not incorporated substitutionally may also form new acid sites. A rather detailed study of this latter system has been reported for Fez+ and Fe3+ ions deposited on MgO, TiO2. A1203 and SO2, and Zn2+, Ga3+, Fe3+, M3+, Fez+, Sc3+, and Mg2+ deposited on Si02. Mossbauer spectroscopy and pyridine adsorption have been used to characterize these samples [12-141. The dependence of the amounts of acid sites in these systems on the pretreatment temperature indicates that BrQnsted acid sites may be removed by high temperature evacuation of the oxide, that is, dehydroxylation of the oxide. Lewis acid sites are present on samples pretreated by evacuation at low temperatures, and they persist to much higher temperatures.

A model has been proposed by Connell and Dumesic [13,14] to explain the formation of Lewis acid sites for ions deposited on Si02. The model assumes that the formal charge on the deposited cation is balanced by the coordinating surface lattice oxygen ions, and that a cation is coordinatively saturated if' it has a coordination number four as Si ions have in SO2. If the number of coordinating oxygen ions is less than four, the deposited cation may be coordinatively unsaturated and act as a Lewis acid site. It is further assumed that since oxygen ions in the Si02 surface have a coordination number of two, the formal charge of oxygen along each cation to anion bond is negative one. A +2 cation on the Si02 surface requires two coordinating surface oxygen for charge neutrality. For a +3 cation, this number is three. In both cases, the number of coordinating surface oxygen is less than four, and Lewis acidity is expected to develop at these cation sites. This model successfully explains the data presented. In addition to the requirement of coordinative unsaturation of the deposited cation, the formation of a new Lewis acid site also requires that the matrix oxide is not basic. Finally, the strength of the acid site is determined by the electronegativity of the cation: more electronegative cations result in stronger acid sites.

5.3 DETERMINATION OF ACIDITY

Since surface acidity is an important property that often determines the surface chemistry, various methods have been developed to mcasure its presence and strength. The most commonly used method involves spectroscopic investigation of adsorbed probc molecules. This method is perhaps experimentally the most convenient. With the much improved spectroscopic techniques available today, rapid determination of the presence of surface acid sites is possible. The sccond method employs the dctcrmination of

Three methods will be described.

SURFACE ACIDITY 81

isoclcctric point and the third method, titration with indicators.

Adsorption of Ammonia and Pyridine Infrared spectroscopic investigation of adsorbed probe molecules has been

employed routinely to characterize various oxides [15]. The well developed technique of spectral subtraction, made convenient with computers and Fourier transform infrared spectroscopy, has made possible the investigation of colored samples including iron oxide, and samples of low surface areas. Because this is a dry tcchnique, there is a wide latitude in the pretreatment of the sample, and thus it is convenient to study the effect of sample pretreatment on the amount and strength of surface acidic sites.

The most common probe molecules are pyridine and NH3, although other amines are sometimes used. NH3 and pyridine are bonded to the surface in three different modes. In the first mode, the molecules are adsorbed abstractively: it is protonated by a proton from a surface hydroxyl group. Therefore, it probes surface Brdnsted acid sites.

C ~ H ~ N + HM -+ C ~ H ~ N H + + 0-4 (5-5)

In the second mode, the electron lone pair of the nitrogen atom adds to the (cus) cation of the oxide which acts as a Lewis acid. Such bonding involves o-donation and requires an exposed surface coordinatively unsaturated cation. Hydrogen bonding is the third mode. It is the weakest mode of interaction. and has not been uscd to measure the acidity of surface hydroxyl groups, although it can be done in principle.

The presence of BrBnsted acidity is revealed by the formation of protonated ammonium or pyridinium ions. These ions possess characteristic vibrational bands. Likewise, ammonia or pyridine molecules bonded to Lewis acid sitcs or hydrogen-bonded to the surface also possess characteristic bands. The characteristic band for a N&+ ion is located at 1400-1480 cm-' [16]. As a rcfercnce, a detailed table prepared by Tsyganeko [ 171 of NH3 infrared absorption bands in solutions of various proton-accepting and proton-donating solvents can be uscd to help identify the bands. NH3 bonded via the electron lone pair at the nitrogcn atom to a Lewis acid site possesses an asymmetric and a symmetric stretching frequency of v,, = 3330-3380 cm-' (compared with 3444 cm-' for a molccule in the gas phase) and v, = 3260-3280 cm-' (3336 cm-' in the gas phase) [ 31. Additional hydrogcn-bonding intcraction with neighboring anions would lower these frcqucnces [ 181. Whcn hydrogen-bonding to a surface hydroxyl group is the only interaction, the NH3 infrared bands appear near 3400 and 3200 cm-' [3].

For pyridinc, according to Knozinger [ 191, the ring-vibration modes ( 1 9b and 8a modes) arc most affected by the nature of intermolecular interaction via the nitrogcn atom. These two modes are observed at 1440-1447 and 1580-1600 cm-', rcspectively, for hydrogcn-bonded pyridine, at 1535-1550 and about 1640 cm-' for a pyridinium ion, and at 1447-1464 and 1600-1634 cm-' for pyridine

82

coordinatively bonded to Lewis acid sites. Although a general classification of the mode of bonding is readily possible,

it has nor becn possible using this method to obtain quantitative information on the strengths of the acid sites as measured by the proton affinities of the conjugate bases formcd or the pK,'s of the acid sites. Attempts to determine the strength of adsorption by temperature programmed desorption of adsorbed NH3 or pyridine often fail because of oxidation of these molecules by the transition metal oxide. It has not been possible either to correlate the vibrational band frcqucncics with the nature of the metal ion or its coordination [20,21].

Athough ammonia and pyridine have been used to probe similar surface propcrties, they are not identical. One difference is their sizes. The molccular cross-sectional area of NH3 is 0.127 nm2, and of pyridine is 0.313 nm2. Thus it may be possible that more ammonia is adsorbed than pyridine. The larger size of pyridine may also result in weaker bonding to a Lewis acid cation if steric crowding is a factor, especially if the cation Lewis site is recessed into the surfacc. This latter effect has been used to explain the smaller amount of pyridine adsorbed than NH3 on Cr2O3 [ 181.

Another difference betwecn the two molecules is their relative basic strcngth. In an aqueous solution, ammonia is a stronger base with a pKb of around 9 compared to a pKb of about 5 for pyridine [221. However, the basicity of pyridine in the gas phase is significantly higher [23,24]. It has been pointed out in section 5.1 that the polarizable surface oxygen ions make molecules on an oxide surfacc behave as if they were in an aqueous solution. Thus the rclativc basicity in aqueous solutions should be more appropriate to describe the relative Brdnstcd basicity of these two molecules. Indeed, on the oxidized and the reduced molybdena/alumina, N h + ions are detected. On the other hand, pyridinium ions are dctected only on the oxidized form [25]. The relative Lewis acidity is different. For example, pyridine is desorbed more slowly than ammonia from thc Lewis acid sites of F%03, which suggests that it is adsorbed morc strongly than

Most of the first row transition metal oxides and a fcw of the othcrs have bccn studicd with these two moleculcs. As shown in Table 5-3, Lewis acid sitcs arc found on practically all of the oxides studied that have been prctreatcd by hcating to 400°C or higher in vacuo. On the other hand, Brbnstcd acid sitcs are found only on V205, Nb205, Moo3, W03 , Re207, and Cr2O3/AI2O3 [ l S ] .

The density of acid sites dcpends on the sample prctreatmcnt. Figurc 5-3 shows an example of this effect on Nb205 as dctcrmined by pyridinc adsorption [27]. The sample dried at 100°C possesses a rather large amount of Brdnstcd and Lewis acid sitcs. Hcating Ihe sample to 300°C in vacuo results in a decrcasc in the numbcr of BrBnsted acid sites but an increasc in Lewis acid sitcs. Thus thcrc is an intcrconversion upon dehydration. Further heating to 500°C in vacuo rcsults in sharp dccreascs in both typcs of acid sitcs, which are due mostly to a rapid loss in surface arca. The cffcct of evacuation aftcr pyridinc adsorption is also shown in the figurc. As expcctcd, a highcr evacuation temperature rcsults in smallcr amounts of pyridinc adsorbed. I t is interesting to note that thc dccrcasc in the amount of pyridinc adsorbcd is qiiitc gradual with the incrcasc in evacuation

NH3 [26].

SURFACE ACIDITY 83

Table 5-3 Acid Sites on Various Transition Metal Oxides Determined by Ammonia or Pyridine Adsorption

Mode of Bonding Oxidea H-bonding Lewis Acid Brbnsted Acid

+ +

+

+

+ + + + + + + + + + + + +

+ +

Footnote: a) Pretreated by evacuation at 400°C or higher.

temperature. This suggests that there is a broad range of strength of interaction of pyridine with the surface, which implies the presence of sites of a broad range of acid strength.

Isoelectric Point A second method to determine the presence of acid sites is by determining the

isoelecuic point of an oxide. This method makes use of the phenomenon that solid oxide particles in aqueous suspensions are often electrically charged. Charged particles are formed when there is an imbalance between the densities of adsorbed H+, OH-, and ionized surface OH groups. They may also be formed by adsorption of metal hydroxo complexes derived from the hydrolysis products of material dissolved from the solid, i.e., [Mz+(OH),,]zA species. The presence of a net charge on the oxide particle can be observed in electrophoresis experiments. The magnitude of the surface charge depends on the pH of the solution, and there exists a pH at which there is no net surface charge. This pH is called the isoelectric point, It is equivalent to h e point of zero charge, and is the pH at which suspended particles in H 2 0 do not move in an electric field.

Since the oxide particles arc in an aqueous solution during the determination of isoelccuic point, much of the variation of surface acidity due to different extents

84

0 200 400 600

Pretreatment Temp. C

Figure 5-3 Relative amounts of a: Lewis acid; b: Brdnsted acid sites of a sample of Nb2OS.nHzO as a function of pretreatment temperature as determined by the IR absorption peak intensity of adsorbed pyridine after evacuation at: 1, room temperature; 2. 100°C; 3, 200°C; and 4, 300°C. The surface areas of the sample after pretreatments at various temperatures are : lOO"C, 164 m2/g; 300"C, 126 m2/g; and 500°C. 42 m2/g. (From Bull. Chem. SOC. Jpn., 56, 2927 (1983), copyright Chemical Society of Japan).

of hydroxylation cannot be studied. However, the isoelectric point is still a rather sensitive function of the history of the sample. This is because the density of charges on any oxide surface is quite low. Thus the surface charge could be significantly altered by the presence of impurity cations or anions on the surface, as well as surface imperfection such as vacancies and dislocations. Its value could also depend on the nature of the ions in the aqueous phase that are used to vary the pH due to the possibility of specific adsorption of the ions. These factors have led to wide variations in the reported values of the isoelecuic points of oxides.

Isoelecuic points can be determined in an elcctrophoresis apparatus where the zeta-potential of a solid is measured as a function of the pH of the solution. Since zeta-potential is the potential between a charged surface and the electrolyte solution, its value is zero when the net charge of the surface is zero. Therefore, the pH at which the zeta-potential is zero is the isoelecuic point, assuming that there are no adsorption of charged species other than protons and hydroxide ions on the

SURFACE ACIDITY 85

Figure 5-4 Zeta-potential at 22.5'C as a function of pH. Curve 1: NiAl204 and CoA1204; curve 2: C0304; curve 3: A1203; and curve 4: NiO. (From J. Catal.. 83, 225 (1983). copyright Academic Press).

surface. Figure 5-4 shows some sample data for a number of oxides [28]. The isoelectric points for a large number of oxides have been determined.

Table 5-4 is an updated summary of the comprehensive data collected by Park [29,30]. From these data, Park proposes a generalization presented below based on the stoichiometry of the oxide:

Isoelectric point (IEP) M 2 0 IEP > pH 11.5 MO 8.5 <IEP < 12.5 M2°3 6.5 < IEP < 10.4

0 < IEP < 7.5 M205, M03 IEP < 0.5 MO2

Although the expcrimental values of isoelectric point dcpcnd on the hydration state, purity, and surface imperfection, they are fundamental properties of an individual oxide whcn thcsc complications are eliminatcd. It becomes possible to use isoclcctric points to determine tile fraction of each oxide exposed on the surface in a system containing two oxides. For example, the isoelectric point of Moog is 6.25, and of alumina is 8.80. It has been found that Moo3 supported on A1203 shows an IEP intermediate between these two values, and there exists a onc- to-one correlation between the value of IEP and the Mooj surface area detcrmincd by oxygen chemisorption [31]. I t has now been shown with Co [28], Mo [32], and V [33] oxides supported on alumina that a supportcd oxide often shows an IEP

86

Table 5-4 Isoelectric Points of Transition Metal Oxides and Hydroxides

7.0

7.3

7.6

11 f 0.2

9.5 f 0.4

12 ? 0.5 6.5 f 0.2 5.2 to 8.5 6.7 rt 0.2

4 to 4.5 7.0 6.2

8.3 to 10.3 11 to 12 4.5 to 6.3

1.4 <0.5

8.7 to 9.2

Table 5-5 Isoelectric Points of Oxides Supported on Alumina

Co-Alb pure

% Coa pure y-Al2o3 0.6 1.0 1.6 4.4 5.9 Co304 CoA1204 IEP 8.80 8.30 8.05 7.65 7.95 8.05 7.30 2.65

Ni-Alb

% Nia pure y-A1203 2.3 4.5 6.0 9.0 pure NiO NiAI2O4 IEP 8.80 8.15 8.00 7.80 7.60 8.30 2.65

MO-AI' .

% Moa pure y-AI2O3 1.1 4.1 8.8 11.4 13.8 pure Moo3 IEP 8.80 8.65 8.50 8.30 8.00 7.85 6.25

Footnotes: a) g oxide per 100 g y-Al203. b, From J. Catal., 83, 225 (1983). ") From Appl. Catal., 4, 371 (1982).

SURFACE ACIDITY

intermediate between the component oxides. Furthermore, the formation of the aluminates can be easily detected as they have very different isoelectric points from the component oxides. Table 5-5 shows the data for Co, Ni, and Mo supported on alumina. For supported vanadia, the EP varies with increasing vanadium content from 8.25 to 6.80 on a y-alumina support, and from 4.5 to 1.95 on a titania support. The values for pure alumina, vanadia, and titania are 8.80, 1.40, and 6.30, respectively. Among these examples, one may note that the NiO- A1203 system seems to be anomalous in that the IEPs of the mixed oxides fall outside the range bracketed by the component oxides. Nonetheless, there is a monotonic trend as observed in the other oxide systems.

Titration with Indicators A third method to determine surface acidity is by titration with indicators.

Conceptually the method involves first adsorbing onto the solid a particular indicator which changes color on protonation. It is important that the adsorption process is completed without exposure of the pretreated sample to moisture. Thus careful exclusion of air and an anhydrous organic solution of the indicator is used. If the surface possesses hydroxyl groups more acidic than the pK, of the indicator, the indicator becomes protonated and undergoes a change in color. The solid is then titrated with a strong base such as n-butylamine to deprotonate the indicator. The amount of n-butylamine needed to completely deprotonate the indicator, as indicated by a color change, is the amount of surface hydroxyl groups more acidic than the pK, of the indicator. By using different indicators of different pK,'s, the distribution of acid strength of an oxide can be determined.

This method has a number of experimental difficulties. The most severe one is the long time required to attain equilibrium between the surface and the indicator. During this equilibration time, extreme care is needed to prevent the oxide from exposure to air, especially moisture and COz. Another problem is that accurate pK, values of indicators that should be used under these experimental conditions are not available. Aqueous pK, values are often used as reasonable approximates.

This method has been extensively employed by Tanabe and coworkers [7,8]. The common indicators used include Methyl Red (pK,=+4.8), 4-dimethylamino- azobenzene (+3.3), 4-anilinoazobenzene (+1.5), Crystal Violet (+0.8), dicinnamylidene-acetone (-3.0), benzylideneacetophenone (-5.6), and anthraquinone (4.2) . Table 5-6 shows some exemplary data using this method. Among these examples, A1203, Ti02 and Zr02 all show a rather narrow distribution of acid strength. There is a sharp drop in the amount of acid sites of pK, more acidic than 3.3. On the other hand, mixed oxides show much broader distributions of acid strength. Most interestingly, new acid sites are developed in the mixed oxides that are much stronger than those found in the component oxides. The catalytic importance of strong acid sites has been well documented for silica- alumina. Acidic mixed oxides involving transition metals may eventually find technological applications.

88

Table 5-6 Distribution of the Amounts of Acid Sites of Different Acid Strengths a

Surface Area

Acid Amounts (mmoVg) at different pK,'s

Oxide mZ/g +4.8 +4.0 +3.3 +1.5 -3.0 -5.6 -8.2

A1203 190 - 0.285 0.075 0 Ti02 38.5 0.057 0.057 0 Zfl2 72.0 - 0.280 0.060 0.060 0 Ti02-Zr02b 230 - 0.475 0.380 0.350 0.375 0.125 0.050 TiO2-Al2O3; 204 0.422 0.422 0.337 0.250 0.220 0.060 0 Zfl2-Al203 320 - 0.590 0.205 0.205 0.045 0.045 0

Footnotes: a) Taken from Bull. Chem. SOC. Jpn., 46, 2985 (1973), copyright Chemical Society

b, The molar ratios are about 1:1. of Japan.

5.4 ROLE OF ACID SITES IN CATALYTIC REACTIONS

The importance of strong Brbnsted acid sites in hydrocarbon conversions has been well established [34]. Transition metal oxides are no exceptions. Those with strong Brbnsted acid sites catalyze cracking and skeletal isomerization of alkanes. These include Nb205 and TazO5 that are not readily reduced. They also catalyze carbon-carbon double bond isomerization and alcohol dehydration. These latter reactions may proceed by other mechanisms not involving acid sites as well. They are separately discussed in later chapters.

Perhaps the reactions most uniquely catalyzed by transition metal oxides are metathesis and selective oxidation reactions. The current mechanism of metathesis does not require acid sites. On the other hand, both Brujnsted and Lewis acid sites may participate in selective oxidation reactions. For example. it has been proposed that the oxidative cleavage of butcne to acetic acid and acetaldehyde proceeds via the formation of carbenium ions [35]:

SURFACE ACIDITY 89

The role of Lewis acid sites in the selective oxidation of alkenes has been suggested [36-381. The activation of the alkene molecule is assumed to be the adsorption onto a Lewis acid site. Details of these reactions are discussed in Chapters 11 and 12 on selective oxidation.

REFERENCES

1. R. G. Pearson, "Hard and Soft Acids and Bases," Dowden. Hutchinson, and Ross

2. R. Spitz, J. Burton, M. Barteau, R. Staley, and A. Sleight, J. Phys. Chem.. 90,

3. T. Morimoto, H. Yanai. and M. Nagao. J. Phys. Chem., 80,471 (1976). 4. K. Segawa, and W. Hall, J. Cutul., 76, 133 (1982). 5. R. Groff, J. Cutul., 86, 215 (1984). 6. T. Kantoh, and S. Okazaki, Bull. Chem. Soc. Jpn., 54.3259 (1981). 7. K. Tanabe, "Solid Acids and Bases, Their Catalytic Applications," Academic Press,

8. K. Tanabe, in "Catalysis Science and Technology", J. R. Anderson and M. Boudart ed.,

9. K. Tanabe, T. Sumiyoshi. K. Shibata, T. Kiyoura, and J. Kitagawa, Buff. Chem. Soc.

10. K. Shibata, T. Kiyoura, J. Kitagawa, T. Sumiyoshi, and K. Tanabe, Bull. Chem. Soc.

11. H. Kung, J. Solid Stute Chem., 52, 191 (1984). 12. G. Connell, and J. A. Dumesic, J. Cutul., 101. 103 (1986). 13. G. Connell, and J. A. Dumesic, J . Cutul.. 102, 216 (1986). 14. G. Connell, and J. A. Dumesic. J. Cutul., 105, 285 (1987). 15. M. C. Kung. and H. H. Kung, Cutul. Rev. Sci. Eng., 27,425 (1985). 16. L. H. Little, "Infrared Spectra of Adsorbed Species," Academic Press, New York. 1966. 17. A. Tsyganeko, D. Pozdnyakov, and V. Filimonov, J. Mol. Struc., 29,299 (1975). 18. K. Morshige. S. Kittaka, S. Katsuragi, and T. Morimoto, J. Chem. Soc. Furuduy

19. H. Krozinger, Adv. Cutul., 25. 184 (1976). 20. N. Gill, R. Nuttall, D. Scaife, and D. Sharp, J. Inorg. Nucl. Chem., 18.79 (1961). 21. A. Zecchina, E. Guglielminotti, L. Cerruti, and S. Coluccia, J. Phys. Chem.,

22. E. P. Parry, J. Cutul., 2, 371 (1963). 23. W. B. Jensen, 'The Lewis Acid-Base Concepts, An Overview," Wiley, New York,

24. M. Taagelpera, W. Henderson, R. Brownlee. J. Beauchamp. D. Holtz. and R. Taft,

25. J. Valyon, R. L. Schneider, and W. K. Hall, J. Cutul., 85, 277 (1984). 26. C. Rochester, and S. Topham. J. Chem. SOC. Faruduy T r m . I , 75, 1259 (1979).

Publ.. Stroudsburg, PA, 1973.

4067 (1986).

New York, 1970.

Springer-Verlag, New York/Berlin, vol. 2. p. 231, 1981.

Jpn.. 47, 1064 (1974).

Jpn., 46. 2985 (1973).

Trans. I , 78, 2947 (1982).

76, 571 (1972).

1980.

J . Amer. Chem. SOC., 94, 1369 (1972).

90

27. T. Iizuka, K. Ogasawara, and K. Tanabe, Bull. Chem. SOC. Jpn., 56, 2927 (1983). 28. F. Gil-Llambias. and A. M. Escudey-Castro, J. Cataf., 83, 225 (1983). 29. G. A. Park, Chem. Rev., 65, 177 (1965). 30. G. A. Park, Adv. Chem. Series, 61. 121 (1967). 31. F. Gil-Llambias, and A. Escudey-Castro, J . Chem. SOC. Chem. Commun., 478 (1982). 32. A. Escudey-Castro, L.B. McLeod, and F. Gil-Llambias, Appl. Cataf., 4, 371 (1982). 33. F. Gil-Llambias. A. Escudey-Castro, J. Fierro, and A. Lopez-Agudo, J. Cataf.,

34. H. Pines, 'The Chemistry of Catalytic Hydrocarbon Conversions."

35. Y. Takita, K. Nita, T. Maehara, N. Yamazoe, and T. Seiyama, J. Catof.,

36. D. Dadyburjor, S. Jewur, and E. Ruckenstein, Cutaf. Rev., 19, 293 (1979). 37. H. Kung, Ind. Eng. Chem. Prod. Res. Dev., 25, 171 (1986). 38. M. Ai, J . Catal., 52, 16 (1978).

95, 520 (1985).

Academic Press, New York, 1981.

50, 364 (1977).

Chapter 6

REDUCTION OF OXIDES

6.1 INTRODUCTION

Reduction of an oxide can be accomplished by removal of lattice oxygen, or by dissolution of the reductant into the lattice. The former is common to all oxides, while the latter occurs only in very selected systems. An example of the latter is the dissolution of hydrogen in WO3 to form H,W03, tungsten bronze.

Removal of lattice oxygen can be achieved by many different reducing agents. The common ones include hydrogen, carbon monoxide, ammonia gas, and hydrocarbons. Usually the thermodynamic driving force is the formation of water and carbon dioxide. However, while the use of different reducing agents may involve nearly equivalent thermodynamic driving forces, the kinetics and the rncchanism of reduction can be very different. The rates of reduction may depend strongly on the presence of surface defects and bulk grain boundaries, the orientation of the exposed surface planes, the nature of the support if any, the presence of hydroxyl groups, and the presence of other metals. When the bulk structure is anisotropic, the rate of reduction in the bulk is likely anisotropic. For example, molybdenum oxide has a layer structure, and shear planes are more readily formed along some directions (see Chapter 2). Reduction along the shear planes is much more rapid than along other directions.

Reduction in the surface region may be accomplished also by ion sputtering which usually preferentially removes oxygen atoms. The extent of reduction obtained by this process depends on the orientation of the surface plane, the nature of the oxide, and the sputtering conditions [la].

Because of the large number of significant variables, the surface properties of samples of partially reduced oxides may be quite different and dependent on the dctails of the reduction process, even for the same overall degree of reduction. For example, the partial reduction of iron oxide, Fe203, by CO usually results in a lower oxide in the surface region of a grain and a surface carbonate. Under some conditions, a carbide may be formed. Reduction by hydrogen, on the other hand,

91

92

results in a hydrogenated surface of reduced oxide. Reduction by hydrocarbon could lead to a lower oxide with a surface C ~ ~ ~ ~ M C ~ O U S layer, as well as a hydroxylated surface and surface carbonate. Another example is the reduction of bismuth molybdate by hydrogen, butene or propene. The catalytic activity of the partially reduced oxide for propene and butene oxidation has been found to depend on the reducing agent used (see Chapter 11) [lb].

6.2 THERMODYNAMICS OF REDUCTION

The reaction between a metal oxide MO, and hydrogen to form a lower oxide MO,-l and water vapor can be represented by the equation:

The Gibbs frce energy change, AG, for this reaction is related to the standard value, AGO, by:

AG = AGO+ RT tn(PH20/PH2) (6-2)

If the oxide MO, is reduced totally to the metal M, the last term of this equation becomes nRT tn(PH20/PH2), and the standard Gibbs free energy change now refers to the total reduction reaction.

Inspection of the stoichiometric ratios of this reaction in Eq. (6-1) shows that there is no change in the number of gas phase molecules between the reactants and the products. Thus the AS contribution to AG is small.

Thermodynamic treatments are quite useful for estimating the extent of reduction of bulk oxides and large oxide crystallites. Their usefulness is more limited for well-dispersed oxides on a support or for amorphous oxides, because the AGO values for these oxides are usually not available, cannot be estimated reliably from bulk metal-oxygen bond energy, and vary from one sample to the ncxt. The metal-oxygen bond lengths and bond angles in these oxides may be quite different from bulk oxides. For example, it has been mentioned in Chapter 3 that the lattice of small a-Fe203 crystallites dilates. The energy associatcd with the interaction with a support cannot be evaluated accurately either.

If the ratio of the partial pressure of water to hydrogen is known (or equivalently the gas phase oxygen activity), equations (6-1) and (6-2) permit the calculation of the MO,-l to MO, ratio. That is, the bulk stoichiometry of the oxide can be evaluated. However, this does not imply that the surface stoichiometry is known. At thermodynamic equilibrium, the thermodynamic activity of the oxygen ion in the surface is the same as that in the bulk. However, the different electrostatic potcntial and the different numbers of ncighboring ions at a surface site versus a bulk site make the energy of the surface ions diffcrent from thc bulk ions. Furthermore, there may be reconstruction of the surface structure. The result is that at equilibrium, the surface stoichiometry may be somewhat diffcrent from the bulk stoichiometry. The extent of this possible difference is

REDUCTION 93

currently unknown.

6.3 KINETIC MODELS

The first step in reduction is the activation of the reducing agent. If carbon monoxide is used, it is most likely first adsorbed onto a coordinatively unsaturated surface metal ion of the oxide. This is followed by its reaction with the lattice oxygen to form a surface carbonate with reduction of the metal cations, which decomposes to carbon dioxide. If hydrogen is used, it is adsorbed by heterolytic dissociation. A surface hydroxyl group is formed in this process (see Chapter 4). Desorption of water that is formed by the reaction of a hydroxyl group and a hydride would results in reduction of the oxide.

Since adsorption of hydrogen and/or carbon monoxide might be enhanced by the presence of reduced metal oxide which has a higher density of coordinatively unsaturated metal cations of lower oxidation states, the reduction process could be autocatalytic, depending on the method of growth of the reduced portion of the oxide.

Mathematical description of the reduction kinetics is available for bulk oxide particles which can be modeled as isolated spheres in the absence of any effect due to support, presence of impurities, or anisotropy. There are two models, the nucleation model and the contracting sphere model.

In the nucleation model, surface oxygen ions are removed from the lattice by reduction, leaving behind an anion vacancy. When the concentration of vacancies reaches a critical value, the vacancies are annihilated by rearrangement of the lattice with the eventual formation of small grains of lower oxide or metal (see Fig. 6-1). Such rearrangement can be readily accomplished if the oxide tolerates the presence of shear planes. The small grains of reduced oxide grows by inward diffusion of the reduced metal ions and/or outward diffusion of the oxygen ions. Eventually, these grains of reduced oxides or metal grow to the extent that their boundaries overlap, resulting in a particle of an oxide core with a shell of reduced oxide. Further reduction of this particle follows the contracting sphere model.

The time dependence of the degree of reduction, a, in the initial stage of isothermal reduction by the nucleation model is described by eq. (6-3) [2,3]:

where Vfinal is the final volume of the grain of reduced oxide, C1 and p are constants related to the time dependence of the volume of a grain, t l is the induction time, and C2 and q are constants related to the time dependence of the number of grains.

Figure 6-2 shows the fraction of an oxide reduced with respect to time for the nucleation model. There are two characteristics for the kinetics of reduction according to this model: the presence of an induction period, and the possibility

94

nucleus of metal or reduced metal oxide

metal oxide

Figure 6-1 Initial stage of reduction of metal oxide by the nucleation model.

Time

Figure 6-2 Fraction of oxide reduced as a function of time for the contracting sphere model (I) and the nucleation model (n).

Figure 6-3 Initial stage of reduction of metal oxide by the contracting sphere model.

REDUCTION 95

of autocatalysis. Both of these characteristics are results of the fact that in the early stage of reduction, the grain-oxide interfacial area increases with the size of the grain, and that a grain of reduced oxide activates the reductant more readily than the fully oxidized oxide. These result in an increase in the rate of supply of reductant to the oxide, and the corresponding increase in the rate of reduction as the reduction proceeds. Thus the fraction of oxide reduced as a function of time is concave upward initially. Eventually when the grains of reduced oxide coalesce, and the reduction follows the contracting sphere model, the rate of reduction decreases with time. Thus the reduction curve has a sigmoidal shape.

This nucleation model describes the reduction of NiO with hydrogen. An induction period is observed when the NiO (100) plane is reduced. As the surface becomes severely reduced, Ni metal islands exposing primarily the (100) plane are formed with Ni <010> and <001> directions along the NiO cOlO> and <001> directions, respectively. The rate of oxygen removal is first order in surface oxide concentration, and first order in the hydrogen pressure [4]

The contracting sphere model is an extreme case of the nucleation model in that it is assumed that the number of reduced oxide grains formed on the surface of the sphere is so large that the boundaries of the grains overlap when the diameters of the grains are still small versus the radius of the sphere. This situation can be accurately modeled as a rapid formation of an uniform layer of reduced oxide or metal on the sphere (Figure 6-3). The thickness of this reduced layer grows uniformly, resulting in a spherical core of oxide that shrinks with time. Since the rcduction of the core oxide is accomplished by diffusion of ions and reductant across the oxide-reduced oxide interface whose area decreases with time, and the distance of this interface from the surface of the sphere increases with time, the rate of reduction decreases with time. This results in a characteristic curve for the extent of reduction shown in Figure 6-2.

The time dependence of the degree of reduction for isothermal reduction by the contracting sphere model is described by eq. (6-4) [2]:

k rodo (‘O -

P c,,>t = [ l - a - - ( 1 - 3

where k, is the rate constant of reduction per unit area at the oxide-reduced oxide interface, kd is the diffusivity of reductant through the reduced layer, ro is the radius of the entire sphere, and C,, C,, are the actual concentration of reductant at the outer surface of the sphere and the equilibrium concentration. The expression assumes that the chemical rcaction at the interface is first order with respect to Cl-C,,, where C , is the reductant concentration at the interface.

Another situation is possible that is in between the nucleation and the contracting sphere models. This situation is as follows. Small grains of reduced oxide are first formed on the surface of an oxide particle as in the nucleation model. However, these grains of reduced oxide are less active in activating the reductant molecules than the fully oxidized oxide. Therefore, their presence on the oxide surface inhibits the rate of reduction bccause it reduces the area of exposed fully

96

oxidized oxide. These grains grow as reduction continues, eventually cover up the oxide particle. Whence the oxide is reduced by the contracting sphere model.

In this situation, the rate of reduction decreases continuously with increasing extent of reduction, and the fraction of reduction-versus-time curve should be qualitatively the same as the one for the contracting sphere model.

The descriptions so far assume that the cations and anions in an oxide have homogeneous properties. However, the surface chemistry and catalytic properties depend heavily on the surface condition. In section 4.1, it is shown that a surface may possess more than one type of surface lattice oxygen. A Ti02 (1 10) surface is one such example (Fig. 4.1) where there are two types of surface lattice oxygen ions. Type one is bonded to two six-coordinated Ti cations, and sits above the plane of the other ions. The other type is bonded to one six- and two five- coordinated Ti cations and is in the surface plane. Their different positions and bondings should make them removable with different degrees of difficulty. Removal of the first type would involve breaking two Ti-0 bonds. Removal of the second type would involve breaking three Ti-0 bonds, and should be more difficult than the first type. Indeed, this has been confirmed on the Sn02 (1 10) surface [67]. Even within each type, it is expected that there will be a gradual increase in the difficulty in removing the oxygen ions as the reduction increases.

This consideration suggests that unlike the kinetic treatment above which applies well to bulk reduction, reduction of the surface is likely to proceed with a rate constant that depends on the orientation of the surface plane, the extent of reduction, and the density and nature of surface defects. Unfortunately no quantitative experimental data are available.

6.4 MECHANISM OF REDUCTION

The most heavily studied system is perhaps the reduction of well-dispersed molybdenum oxide on a support. In the reduction of molybdcnurn oxide supported on alumina, it has been found that one molecule of water is relcased for every two hydrogen molecules consumed at low extents of reduction. Thus some hydrogen is retained by the partially reduced oxide [51. The Mo(V) specics has been detected by EPR, but its concentration is low compared with thc extent of reduction [6]. A model of the reduction process that is consistent with these results is shown in Figure 6-4 [7].

In this model, the well-dispersed supported molybdenum oxidc is present as an oxide chain. The initial step of reduction is the heterolytic adsorption of hydrogen molecules followed by reduction of Mo cations and migration of H from a cation to an oxide ion to form a p-hydroxl group. Thus two p-hydroxyl groups are formed for one hydrogen molecule adsorbed. The two p-hydroxyl groups may bridge the same Mo-Mo pair or diffcrcnt pairs as shown in structure 111, and these two modes probably occur with equal likelihood. On the other hand, dehydration probably occurs more readily if the p-hydroxyl groups span the same Mo-Mo pair. This would explain the observation of one molecule of water released for every two molecules of hydrogen consumed at low extents of reduction. This i s

REDUCTION 97

I\ / I - o-o-o o-o-o-

H 0

-1- H2°

(VI)

- o-o-o o-o-o-

Figure 6-4 A possible mechanism for the reduction of supported MOO,

98

illustrated by the conversion of III -t IV. Reduction continues by further adsorption of hydrogen and dehydration (IV + V + VI).

If the molybdenum species is part of a crystallite instead of the monolayer shown, diffusion of lattice oxygen from the bulk to the surface occurs. This is equivalent to the diffusion of anion vacancies from the surface into the bulk, and surface molybdenum oxide can then undergo another cycle of reduction. This process results in bulk reduction.

The mechanism and the extent of reduction determine whether the reduced oxide can be rapidly reoxidized. In general, if the reduction preserves the essential structural feature of the oxidized form, rapid oxidation may be expected. A rapidly reversible system has been demonstrated for CaMn03 and CaMnO2.5, and for CazMn04 and Ca2Mn03.s [8]. The reduction of M a 3 to the lower oxides by the formation of shear planes is also expected to be rapidly reversible.

When there are a number of intermediate oxides between the fully oxidized state and the metallic state, then during reduction, reduced oxides are normally formed in the sequence of increasing degree of reduction. For example, in the reduction of large crystallites of CuO at 250°C in 2% H2/Nz, the oxide is almost totally reduced to CuzO before metallic Cu appears [9]. However, depending on the system and the reduction conditions, some intermediate oxides are further reduced immediately upon formation. For example, one might expect the reduction of F@03 to proceed via Fe304 and then FeO en route to metallic Fe. However, when Fez03 is reduced in Hz, F%04 is lirst formed, and Fe304 is then reduced directly to Fe metal. Little or no FeO is detected [lo].

Even with a given oxide, the reduction sequence can depend on the crystallite size, the crystallographic form, and the nature of the support. In the example of CuO mentioned above, much less Cu20 is detected during reduction of small crystallites, and a majority of CuO appears to be directly reduced to Cu metal [9]. The dependence on crystallographic form is illustrated by N b z 0 5 . The reduction of p-Nb205 proceeds readily in H2 at 750°C to form Nb02 [ l l ] . However, a- Nb205 is reduced through a series of complex intermediates, but no Nb02 is formed [12]. The effect of the support is quite complicated. It is separately discussed in the next section.

6.5 EFFECT OF SUPPORT

The interaction of an oxide with a support could affect both the thermodynamics and the kinetics of its reduction. The effect arises partly because a support can act as a dispersing agent, and small crystallites may be reduced differently than large crystallites, such as the example of CuO mentioned above. A support can also interact chemically with the oxide. Often, but not always, such a chemical interaction increases the resistance of the oxide to reduction.

There are quite a number of examples of increased resistance to reduction when an oxide is supported. W 0 3 supported on A1203 is more difficult to reduce than unsupported W 0 3 [13,14]. Fe2O3 supported on y-AI2O3 [2], and Pt oxide supported on y-AI2O3 [15] are similarly less reducible than their unsupported

REDUCTION 99

oxides. It has been reported that a F't+ signal is detected by EPR on a I"fl-Al203 sample even after reduction at 623 K [16].

There are disagreements for some other systems. One report mentions that V205 supported on Ti02 is less reducible than unsupported Vz05 [17]. Another lists the ease of reduction as decreasing in the sequence Vh-Al203, VDiO2, V2O5 [18]. In this case, the influence of impurities present in the support may be important.

It has been reported that in temperature programmed reduction, carefully oxidized Rh crystallites on A1203 or Ti02 that contain only a surface layer of Rh oxide are reduced more readily than the bulk oxide [19]. In another report, it is mentioned that RhIy-Al203 that is oxidized at a temperature above 873 K is reduced less readily than a similarly pretreated bulk oxide [20]. It has been proposed that a high temperature of oxidation is necessary to induce the strong interaction with the support that leads to lower reducibility. It is interesting to mention that temperature programmed reduction of a Rh/Ti02 sample fully oxidized at low temperature shows two distinct reduction peaks. The lower temperature peak is due to well-dispersed Rh2O3 on the support, and the higher temperature peak is due to large Rh2O3 crystallites [19]. It appears that in the supported Rh203 system, different pretreatments lead to different effects of dispersion and chemical interaction on the rates of reduction.

In some cases, the apparent support effect may be actually due to the reaction between the oxide and the support to form a new compound. In this case, the increased resistance to reduction is due to the different properties of the new compound. This is the case for NiO/A1203. The reduction of this oxide often shows portions of different rates. It is now understood that the different rates correspond to the reduction of crystalline NiO, Ni2+ bound to alumina, and nickel aluminate [21-241. The increased resistance to reduction of V205/Mg0 is likewise due to the formation of Mg3(V04)2 [25].

The reduction of a metallic ion in a zeolite depends on the location of the ion. It has been reported that a Ni2+ ion in a hexagonal prism of faujasite is more difficult to be reduced than a Ni2+ ion in the sodalite cage or supercage [26]. Presumably, the lattice oxide ions of the hexagonal prism stabilize the Ni2+ ion rather effectively.

It should be noted that while a support affects the reducibility of an oxide, the presence of an oxide may also cause changes in a support during reduction. For example, during temperature programmed reduction of a V205/Ti02 (anatase) sample, reduction of V2O5 and transformation of anatase to rutile are simultaneously observed [27-291.

When an oxide strongly interacts with a support, the support could determine the structure of the oxide and its reduction behavior. For example, reduction of vanadia well-dispersed on anatase Ti02 removes an average of 0.85-0.90 oxygen atoms per vanadium ion [30]. Reduction of vanadia well-dispersed on A1203 removes an average of 0.6 oxygen atoms per vanadium. Reduction of large crystallites of V205 removes 0.55 oxygen atoms per vanadium.

The rhodium oxide system is interesting.

100

6.6 EFFECT OF OTHER COMPONENTS

The presence of other components may enhance the rate of reduction of an oxide by promoting the rate of activation of the reductant. For example, the presence of a noble metal may enhance the dissociation of hydrogen. The activated reductant then migrates to the reduction site and reduction is enhanced. In some cases, migration is by the spillover phenomenon in which the activated reductant migrates over the support. This mechanism has been used to explain the results from a physical mixture of F~03&-Al203 and ptly-AlzO3. Reduction of FezO3/y-AIzO3 commences at about 430 K in this mixture, much lower than 640 K without F't [2].

The reduction of Re-Pt/Alz03 is another well studied example. The addition of Pt considerably lowers the temperature required to reduce Rez&. This has been explained by the surface migration of R%07 to Pt [31,32]. Such migration results in physical contact between the two components, and hydrogen atoms dissociated on Pt can migrate readily to R%07. A similar enhanced reduction by the addition of Pd to R%07/A1203 is also observed [33].

The reduction of Pt is also affected by the physical contact between Pt and RezO7. In a temperature programmed reduction experiment, the reduction peak for Pt is moved to a higher temperature in the presence of Rez& than in its absence WI.

The formation of hydrogen bronze is much facilitated when the oxide is in contact with a metal that dissociates hydrogen molecules. It has been reported that HxW03, H,M0O3, and HXV2O5 can be readily prepared by exposing the oxide deposited with platinum to hydrogen gas [64-661. It is interesting that only part of the hydrogen incorporated into the bronzes can be removed easily, such as by evacuation or by consumption in the hydrogenation of ethene [65,66].

6.7 REACTIVITY OF REDUCED SURFACES

Reduced surfaces often possess chemisorptive and catalytic properties different from stoichiometric surfaces. It has been shown that NO is adsorbed strongly on Fez+ but only weakly on Fe3+ [35-371. and strongly on partially reduced tungsten oxide but not on fully oxidized W6+ [38]. CO has been shown to adsorb on Cu+ strongly, but on Cu2+ weakly [39]. The reason for the difference between the oxidized and the reduced ions is probably the different electronic configurations of the ions, and the lower degree of coordinative unsaturation in the fully oxidized oxides.

Together with the presence of surface cations with a high degree of coordinative unsaturation, a reduced surface also has cations of a lower oxidation state and anion vacancies. Some of these factors lead to the enhanced activity of the surface to dissociatively adsorb molecules that contain oxygen atoms. The dissociative adsorption often results in reoxidation of the surface. The data in Table 6-1 illustrate this effect. In these experiments, the fully oxidized (i.e., stoichiometric) surfaces are prepared either by extensive annealing of the surface

REDUCTION 101

Table 6-1 Chemical Properties of Oxidized and Reduced Single Crystal Surfaces

Oxidized" or

Surface Reduced Adsorption Ref.

v203(W'

SrTi03( 1

Reduced

1 Oxidized

Reduced

1) Reduced

SrTi03(100) Oxidized

Reduced

Ti02( 1 10) Oxidized Reduced

Ti02(100) Oxidized (1x3) Reduced

Ti02( 100) Reduced (1x7)

Ads. O2 readily as 02-; Ads. H 2 0 molecularly.

Ads. O2 readily as 02-; Ads. H 2 0 dissociatively.

Ads. O2 readily as 02-; Ads. H 2 0 dissociatively. Ads. O2 readily; Ads. H 2 0 dissociatively.

Ads. H 2 0 reoxidizes the surface.

Ads. H 2 0 molecularly; Ads. small amount of 02. Ads. O2 reoxidizes the surface; Ads. H 2 0 dissociatively.

Does not ads. CO or H2. Ads. O2 dissociatively and reoxidizes the surface;

Ads. H20 dissociatively.

Ads. H 2 0 molecularly. Ads. H 2 0 dissociatively.

Ads. H 2 0 dissociatively.

a b

a b

C

C

d

j e e f

g g h

f

I

1

1

Footnote: a) Fully oxidized sufaces are stoichiometric surfaces.

References: a) R.L. Kurtz, and V.E. Henrich. Phys. Rev. B. 25. 3563 (1982). b) R.L. Kurtz, and V.E. Henrich. ibid. 26, 6682 (1982). c) R.L. Kurtz, and V.E. Henrich, ibid, 28, 6699 (1983).

102

Table 6-1 continued

d) S. Ferrer. and G.A. Somorjai, Surface Sci.. 97. L304 (1980). e) V.E. Henrich, et al., J. Vac. Sci. Techno]., 15, 534 (1978). f) V.E. Henrich. et al., Solid State Commun., 24. 623 (1977). g) W. Gopel, et al., Phys. Rev. B, 28, 3427 (1983). h) V.E. Henrich, et al., Phys. Rev. Lett., 36, 1335 (1976). i) W. Lo. et al., Surface Sci., 71, 199 (1978). j) by HREELS. R.G. Egdell and P.D. Naylor, Chem. Phys. Lett., 91, 200 (1982).

or by cleavage of a single crystal. Reduction of the surface is achieved by ion- sputtering or H2 reduction. Adsorption and dissociation of the molecules are monitored with UPS.

The results in Table 6-1 are supported by other measurements such as temperature programmed desorption. When adsorbed CI8O2 is desorbed by heating from a partially reduced TiOz surface, the amount of C1*0 desorbed increases with the extent of reduction 1401.

Reduction of a surface does not always lead u, stronger interaction with molecules. The peak temperature in the temperature programmed desorption of CO from Ti02 does not change with increasing degree of reduction, although the amount increases [40]. On ZnO, the differential heat of adsorption of CO is much lower on a reduced (7 kJ/mole) than on an oxidized surface (44 kJ/mole) [41]. In contrast to NO which is adsorbed strongly on Fez+ but weakly on samples containing only Fe3+ [35-371, pyridine is adsorbed more strongly on samples containing Fe3+ than Fe2+ [42]. It is interesting to note that in this study, Mossbauer spectroscopy detects direct interaction between pyridine and Fe2+ of low coordination, but not that between pyridine and Fe3+. However, pyridine is shown by infrared spectroscopy to adsorb on Lewis acid sites. Since Fe3+ is a harder and smaller cation than Fez+, it is a stronger Lewis acid.

6.8 INFLUENCE OF REDUCED OXIDES ON THE PROPERTIES OF TRANSITION METALS

It was first reported about ten years ago that certain transition metal oxides can have a dramatic effect on the chemisorptive and catalytic properties of metals when they are in close contact with the metal. This effect was originally tcrmed Strong Metal-Support Interaction, but is currently referred to as decoration effcct. The first report of this effect was on the suppression of hydrogen and carbon monoxide adsorption capacity of noble metals. It has been found that when a noble metal is supported on Ti02, the material chemisorbs H2 or CO at a stoichiometry of roughly one H atom or CO molecule per surface exposed noble metal atom, if the material is reduced at 200°C. If the material is reduced at 50O0C, its ability to chemisorb H or CO is essentially completely suppressed [43,44]. This is illustratcd with typical data in Table 6-2. It was later found that

REDUCTION 103

I I 100 300 500 700

C TA

Figure 6-5 Hydrogen chemisorption on iridium supported on various oxides as a function of activation in hydrogen for 1 h at each of various temperatures. T, is the activation tcmperature. and H/M is the atomic ratio of hydrogen adsorbed to iridium in the catalyst. (From Science, 211, 1121 (1981), copyright American Association for the Advancement of Science).

Table 6-2 Suppression of H2 Adsorptive Capacity due to Decoration Effect (From Scicnce, 211, 1121 (1981), copyright American Association for the Advancement of Science).

2% Metal on Ti02 Support Reduction at 200°C Reduction at 500°C

H atom adsorbed/rotal metal atoms

Ru 0.23 Rh 0.7 1 Pd 0.93 0 s 0.2 1 Ir 1.60 Pt 0.88

0.06 0.01 0.05 0.1 1 0.00 0.00

this effect is not unique to Ti02 as the support. Fig. 6-5 shows the extent of suppression of the hydrogen chemisorptive capacity as a function of reduction temperature for various oxide supports. In general, oxides that are readily reducible at intermediate temperatures all show this effect. The effect is reversible. That is. reoxidation with oxygen followed by a low temperature reduction at 200°C restores most of the adsorptive capacity.

The influence of the decoration effect on the catalytic properties of the metal depends on the reaction. Some typical examples are shown in Table 6-3. It can be seen that the decoration effect suppresses the activity of R for benzene hydrogenation and cyclohexane dehydrogenation, but enhances the activity in CO hydrogenation. The activity of the Fe catalyst in ammonia synthesis is slightly decreased, but the activation energy is greatly increased. In the case of butane hydrogenolysis, the selectivity and the activity are both altered.

In addition to these examples, it has been shown that CO hydrogenation on Ni catalysts is also enhanced by the decoration effect, although the extent of enhancement may vary from very little [48] to rather substantial [49]. In ethane hydrogenolysis on titania-supported Rh, it has been found that the activity decreases rapidly with increasing reduction temperature (Fig. 6-6), whereas the cyclohexane dehydrogenation activity over the same catalyst remains almost unchanged [ 501.

It is now rather well established that this effect is not due to the formation of alloys (e.g. R-Ti alloy), the encapsulation of the metal by the support, sintering of the metal, poisoning of metal by impurities in the support, or simply electron transfer between the bulk of the support and the bulk of the metal crystallite. These conclusions follow from a variety of experimental observations. For example, transmission electron microscopic studies as well as X-ray diffraction show no evidence of sintering [45,51], and Mossbauer spectroscopy shows that the bulk of the iron crystallites is the same whether or not the sample is exhibiting the decoration effect [47], even though the titania support near the metal crystallite is reduced from Ti02 to TbO, [51,52].

The current picture is that the origin of the effect is the migration of small particles of reduced titania (or other reduced support) onto the metal crystallites to "decorate" the metal surfaces. These decorating reduced oxide particles may partially block the metal surfaces from gas molecules, affect the electronic structure of the neighboring metal atoms, or provide an oxide-metal interface for interaction with molecules. Depending on the reaction, one or more of these effects may participate to affect the observed characteristics of the reaction.

That decoration is the physical picture was first suggested by Dumesic [471 on the Fe/Ti02 system. Using Mossbauer spcctroscopy, it has been observed that the bulk properties of Fe crystallites are the same whether they are in the decorated state or not. Thus the effect must be a surface phenomenon. It is then proposed that titania species cover the iron crystallites. Such a decoration model would suggest that the extent of the effect should depend on the time allowcd for the reduced oxide particles to migrate onto the metal crystallites and thc interface between the metal and the oxide. These have been confirmed. It is observcd that using the rate of ethane hydrogenolysis as a measure, the extent of the decoration

REDUCTION 105

Table 6-3 Effect of Decoration on the Catalytic Properties of Noble Metals

Reaction Other Reaction Catalyst Rate Effects Ref

C6H6 hydrogena- 4.8% W i O 2 tion at 288K 523 K reduced

773 K reduced

C6HI2 dehydroge- 2.7% IrKi02 nation at 523 K 523 K reduced

773 K reduced

CO hydrogenation 1.9% Pt/Ti02 at 524 K 473 K reduced

773 K reduced

NH3 synthesis 1.14 % Fe/Ti02 at 673 K 713 K reduced

798 K reduced

C4Hlo hydro- 4.8% Pfli02 genoly sis 623 K reduced at 623 K

773 K reduced

45 40 mmole/h-g cat. 3.5 mmole/h-g cat.

1400 mmole/h-g cat. 304 mmole/h-g cat.

0.01 11 molecules/s-Pt, 0.076 molecules/s-% (0.0195 if assumed the same dispersion as low- temperature reduced sample)

45

46

47 0.03 1 ks-l

0.01 1 ks-'

Eact = 100 kJ mole-'

Eact = 220 kJ mole-'

45 Relative product formation rate:

C, =0.65, C2= 1.1, C3=0.7, i-C4=0 c1=35, C2=47, C,=25, i-C4=88

effcct depends linearly on the square root of the reduction time [501, which is characteristic of diffusion processes. For a given reduction time, the hydrogenolysis activity decreases as the inverse of the particle diameter. Finally, the increasing suppression of the hydrogenolysis activity as the reduction time increases is found to parallel a similar suppression by the addition of copper to a nickcl catalyst, which is interpreted by the braking-up of nickel ensembles on the surface by copper atoms. Thus the data are consistent with the model that the surface rnctal ensembles are broken up by reduced titania particles decorating the surface.

106

Reduction Temp- K

Figure 6-6 Ethane hydrogenolysis and cyclohexane dehydrogenation on Rh/Ti02 catalyst as a function of catalyst reduction temperature. (From J. Catal., 82, 279 (1983). copyright Academic Press).

Recently, it has been further proposed that the metal-oxide interaction occurs through the interaction of metal atoms with oxygen ion lattice vacancies in the reduced oxide. At temperatures sufficiently high to induce the decoration effect, the oxide support is reduced so that it has a high concentration of anion vacancies. The high concentration of anion vacancies enhances the diffusion of metal atoms into the near-surface region of the bulk, and results in the formation of a raft-like metallic cluster covered by a thin (atomic) layer of the support. When the support is reoxidized, the anion vacancies are filled, driving the metal atoms back to the surface [53].

When titania particles are deposited on model catalysts of a Ni (1 11) single crystal surface 1541 or a F’t foil [55,56], a similar suppression of the H2 or CO chemisorption capacity is observed which is similar to that resulting from high temperature reduction of supported metals. Enhanced catalytic activity in CO hydrogenation has also been observed on these low surface arca catalysts [54,56]. In the case of Pt foil, the activation energy is reduced from 126 to 80 W/mole. Furthermore, a small amount of deposited TiO, particles is found to enhance the methanation activity of Ni. At an optimum coverage of 8%, (recently there is doubt about this number because of questions about the calibration method employed), the activity is four times that of a clean Ni surface. On the other hand, complete suppression of the chemisorptive capacity rcquires complete coverage of the metal surface by the oxide particles 156,571.

REDUCTION 107

The extent of election transfer between the metal and the decorating particles is not established. For example, a study of Pt crystallites supported on a Ti02 single crystal surface by XPS and AES suggests electron transfer from Ti02 to F't [%I, while a XANES study suggests electron transfer from Pt to Ti02 [591.

The influence of the decorating effect on the heat of adsorption depends on the system. On Pt supported on titania, the effect results in a decrease in the initial heat of adsorption of H2 from 92 to 82 kJ/mole, but no change for CO adsorption [60]. The integral heats of adsorption of both CO and H2 are substantially reduced [61]. However, the integral heat of adsorption of CO or H2 on Pd is not affected [62,63].

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REDUCTION 109

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Chapter 7

OXYGEN ON OXIDES

7.1 NATURE OF ADSORBED OXYGEN

Depending on the sample history, it is possible that there are oxygen atoms, either neutral or charged, on an oxide surface that are in positions different from the positions of surface lattice oxygen ions. These oxygen atoms and ions may have different charges than the lattice oxygen, have different energics of binding (or adsorption), and be desorbed at different temperatures. Adsorbed oxygen on a stoichiometric surface of a fully oxidized oxide is readily identifiable. It is ususally desorbed at temperatures lower than the sublimation temperature of surface lattice oxygen. On a partially reduced surface, adsorbed oxygen may result in reoxidation of the surface cations to different degrees, depending on the extent of charge transfer between the reduced center and the adsorbed oxygen. If the charge transfer is such that an adsorbed oxygen atom acquires the same elcctron density as a surface lattice oxygen ion, and it occupies a lattice site, the surface is reoxidized and the adsorbed oxygen becomes a lattice oxygen. If the charge transfer is less extensive and/or the oxygen species occupies a site diffcrcnt from a surface lattice site, it is an adsorbed oxygen.

Adsorbed oxygen may be present as atomic or molecular species with various charges. On transition metal oxides, the most common species are O,O-, 02, and 0 2 , On some basic alkali and alkaline earth oxides, 022- has been reported [l]. On UV-irradiated TiOz, 03- species has been observed with EPR [2]. Various review articlcs discussing these and other species such as 03- havc appeared in recent years [2-41.

In general, atomic oxygen spccics are adsorbed more strongly than molccular species. This is because strong surface-atomic oxygen bonds are needcd to compensate for the energy required to break the double bond of thc oxygen molecule. The rate of dissociativc adsorption of oxygen to atomic species is expcctcd to be lower than that of molccular adsorption. This is because a pair of neighboring surface sites must be available for the former process. Otherwise

110

SURFACE OXYGEN 111

lattice or surface diffusion of mononuclear oxygen ions (or lattice anion vacancies) is needed before dissociation of the oxygen molecule can be achieved. This is consistent with the fact that a saturation coverage of adsorbed atomic oxygen corresponds to a few atoms per nm2 of surface (see Table 7-1). On the other hand, molecular adsorption of oxygen can take place on an isolated surface site.

This requirement of the availability of a pair of surface sites can be overcome by using species such as N20 instead of 02, the decomposition of which would leave an atomic oxygen on the surface. Indeed, decomposition of N20 at low temperatures resulting in an adsorbed atomic oxygen (usually detect& as 0-) has been observed on many oxides [2-41.

7.2 DETECTION OF ADSORBED OXYGEN

One method to detect the presence of adsorbed oxygen is by temperature programmed desorption after exposing the oxide to 02. Presumably, adsorbed oxygen is bonded differently on the surface than lattice oxygen, and would be desorbed at temperatures different from the vaporization temperature. The temperature programmed desorption profiles of a number of oxides have been reported [5-81. The general feature of these profiles is schematically shown in Fig. 7-1. Upon heating of an oxide with adsorbed oxygen from room temperature, three types of desorption peaks are observed. Type I occurs at a relatively low temperature and it represents the most weakly adsorbed oxygen. Normally this is assigned to adsorbed molecular oxygen which usually desorbs below 300°C. Type I1 which occurs at an intermediate temperature usually results from adsorbed atomic oxygen, and desorbs below 600°C. The amounts of oxygen desorbed in type I and I1 are small, of the order of a few percent of a close-packed monolayer. Type 111 occurs at high temperatures and at a rate which may increase continuously with increasing temperature. The amount of type 111 may be much larger than those of the other two types. This type results from vaporization of lattice oxygen, and desorption of metal atoms occurs simultaneously, that is, the metal oxide is subliming at these temperatures.

Table 7-1 summarizes the desorption temperatures of oxygen from a number of oxides [51. Except for those noted, oxygen is adsorbed by cooling the oxide in oxygen from 600°C to 10°C. In this manner, adsorption is achieved even for species whose adsorption is activated. It is clear from the table that the schematic desorption profile shown in Fig. 7-1 is oversimplified. For many oxides, there are more than two desorption peaks below sublimation temperature. For example, there are three desorption peaks for iron oxide. From their temperatures, one is probably adsorbed molecular oxygen and two are adsorbed atomic oxygen. Titanium dioxide appears to have three forms of adsorbed oxygen. The different forms of adsorbed oxygen reflect interaction with different environment of the surface. At present, there is little detailed understanding of the differences.

Attempts to gain further understanding of the adsorbed oxygen species is complicated by the fact that the extent of interaction of the adsorbed species with the surface (and perhaps also with each other) may depend on the adsorption

112

Table 7-1 Desorption Temperatures and Total Amounts of Oxygen Desorbed from Metal Oxides a

Desorption Volume Desorbed Oxides Temp. C" ml(STP)/m2

v205 M003 Biz03

Biz03 .2Mo03 wo3

& 0 2

Fez03 ( 3 3 0 4 NiO CUO

S iOz Ti0 (anatase)b ZnO Sn02

Cr203

A120zb

?b

450 50,270, 360,540 55, 350,486 30, 165, 380 35,335,425,550 125, 390 65 100 125,1190, 250 190, 320 80, 150

0 0 0 0 0

2.13 x lop2 6.54 x lop2 4.05 x 10-3 3.30 x lop2 1.12 x 10-2 1.42 x 10-1 2.05 x 1 0 4 2.99 x 10-5 5.52 x 10-5 2.45 x 1 0 4 2.11 x 10-3

Fo o motes: a) From Iwamoto. et al., J. Phys. Chem., 82, 2564 (1978), coyright American

Chemical Society. Oxygen is adsorbed by cooling the sample in oxygen from 600 to 10°C.

b, Oxygen is adsorbed at room temperature. ') Obtained with a heating rate of 20°C/min.

conditions. For example, the temperature programmed desorption profile of oxygen adsorbed on a-Fe203 has been shown to depend on the adsorption temperature. The profile in this figure can be explained as follows. There are three types of adsorbed oxygen on a-Fe203. The fist type produces a desorption peak at about 490°C. The second type produces a peak at about 370°C. The peak temperatures of these two types do not vary with the adsorption temperature. However, adsorption into these forms is activated and temperatures close to the desorption temperature are needed for adsorption to occur. The desorption of the 370°C species is second order in surface coverage, suggesting that it may be an atomic species [8]. The third type produces a peak below 200°C whose exact temperature depcnds on the adsorption temperature. The desorption temperature suggests that this peak is due to adsorbed molecular

This is shown in Fig. 7-2.

SURFACE OXYGEN 113

T- Figure 7-1 Schematic temperature programmed desorption profile of oxygen born an oxide. I: from adsorbed molecular oxygen; II: from adsorbed atomic oxygen; III: from sublimation of lattice oxygen.

0)

m

C 0

w

a

.- c,

L n 0 VJ Q)

a

I I I I I

100 300 500

Temperature, C

Figure 7-2 Effect of adsorption tcmperature on the TPD profile of adsorbed oxygen on a- Fe203. Adsorption tcmperatures: a. Sample cooled from 600 to 10°C in 0,; b. 10°C; c. 110°C; d. 250°C; e. 400°C. (From Bull. Chem. SOC. Jpn., 51. 2765 (1978), copyright Chemical Society of Japan).

114

Figure 7-3 Room temperature EPR spectrum of adsorbed 0 2 - and Zn' on ZnO. (From J. Phys. Chem., 82. 2564 (1978), copyright American Chemical Society).

oxygen (type I). The variable temperature of the desorption peak suggests that there is a broad distribution of adsorption sites. Similar dependence of the temperature desorption profiles on the adsorption conditions has been observed also on NiO [91 and Mn02 [51.

The assignment of adsorbed atomic or molecular oxygen species to the various desorption peaks has been confirmed by EPR measurements on some oxides. On ZnO, adsorption of oxygen generates an anisotropic signal with g,=2.052, g2=2.009, and g3=2.003 (Fig. 7-3). This signal is assigned to 02-. The signal disappears upon evacuation at 20O0C, which coincides with a desorption peak at 190°C. It appears that the EPR signal and the desorption peak are from the same species [ 5 ] . It is interesting that this 02- signal and the Zn' signal are complementary with each other, showing an electron transfer from Zn+ to oxygen:

On anatase TiOz, oxygen adsorption results in three EPR signals that have been assigned to 0 2 - on three different sites. These three signals disappear upon heating to temperatures that are coincident with the three temperature programmed desorption peaks [5 ] . Thus these three desorption peaks are from adsorbed molccular oxygen.

Direct spectroscopic observation of adsorbed oxygen has been made using IR and EPR. Because of the lack of a dipole moment, the oxygen molccule in the gas phase is IR-inactive. However, it is Raman active with a vibrational stretching frequency of 1552 cm-'. Adsorbed molecular oxygen will have a lower symmetry than in the gas phase because of perturbation by the surface, and some weak IR bands have been assigncd to adsorbed 02. Some of these assignments are summarized in Table 7-2. The assignment of the charge on the adsorbed oxygen molecules is generally made by comparison with coordination complexes.

Examples of adsorbed superoxide ion (023 detected by EPR have been given above for ZnO and Ti02. Adsorbed superoxide ion has been detected on many

SURFACE OXYGEN 115

Table 7-2 IR Detection of Adsorbed Oxygen Molecules on Some Transition Metal Oxides

Stretching Oxide Adsorbed Species Frequency, cm-l Ref.

gas phase 0 2

0 2 - 0 2 2 -

NiO singlet O2 0 2 -

Ti02 0 2 - 0 2 h 0 2

cr203 0 2 2 -

~ - F Q O ~ 0 2 - 0 2 2 -

singlet O2

1552 a-c 1140 850

1500 d 1140,1060 d 1180-1060 e 16GO- 1580 f 1600-1 700 e

985 g 1460 d

1350,1325,1300,1270 i 1100-900 i

References: a) J. Shamir. et al. J. Amer. Chem. SOC., 90. 6223 (1968). b) K. Nakamoto, "IR and Raman Spectra of Inorganic and Coordination Compounds,"

c) N. Sheppard, in "Vibrational Properties of Adsorbates." R. F. Willis, ed..

d) A. Tsyganenko, et al., Spect~os. Lett. 13, 583 (1980). e) A. Davydov, et al., Kinet. Catal., 14, 1342 (1973). f) A. A. Davydov, et al., Symposium on "Adsorbirovanny Kislorod v Katalize,"

Wiley-Interscience, NY, 3rd edition, 1978.

Springer-Verlag, Berlin, 1980.

1972, Institute of Catalysis, Acad. of Science, USSR, Novosibirsk. Preprint No. 19.

g) A. Davydov, et al., Kinet. Catal., 13, 980 (1972). h) F. Al-Mashta. et al., J. Chem. SOC. Faraday Trans. I, 78, 979 (1982).

oxides including ZnO [5,10-121, Ti02 (5,131, supported V205 [14], supported Moo3 1151, mixed oxides of COO-MgO [161, and nontransition metal oxides [3,41. This species is characterized by an anisotropic signal of g1=2.015 to 2.077, g2=2.002 to 2.012, g3=2.001 to 2.011 [3]. There is a correlation between the magnitude of g1 and the formal oxidation state of the metal ion: the higher is the oxidation statc, [he lower is the value of gl because of larger crystal field inlcractions [3,17]. Interaction of the superoxide ion with the nuclear magnetic moment of the metal ion may result in hyperfine splitting of the EPR signal [14].

116

Because of the electronegativity of oxygen, electron transfer from the oxide to adsorbed oxygen commonly occurs, although it may not be a necessary condition for chemisorption. Therefore, adsorption of oxygen becomes much more facile if the oxide is reduced. It has been shown that chromium oxide reduced at 500°C and evacuated (to dehydroxylate the surface and make some surface chromium ions coordinatively unsaturated) can adsorb oxygen strongly at as low as -195°C [MI. Similarly, molybdena reduced at 500°C and evacuated chemisorbs oxygen at -78°C and below [19]. In these two cases, adsorption is believed to be on Cr203 and M a 2 . The facile adsorption has made it possible to use oxygen chemisorption to determine the surface area of supported chromia and molybdena [20,21]. However, the stoichiometry of adsorbed oxygen and surface cation is not unity, and an assumption has to be made that this stoichiomeuy is the same for the supported and the unsupported sample whose surface area can be independently determined. The nature of the adsorption site is still a subject of investigation [22]. There is indication that the adsorption site on molybdenum oxide is a surface Mo2+ center based on the competitive behavior between O2 and NO adsorption ~231.

7.3 REACTIVITY OF ADSORBED OXYGEN

Adsorbed oxygen species vary greatly in their reactivities depending on their nature and the nature of the oxide. Among species that have been identified with EPR, 0- has been shown to be very reactive. It abstracts an H atom from an alkane molecule at as low as 77 K. 0 2 - is less reactive, and it forms complexes with adsorbed alkene molecules. The reactivities of those species not detected by EPR vary. On iron oxide, for example, the strongly adsorbed oxygen appears to be unreactive, in contrast to the weakly adsorbed oxygen which degrades adsorbed alkenes rapidly to combustion products [24].

Most of the studies on the reaction of adsorbed oxygen concerns the EPR- active 0- species. The common method to generate this species is either by irradiation of an oxide with y-ray or UV light, or by decomposition of N20 on the oxide. There does not appear to be any significant difference in the species produced by either method.

Adsorbed 0- is a very reactive species. It reacts readily with alkanes at low temperatures on a number of oxides including ZnO, vanadia, and molybdena. Cleavage of a C-H bond and formation of OH and an alkyl radical is a common first step of the reaction. In somes cases, partial oxidation products of alkanes are formed. Examples of these reactions are described below.

Figure 7-4 shows such a reaction on ZnO. Adsorbed 0- is generated by irradiating ZnO with UV light at 90°K. The presence of this species is detected by EPR (curve b). On exposure of the sample at 90 K to methane, a new EPR signal of .CH3 is detected (curve c). Thus 0- abstracts an H atom from methane according to the equation:

O-(ad) + CH4 -+ .CH3(ad) + OH-(ad) (7-2)

SURFACE OXYGEN 117

___) A H

Figure 7-4 EPR spectra taken at 90 K showing reaction of C& with adsorbed 0- on ZnO. a. ZnO pretreated in 0,; b. Adsorbed 0- on ZnO generated by UV radiation; c. Spectrum of C H 3 after exposure to C&. (From React. Kinet. C a d . Lett., 18, 243 (1981). copyright Elsevier Scientific F'ubl.)

The same reaction has been detected on silica-supported vanadia, molybdena, and tungsta [26,27]. When the reaction is conducted at room temperature on V/Si02, a small amount of C2H, formed by coupling of methyl radicals is detected in the gas phase. Upon heating the oxide, large quantities of C2H, and CO are desorbed together with small quantities of CH, and C02. Similar observations have been made on UV-irradiated Ti02. It has been proposed that the methyl species is adsorbed as a surface methoxide, at least above ambient temperature [27]:

CH3 + 02- + CH30- + e- (7-3)

or .CH3 + 0- + CH30-

The formation of ethane, however, should result from the coupling of two methyl radicals.

When the adsorbed 0- on V/Si02 is exposed to a mixture of methane and oxygen at room temperature, the 0- EPR signal is immediately replaced by an 03- signal. Formaldehyde becomes the major reaction product instead of ethane. On TiOz, CO and C02 are the major products. It is interesting that if ' * 0 2 is present in the gas phase, the oxygen in the formaldehyde formed is not labeled [27].

Adsorbed 0- reacts readily with ethane also. On Mo03/Si02, the reaction is instantaneous and proceeds to almost completion at room temperature [28]. Upon heating, C2H, is desorbed. However, the amount of ethene observed is more than the amount of 0- detected by EPR. This may indicate the presence of undetected

118

0- which is EPR inactive because of strong magnetic dipolar interaction with the solid. 0- adsorbed on V/SiO2 also reacts readily with C2H.j. Upon heating, CI&, C2&, and carbon oxides are the major desorbed species. No coupling product, butane, is observed [271.

Upon heating, ethene and methane are the major products. Interestingly, the total amount of ethene and methane formed is about five times the amount of 0- detected, as in the case of Mo03/Si02 mentioned above. The reaction mechanism has been proposed to be:

Adsorbed 0- on Co-MgO also reacts readily with ethane [29].

O-(ad) + C2& + C2HS(ad) + OH-(ad) (7-4)

C2Hs(ad) + 02-(lattice) + C2HSO-(ad) + e-(s) (7-5)

C2HsO-(ad) + 2 02-(lattice) + CH3COO-(ad) + H20 + 4 e-(s) (7-6)

CH,COO-(ad) + OH-(ad) + C& + C032-(ad) (7-7)

In addition to alkanes, adsorbed 0- also reacts with HZ, D2, and alkenes. Reactions of adsorbed 0- with H2 or D2 are also rapid, as is that with various alkenes. At 10°C the reactivity of 0- with various molecules shows the order D2 < H2 < C21& < CO < CH, [26]. The reactivity also depends on the nature of the oxide. For H2, D2, and CI&, it decreases in the sequence V205/Si02 > Mo03/Si02 > W03/Si02 [26,31]. The reaction with H2 results in the formation of adsorbed OH-. The rate shows a normal deuterium kinetic isotope effect of about four at -lOO°C. and of about two at 10°C. The reaction with CO results in the formation of C02- [26,32]. A similar associative reaction of 0- with 0 2 to form adsorbed 03- has been reported [27.29,30]. On warming and evacuation, the reaction is reversible and 0- and 02(g) are formed.

The reaction of adsorbed 0- with Cz& is more complex and the detail depends on the oxide. On Mo03/Si02, exposure of adsorbed 0- to C2H4 at 110 K results in a new EPR signal that has been assigned to a linear CH2CH20- species. On warming, the signal is replaced by one that has been assigned to a CHCH2 radical [33]. The assignment to a linear .CH2CH20- species is supported by deuterium and 13C labelling studies [34]. On W03/Si02, the same reaction at low temperatures leads to the appearance of a EPR signal that has been assigned to a cyclic (-CH2CH03 epoxide-like species that has not been clearly identified. On warming to 90 K, the -CHCH2 species is also formed 1351.

When a sample of Mo03/Si02 is irradiated with W light in the presence of C2&, 0- is presumably formed by the reaction: Mo6+a2- + Mo5+4-. An EPR signal assigned to an ethylene oxide species has been observed at 77 K [361 that is formed by:

C2& + 0-(ad) - (H2C- H2)- (7-8) \/= 0

SURFACE OXYGEN 119

Under this condition, a small amount of propene is also detected in the gas phase. On warming, more propene as well as some 1- and 2-butene and formaldehyde are detected. It is proposed that the adsorbed ethylene oxide decomposes on warming to form Mo-methylene species, which participate in metathesis at the elevated temperatures to produce the higher hydrocarbons.

Other surface species have been observed in the reaction of ethene with adsorbed 0-. On Co-MgO, the reaction at 25°C results in a EPR signal assigned to .CCHp--OH- [25]. On ZnO at 90 K, a polymeric radical species of R<H2 has been reported. The EPR signal of this species disappears on exposure to oxygen or evacuation [371.

Adsorbed 0- also reacts with other alkenes. The reactions with propene and 1-butene result in the formation of methyl radicals and linear -CH2CHO- radicals. This indicates cracking of the hydrocarbon [34].

There are few studies on transition metal oxides on species of adsorbed oxygen other than 0-. The associated reaction mechanisms are not well understood. It has been reported that adsorbed 02- reacts with propene at room temperature to form a propene-oxygen complex [38]. On desorption. combustion products are observed from ZnO [39] and copper oxide [40]. The 02- species on Co-MgO is inactive towards ethane [29].

Adsorbed 03- species on Co-MgO reacts with ethane. On heating, ethene is desorbed as a product [29].

REFERENCES

1 . F. Blunt, P. Hendra. and I. Mackenzie, Chem. Commun., p. 278 (1969). 2. M. Che, and A.J. Tench, Adv. Cafal., 31, 77 (1982). 3 . J. Lunsford, Adv. Catal., 8, 135 (1973). 4. A. Bielanski, and J. Haber, Adv. Catal., 19. 1 (1979). 5. M. Iwamoto, Y. Yoda, N. Tamazoe. and T. Seiyama, J. Phys. Chem., 82, 2564

6. B. Halpem, and I. Germain. J . Catal., 37. 44 (1975). 7. M. Iwamoto, Y. Yoda. N. Yamazoe. and T. Seiyama, Bull. Chem. SOC. Jpn.. 51,

8. B. Yang, and H. Kung, J. Catal., 75. 329 (1982). 9. M. Iwamoto, Y. Yoda. N. Yamazoe, and T. Seiyama. J. Phys. Chem. Solid,

(1978).

2765 (1978).

80, 1989 (1976). 10. R . Kokes, Proc. 3rd Intern. Cong. Catal., Amsterdam, 1964. p. 484. 1 1 . J. Lunsford and J. Jayne. J. Chem. Phys. 44. 1487 (1966). 12. A. Tench and T. Lawson, Chem. Phys. Lett., 8, 177 (1971). 13. C . Naccache. P. Meriandeau, M. Che, and A. Tench, Trans. Faraday SOC., 67.

14. V. Shvets, V. Vorotyntsev. and V. Kazansky. Kinet. Catal., 10. 356 (1969). 15. K. Seshadri, and L. Petrakis. J. Phys. Chem., 74. 4102 (1970). 16. K. Dyrek, Bull. Acad. Pol. Sci.. Ser. Sci. Chim., 21, 675 (1973). 17. M. Dufaux, M. Che. and C. Naccache, C.R. Acad. Sci., Paris, IS]. 268, 2255 (1969).

506 (1971).

120

18. D. S. McIver, and H. Tobm, J. Phys. Chem., 64. 451 (1960). 19. H.-C. Liu, L. Yuan, and S. W. Weller. J. Catal.. 61, 282 (1980). 20. S. W. Weller, k c . Chern. Res.. 16. 101 (1983). 21. A. Lopez Agudo, F. Llambias, P. Reyes, and J. Fierro, Appl. Cat& 1, 59 (1981). 22. N. K. Nag, J. Catal., 92. 432 (1985). 23. J. Valyon, and W. K. Hall. J. Catal., 84, 216 (1983). 24. B. Yang, and H. Kung, J. Catal., 77, 410 (1982). 25. A. M. Volodin. and A. E. Cherkashin, React. Kinet. Catal. Lett., 18, 243 (1981). 26. N. I. Lipatkina, V. A. Shvets, and V. B. Kazanskii, Kinet. Catal., 19. 979 (1978). 27. S. Kaliaguine, B. Shelimov, and V. Kazansky, J. Catal.. 55, 384 (1978). 28. M. B. Ward, M. J. Lin, and J. H. Lunsford, J . Catal., 50, 306 (1977). 29. K. Aida. M. Tajima, M. Isobe, and T. Onishi, Proc. Intern. Cong. Catal.,

30. V. B. Kazanski, V. A. Shvets, M. Ya Kon, V. V. Nikisha, and B. N. Shelimov,

31. V. A. Shvets. and V. B. Kazansky. J. Catal., 25, 123 (1972). 32. A. R. Gonzalez-Elipe, C. Louis, and M. Che, J. Chem. SOC. Farad. Trans. 1,

33. V. B. Sapozhnikov, V. A. Shvets. N. D. Chuvyllcin. and V. B. Kazansky.

34. J. F. Hemidy, and A. J. Tench, J. Catal., 68, 17 (1981). 35. V. A. Shvets, V. B. Sapozhnikov, N. D. Chuvylkin. and V. B. Kazansky.

36. M. Anpo, and Y. Kubokawa, J. Catal., 75, 204 (1982). 37. A. M. Volodin, and A. E. Cherkaskin. React. Kinet. Catal. Lett., 20, 347 (1982). 38. T. Tabasaranskaya, A. A. Kadushin, K. N. Spiridonov, and 0. V. Krylov.

39. A. A. Davydov, A. A. Yefremov, V. G. Mikhalchenko. and V. D. Sokolvoskii,

40. V. Mikhalchenko, V. Sokolvoskii, A. Filippova, and A. Davydov,

8th. vol. ID, 335 (1984).

Proc. 5th Intern. Cong. Catal., (1972), 2, 1423 (1973).

78, 1297 (1982).

Kinet. Catal.. 17. 1251 (1976).

J . Catal.. 52, 459 (1978).

Kinet. Katal., 13, 1370 (1972).

J . Catal., 58, 1 (1979).

Kinet. Katal., 14, 1253 (1973).

Chapter 8

PREPARATION OF OXIDES

8.1 GENERAL CONSIDERATIONS

Oxides can be prepared in the form of single crystals, or polycrystalline or amorphous samples. For studies where surface atomic arrangements need to be known, single crystal samples are used from which well-ordered crystallographic surface planes can be prepared. There are many well established methods to grow single crystals. The most common method is growth from the melt. In this method, very high temperatures are required, and the oxygen partial pressure (or oxygen fugacity) above the melt must be carefully controlled to ensure that the crystal attains the desired stoichiometry. When the melting point of the oxide is so high that the oxide decomposes, other lower temperature methods may be employed, such as vapor transport. Many books are available on the various methods of crystal growth [l]. In this chapter, we concentrate on methods to prepare powder samples.

For most practical purposes where a large specific surface area is essential, polycrystalline or amorphous samples are desirable. Experience has shown that in general, low temperature processes are necessary to obtain oxides of large surface areas or small particle sizes. However, methods can be developed such as in the preparation of aerogels in which high temperatures are used [2a]. A low temperature treatment does not necessarily lead to the formation of the thermodynamically most stable phase. Instead, depending on the details of the preparation procedure, metastable phases and/or amorphous samples may be obtained. Unfortunately, there is still little understanding of the reasons behind the formation of a particular phase, and each case has to be treated individually. A low temperature treatment may also lead to the formation of hydrous oxides from precursors obtained by aqueous precipitation. The water in the hydrous oxides can be removed by treatment at elevated temperatures.

Broadly speaking, polycrystalline or amorphous samples of oxides are often prepared from a solution containing the appropriate precursor compounds. There

121

122

SOLUTION CONTAINING PRECURSOR CATIONS

/-2 \ \ COPRECIPITATION COMPLEXATION GELFORMATION GALAXING WITH

1 (acid, alcohol addn) I OXIDE PRECURSORS (AGING) 1

1 DRYING

I WASHING

(AGING) 1

DRYING .1 1 THERMAL

THERMAL ACTIVATION ACTIVATION (decomposition)

\ L /

1 (AGING)

1 DRYING

REMOVAL OF 1 VOLATILE COMPDS EXTRUSION OR

1 BALL MILLING

1 (AGING) 1

THERMAL THERMAL ACTIVATION ACTIVATION

/ (solid-state rxn) ADDITION OF

3THER ELEMENTS 1

impregnation malaxing

ball milling 1

(aging) 1

1 DRYING

THERMAL ACTIVATION

FORMING PROCESS (optional) \

(tabletting. extrusion) 1

(aging) 1

1 DRYING

THERMAL ACTIVATION

FINAL OXIDE 1

Figure 8-1 Common methods for the preparation of mixed oxides. (From "Preparation of Catalysts. III", G. Poncelet, P. Grange, and P. Jacobs, ed., Elsevier Publishers, 1983, p.185. copyright Elsevier Scientific Publishers).

OXIDE PREPARATION 123

are many methods to transform the metal ions in the precursor solution into the desired oxides. Fig. 8-1 illusuates the most commonly employed methods. Whether by precipitation, gelation, or complexation, a solid containing the metal ion of the oxide is first separated from the solution. This precipitate is then washed and dried, and finally calcined and perhaps reduced to form the desired oxide. Sometimes other components are added along the process. The oxide at this point is a powder. It can be formed into pellets or tablets if needed.

However, different techniques of forming an oxide may result in samples of very different surface properties. This is because the crystallographic form of the oxide obtained, its morphology, its composition, and its surface structure may vary depending on the preparation procedure. Any one of these factors can have a significant effect on the surface property. In this chapter, examples to demonstrate this will be presented.

The preparation procedure may also determine the pore volume and pore size and structure of the oxide particles. Although no detailed study is available for transition metal oxides in general, it is expected that the applicable principles should be similar to those for alumina, which has been recently reviewed [2]. It is mentioned that the sedimentation and flocculation of the oxide colloidal particles, the procedure of drying the wet precipitate agglomerates, the aging process, the peptising agent, and the nature of any additives are among the factors that affect the pore size of the oxide. Therefore, attention to details is important in the preparation.

8.2 PREPARATION OF UNSUPPORTED SINGLE COMPONENT OXIDES

Single component oxides may be prepared by thermal oxidation of the metal. Usually, the metal samples arc available in forms that have low surface areas such as pellets, course powder, or foils. The kinetics of oxidation is slow because lattice diffusion of metal and oxygen ions is necessary. Therefore, high temperatures are used to obtain reasonable rates. The result is that the surface areas of the oxides obtained are low.

For many practical purposes where the available surface area is critical, the oxides are commonly prepared by the decomposition of precursor compounds containing the desired metal ion in air or oxygen. Usually, preference is given to precursor compounds that decompose at low temperatures to minimize sintering of the resulting oxide which causes low surface area. Carbonates, bicarbonates, hydroxides, nitrates and oxalates are usually preferred over sulfates for this reason. The decomposition may also be carried out in vacuo to lower the decomposition temperatures. The desired precursor compounds can usually be obtained by precipitation from aqueous solutions of the readily available compounds. When the readily available compounds of some elements (such as titanium) are not soluble in water, hydrolysis of nonaqueous solutions or direct vapor phase oxidation of some compounds may be employed.

The method to prepare aerogels has also been successfully applied to prepare

124

some high surface area transition metal oxides, notably chromia. In this method, a hydroxide precursor is first precipitated from an alcoholic solution by hydrolysis of an akoxide, for example, with a controlled amount of water. The alcohol solvent is then removed by evacuation at a temperature above the critical temperature of the solvent. An alcoholic solvent is used instead of water because of its lower surface tension so that the gel structure can be retained from collapsing during drying. If the hydroxide precursor is prepared by hydrolysis in an aqueous solution, the water should be first replaced by an alcoholic solvent before evacuation. Otherwise, dissolution of the oxide in water at the high temperature may occur [2a].

Crystallographic Phases It is rather common that different preparation methods result in oxides of

different crystallographic phases. One example that is of significant industrial interest is the preparation of titanium dioxide (Ti02). Titanium dioxide exists in three crystallographic forms: rutile, anatase and brookite. The brookite phase is seldom formed. Although the rutile phase is generally considered the most stable form, there are thermochemical data which indicate that anatase may be more stable than rutile [3].

Common titanium compounds are not soluble in water. Very pure Ti02 can be prepared by hydrolysis of prepurified TiCb. Ti(1V) sulfate, or Ti akoxides. The resulting solid is best regarded as hydrous titanium oxide. Subsequent calcination in air produces Ti02. Normally, rutile is the predominant form, but anatase is often present in minor amounts. The anatase form can be produced by hydrolysis of Ti halides at temperatures around 600°C, or via low temperature calcining (-700°C) of precipitated titanic acid H2Ti03 .nH20 [4].

Ti02 can also be prepared by vapor phase oxidation of TiC12. TiH2 can also be used [5].

ZrO2 crystallizes in the monoclinic and tetragonal forms. The absorption of mechanical energy by the Zr02 particles finely dispersed in A1203 when they transform from the tetragonal to the monoclinic form is the reason for the enhanced fracture toughness of such alumina. Davis has investigated the final crystallographic phase of ZrOz as a function of the pH of the precipitation mixture [6]. The results are shown in Fig. 8-2. The hydrous Zr02 precipitates are prepared by adding W O H . NaOH or KOH rapidly to a zirconyl nitrate solution to bring it to a desired pH. The gel is then calcined to 600°C to yield Zr02. Precipitation at a pH of 3-4 or 13-14 results in a nearly 100% tetragonal phase. On the other hand, at pH 7-10, over 90% of the ZrOz is in the monoclinic phase.

The phases obtained by hydrothermal treatment of amorphous hydrous z102 gel depends on the other salts present in the mixture. Only the monoclinic phase is produced using KF or NaOH, while a mixture of both phases is produced using LiCl, KBr or no salt [7]. Monoclinic Zr02 is also obtained by heating stabilized colloidal zirconyl nitrate in concentrated nitric acid [8].

Molybdenum trioxide also exists in two different crystallographic forms, orthorhombic and hexagonal. If the oxide is prepared by th addition of HN03 to an aqueous solution of ammonium heptamolybdate to form a precipitate which is


.I NHiOH "1 dKOH NaOH

0 8 i - z 0 0,

0,

C 0

' 0 4

06

c

- c

P LL

0 2

2 4 6 8 1 0 4 2 4 4

PH

Figure 8-2 Fraction of tetragonal zirconia in materials obtained by calcining at 6OO0C gels precipitated from solutions with NH40H. NaOH, or KOH versus h a 1 pH of the solution. (From J. h e r . Ceram. Soc., 67, C-168 (1984). copyright American Ceramic Society).

heated to above 300"C, the orthorhombic form is produced predominantly. If the precipitate is heated to below 250"C, hexagonal Moo3 is obtained predominantly [91.

Morphology of Oxide Crystallites The preparation method determines the morphology of the final product as

well as its crystallographic phases. In general, the higher the temperature of calcination, the coarser and the more crystalline is the final powder. Some examples of preparing different morphologies are described bclow.

ZnO is the first example [lo]. If ZnO is prepared at high temperature by ignition of zinc in an oxidizing atmosphere, a three dimensional crystalline sample of a low surface area is obtained. Electron microscopy shows that thc sample consists of crystallites of about 700 nm long and 300 nm wide, exposing the prismatic face on the long side, and the polar faces on the short side. Small crystals grown from vapor transport are needles of hexagonal cross-section, about 0.1 cm long and 0.001 cm in diameter. The long side exposes the prismatic face. A powder preparcd from ZnC03 precipited from a nitrate solution with sodium carbonate followed by washing and conversion to oxide by calcining has a high surface area about 36 m2gp1. The resulting oxide contains hexagonal platelets about 15 nm across. The polar face represents about 33% of the exposed surface. A powder prcpared by the decomposition of zinc acetate contains particles grown

126

mostly along the (0001) orientation, while that by the decomposition of zinc formate mostly contains particles grwon along the (1 120) direction [ 111.

The surface area of ZnO prepared by the decomposition of ZnC03 precipitated from a nitrate solution depends on the calcination procedure. If the water trapped in the prccipitate is removed efficiently during heating of the precipitate (ZnC03), a high surface area sample (above 30 m2g-') is obtained. Otherwise, the surface area is much lower (10 m2gA1 or lower),

Haematite a-Fe203 is another example [12]. It is reported that the dish- bution of the exposed planes depends on the precipitation method and the temperature of oxygen pretreatment. Precipitation at high pH from aqueous ferric nitrate leads to the formation of geothite (a-FeOOH) with a smaller amount of haernatite. Mainly haematitc is formed at low pH [13-151. Precipitation from a ferric chloride solution leads to the formation of P-FeOOH, particularly when the precipitate is aged in contact with a solution containing chloride ions [14-171. Subsequent heating of these different precipitates results in a-Fe203 particles that possess different proportions of hydroxyl groups of different IR absorption frequencies. The latter is an indication that different amounts of various crystal faces result from the various preparations [ 121.

In addition to preparing oxides with the cxposure of different crystallographic planes in the crystallites of different samples, it is also possible to prepare oxide particles of different macroscopic shapes by careful aging of a solution of a particular composition at a particular temperature and for a particular time [18]. Table 8-1 summarizes the examples in the literature which show that depending on aging conditions, spherical, cubic, spindle or disk-like a-Fe203 can be obtained. It is noted, though, that each of the uniformly shaped macroscopic particles consists of many subparticles. That is, the macroscopic particles are not single crystals. At this point, except for the fact that slow aging is required, there is little understanding of why different but uniform shapes are obtained under different conditions. In general, if the hydrolysis products consist of discrete well-defined ionic complexes, crystal growth occurs, yielding particles of fixed stoichiometry and most often of well-defined crystal habits. If the hydrolysis results in the formation of polymeric metal complexes, spherical particles are produced.

Indeed, low temperature hydrolysis has been a method to produce high surface area oxide gels. For example, chromia gel has been produced by the hydrolysis of chromium nitrate in an urea solution. Heating the solution to decompose the urea permits slow and uniform hydrolysis throughout the solution. The gel precipitates after boiling the solution for a few hours [19]. Drying at low temperatures results in a high surface-area gel. Evacuation may help the drying. Use of organic instead of aqueous solvents results in gels of larger pores because of the lower surface tension of organic solvents. Other hydrolysis methods can also be used. For example, hydrolysis of tris(2,4-pentanedionate) chromium (Cr(a~ac)~) followed by precipitation with ammonium hydroxide, or hydrolysis of chromium ethoxide by water and ammonium hydroxide has been successfully attempted [20]. Such chromia gel is mostly amorphous to X-ray diffraction. Only a small fraction is crystalline.


Table 8-1 Examples of the Dependence of the Morphology of an Oxide on the Preparation Method.

Material and Aging solution Aging temp. Reference1 Morphology ture and time

spherical 0.032 M FeClz 100"C, (- 0.1 to 0.5 p) hematite (a -Fq03) sol

+ 0.005 M HCl 2 weeks

cubic (- 1 pn) FeC13 in HzO. 100°C. days a- Fez03 HzO-ethanol, or

HzO-methanol soh.

disklike 0.0040 M Fe(NQ)3 50°C. 1 h (- 1 x 10 pm) a-FqO, + 1.20 M NaOH

+ 0.20 M triethanolamine

+ 0.50 M Hz02

spindle a-FeZO3 Fe3+ + N a H z Q 100°C. days

spherical amor- Crz(S04h 75°C phous chromium hydroxide

Cr3+ + KHZPO4 but 75°C not in the presence of ( C l , NO3- or acetate)

spherical rutile TiC14, S04z- 98"C, days

spherical (< 1 pm) Soh. of FeS04 and Ni,Fq-x04 KOH in OZ-free soh.

well-formed 0.02 M Fe(NQh 25OoC, 2 h crystallites + 0.40 M TEA Fc304 + 2.4 M NaOH

+ 0.85 M NzH4

E. Matijevic and P. Scheiner, J.C.I.S. 63, 509 (1978).

S. Hamada, E. Matijevic, J.C.I.S.. 84, 274 (1981).

E. Matijevic, Acc. Chem. Res., 14, 22 (1981).

M. Ozaki, S. Kratohvil, and E. Matijevic J.C.I.S. 102, 146 (1984).

R. Demchak. and Matijevic. J.C.I.S. 31, 257 (1969).

E. Matijevic. A. D. Lindsay, S. Kratohvil. M. E. Jones, R. I. Larson, N. W. Caycy, J.C.I.S., 36. 273 (1971).

E. Matijevic, M. Budnik, L. Meites, J.C.I.S., 61, 302 (1977).

A. E. Regazzoni, and E. Matijevic. Corrosion, 38, 212 (1982).

R. S. Sapieszko. E. Matijevic J.C.I.S.. 74, 405 (1980).

128

Table 8-1 continued

polyhedra

CUO

needle ZnO (- 1 x 50 pm)

(- 20 Pm)

VzOs leaflets

crystalline cobalt oxide

crystalline Zr02 (16 nm) mixture of monoclinic and tetragonal

monoclinic 6 nm ZrOz

spherical (-1 pm) A1 hydroxide sol

ferrous hydride gel containing excess

FeS04

0.040 M CU(NO:,)~ + 0.20 M HEDTA + 1.20 M NaOH

0.040 M Zn(NQ)z + 0.20 M TEA + 1.2 M NaOH + 0.85 M N2H4

0.040 M Na3V04 + 0.20 M HEDTA + 1.20 M NaOH + 0.85 M N2H4

acetate

amorphous hydrous Z r O z in distilled water with or without added KF, NaOH. LiCI. or KBr

90°C

100"C, 70 min

25OoC. 2 h

25OoC, 3 h

200°C

stabilized colloidal 150°C. zirconyl nitrate in 14 days HN03

0.002 M AIz(SO3)3 97"C, 48 h

E. Matijevic, J.C.I.S., 58, 374 (1977).

E. Matijevic, Am. Chem. Res.. 14, 22 (1981).

ibid.

ibid.

T. Sugimoto. E. Matijevic J. Inorg. Nucl. Chem., 41, 165 (1979).

E. Tani, M. Yoshimura, S . Somiya, J. Amer. Ceram. SOC., 66, 11 (1983).

P. Morgan, J. Amer. Ceram. SOC., 67. C-204 (1984).

R. Brace and E. Matijevic, J. Inorg. Nucl. Chem., 35, 3691 (1973).

Footnote: a) J.C.I.S. stands for J. Coll. Interface. Sci.


8.3 PREPARATION OF SUPPORTED OXIDES

There are a number of reasons to use supported oxides instead of unsupported ones. One reason is to increase the surface area. Often oxides are prepared by calcination at relatively high temperatures that result in significant sintering or loss of surface area. A suitable support can reduce sintering. Sometimes a support is used as a heat conduction medium. This may be particularly important in oxidation reactions which have large heats of reaction. A support of high thermal conductivity helps to remove the heat from the reaction site, and reduces hot spots which usually degrade selectivity. Silicon carbide has been used for this purpose. In a number of systems, the support provides a template for the oxide to form in a certain desirable morphology, such as the exposure of one predominant plane. V2O5 supported on anatase Ti02 is a well known example [21,22]. V205 forms a layer structure on TiO2, preferentially exposing the (010) plane which contains the V=O groups believed to be important in oxidation reactions. In this regard, the exact nature and structure of the support can be critical. As described in Chapter 2, V2O5 forms a layer structure much more readily on anatase TiOz than Si02 [21]. Similarly, Moo3 forms a monomolecular layer structure on A1203 but not on SiO, [231.

It is common that an oxide support interacts strong with the oxide. The formation of a layer structure on the support is a result of such strong interaction. The reducibility of the oxide is often affected by a strongly interacting support (see Chapter 6). The strong interaction may also cause morphological changes in the support. The surface area of V205 supported on anatase Ti02 is much smaller than that of the original support, whereas that of Moo3 supported on the same Ti02 has the same surface area [24].

Supported oxides are often prepared by impregnation or ion-exchange of a precursor compound onto the support. Ion-exchange can be used for oxides that have precursor compounds that dissolve in water and form ionic complexes that contain the desired metal. This requirement usually excludes the preparation of supported Ti or Zr oxides by ion-exchange. Schematically, the process involves ionization of surface hydroxyl groups of the support in the aqueous solution. The desired ion is then attached to the ionized hydroxyl group by electrostatic forces. The process for high pH is:

S-OH + OH- + S - 0 - + H20 (8-1)

The corresponding process at low pH is:

S-OH + H+ 4 S-OH,' (8-3)

S 4 H 2 + + (M)"- --+ S-OH,+(M)"- (8-4)

In these equations, S-OH stands for the hydroxyl group of the support, and

130

(I@"+ Or "- is the ionic complex. For most metals, cationic amine complexes are readily available. For Mo and V, the anionic 0x0 complexes are often used [25]. Subsequent calcining of the ion-exchanged material produces the oxide. This method produces highly dispersed oxides because electrostatic repulsion keeps the complexes apart. The loading obtainable is also low for the same reason.

In impregnation, a solution of the salt containing the desired cation is added to the support. If the volume of the solution is just enough to fill the pore volume of the support, the method is called impregnation by incipient wetness. Excess solution may also be used. The solvent is eventually removed by drying, and the oxide is formed after calcination.

Impregnation is the most widely used method because it is applicable to practically all oxides. The possible use of either aqueous or nonaqueous solvent is the major factor of the versatility of the method. The loading can be varied over a wide range by changing the concentration and the amount of solution used. The loading as well as the details of the drying and calcination conditions affect the dispersion of the final oxide.

The pH of the ion-exchange or impregnation solution is one of the factors that determine the nature of the final oxide. In addition to changing the extent of protonation or dissociation of surface hydroxyl groups, different pH's may also change the species in the solution. For example, at pH c 1, vanadyl cation (VO,') is the major species in solution. At pH 4, the major species is a three-dimensional decavanadate, [HVI 0028]s-. At pH 7, two-dimensional tetra- and trivanadate species are present, and at pH 9 and 12, divanadate and monovanadate species are formed, respectively [26]. At high pH's, dissolution of the support may occur.

Special techniques are sometimes used for better wetting of the support by the oxide precursors. For example, spraying in vacuo a solution of the precursor onto the support has been attempted [27] to prepare silica-supported Fe203-Mo03. Precursors may also be introduced in the gas phase. For example, alumina- supported V205 has been prepared by treating the A1203 with a stream of N2 containing VOC13. The VOC13 is then hydrolyzed with steam diluted with N2 [21]. The process can be repeated to increase loading. Organometallic compounds may also be used in the gas phase method. For example, cyclopentadienyl zirconium can be used as precursor for Zr02, and x-ally1 chromium for Cr203. In these cases, the Zr or Cr precursors are anchored onto the surface by the hydrolysis reaction of the organometallic complexes with the surface OH groups. Since this hydrolysis reaction involves specific interaction between the compound and the surface OH groups, the compound may also be introduced in an inert solvent. Contacting a solution (normally a thoroughly dehydrated hydrocarbon or ether solution) of the compound with the solid support often works well. With this technique, Ti02 supported on A1203 can be prepared using a heptane solution of TiCI4 or Ti(OC3H7)4 [28], and Nb205 supported on Si02 using a hexane solution of Nb(OC2Hs)s [291.

Repeated ion-exchange, repeated impregnation, or a combination of ion- exchange and impregnation are sometimes employed to increase the final loading. It may be noted that because of the pH of the impregnation solution, sometimes ion- exchange occurs during impregnation. This happens when alumina is impreg-


nated with a solution of ammonium heptamolybdate of low pH. A highly dispersed molybdenum oxide is thus obtained [30].

It is important to recognize the possibility that the support, which is almost always another oxide, may react with the oxide. Among the common supports, Si02 may form silicate with many oxides, especially those in the first transition period, A1203 may form aluminates, and Ti02 titanates. In fact, the formation of new compounds is often thermodynamically favorable but kinetically limited. Thus high temperature calcination enhances their formation. In some cases, such as supported NiO, this leads to the apparent increase in the resistance to reduction of the oxide [31].

For TiO,, there is the additional complication that phase transformation of titania may be enhanced by the presence of a second oxide. It has been reported that the transformation of anatase to rutile is enhanced by the presence of vanadium oxide [32]. In another study, vanadia-titania catalysts were prepared by coprecipitation. It was shown that the different precipitation conditions and different subsequent treatments led to widely different fractions of titania being in the rutile form [33].

Like unsupported oxides, the crystallographic form of a supported oxide may depend on the preparation method. This has been reported to be the case for a silica-supported iron oxide. A sample prepared by hydrolysis of a mixed solution of ethyl silicate and iron (111) nitrate in ethylene glycol consists of very small crystallites of y-Fe203 on the SiOz support, but one prepared by impregnation of silica gel with aqueous iron (111) nitrate consists of a-F@03 particles with a broad size distribution [34].

8.4 PREPARATION OF MULTICOMPONENT OXIDES

Since the number of possible compounds that can be formed increases rapidly with the number of components, multicomponent oxide systems can be very complex. Since catalysts are often prepared at as low a temperature as feasible to preserve the surface area, the formation of compounds is often kinetically controlled instead of thermodynamically controlled (that is, the compounds observed may not be the most thermodynamically stable ones). Thus the overall chemical composition of the oxide can be used only as a guideline for the possible compounds present. It is not uncommon to find in a mixture of compounds some of stoichiometries quite different from the overall composition. However, this may be desirable in some cases when a multi-phase is necessary for the desired selectivity and activity of a catalyst. Nonetheless, this behavior makes it more difficult to reproduce the properties of a catalyst, whose behavior is vulnerable to small changes in the preparation procedure.

In general, the various methods shown in Fig. 8-1 are used to prepare multicomponent oxides. The most desirable method would be one which results in a solid of a reasonable surface area, the desired stoichiometry, the desired number of phases, and, very importantly, the desired spatial distribution of these phases. The latter is important in catalytic reactions where the diffusion of

molecules is critical to selectivity. There is not yet a universally most desirable preparation method for all oxides.

The traditional ceramics preparation technique of ball-milling the appropriate ratio of the component oxide powders followed by high temperature calcination to bring about compound formation is seldom used in the preparation of high surface area samples. This is because the reaction rates between component oxides are usually small. Thus a high temperature and a long heating time are needed which cause the formation of low surface area compounds. To overcome this problem, instead of mixing component oxides, precursor compounds that contain the desired metals and decompose at low temperatures are used. Carbonates and oxalates are among the precursor compounds. For example, CaMn03 and Ca2Mn04 can be prepared by the oxidative decomposition at 1000°C of CaMn(C03)2 and Ca2Mn(C03)3, respectively [35]. High surface area solid solutions of NiO-MgO and Coo-MgO are obtained by impregnating Mg(OH)2 with Ni or Co nitrate. The mixture is then thermally decomposed and annealed in vacuo at 1000°C [36]. Intermetallic compounds have also been used as precursors.

Although the ion-exchange method may be used to prepare multi-component oxides, the need to know the ion adsorption coefficients so as to control the stoichiometry of the final product makes the method less convenient to use. Impregnation does not have this limitation. A solution containing the desired ratio of the compounds can often be made. By virtue of the method, all solute deposits onto the support. Repeated ion-exchange, repeated impregnation, or a combination of ion-exchange and impregnation can also be used.

In these methods, the desired final compound is formed without the necessity of very severe heating because the metal ions are already brought into intimate mixing in the precursors. To achieve good intimate mixing of the metal ions in the precursors is also the goal of many other techniques. Freeze-drying, for example, involves spraying a solution that contains decomposable compounds of the desired metal ions into liquid nitrogen. Because of the rapid freezing, the precursor compounds form very small particles and little segregatory crystallization takes place. Upon drying in vacuo such that the solvent is removed without melting the solid (which sometimes may exist as a glassy solid), the intimate mixing of the precursor compounds is preserved. Among many examples, Zn ferrite has been prepared with this technique [37].

Spray-drying employs a similar principle as freeze-drying. Here, a solution of precursor compounds is sprayed into a hot chamber. The solvent is rapidly evaporated, and the precursor compounds form fine particles. Again, because of the speed of the drying process, segregation of the components is minimized. Both spray-drying and freeze-drying are commonly employed techniques. They are applicable to systems that have precursor compounds that are soluble in a common solvent.

A variation to freeze-drying and spray-drying is to make use of complexation. In the preparation of a Cu-Co-Al-Zn oxide, an organic acid (citric, malic, tartaric, glycolic, or lactic) is first added to the metal nitrate solution. The solution is then evaporated in vacuo to yield a glassy amorphous material. This is then calcined WI.


1 I I I 1 I I 0 0.4 0.8

2- Amount of CO Added X i [ M " + ] 3

Figure 8-3 Neutralization curve of a nitrate solution of Cu", A13+, and ZnZf (0.35 M) with 0.4 M sodium carbonate. (From the same source as Figure 8-1, copyright Elsevier Scientific Publisher).

One advantage of the techniques mentioned thus far, except ion-exchange, is that all of the starting material is recovered in the solid. Thus the overall stoichiometry of the final compound is known with certainty during preparation. Such is not the case with another commonly employed technique of coprecipitation.

In coprecipitation, a precipitation agent which is often hydroxide or carbonate of ammonium or sodium is added t~ a solution containing the precursor compounds. Hydroxides or carbonates of the metal ions are precipitated out of the solution. They are filtered, washed, and then calcined to form the oxide. The homogeneity and the degree of atomic mixing of the metal ions in the precipitates depend on the solubility products of the compounds involved. As an example, a Cu-Zn-A1 oxide catalyst is prepared by coprecipitation of a solution containing Cu, Zn, and A1 nitrates with sodium carbonate. The precipitation curves of Cu, Zn, and A1 hydroxycarbonates are shown in Fig. 8-3 [39]. It can be seen that the precipitation curves are so differcnt that if Na2C03 is added slowly to the solution of nitrate mixture, Al, Cu, and Zn hydroxycarbonates would precipitate out quite independently. In such a situation, it is advisable to add the nitrate solution slowly into a sodium carbonate solution. It is also interesting to note that the precipitation point for A1 and Cu hydroxycarbonates shown in Fig. 8-3 are much lower than those if the hydroxycarbonates are precipitated separately from a solution containing only one metal salt.

When properly performed, coprecipitation can result in a fairly well mixed

134

precursor mixture. This is particularly m e if the precipitate compound contains a mixture of metal ions. In the example of the Cu-Zn-A1 oxide above, the compounds (Cu,Zn)2(OH)2C03, (Cu,zn)dOH)6(CW2, and (Cu,Zn)gA12(OH)16- C03 can be formed [39,40]. These compounds provide intimate mixing of the metallic ions. Upon calcination, well mixed CuO and ZnO, as well as possibly Cu+ dissolved in ZnO are formed.

One more example will be mentioned to demonstrate the importance of the pH in the stoichiometry of the final oxide obtaincd. In the preparation of the NiMo04-Mo03 system, the amount of Moo3 in the final oxide depends on the pH of the solution during the addition of ammonium heptamolybdate solution to the nickel nitrate solution. If the pH is 6, stoichiometric NiMo04 is formed. If the pH is 5 or lower during the addition, and then raised to 5.5 at the end of the addition, excess Moo3 is formed. The amount of excess Moo3 increases with decreasing pH [41].

There are variations to the simple coprecipitation technique. For example, coprecipitation can be performed in the presence of a suspension of another precursor. This is used as one of the methods to prepare a Ag-Mn-Co oxide [42]. Thus a Ag and Co carbonate mixture is precipitated out in a solution containing a MnOz suspension. Other variations such as precipitation of Ag-Mn oxide in a C0C03 suspension can also be used.

REFERENCES

1. M. M. Faktor. and I. Garrett, "Growth of Crystals from the Vapor," Chapman and Hall Ltd., London. 1974; D. Elwell, and H. J. Scheel." Crystal Growth from High-Temperature Solutions," Academic Press Publ., New York, 1975.

2. D. L. Trimm. and A. Stanislaus. Appl. Cafal., 21, 215(1986). 2a."Aerogels", ed. J. Fricke. Springer-Verlag Publ.. Berlin, 1986. 3. F. A. Cotton and G. Wilkinson, "Advanced Inorganic Chemistry," 4th Edition,

John Wiley & Sons, N.Y. 1980. p. 695. 4. "Handbook of Preparative Inorganic Chemistry", Ed. by G. Braucr (Engl.

transl.) Vol. 2, 2nd edition, Academic Press, 1965. 5. V. A. Lavrenko, V. Zh. Shemet, S. K. Dolukhanyan, CA 98382827k 6. B. H. Davis, J. Amer. Ceram. Soc., 67. C-168 (1984). 7. E. Tani, M. Yoshimura. S . Somiya, J. Amer. Ceram. Soc., 66, 11 (1983). 8. P. E. D. Morgan, J . Amer. Ceram. Soc.. 67, C-204 (1984). 9. J.-M. Tatibouet and J. E. Germain, C. R. .4cad. Sci. Paris C. 290. 321 (1980).

10. M. Bowker, H. Houghton. K. C. Waugh, J. Cafal., 84, 252 (1983). 11. G. Djega-Mariadasson and L. Davignon. J. Chem. SOC. Faraday Trans. 1.

12. C. H. Rochester and S. A. Topham, J. Chem. Soc., 1073 (1979). 13. R. J. Atkinson. A. M. Posner. and J. P. Quirk. J. Inorg. Nuclear Chem.,

14. P. J. Murphy, A. M. Posner. and J. P. Quirk, J. Colloid Interface Sci.,

78, 2447 (1982).

30, 2371 (1968).

56, 270, 284, 312 (1976).


15. T. G. Spiro, S. E. Allerton, J. Renner, A. Terzie, R. Bils, and P. Saltman,

16. A. L. Mackay, Mineral. Mag., 32, 545 (1960). 17. K. Kauffman, F. Hazel. J . Coffoid Inferface Sci., 51. 422 (1975). 18. E. Matijevic, Acc. Chem. Res., 14, 22 (1981) and reference within. 19. R. L. Bunvell, Jr., A. B Littlewood. M. Cardew. G. Pass, C. R. H. Stoddart,

20. K. C. Taylor, PhD Thesis, Northwestern Univ.. 1968. 21. Y. Murakami, M. Inomato, K. Mori, T. Ui, K. Suzuki, A. Miyamoto, T. Hattori.

Proc. 3rd Intern. Symp. on Scientific Basis for the Preparation of Heterogeneous Catalysts, G. Poncelet, P. Grange and P. A. Jacobs, ed., 1983. p. 531.

22. G. C. Bond and P. Konig, J. Cafaf.. 77, 309 (1982). 23. N. Giordano, A. Castellan. J. C. J. Bart,, A. Vaghi. and F. Campadelli.

24. M. Del Arm, M. J. Holgado, C. Martin, and V. Rives, J . Catal.. 99. 19 (1986). 25. L. Wang and W. K. Hall, J. Card., 66, 251 (1981). 26. M. Pope, and B. Dale, Quarr. Rev. (London). 22, 527 (1968). 27. M. Carbucicchio and F. Trifiro, J . Cataf., 62. 13 (1980). 28. G. B. McVicker and J. J. Ziemiak. J. Cafaf., 95, 473 (1985). 29. E. J. KO, R. Bafrali, N. T. Nuhfer, and N. J. Wagner, J. Cafal., 95. 260 (1985). 30. M. Houalla, C. L. Kibby, L. Petrakis, D. M. Hercules, J. Caraf., 83, 50 (1983). 31. M. Houalla, Proc. 3rd Intern. Symp. on Scientific Basis for the Preparation

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32. G. C. Bond, and A. Sarkany, J . Cafaf., 57, 476 (1979). 33. W. E. Slinkard and P. B. DeGrooS J . Calaf.. 68. 423 (1981). 34. T. Ida, H. Tsuika. A. Ueno, K. Tohji, Y. Udagawa, K. Iwai. and H. Sano.

35. H. S. Horowitz and J. M. Longo. Mat. Res. Buff., 13, 1359 (1978). 36. A. P. Hagan, M. G . Lofthouse, F. S. Stone and M. A. Trwethan, in "Scientific

Bases for the Preparation of Heterogeneous Catalysts, II", B. Delmon, P. Grange, P. Jacobs, and G. Poncelet, ed., Elsevier Scientific Publ., 1979, p. 417.

84, 382 (1980).

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G. Poncelet, P. Grange, and P. A. Jacobs, ed., Elsevier Scientific Publ., 1983, p.225.

Chapter 9

METATHESIS AND ISOMERIZATION

9.1 METATHESIS

In fro ductio n When a mixture of two alkenes is exposed to a metathesis catalyst, the

substituents of the two alkenes are interchanged to form new alkenes according to the reaction:

Usually, both the E and the Z isomer of each of the four products shown are obtained. These products can be predicted by assuming that each alkene molecule is dissociated into two (R)2C fragments by breaking the C=C bond. These fragments recombine with other fragments to form the products. One must emphasize that the actual reaction mechanism is quite different from this picture. Since the discovery of highly selective catalysts for metathesis, this reaction has been used to prepare specific alkenes, particularly internal alkenes, from more commonly available ones.

Metathesis does not always result in the production of new alkene molecules with a different carbon chain length or with a diffcrent position of the carbon- carbon double bond. When this happens, the reaction is dcgenerative (nonproductive). Reactions (9-2) and (9-3) are two examples of degenerative mctathesis:

136

METATHESIS AND ISOMERIZATION &37

CH2=CHCH3 + CD2=CDCD3 ---+ CH2-CDCD3 + CDZ=CHCH2 (9-2)

(9-3)

Concurrent with reaction (9-2), productive metathesis occurs that produces cthene and butene as products:

CH2=CHCH3 + CD2XDCD3 + CH2=CD2 + CD3CMHCH3 (9-4)

It can be seen that according to reaction (9-3), metathesis is also a route for cis- trans isomerization.

Co-Mo oxide supported on A1203 is the first transition metal oxide heterogeneous metathesis catalyst reported [ 11. Since then, many other oxidic catalysts have been discovered [2,3], including MoO3-Cr2O3/AI2O3. Mo03/Si02, W03/Si02, W03/A1P04, Re207/A1203, Ta20s/Si02, Nb20s/Si02, and Te03/Si02. Among these, those based on Moo3, W03 and R%07 are the most active and selective, with Re207 being active at close to room temperature. Progress on this reaction has been reviewed [2,3]. Some nonoxidic catalysts are also very active, especially those prepared from molybdenum hexacarbonyl.

Mechanism It is generally accepted that metathesis proceeds via a metallocarbene

intermediate. For example, consider the metathesis of a mixture of propcne and propene-d6. The reaction of propene-d6 with the catalyst results in the metallocarbenes I and 11:

M q D 2 MXDCD3

Addition of propene to I generates a metallocyclobutane the decomposition of which leads to cithcr dcgencrative (eq. 9-5) or productive (eq. 9-6) metathesis, depending on the orientation of the propene molecule. Analogous reactions can be written for species 11.

Dcgencrative:

CH3 I

138

Productive:

One characteristic of this mechanism is the retention of the identity of the CR2 fragments. Indeed, when the center carbon of propene is labeled, the label is found in the product butene but not in ethene [4]. A mixture of C2H, and C2D4 yields exclusively C2HzD2 as the product. No isotopic mixing is detected [5]. Similarly, the products from a mixture of propene-d,, and propene-4 are almost exclusively ethene-d,,, d2 and 4, propene-d,,, d2, 4, and 4, cis-2-butene-4, 4, and d8, and trans-2-butene-d,,, 4, and ds [6]. Thus the reaction proceeds via transalkylidenation and not transalkylation.

Another possible mechanism, which involves a cyclobutane intermediate (eq. 9-7) has not received much support. Cyclobutane only reacts very slowly over a metathesis catalyst.

H2C=CH2 + D2C=CD2 + H2yTDZ --+ 2 H2C=CD2 (9-7) H2C-CD2

There are other evidence that support the involvement of some organometallic species (such as I or II) in this reaction. It has been observed that prereduction of the catalyst by propene or butene produces a very active catalyst, otherwise an induction period is observed before the catalyst becomes active. Strong evidence for the metallocarbene intermediate is provided by homogeneous catalysis where the intermediate is isolated and shown to be an active catalyst [7,8]. It appears that the carbene-intermediate mechanism can explain all of the available data on metathesis, including that of cyclopropane which may proceed by first isomerizing to propene [9,10], and stereospecific metathesis of terminal alkenes [ 131.

The rate of metathesis depends on the substituents around the double bond due to different degrees of steric hindrance and different polarities of the intermediates. The sample data shown in Table 9-1 show that on reduced M003/A1203, butenes and propene react much faster than ethene [l l] . It has also been found that degenerative metathesis of' propene and 1-butene is about ten timcs faster than productive metathesis [6], and the rate of degenerative metathesis of cis-2-butene is slower than that of 1-butene [ll]. This latter observation is different from those reported in homogeneous systems where the activity sequence in decreasing ordcr is: degenerative metathesis of terminal alkene > cross metathesis of terminal and internal alkenes > metathesis of internal alkene > and productive metathcsis of terminal alkene [12]. In a mixture of alkenes, the alkene with a higher metathesis rate tends to suppress the metathesis of the alkenc of a

METATHESIS AND ISOMERIZATION 139

Table 9-1 Relative Rates of Metathesis and Isomerization on Reduced M 0 0 3 / y - A 1 ~ 0 ~ ~

Relative Initial Rates Catalyst Productive Degenerative

Reactant Activated byb Metathesis Metathesis Isomerization'

ethene none 0 1.2 12.7

CZH, c-2-butene

propene none 1.3 0 c3H6 9.4 53.5

c-2-butene 6.2 47.5

1 -bu tene none 0 0 64.6 1-butene 4.0 32Sd 24.7

c-2-butene 1.7 26.0d 27.4

c-2-butene none 0 29.8 c-2-butene 0 81.9" 1.9

Foomotes: a) From J. Engelhardt. et al., J. Catal., 70, 364 (1981). copyright Academic Press. b, Catalyst has been preoxidized and then reduced by the hydrocarbon indicated. ') By mechanism other than metathesis. d, Measured by the formation of 1-butene-dz and 4 from a mixture of 1-butene-4

e, Measured by the formation of c-2-butene-d4 and t-2-butene from a mixture of and de.

c-2-butene-6 and d8.

slower rate [ 1 I].

Nature of Catalyst Oxides of molybdenum, tungstcn, and rhenium are among the most active

oxide catalysts for metathesis. However, the oxide must be reduced first before it is active as is illustrated in Table 9-1. A fully oxidized catalyst has very low activity, but the activity increascs with the time of contact with the alkene. In the case of molybdenum oxide, reduction of the catalyst by alkenes has been directly confirmed by the color change of the catalyst and the detection of the EPR signal of Mo5+ [ l 11. Indeed, molybdenum oxide on a variety of supports but contains only Mo6 is inactive for metalhesis 1141. Similar observations have been made on + . .

140

W03/Si02 [3]. Exposure of the catalyst to oxygen which reoxidizes the catalyst also poisons it [23.

The induction period of the fully oxidized catalyst can be eliminated by prereduction with alkenes. It has been found that propene and butenes are much more effective activation agents than ethene [11,15]. This is because ethene does not reduce the oxide as effectively as the other alkenes. For example, a Mo0,/A1203 catalyst possesses a Mo5+ spin concentration of 0.09 x 1014 after reduction with ethene, and 1.0 x 1019, 2.0 x 1019, and 2.4 x 1019, respectively, after reduction with propene, 1-butene, and c-2-butene at 60°C [ l l ] . Reduction by other agents such as HZ, CO or NH3 is also possible but not as effective as propene or butene [15]. Perhaps this is because of the fact that the active catalyst is a surface metallocarbene which can be formed during the reduction by alkenes. It is known that carbonaceous residue is retained by the catalyst after reduction [9]. However, only a small fraction of the adsorbed CH3 and CH2 groups after reduction are associated with the active sites. It has also been reported that treatment with halogenated alkenes also greatly enhances the metathesis activity [161.

Although reduction of the catalyst is important for the activity, the activity does not necessarily correlate with the extent of reduction. It has been found that on a Md3/A1203 catalyst, the metathesis activity is independent of the extent of reduction beyond an equivalent of 0.3 e/Mo [17]. Furthermore, poisoning of the catalyst by pyridine and water shows that the amount of active sites is an order of magnitude less than the number of anion vacancies generated by reduction or the number of hydrogen atoms retained in the reduction step.

The activity also depends on the type of support. For molybdena, using A1203 or Ti02 as support generates active catalysts, using SO2 , Zr02, Cr2O3, or ZnO generates less active catalysts, and using Co304, NiO, MgO, Ge02, or SnO2 generates inactive catalysts [141. In the case of Co304, NiO, and SnOz, prereduction reduces the support but not Moo3. In the case of MgO, a magnesium molybdate compound is probably formed in which the Mo6+ is more difficult to reduce. Attempts to correlate the activity with the acidity of the catalyst have not been successful [ 101.

9.2 ISOMERIZATION OF ALKENES

Introduction Isomerization is perhaps one of the least demanding reactions in terms of the

requirements for the active site. Thus many oxides and metals are active isomerization catalysts. Indeed, practically all transition metal oxides possess isomerization activity. However, the activity and the selectivity vary among different oxides. In general, there are two types of isomerization reactions involving alkenes: double bond migration and cis-trans isomerization.

There are four well recognized mechanisms for isomerization: by metathesis, via carbanion intermediate, via carbenium ion intermediate, and via a radical-like (i.e. nonionic) alkyl intermediate. In the case of double bond


migration, sometimes the product can be either a cis or trans-alkene. The cis/trans ratio differs substantially between the carbanion and carbenium ion mechanism, being ten or higher for the former, and close to unity for the latter. Double bond migration cannot be achieved by metathesis. Except for the charge on the alkyl intermediate, the mechanism via a free radical-like alkyl intermediate is similar to that for the carbenium ion intermediate. Therefore, this mechanism will not be discussed further.

The cis/trans ratio is often used as an indication of the reaction mechanism. It should be noted that the pretreatment condition may alter this ratio for some oxides. For example, the cis/uans-2-butene ratio from 1-butene on a Cr2O3 catalyst pretreated by evacuation ranges from 2 for evacuation at 300-45OoC, which suggests the involvement of carbenium ion intermediates formed with surface hydroxyl groups, to as high as 50 for evacuation at 750°C. which suggests involvement of carbanion intermediates when the surface is dehydroxylated [18,26]. The cis-2-hexene/trans-2-hexene ratio from 1-hexene is 0.5 if the catalyst is activated with hydrogen, which may leave adsorbed hydrogen on the surface to form BrBnsted acid sites for the cationic reaction mechanism [27].

The effect of prereduction of a M a 3 catalyst is different from a Cr2O3 catalyst. On a reduced Md3/AI2O3 catalyst, the metathesis activity is high and isomerization by metathesis is the major mechanism for cis-trans isomerization [11,19]. On a fully oxidized Mo03/A1203 catalyst, isomerization proceeds mainly by a cationic intermediate. In this latter case, the contribution of A1203 to the isomerization activity has not been identified.

Isomerization by Metathesis In section 9-1, it is pointed out that degenerative metathesis may lead to cis-

trans isomerization, and that the CR2 fragment is retained in the process. The latter is a characteristic of isomerization by metathesis which is illustrated by the data in Table 9-2. Here, a mixture of cis-2-C4Ds and cis-2-C4H8 is passed over a reduced molybdena catalyst. The deuterium content of the product mixture is analyzed. The product trans-2-butene contains essentially only &, 4 and ds species. This can be readily understood if the CR2 fragment is retained during isomerization as in metathesis. In contrast, the formation of 1-butene which involves double bond migration must involve a different mechanism. The deuterium distribution in this molecule is much broader [19]. The much broader deuterium distribution in all products from an oxidized molybdena catalyst also points to a different reaction mechanism. It will be shown later that these reactions proceed via carbenium ion intermediates.

The activity for isomerization by metathesis increases drastically when the catalyst is reduced, as is discussed in the last section. However, reduction also suppresses isomerization by other rncchanisms. This is shown in Table 9-1. For example, a fully oxidized MoO3/y-AI2O3 catalyst is very active in isomerizing cis- 2-butene to I-butene and trans-2-butene. Reduction of the catalyst greatly suppresses the formation of I-butene, and the formation of trans-2-butene is now primarily by metathesis. The cis to trans-2-butene isomerization is suppressed by the presence of propene or l-butene, and less effectively by ethene [ll]. The

142

Table 9-2 Deuterium Distribution in the Isomerization of a Mixture of cis-2-C4H8 and trans-t-CdH8'

Deuterium Distribution %

cis-2-butene 47.1 0 0 0 0 0 1.1 5.9 45.9 (reactant)

oxidized cis-2-butene 55.8 9.1 1.2 0.4 0.8 0.3 0.9 5.7 25.9 Moo3/ trans-2-butene 28.3 18.6 3.2 1.1 2.2 0.6 3.4 18.6 24.2

y-A1203 l-butene 30.2 23.6 4.3 1.0 0.7 1.0 4.7 16.6 17.9

reduced cis-2-butene 28.3 1.9 0.6 2.0 28.3 1.5 0.7 3.7 33.1 Moo3/ trans-2-butene 26.6 1.7 0.7 3.1 38.9 2.3 0.5 2.9 23.3

y-Al2O3 l-butene 59.5 7.1 2.4 2.4 9.5 2.4 2.4 4.8 11.9

Foomote: a) From J. Goldwasser, et al., J. Catal., 70. 275 (1981), copyright Academic Press.

reaction is poisoned by H20 and pyridine, and partially by NH3 [17].

Isomerization via Carbenium Ion Intermediate In this mechanism, a proton is added to one carbon of the double bond of an

alkene as in abstractive adsorption to form an adsorbed carbenium ion. Loss of a different proton from the same carbon may result in cis-trans isomerization, and loss of a proton from adjacent carbon may result in double bond migration. Reaction (9-8) illustrates this mechanism:

+ - R1 CH=CH-CH2R2

I * +H

Since a proton needs to be supplied by the catalyst, the catalyst must possess BrQnsted acid sites (see Chapter 5 for a discussion of acid sites). Practically all acidic catalysts that possess BrQnsted acid sites can catalyze isomerization by this mechanism.


According to the mechanism, if the reactant is a perdeuterated alkene, one deuterium atom in the molecule is equilibrated with the surface proton pool each time it undergoes isomerization. Multiple exchange is also possible if the molecule undergoes more than one reaction cycle before desorption. This would result in a broad isotopic distribution in the product. This is illustrated in Table 9-2 by the products from a fully oxidized MoO3/AI2O3 catalyst and by 2-butene from the reduced molybdena catalyst. The formation of a product alkene-d, requires at least one reaction cycle, alkene-d2 at least two reaction cycles, and so on. Thus the larger is the difference between the deuterium content of the reactant and the product, the larger is the number of reaction cycles the molecule has undergone. In support of the fact that these product distributions are results of a cationic intermediate mechanism, 1 -butene yields only cis- and trans-2-butene as products on the fully oxidized Mo03/A1203 with a cis/trans ratio of 1.9 which is characteristic of such a mechanism.

The isomerization of 1-butene to 2-butenes on TiOz-supported W03 has been shown to proceed via a carbenium ion intermediate. The reaction rate parallels the density of acid sites as determined by titration with indicators. The cis/trans ratio of the product 2-butene is 1.7 [6], which is a typical value for this mechanism. The isomerization of cyclopropane to propene also shows a qualitative correlation with the acidity among various supported W03. Thus W03/Si02 and W03/Mg0 which have low acidity are inactive.

Since an acid site is required, its removal should render the catalyst inactive. Indecd, molybdena catalysts are poisoned by H20, NH3 and pyridine which interact strongly with the acid sites, while NO, CO, or CO2 are not poisons [17]. Reduction of the molybdena catalyst reduces the activity for isomerization by this mechanism, in agreement with the decrease in the amount of pyridine adsorbed (that is, acidity) [ l l ] . On the other hand, preadsorbed hydrogen on a reduced catalyst enhances the activity two to five times. For comparison, the metathesis activity is not affected by the presence of preadsorbed hydrogen [20], but is poisoned by CO [21.

Preadsorbed hydrogen also greatly enhances the isomerization activity of Co304. If preadsorbcd deuterium is used, it is incorporated into the isomer product to form a dl species. However, its incorporation is inefficient, and the preadsorbed deuterium is preferentially released back to the surface after isomerization. Exactly how this is accomplished is not clear. The enhanced activity cannot be maintained as the preadsorbed hydrogen is consumed to hydrogenate the alkene [21,22].

Isomerization via Carbanion Intermediate Double bond migration can also be achieved by first abstraction of an allylic

hydrogen from the alkene molecule to form a x-ally1 species. Addition of hydrogen back to the molecule but at the opposite end of the allylic bond results in isomerization. This process is shown in reaction (9-9) in which the lost hydrogen is shown as a proton:

144

The heterolytic dissociative adsorption of alkenes has been described in Chapter 4. In this step, the catalyst provides a basic coordinatively unsaturated surface lattice oxygen ion to accept the proton from the molecule. The x-ally1 anion is bonded to a surface coordinatively unsaturated cation. This mechanism has been shown to be operative on ZnO and Cr2O3 pretreated by evacuation where the cis-2- butene/trans-2-butene ratio from 1-butene is 10, and 2 to 50, respectively [18].

The interaction of alkenes with ZnO has been studied in detail (see Chapter 4). Infrared spectroscopy shows that propene is adsorbed both as a x-complex with a C==C stretching band at 1620 cm-' and as a x-ally1 species with a C 4 4 antisymmetric stretching band at 1545 cm-' [23]. The assignment of the bands has been confirmed using various isotopically labeled propene molecules. The formation of x-ally1 is accompanied by the formation of surface OH groups. These adsorbed species can be removed readily by evacuation. The idea of x- ally1 intermediates based on propene adsorption data has been used to interpret data in the reactions of butenes, although some of the adsorbed butenes cannot be removed by evacuation, and behave kinetically differently from the isomerization reaction [ 181.

It is interesting to note that reaction (9-9) may proceed either by intramolecular or intermolecular H-transfer. The relative contributions of these two pathways have been evaluated on Cr2O3 [24]. Using CD2=CH-CH3 and CD2=C(CH3)2 as reactants, intramolecular H-transfer would yield CD2H<H=CH2 and CD2H-C(CH3)=CH2, respectively, as the only initial products, whereas intermolecular H-transfer would yield propene-dl and d3 and isobutene-dl and d3 products. It is found that the ratio of intrdintermolecular H- transfer is 1.4 for amorphous Cr2O3 and 3.9 for crystalline Cr2O3 for propene. Both inter- and intramolecular H-transfer take place on ZnO also [25]. It has been further suggested that the intermolecular pathway proceeds via the x-complex and not the x-ally1 intermediate.

REFERENCES

1. R. L. Banks, and G. C. Bailey, Ind. Eng. Chem. Prod. Res. Develop., 3, 170 (1964). 2. G. C. Bailey, Cafal. Rev., 3, 37 (1969). 3. J. C. Mol, and J. A. Moulijn. Adv. Catal.. 24. 131 (1975). 4. J. C. Mol, J. A. Molijn. and C. Boelhouwer, Chem. Commun., 663 (1968). 5 . J. C. Mol, F. R. Visser, and C. Boelhouwer, J . Cafal.. 17. 114 (1970).


6. K. Tanaka, K.-i Tanaka, and K. Miyahara, J . Chem. Soc. Chem. Commun.,

7. D. J. Cardin. M. J. Doyle. and M. F. Lappea J. Chem. Soc. Chem. Commun.,

8. R. H. Grubbs, and C. R. Hoppin, J . Amer. Chem. Soc., 101, 1499 (1979). 9. M. Lo Jacono, and W. K. Hall, J . Colloid Interfnce Sci., 58, 76 (1977).

10. T. Yamaguchi, Y. Tanaka. and K. Tanabe. J. Catal., 65, 442 (1980). 11. J. Engelhardt. J. Goldwasser, and W. K. Hall, J. Catal., 70, 364 (1981). 12. J. McGinnis, T. J. Katz. and S . Hurwitz, J . Amer. Chem. Soc., 98, 605 (1976). 13. C. P. Casey, and H. E. Tuinstra, J . Amer. Chem. Soc.. 100, 2270 (1978). 14. K. Tanaka, K. Miyahara. and K.-i Tanaka, Proc. 7th Intern. Cong. Catal.,

15. A. A. Olsthoorn, and C. Boelhouwer. J. Catal., 44, 207 (1976). 16. E. S. Davis, D. A. Whan, and C. Kambell, J . Chem. Soc. Chem. Commun.,

17. E. A. Lombardo, M. Lo Jacono, and W. K. Hall, J . Catal.. 64. 150 (1980). 18. C. C. Chang, W. C. Connor. and R. 1. Kokes, J . Phys. Chem.. 77. 1957 (1973). 19. J. Goldwasser. J. Engelhardt. and W. K. Hall, J . Catal., 70, 275 (1981). 20. J. Engelhardt, and D. Kallo, J . Caul., 71, 209 (1981). 21. T. Fukushima, and A. Ozaki, J. Catal., 32, 376 (1974). 22. T. Fukushima, and A. Ozaki, J. Catal., 41, 82 (1976). 23. A. L. Dent, and R. J. Kokes, J . Amer. Chem. Soc., 92, 6709 (1970). 24. G. L. Haller, and C. S . John, Proc. 7th Intern. Cong. Catal.. Tokyo, Japan,

25. A. L. Dent, and R. J. Kokes, J . Amer. Chem. Soc.. 92, 6718 (1970). 26. N. E. Cross, and H. F. Leach, J. Catal., 21, 239 (1971). 27. R. L. Burwell, Jr., G . L. Haller, K. C. Taylor, and J. F. Read, Adv. Catal.,

314 (1979).

927 (1972).

Tokyo, Japan, 1980, paper B47.

1202 (1971).

1980, paper B20.

29, 1 (1969).

Chapter 10

DECOMPOSITION, HYDROGENATION AND

RELATED REACTIONS

10.1 DECOMPOSITION OF ALCOHOLS

There are three common reactions of alcohols on transition metal oxides: catalytic dehydration to the corresponding alkenes, catalytic dehydrogenation to ketones or aldehydes, and oxidation to carboxylates or carbon oxides with reduction of the oxide. In some cases, formation of hydrocarbon by coupling and bimolecular dehydrogenation is the major catalytic reaction.

Dehydrogenation may be carried out both oxidatively in the presence of oxygen or other oxidants, and nonoxidatively. In the nonoxidative route, hydrogen is the other product. In the oxidative route, water is the byproduct. Examples of the oxidative route include the highly selective oxidative dehydrogenation of methanol to formaldehyde (see chapter 12) and ethanol to acetaldehyde [l] on molybdate catalysts. Ordinarily, an oxidant is needed to prevent reduction of the oxides. In this chapter, the nonoxidative route is discussed.

Methanol

and dissociatively to form a surface methoxide: It was discussed in Chapter 4 that methanol is initially adsorbed heterolytically

CH3O H I I

CH30H + M a + M -- 0 (10-1)

Adsorbed methoxide has been shown to result from adsorption of methanol on ZnO at room temperature by use of infrared spectroscopy [2,31 and XPS (41. Molecularly adsorbed methanol may also be present. On heating, the undisso- ciated methanol and possibly some methoxide are desorbed as unreacted methanol,

146

DECOMPOSITION AND HYDROGENATION 147

the latter via the reverse of reaction (10-1). The remaining methoxide reacts via two pathways:

In the lower temperature catalytic pathway (10-2), adsorbed methoxide reacts to form formaldehyde, CO and H2. In the higher temperature pathway (10-3), it is oxidized by the zinc cation to an adsorbed formate which is decomposed to CO, C02, H2 and H20 [4-61. In this pathway, the ZnO surface is reduced which results in the evolution of Zn atoms into the gas phase at a temperature lower than the ordinary sublimation temperature of ZnO [7]. These two pathways are readily observed in a temperature programmed decomposition experiment (Figure 10- 1).

These two pathways are also observed under (quasi-)steady state conditions. Methanol is decomposed on ZnO into H2 and HCHO in the temeprature range 453- 513 K, and into HCHO, CO, C02, and H2 over 563-613 K [8]. Reduction of ZnO during the reaction is detected, and the formation of a "metallic silver mirror'' on the reactor used for methanol decomposition has been mentioned.

In this reaction, formaldehyde is formed by C-H bond dissociation in the adsorbed methoxide. Indeed, when CH3180H is used, only HCH180 and C l 8 0 are detected in the lower temperature pathway [7]. On the other hand, the formation of an adsorbed formate in the higher temperature pathway results in scrambling between the oxygen atom in methanol with those in the surface, and a mixture of C1602, Cl60l8O, Cl60 and C l 8 0 is observed [7,8]. Since C02 is not formed by the reaction of CO with the surface, it must be a primary product in methanol decomposition [8]. The formate is formed above 400 K and can be detected by IR [2] and X P S [4], but it decomposes at a much higher temperature [4,6]. C-H bond breaking is involved in the overall kinetics of methanol decomposition. It has been observed that there is no H/D kinetic isotope effect for CH30H and CH30D, but there is a normal kinetic isotope effect for CH30D and CD30D [8]. Finally, it has been observed in temperature programmed decomposition studies that this surface formate, whcther it is formed by the oxidation of methanol or formaldehyde, or by the dissociative adsorption of formic acid, decomposes at the same temperature [4,6].

The fraction of adsorbed methanol on ZnO that is adsorbed weakly depends on the extent of surface dehydroxylation [9]: it increases with increasing hydroxyl contcnt of the surface. The density of strongly adsorbed methanol increases proportionally with decreasing hydroxyl content, and reaches a limiting value of four molecules per nm2 on a totally dehydroxylated surface. The fraction of strongly adsorbed spccics that reacts by dehydrogenation versus that which reacts by the oxidation pathway dcpends on the condition of the surface. A more

148

a Z

8

0 100 200 300 400 500

TEMP C

Figure 10-1 TPD spectrum of CH30H decomposition on a Zn-polar surface. (From J. Phys. Chem., 90, 3184 (1986), copyright American Chemical Society).

reduced surface favors dehydrogenation [6]. The activity of the ZnO surface towards decomposition of methanol under

ultra-high vacuum conditions depends on the surface crystallographic orientation Temperature programmed decomposition studies show that the Zn-polar surface is the most active. The prismatic surfaces containing step defects or anion vacancies are also active, but the 0-polar surface is very inactive [4,6,10,11]. It has now bcen demonstrated that the catalytic activities of the Zn-polar surface for methanol decomposition at low (mtorr) and near atmospheric pressure are also much higher than the other ZnO surfaces [98,99].

The methanol decomposition pathways are somewhat different on other oxides. On NiO, CO, H2 and H20 are the only products detected [12]. On a highly dehydroxylated anatase TiOz surface, adsorbed methanol decomposes in TPD studies to form dimethyl ether, ethane and methane as the major products at high methoxide coverages, and mostly dimethyl ether and ethane at low coverages [13]. Dimethyl ether is formed by bimolecular dehydration, which is proposed to occur between methoxy groups on adjacent Ti cations:

YcH3 YCH3 + p, + H3COCH3 T i T i Ti Ti

(1 0-4)

Ethane is proposed to be formed from two Ti-CH3 species the presence of which is detected by IR. Methane is formed from isolated Ti-CH3 species. If a Ti02


surface is first covered with methoxy species, exposure to ethene yields propene, 1- butene and t-2-butene, while exposure to ethanol yields propene, c-2-butene. methyl ethyl ether, and C5 oligomer. The appearance of these products is consistent with the presence of TiXH3 groups on the surface. In this aspect, the surface of Ti02 is similar to a Ziegler-Natta oligomerization catalyst. It is possible that the dissociative adsorption of methanol to form T iXH3 and OH(ad) occurs on an oxygen vacancy and a surrounding reduced Ti ion. It should be mentioned that methane has been reported as a minor product of reaction of methanol on ZnO [6] and Cr2O3 [14], suggesting the possible presence of surface methyl species.

Oxidation of surface methoxide to formate is also found on chromium oxide [14]. Decomposition of formate then yields CO, COz, and HZ. Methane and dimethyl ether are the other products.

Catalytic decomposition of methanol on CuO and Cu20 yields C02, H20 and H2 as initial products and methyl formate at a later stage. Methyl formate, CO, and H2 are also products on CuCr204, but in this case, reduction of the oxide to metallic copper dispersed on Cr2O3 occurs [15].

2 -Propano1 The decomposition of 2-propanol is another well-studied reaction. The two

most common reaction pathways are dehydrogenation (10-5) and dehydration (10- 6):

CH3CHOHCH3 CH3COCH3 + H2 (10-5)

(10-6) + H+ CH3CHOHCH3 ------) [CH3,HCH3]+

OH2

Sincc dehydrogenation is the preferred pathway on basic oxides, while dehydration is preferred on acidic (BrQnsted) oxides [16], this reaction has been used as a probe for the acidicbasic properties of an oxide [17]. Indeed, the predominant reaction is dehydrogenation to acetone on oxides such as CaO and MgO, and dehydration to propene on oxides such as A1203 and V205. However, the situation is less definite on transition metal oxides. It has been shown that dehydrogenation is predominant on Cr2O3, Mn02, FezO3, C0304. NiO, CuO. and ZnO some of which are considered acidic (see Chapter 5). Furthermore, anatase and rutile Ti02 give different selectivities: dehydrogenation is predominant on rutile, and dehydration on anatase [611.

The detailed mechanism of the reaction on ZnO has been studied. Infrared spectroscopy shows that 2-propanol is adsorbed heterolytically as 2-propoxide with the simultaneous formation of surface hydroxyl groups [18]. At 363 K,

transformation of the surface alkoxide to an adsorbed acetone occurs. Further- more, a surface enolate species is detected suggesting a rapid interconversion between acetone and enolate:

CH 3

(10-7)

The dehydrogenation to acetone is the predominant reaction under catalytic conditions, accounting for 90% of the products.

The interconversion of acetone and enolate on ZnO has been independently confirmed with adsorbed acetone [l8,19]. It has also been observed on NiO [20,21]. This process has also been confirmed with deuterium-labeling studies. When a mixture of 2-propanol-& and dg is decomposed on ZnO, H/D mixing does not occur in 2-propanol, but it occurs in the product acetone [19].

At 160°C, the normal H/D isotope effect at the secondary C-H bond suggests that the breaking of this secondary C-H bond is rate limiting in 2-propanol decomposition. In contrast, deuteration of the methyl groups does not affect the reaction rate [19]. Use of a mixture of C3H70D and C3D70D with monitoring the evolution of D2, HD, and H2 has shown that the product H2 is formed exclusively from the a-C-H and &H of 2-propanol [22].

The dehydrogenation of 2-propanol to acetone on ZnO is a totally reversible reaction, and acetone is readily hydrogenated to 2-propanol.

At low pressures, dehydrogenation is first order in 2-propanol, and becomes zeroth order at higher pressures (2 P a and above) [22,23]. At low pressures, the reaction rate depends on the orientation of the exposed surface plane. At about 2OO0C, the rate of reaction is much faster on the Zn-polar surface than the 0-polar surface (Fig. 10-2) [23,24]. Interestingly, the activation energy is higher on the Zn-polar than the 0-polar surface. The reaction has also been studied on a stepped prismatic surface. The rate is higher if the alcohol molecules impinge on the surface in the direction up the step than down the step. At a higher pressure of alcohol and using powder samples. a different conclusion has been made that the reaction rate is proportional to the number of surface coordinatively unsaturated Zn ions independent of the crystallographic plane [25]. At the moment, the reasons for the differences between the low-pressure single-crystal study and the higher- pressure powder study are not understood. It is possible that a reaction that is structure sensitive at low pressure, i.e., low surface coverage, may become structure insensitive at high coverage, i.e., high surface coverage [26].

The high selectivity for dehydrogenation observed under catalytic conditions on ZnO is not observed in temperature programmed decomposition studies [27-291 where a selectivity as low as 60% dehydrogenation has been observed. It is further found that the ratio of dehydrogenation to dehydration depends on the pretreatment of the ZnO sample.

A Zr(OH), The selectivity for this reaction depends on the oxide used.


M I 0 7

v; 3 . 5 5 I

W +J

v1 \ W

.-

2 .84 U W

0 E

d

c 2 . 1 3 0

c, u u 0

.4

a

1 . 4 2 W c 0 c, W U

0.71 L 0

Q-

L L

2

I

/ d 100 1 80 260 340

Temperature C

Figure 10-2 Rate of reaction measured as turnover frequency for acetone production as a function of temperature for the Zn-polar ( 1 ) and the 0-polar surfaces (2) of ZnO.

sample treated with H2Te04 is 100% selective for acetone production [30]. On the other hand, TiOz [31,341 and Se-treated Ti02 [30] samples are quite selective for dchydration. The activation energy of the reaction depends on the promoters. For example, on an alkali-promoted VzOs catalyst, the activation energy varies from 1.6 to 38 kcal/molc depending on the alkali ion used [32]. The interaction of the hydroxyl group of the alcohol with the alkali ion was used to explain this obscrvation. In other instances, the band theory of oxides [I61 or the geometric rcquircmcnt of surface atoms [331 have been invoked to explain similar obscrvations.

Rcccntly, it has been reported that on the Cu-CuO system, dchydrogcnation of 2-propanol docs not take place on Cu metal, Cu20, or CuO. However, the activity on a prcrcduced Cu20 or prcoxidizcd Cu is very high. It is also high for a physical mixture of Cu and CuO. The results suggest that there is a synergistic cffcct bctween oxidized and metallic Cu that lcads to high activity [42], but Lhc exact cause has yet to be found.

152

Other Alcohols The decomposition of other alcohols has been less extensively studied. The

patterns of their reactions are similar to 2-propanol. On anatase Ti02, dehydration of ethanol to ethene and 2-methyl-2-propanol (tert-butyl alcohol) to 2- methylpropene (isobutene) are the major reactions r31.351. In the reaction of ethanol, small amounts of ether and butenes are detected, which indicate bimolecular dehydration and oligomerization of ethene. The latter is due to the formation of Ti-C2H5 species [13]. Infrared spectroscopy detects the presence of surface alkoxides.

Dehydration of ethanol to ethene is also the major reaction on Cr2O3 [36] and Mo03/Si02 [37]. Small amounts of C€&, C2H6 and C3€& are also produced on Mo03/Si02. This may be due to the formation of Mwalkyl or Mmarbene species after some of the surface Moo3 is reduced by the reaction with the alcohol. On niobia, the products from ethanol decomposition at 400-550 K range from 50% diethyl ether and 50% ethene on Nb.L05/Si02 prepared by impregnation, to 50% ether and acetaldehyde and 50% ethene on unsupported Nb2O5, to over 99% ethene on a thin layer of niobia on Si02 [43].

The decomposition of various alcohols has also been studied on ZnO using temperature programmed decomposition technique. As mentioned earlier, for reasons not yet fully understood, the dehydration pathway is more prominent using this technique than under catalytic conditions. It has been observed that 90% of ethanol is decomposed to ethene, about 3040% of 1-propanol, 2-propanol, and 1- butanol to the corresponding alkenes, and 17% of 2-bum01 to butenes [38]. The remaining products are the corresponding aldehydes and ketones. A high selectivity for dehydration of ethanol has also been observed on ZnO single crystal surfaces [39]. In contrast, catalytic decomposition of ethanol on ZnO powder is predominantly dehydrogenation to acetaldehyde [401.

In the temperature programmed decomposition studies, the dehydrogenation products (aldehydes and ketones) are evolved about 10-15°C lower than the dehydration products (alkenes). This indicates higher activity for dehydrogenation on ZnO, which may account for the much higher selectivity for dehydrogenation under catalytic conditions. It is also observed that branched alcohols decompose about 50°C lower than straight chain alcohols. I f this implies that the adsorbed species from a branched alcohol is less stable, it would indicate that the transition state in the dehydrogenation reaction of surface alkoxides on ZnO might have a carbenium ion character [39].

2-butanol decomposition has been studied on mixed iron oxides. The selectivity depends on the oxide components. Dehydrogenation to methyl ethyl ketone is the only reaction observed on F Q O ~ . It is also the predominant reaction on Fe203-Zn0. However, dehydration to butenes is the predominant reaction on Fe203-Ti02, while both reactions proceed about equally on F+03-Zr02. Reduction of these catalysts results in a much enhanced production of butane, presumably formed by nucleophilic substitution of OH- in butanol by H- that is lcft on the surface from the dchydrogenation of butanol [411.


10.2 DECOMPOSITION OF NITROUS OXIDE

The decomposition of nitrous oxide produces only nitrogen and oxygen:

2 N 2 0 --+ 2N2 + 0 2 (10-8)

Because of its apparent simplicity, this reaction has been used as a probe for the properties of solids. However, an unequivocal correlation between reaction kinetics and properties of the catalyst depends on a firm establishment of the reaction mechanism for which more than one has been proposed over the years.

One of the most popular mechanism is as follows [16,44]:

N20(g) + V,- N20-(ad) (10-9)

N20-(ad) - N2(g) + 0-(ad) (1 0- 10)

2 0-(ad) --+ 02(g) + 2 V,- (10-11)

In this mechanism, nitrous oxide is first adsorbed on a surface anion vacancy V,- with a trapped electron. This step involves electron transfer from the solid to the adsorbate. The adsorbed N20- decomposes quite readily to gaseous N2, leaving an adsorbed 0- on the surface. Recombination of 0- to form gaseous oxygen and two anion vacancies completes the cycle.

There is experimental evidence to support this mechanism. Stoichiomeuic decomposition of N20 to N2(g) and adsorbed oxygen can be accomplished at a temperature lower than the catalytic decom osition temperature [45,46]. The species 0- has been detected by EPR on Md or V205 supported on SO2 [47,48] and UV-irradiated Ti02 [49] and ZnO [50] upon decomposition of N20. This evidence further suggests that the recombination of adsorbed oxygen is the slow step in the reaction. Indeed, the rate expression for a large number of oxides, including NiO, ZnO, Fe2O3, Cr203, Ti02, CuO, Rh2O3, Mn02, Ir02 , and many nontransition metal oxides has the general form [44,51]:

P

(10-12)

which can be derived from the reaction sequence assuming that the step of N20-(ad) decomposition is irreversible, but the adsorption of N20 and desorption of oxygen are reversible. This expression implies that the reaction is inhibited by gaseous oxygen as is observed.

This reaction mechanism suggests electron transfer between the adsorbates and the solid. Electron transfer is involved in the formation of 0-(ad) which has been mentioned above. For semiconducting oxides, electron transfer affects the electrical conductivity. It has b u n observed that the electrical conductivity of ZnO changes when subjected to the reaction condition, which supports the presence

154

of electron transfer [62]. Changes in electrical conductivity during oxygen chemisorption are also well known. In these cases, the conductivities are strong functions of the oxygen partial pressures which may reflect changes in the concentration of anion vacancies.

The involvement of electron transfer in the reaction mechanism suggests possible correlations between the semiconducting properties of oxides and their activities in this reaction. Indeed, the most active oxides which include Cu20, COO, Mn2O3, and NiO are all p-type semiconductors, the least active oxides which include ZnO, CdO, Ti02, Fe2O3, and Cr2O3 are n-type semiconductors, and the oxides that are intermediate in activity which include CuO, MgO, CaO, and Ce02 are insulators [63]. Essentially the same classification is obtained if activation energy is used as the criterion instead of activity [HI: the p-type semiconducting oxides have the lowest activation energy, and the n-type semiconducting oxides have the highest activation energy. A linear correlation between the activation energy and the electrical conductivity has been reported [ 5 5 ] .

The higher activity of p-type semiconducting oxides is explained by the fact that oxygen recombination and desorption is a slow step. Since electron is transferred into the solid in this step, its rate is enhanced by the availability of holes in the solid. Thus p-type semiconductors are more efficient than n-type semiconductors or insulators. Alternatively, the different positions of the Fermi levels in the two types of semiconductors may be the reason for the different rates of electron transfer in this step [65].

Since the oxygen desorption step is slow, there have been attempts to correlate the kinetic parameters of this reaction with those of isotope exchange between gaseous oxygen and the oxide surface. There appears to be a strong correlation in the preexponential factor and the activation energy for these two processes (Fig. 10-3 and 10-4) [441.

The apparent success in correlating semiconducting and catalytic properties of oxides has led to attempts to study the effect of dopants on the reaction. Studies using doped NiO [MI, doped ZnO [51], and doped Fe203 [52]. have been conducted, yet they were not always successful. Some complications encountcred in these studies include the fact that the activation energy is found to depend on thc outgassing temperature of the oxide, the oxide is oxidized during reaction [531, and a new compound is formed between the dopant and the oxide [52,541. It has also been reported that the activation energy depends on the crystallite size of the oxide. On iron oxide, the value for very small crystallites is lower than for large crystallites [56].

Recently an alternate mechanism involving reduction and oxidation of the oxidc has been proposed on iron-exchanged zeolites [57]:

N20 + O’-(S) N2 + 0 2 + VS-- + C(S) (10- 13)

NzO + V,- + e-(s) --+ N2 + O’-(s) (10- 14)

0 2 + 2 v,- + 2 c-(s) --+ 2 02-(s) (10-15)


68

84

68 -

84 -

0 1 1

84 126 168 kJ/mole

1 1

84 126 168 kJ/mole

Activation Energy for N 0 Decomp. 2 Figure 10-3 Correlation between the activation energies for NzO decomposition and those for the oxygen-exchange reaction on various oxides. (From J. Catal.. 19. 32 (1970), copyright Academic Press).

In this mechanism, the evolution of oxygen is a result of reaction (10-13) in which N 2 0 reacts with a surface lattice oxide ion to yield N2, 0 2 and an anion vacancy. The anion vacancy is filled and the oxide is reoxidized either by N20 (eq. 10-14) or by gaseous oxygen if present. This mechanism is supported by the following observation. When N2I60 is decomposed over a Fe-exchanged Y-zeolite loaded with l 8 0 , the gaseious O2 is composed of 80% 1602 and 20% 180160 [%I. The substantial amount of l8O in the gaseous oxygen product suggests the involvement of lattice oxygen and the redox mechanism. Interestingly, when the same reaction is run over Fe-exchanged zeolite M (mordenite), no l8O is detected in the product. Thus either the number of active sites is very small on this catalyst or the redox mechanism is not operative.

10.3 AND H2-Dz SCRAMBLING

HYDROGENATION, H-D EXCHANGE OF HYDROCARBONS,

Transition metal oxides are generally not very active catalysts for In fact, except for reduced oxides of chromium, vanadium, hydrogenation.

156

MgO Gab / 0 Ca

-;/;s;Hf ' OZ"

Th 0

Sr 1 I

2.0 5.0

' O % o A1

Figure 10-4 Correlation between the preexponential factor of the rate constants for the N 2 0 decomposition Al and the oxygen-exchange reaction AB. (From the same reference as Figure 10-3).

copper, and molybdenum, cobalt oxide, zinc oxide and a few others, other oxides can be considered inactive for hydrogenation when compared with the noble metals. The relatively low activity even among the more active oxides has been advantageously used for selective hydrogenation. For example, copper chromite is used in the selective hydrogenation of nitrobenzene to aniline in which the functional group but not the benzene ring is hydrogenated [59].

Attempts have been made to explain the activity pattern of transition metal oxides. When examining the activities of oxides of the first transidon period for ethene hydrogenation. it is found that the highly active Cr203 and Co304 are separated by the inactive MnO and the nearly inactive Fez03 [60]. (For example, what is in effect a layer of MnO on Si02 will remove O2 from a mixture of H2 and propene without forming any ropane.) This pattern is shown in Fig. 10-5. It is explained by the fact that Mn2' and Fe3+ are in the very stable 3d5 configuration. It is difficult to excite a 3d5 ion into a state that would bind an adsorbate strongly. It should be noted that an active oxide must have surface coordinatively unsaturated ions to adsorb an alkene molecule, dissociatively adsorb hydrogen molecule, and bind the surface alkyl relatively strongly. An inactive oxide may fail because it does not meet one or more of these requirements.

The dependence of hydrogenation activity on the position of the oxide in the

DECOMPOSITION AND HYDROGENATION

60

30

157

-

-

n 0 N

NiO T i02 Cr203 Fe203

"2'3 MnO Co304 ZnO

Figure 10-5 Activity pattern in the hydrogenation of ethene. Figures in parentheses denote reaction temperatures. (From J. Catal., 7, 359 (1967). copyright Academic Press).

Periodic Table as shown in Fig. 10-5 has also been observed in the H2-D2 scrambling [67], dehydrogenation of propane [68], disproportionation of cyclohexene [68], and isomerization of butene [69]. While the correlations and the explanation based on crystal field theory are appealing, the data are for only a limited number of oxides, and the effect of preparations and pretreatments of the oxides have not been taken into account [69]. An oxide may show very different catalytic activity for different degrees of surface hydroxylation, extents of surface reduction, or calcination temperatures, among many variables.

Instead of comparing under one set of conditions, comparison can also be made of the maximum activity of each oxide [69]. It is found that (i) for a given element, the activity of a lower oxide is higher than a higher oxide; (ii) the activity of oxides at their highest common oxidation state increases from right to left within each period; (iii) all trends observed for hydrogenation apply to double-bond migration, exchange of hydrocarbons with deuterium, and H2-D2 scrambling. Unfortunately these conclusions are based on a limited number of oxides, and they should be tested with more data when available.

The mechanism of hydrogenation is believed to be stepwise addition of H atoms to an adsorbed alkene, first to form a surface alkyl species and then an

158

alkane. The sequence is similar to that found on metals. The reversibility of the individual step depends on the oxide, and can be probed by following the incorporation of deuterium atoms in the hydrocarbon molecule upon deuteration.

The most detailed investigation of the hydrogenation mechanism has been conducted on ZnO and CrzO3. For ethene hydrogenation on ZnO, the adsorption and the kinetic data, and the deuterium distribution in ethane lead to the mechanism [ 701 :

H2C=CH 2 I Y z H S I + * C2H4 + Zn-O __L Zn-0 - Zn-0 (10-16)

H I + Zn-O + C2&

+ * -

The adsorption of ethene as a x-complex is confirmed by IR spectroscopy. This step is reversible and the adsorbed ethene can be removed readily by evacuation. Hydrogen is adsorbed dissociatively and heterolytically to form Zn-H+ and &H-. Addition of H to an adsorbed ethene forms a surface a-alkyl. Addition of the second H completes the hydrogenation. These two steps are irreversible because if ethene is deuterated with D2, the product contains 99.9% ethane-d2 [71]. Reversibility of the addition of the first H would lead to deuterium incorporation in the reactant ethene, and reversibility of the addition of the second H would lead to a broad distribution of deuterium content in ethane, neither of which is observed. This mechanism applies to well dehydroxylated Cr203 [71,72] and Co304 1731.

Conclusions concerning the mechanism of ethene hydrogenation are substantiated by results of hydrogenation using H2 and D2 mixtures. On both ZnO and Cr203, the product ethane consists of mostly 4 and d2 species. Small amounts of ethane-dl can be attributed to the simultaneous H2-D2 scrambling reaction that produces HD in the gas phase [71]. On Co304, the production of HD is substantial. Thus the ethane product shows a broad distribuiton of 4, d, , and d2 species which parallels the relative concentrations of H2. HD, and D2 in the gas phase [73].

It is interesting to note that although C2H4 does not exchange with D2 on Co304, intermolecular H-D exchange between C2H4 and C2D4 occurs rapidly. In a mixture of C2H4, C2D4, and H2, ethylene is rapidly randomized to yield a broad distribution of 4. dl, d2, d3, and 4 species. Thus, this intermolecular exchange process is more rapid than hydrogenation [73].

This mechanism and the similarity between ZnO and Cr2O3 do not extend to propene. When propene is deuterated on ZnO, deuterium is incorporated into the reactant propene, and the product propane shows a wide distribution in the deuterium content as shown in Table 10-1 [74]. The data can be explained by the reversible formation of x-ally1 upon dissociative adsorption of propene (eq. 10- 17). If the surface x-ally1 forms propene by picking up a surface deuterium atom instead of a hydrogen atom, deuterium is incorporated into propene. The broad


Table 10-1 Deuterium Distribution in Propene Hydrogenation

ZnO, C3& + D2 at 25°C (ref. a)

propene 84.4 14.1 1.6 0 0 propane 12.2 20.9 55.6 11.3 0

Cr2O3, C3& + D2 at -11°C (ref. b)

propane <4% 94 1.5 0.1

References: a) A. L. Dent, and R. J. Kokes, J. h e r . Chem. SOC., 92, 6718 (1970). b) A. B. Littlewood, and R. L. Burwell. Jr., J. h e r . Chem. SOC.,

82, 6287 (1960).

distribution of deuterium in propane may be due to the rapid reversibility of this step and/or subsequent steps of the formation of surface alkyl.

CH2-CH-CH, H

A C3& + Z n O d - 1 - - - Zn

(1 0- 17)

In contrast to ZnO, deuteration of propene on CrzO3 activated at 450°C leads almost exclusively to propane-d2 [71,72]. Thus the reaction on Cr2O3 probably proceeds as in eq. (10-16) without the formation of n-allyl.

The situation for the hydrogenation of 1 -butene, 2-pentene, cyclopentene, and l-hexene on Cr203 is similar to that of propene. In all cases except pentene, alkane-d2 accounts for 90% or more of the hydrogenated products. For 2-pentene, the product is 78% pcntane-d2 [72]. The deuterium distribution becomes broader when the reaction temperature is raised from about 30" to 130°C. but the amount of alkane-d2 is still higher than that expected from a statistical distribution of deuterium atoms.

The nearly exclusive production of butane-d2 on Cr2O3 permits the determination of whether the two hydrogen atoms are added on the same or opposite sides of the carbon-carbon double bond. It is found that using D2 and cis- 2-butene, meso-2-3-dideuterobutane appears as nearly the only product as shown in eq. (10-18) [75]. Thus D2 is added cis to the double bond.

160

From IR spectroscopy, it has been shown that on ZnO. reversibly adsorbed hydrogen participates in hydrogenation [771. If HD is used as a reactant, it is adsorbed primarily in form I at room temperature and in form 11 below -40°C.

H D I I

Zn 0 D H I 1

Zn 0

I 11

When the selective hydrogenation of buta-1,3-diene to l-butene is carried out at these two temperatures, the same distribution of 75% [3-Dl]but-l-ene and 25% [4- Dl]but-l-ene is obtained in spite of the different forms of adsorbed HD [78]. It is not known whether this is a result of the different reactivity of Zn-H versus O-H and Zn-D versus O-D that happens to compensate each other, or the fact that the catalytically active sites do not distinguish the two forms, or that butadiene is adsorbed in different forms at the two temperatures.

The fact that the deuterium distribution in the alkane product follows closely the D2, HD, and H2 distribution in the gas phase for those oxides that have been carefully studied suggests that either the mobility of adsorbed H (and D) species is very low or the density of active sites is very low. Few quantitative determinations of the density of active sites have been reported. The density is known to depend on the degree of dehydroxylation. A chromia gel activated below 200°C is inactive. Activity is obtained at highcr temperature of activation, and the chromia is very active after 450-500°C activation in an inert atmosphere (see Fig. 4-3b) [72,76]. The deuterium distribution in alkane depends on the activation condition. The amount of alkane-d2 increases with increasing temperature of activation, while the spread of distribution decreases. For example, CrzO3 activated at 215°C yields only 26% hexane-d2 in the deuteration of 1 -hexene, while activation at 400°C yields 90% hexane-d2. The broad distribution resulting from low activation temperatures cannot be explained by rapid isomerization because the ratio of hydrogenation to isomerization does not change significantly with activation temperature. It might suggest that surface hydroxyl groups could exchange with adsorbed deuterium and surface alkyl species readily.

It is most likely that the active sites involve surface coordinatively unsaturated cations. These active sites can be poisoned by water [70,76,77], 02, CO [73,76], and, for Co304, by irreversibly adsorbed H2 [73]. Different alkenes compete for active sites. On C%04, the rate of propene hydrogenation is greatly reduced by the presence of ethene. On a molybdena catalyst, the active sites are the reduced Mo(cus) ions. Interestingly, these sites are poisoned by NO and CO,


but not NH3 or pyridine [91]. Deuterium exchange of alkenes via a x-ally1 intermediate (reaction 10-17) is

known to occur on ZnO. Table 10-1 shows the data that support this. Appar- ently, the reversible formation of x-ally1 is much more rapid relative to hydrogenation on ZnO than on CrzO3. Since negligible amounts of hexene-dl are formed when 1-hexene is passed over chromia with D2 at lower temperatures, although at 200°C small amounts can be detected [72]. The same applies to a mixture of D2 and butene [72]. Isotope exchange between D2 and propene has also been reported on rutile Ti02 [79].

Deuterium exchange of alkanes has been studied extensively on Cr203 [75,76,80]. Chromia gel develops significant activity after activation in N2 or H2 above 300°C to remove higher oxides. The activity is poisoned by exposure to air (presumably moisture) or 0 2 . The predominant reaction is the exchange of one hydrogen atom per period of adsorption. Thus the initial products are alkanes-dl . There is only very minor amounts of multiple exchange the extent of which increases with increasing temperature. The primary carbon in propane and hexane exchanges faster than the secondary carbon. Thus 2.2-dimethylpropane exchanges faster than propane. The faster exchange of primary than secondary carbon can be explained by the involvement of intermediates with some carbanion character.

The situation for cyclic alkanes is less clear. At 200°C on chromia, the rate of exchange per H atom in the molecule follows the order: cyclopropane >> ethylcyclobutane > cyclooctane 2 hexane > cycloheptane > 2.2-dimethylbutane > cyclohexane = cyclopentane = methane > ethane. The order is somewhat different at different temperatures because the values of activation energy are different.

Another reaction related to hydrogenation and deuterium exchange is H2-D2 scrambling:

H2 + D2 + 2 H D (10-19)

which may occur on a surface with the following mechanism:

H H D D M O M 0

H2 + D2 + 2M-O + I I I I + 2 H D + 2M-0 (10-20)

This mechanism is supported by the detection of Zn-H and &H IR vibrational bands when H2 is adsorbed on ZnO [771. On other oxides where such an observation has not been made, such as Cr203, there are suggestions that the reaction may proceed without total dissociation of H2 and D2 [81].

This reaction is known to occur on many oxides including nontransitional oxides, in particular MgO. Across the first transition period, the activity of the oxides shows the two peaks of high activity similar to that shown in Fig. 10-5. except that Co304 and NiO are more active than Cr203 [67]. Reduced vanadia and molybdena are also quite active. In contrast, a reduced anatase Ti02 is less active and the reaction shows a higher activation energy than a reoxidized or an

162

outgassed sample [821. The H2-D2 scrambling reaction is usually suppressed by the presence of

alkenes. For example, butenes suppress greatly this reaction on reduced M a 3 [83]. ethene suppresses this on Co304 [73], CrzO3 and to a lesser extent ZnO [71]. Thus it is likely that hydrogenation and this reaction take place on the same active sites.

10.4 REDUCTION OF NO

NO reduction is an important process because it is one of the reactions in environmental pollutant control. In particular, reduction of NO by NH3 or CO is very desirable as it removes two atmospheric pollutants simultaneously:

2 N 0 +2NH3 + 1/202+ 2N2 + 3H2O (10-21)

NO + CO --j 1/2N2 + CO2 (10-22)

Reduction can also be achieved with H2:

NO + H2 + 1/2N2 + H2O (10-23)

The formation of N2 and H20 are not the only possible reactions in a mixture of NO and NH3. There are at least three competing reactions:

The extent of these competing reactions depends on the reaction conditions such as temperature, composition of the reaction mixture, and the catalyst. For example, in the reaction between NO and NH3 on V205, a higher partial pressure of oxygen greatly enhances the production of N2 over that of N 2 0 so that N2 amounts to over 90% of the nitrogen-containing products [97] (see Fig. 10-6). On copper oxide, the product is 99% N2 when the oxide is fully oxidized, and becomes 75% N2 and 25% N20 when the oxide is partially reduced [92]. The product is 60% N2 and 40% N 2 0 over unsupported Cr2O3. but it is 75% N2 and 25% N2O on a Cr203/A1203 [%I. On Moo3, the ratio depends on the morphology of the oxide particle [90].

A number of oxides has been reported to be active for NO reduction with NH3. CuO [921. A1203-supported Fez03 [931, 0 2 0 3 [941, and VZOS [951, 7 3 0 2 - supported VzOs [84]. Moo3 [851. Fe203 [861, and Nb2Os [87] are reported to be active. Some, such as supported or unsupported V205, are active even for trace amounts of NO and NH3 and in the presence of air or oxygen. The activity of


1.0

0.5

0

O2 Concentration, 96

Figure 10-6 Effect of 0 2 concentration on the production rates of N2 and NZO at 250°C over VzOs. Inlet concentrations of NO and NH3 are lo00 ppm. W/F = 0.33 g-h/rnole. (From J. Catal., 62, 140 (1980), copyright Academic Press).

these oxides depends on the support. V205 supported on Ti02 is more active than on A 1 2 0 3 or Si02 [84]. Moo3 supported on Ti02 is the most active, followed by other supports like Zr02. A1203 and finally SO2 [85]. On Fer03, the activity sequence follows: Si02 > A 1 2 0 3 > Zr02 > Ti02-Zr02 > active carbon > MgO > Si02-Ti02 > Ti02 > Sn02. In this case, the activity difference among the first four supports parallels their different surface areas [86]. For Nb2O5. Ti02 support provides a catalyst that is much more active than A1203 and Zr02, which are more active than Si02 [87]. It has also been reported that sulfate enhances the catalytic reaction for Nb205/Ti02 and Fe2O3/ZrO2 [861.

The activity of the catalyst depends on the oxidation state. The active state of vanadium oxide is V2O5. This state is maintained when there is 1% or more of O2 in the feed under the condition shown in Fig. 10-6 [97]. In the absence of oxygen, the catalyst is reduced to V204. The active state of Cr203/A1203 is also nearly fully oxidized [96]. On the other hand, the activity of copper oxide increases when CuO is reduced [92].

The mechanism of the reaction has been elucidated mostly by use of isotope labeling techniques and infrared spectroscopy. On V205, adsorption of NO only occurs in the presence of O2 with the formation of an adsorbed *-N02- species characterized by an IR band at 1632 cm-'. Adsorption of N H 3 occurs on BrBnsted acid sites to form NI-&+. The adsorbed NO2- and w+ react on the surface to form N2 which is detected in the gas phase [95]. These observations lead to the proposed mechanism:

NO + O-(ad) + N02-(ad) (10-27)

164

C 0 CI .- CI C E ' 0, 0 C

I

(b)

0 20 40 60

Time, s

Figure 10-7 Concentration profiles at the inlet and outlet of the reactor. a. Inlet rectangular pulse of a mixture of NO and NH3. b. Concentration profiles of N2 produced by the reaction of rectangular pulse of the NO and N H 3 mixture with unsupported V205 at various temperatures. (From J. Phys. Chem.. 85, 2367 (1981). copyright American Chemical Society).

*-OH- + NH3 -+ *a2%+ (1 0-28)

*-027W&+ + N02-(ad) - N2 + 2 H20 + *a2- (10-29)

*a2- + H20 + 2 * 4 H - ( 10- 30)

While gas phase oxygen enhances the reaction rate, the reaction proceeds without oxygen. Under such a condition, the V=O groups in V205 are reduced [88]. When NO gas is introduced onto the catalyst with NH3 preadsorbed as W+ at 100°C or higher, w+ disappears and products N2 and H2O are formed [97]. The consumption of V=O in this reaction has been used to determine the surface density of V=O. In this technique, a rectangular pulse of NO and NH3 is passed over a catalyst and the production of N2 is monitored (see Fig. 10-7). The evolution of N2 shows a large peak at the beginning which is assumed to be due to


the reaction with surface V=O (reaction 10-31). This is followed by a long tail which represents regeneration of consumed surface V 4 by lattice oxygen. The reaction sequence is:

V 4 + N H 3 + NO + V-OH ,I- N2 + H2O (10-31)

2V-H + [O] + 2 V 4 + H20 (10-32)

The Occurrence of reaction (10-32) has been questioned recently because it was observed that the same profile of N2 evolution was obtained using a V205Di02 catalyst on which V2O5 presumably forms a monolayer on the support. Since there should not be lattice diffusion of oxygen to reoxidize surface V 4 H groups as in reaction (10-32), the long tail is proposed to be due to reaction on a reduced surface site [89]:

3 V 4 + 2NH3 --+ 3 V 0 + N2 + 3H2O (10-33)

VO + NO -----) V=O + 1DN2 (10-34)

Here, VO represents a vanadium ion associated with an oxygen vacancy. Analogous equations can be written if it represents a V-OH group.

When 15NH3 is used as a reactant with 14N0, isotopically labeled N2 and N20 are produced. On Cr203/A1203, the product consists of 15N14N, 14N14N, and 14N20 but no 15Ni4N0. The ratio 15N14NP4N20 ranges from 1.5 to 3. On CrzO3, 15N14N and 14N20 are the major products, no 14N2 and only a small amount of 15N14N0 are found. The ratio 15N14NP4N20 is slightly below 2 [97]. This ratio is also about 2 on FezO3/Al2O3. The mechanism proposed to account for this is:

2Cr=O + NO + 1 5 N H 3 -+ Cr+NO + Cr4Hl5NHz (10-35)

-+ 15N14N + H20 + C d + Cr-OH

H 2Cr-OH + 2 N 0 --j 2Cr-U-NO (10-36)

--+ 2 0 + N20 + H2O

A combination of reactions (10-35) and (10-36) gives a 15N14NP4N20 ratio of two. The production of 14N2 is assumed to be due to decomposition of l4NzO.

When the same reaction is studied on copper oxide, only 15N14N is produced on fully oxidized CuO. The catalyst is reduced at the same time. Other products begin to appear as the extent of reduction increases. On Cu20, 15N14N 9 14N 2,

15Ni4N0 and 14N20 are formed in the ratio 73:16:9:3 [92].

166

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Chapter 11

SELECTIVE OXIDATION REACTIONS I

11.1 INTRODUCTION

Selective oxidation of hydrocarbons is a very important industrial process that uses catalysts commonly based on transition metal oxides which are often vanadium and molybdenum oxides. Because of their industrial importance and the fact that they pose challenging scientific questions, selective oxidation catalysis has been rather extensively studied. A common feature among these reactions is that the desired products are often not the most thermodynamically favorable ones. For example, selective oxidation of hydrocarbons are usually carried out with oxygen or air. The most thermodynamically favorable products are water and carbon dioxide. The desired products, on the other hand, are alcohols, aldehydes, ketones, acids, anhydndes, or alkenes and dienes. Furthermore, for each hydrocarbon reactant, there are many possible products of various degrees of oxidation before carbon oxides are formed. Depending on the application, one or more of these partial oxidation products are desired. Therefore it is a challenge to understand the factors that account for the activity and the selectivity of a catalyst and to develop a practical catalyst.

The discussions here will concentrate on the chemical factors that determine activity and selectivity, with emphasis on the latter. Readers are urged to recognize that physical factors, such as pore size and heat and mass transfer can be very important. Local heating in a reactor is common because of the high heats of reaction in oxidation. It can lead to a drastic decline in selectivity as carbon oxides are generally favored at higher temperatures. Severe mass transfer limitations in catalyst pores result in concentration of partial oxidation products to be higher in catalyst pores than would be the case in the absence of mass rransfer limitations. Therefore, selectivity to total oxidation is enhanced. Similar enhancement is also observed at high conversions.

If we confine our discussions just to the chemical effects, variations in activity and selectivity among oxide catalysts are due to different chemical

169

170

bondings between the solid and the reactants and the reaction intermediates. There is as yet no universal rule that applies to all selective oxidation processes, perhaps because there are at least two different types of selective oxidation reactions (see next section). It is likely that the rate-limiting and the selectivity- determining steps differ for the two types of selective oxidation reactions. In this chapter, a summary of the current understanding of the role of a catalyst in these reactions is given. This is followed by descriptions of some selective oxidation reactions that are fairly well understood. The latter will be continued in the next chapter.

11.2 TYPES OF SELECTIVE OXIDATION REACTIONS

Selective oxidation reactions can be classified into two types: those that involve only dehydrogenation, and those that involve both dehydrogenation and oxygen insertion into the hydrocarbon molecule. Table 11-1 summarizes the common oxide-catalyzed selective oxidation reactions and the catalysts.

In dehydrogenation reactions, a hydrocarbon molecule is converted into a more unsaturated hydrocarbon by breaking C-H bonds and forming C=€ bonds. Often gaseous oxygen is used as an oxidant to yield water as a byproduct. This provides the thermodynamic driving force for the dehydrogenation process and permits the reaction to be conducted at a lower temperature than a simple dehydrogenation without oxygen. Sometimes other oxidants are used instead of oxygen, including iodine, bromine, and nitrous oxide. Among these, only iodine has had found commerical applications. In these dehydrogenation reactions, the carbon skeletons of the hydrocarbon molecules remain intact.

There are many examples of reactions involving both dehydrogenation and oxygen insertion. Oxygen is needed as an oxidant both for the formation of oxygenates and in the formation of water in the dehydrogenation steps. The general features of these reactions are that C-H bonds are broken, and C 4 bonds are formed. Exceptions to these include the oxidation of ethylene to ethylene oxide in which no C-H bonds are broken, and ammoxidation reactions such as the reaction of propene to form acrylonitrile in which C-N bonds are formed. In some cases, such as the oxidation of benzene to maleic anhydride, the carbon skeleton is broken. In others, the carbon skeleton remains intact. The selectivity is determined in part by the ability of the oxide to catalyze the formation of C-0 bonds without excessive breaking of C-C bonds which leads to combustion.

11.3 FEATURES OF CATALYTIC SELECTIVE OXIDATION

In selective oxidation reactions, gaseous oxygen is consumed in the formation of oxygenates and water, and almost always some carbon oxides. The reaction pathway of oxygen is as follows. Gaseous oxygen is oxidatively adsorbed on the oxide as 0-, 02-, or incorporated as lattice 0’- species. The solid is oxidized in this step, and the electrons acquired by the adsorbed oxygen could be from reduced

SELECTIVE OXIDATION I 171

Table 11-1 Common Oxide-Catalyzed Selective Oxidation Reactions. (From Ind. Eng. Chem. Prod. Res. Devel., 25, 171 (1986). copyright American Chemical Society).

Reaction Catalyst

Oxidative Dehvdrogenation

ethylbenzene -+ styrene isopentane, isopentene -+ isoprene butane. butene --+ butadiene

methanol ----j formaldehyde

Dehvdrogenation and Oxvgen Insertion

butane, butene --+ maleic anhydride propene -+ acrolein @ropene and NH3 -+ acryloniuile)

propene --+ acrolein, acrylic acid,

benzene -+ maleic anhydride o-xylene, naphthalene --+ phthalic

anhydride methane --+ methanol, formaldehyde ethylene -+ ethylene oxide

acetaldehyde

methyl ethyl ketone -+ biacetyl methyl ethyl ketone -+ acetaldehyde,

butane, butene -j acetaldehyde, acetic acid

acetic acid

V-Ti-0 Sn-Sb-0 Bi-Mo-0, promoted Fe-0 promoted V - 0 Fe-Mo-0, M a 3

v-P-0 Bi-Mo-0 Bi-Mo-0, U-Sb-0, Fe-Sb-0, Bi-S b-Mo-0 Co-Mo-Te-0, Sb-V-Mo-O

V-P-0, V-Sb-P-0 promoted V - 0

Mo-0. V - 0 Fe-Mo-0 (also catalyzed by promoted Ag) Co-0 (promoted by Ni, Cu) V-Mo-0

V-Ti-0

surface cations or anion vacancies with trapped electrons. When an oxygen atom is incorporated into a product molecule which is then desorbed, the electrons are returned to the solid. Therefore, in one catalytic cycle, electrons are removed from and then returned to the solid. In other words, the solid has undergone an oxidation-reduction cycle. If the oxygen species that is incorporated into the molecule is a lattice oxygen atom, the sites for adsorption of oxygen and for oxygen incorporation into the molecule may be different, and migration of oxide

172

Figure 11-1 Schematic representation of the redox cycle in a selective oxidation reaction involving lattice oxygen. $'s are anion vacancies.

ions in the solid between the two sites would occur. This situation is shown schematically in Fig. 11-1, which was first proposed by Mars and van Krevelan [l], and is commonly referred to as the Mars and van Krevelan mechanism. One should emphasize that there can be many possible variations from this scheme. For example, alkene adsorption may take place on a partially reduced cation, or the hydrocarbon molecule may adsorb with heterolytic dissociation.

This mechanism which involves lattice oxide ions have a number of consequences. Since some cations in the solid continually undergo alternate oxidation and reduction during the reaction with a consequent change in the local cation- anion ratio in the solid, the catalyst must be able to accommodate these changes reversibly and easily. For example, in molybdenum oxide, this is achieved by the ready interconversion of corner-sharing and edge-sharing M06 octahedra as examplified by the formation of shear planes, and in vanadium phosphorous oxide by the ready interconversion between two crystallographic structures.

Another consequence is the interesting feature that many reactions proceed with the same initial activity and selectivity in the presence or absence of gaseous oxygen. That is, a catalytic reaction and a noncatalytic surface reaction give the same initial rate and selectivity. Table 11-2 summarizes some of the known examples. In the absence of gaseous oxygen, oxidation of the hydrocarbon is by the cation in the solid. Usually, the reaction rate and the selectivity for partial oxidation products decrease with increasing extent of reduction of the oxide, although the rate of decrease depends on the particular oxide. For some oxides, such as a-iron oxide, however, the selectivity is improved initially when the catalyst is reduced before it deteriorates. In this case, the presence of surface anion vancancies may be important.


Table 11-2 Examples Where the Rate andor Product Distribution Have Been Shown to be Similar in the Presence and Absence of Gaseous Oxygen

Reaction Catalyst

Ammoxidation of propene Bi-Mo-0 Oxidation of propene Bi-Mo-0 Oxidation of butene Bi-Mo-0, Fe-SbO, Bi-W-0, Sn-Sb-0

Sb-Mo-0, Sn-P-0, U-Sb-0, Ni-Sn-P-K, V-P-0, Fe-0

CH30H oxidation M a 3

Although the catalysts in Table 11-2 sustain their activities and selectivities even after a substantial extent of reduction of the oxide, some other oxides do not. Cobalt molybdate is an example. When propene is oxidized in the absence of oxygen on this catalyst, acrolein is formed with the consumption of the first surface layer of lattice oxygen. As the degree of reduction increases, the acrylic acid yield as well as the oxidation activity decrease rapidly [2]. Experiments using isotopic gaseous oxygen further show that only a few (3 to 6) layers of lattice oxygen are involved in the oxidation of propene [3].

The incorporation of lattice oxygen into selective oxidation products has now been demonstrated for quite a number of systems. However, it must be recognized that this is not universal. This is illustrated by the selective oxidation of propene over a Sn-Mo oxide catalyst [4]. Passage of a mixture of propene, oxygen-160, and water-I80 over the catalyst yields acrolein, acetone, acrylic acid, acetic acid, and carbon oxides as products. There is a substantial incorporation of in acetone, but much less in acrolein. Thus the sources of oxygen in these two major products are different.

It has been observed that the proximate origins of oxygen atoms used in the formation of water during the oxidative dehydrogenation of butene to butadiene and in the selective oxidation of propene to acrolein are different on a Bi2Mo06 catalyst. Ueda, et al. [5] have shown that the conversion of consecutive propene or butene pulses at 4OO0C decreases in the absence of gaseous oxygen until the catalyst becomes inactive eventually. From the total amount of propene reacted, the extent of reduction required to deactivate the catalyst for the propene reaction can be calculated. If the catalyst is first prereduced to the same extent by the butene dehydrogenation reaction or by hydrogen and then propene pulses are passed over the catalyst, the conversion of propene is found to be essentially the same as for a fresh catalyst. Thus the reduction by butene or H2 does not remove the oxygen important in the oxidation of propene. This conclusion is confirmed by l80 labeling experiments [5] . In these experiments, a catalyst is first reduced by butene pulses and then reoxidized with 1802. The reoxidized catalyst is then

174

10 20 30 40

Reaction Time, min

Figure 11-2 '*O concentration of acrolein formed on "0 labeled BizMoOB catalyst in the oxidation of propene with ''Oz at 400°C. The oxidation of propene was carried out after pretreating the catalyst by propene and "02 (open circles) or 1-butene and I8O2 (closed circles). (From J. Chem. SOC. Farad. Trans. I, 78, 495 (1982), copyright Royal Society of Chemistry).

used for propene oxidation. The l80 content in acrolein is found to increase in the first couple pulses. These results are shown in Fig. 11-2. If the catalyst is reduced with propene instead of butene and then reoxidized, the l80 content in acrolein decreases with the pulse number. Thus the proximate source of lattice oxygen for the formation of water in dehydrogenation must be different from that used for insertion into propene.

Reduction of Bi2Mo06 followed by reoxidation with l8OZ also causes shifts in Raman bands related to the M M bonds at 725, 803, and 844 cm-'. The magnitude of the band shifts are much larger if the reduction is with propene than with 1-butene, consistent with the above conclusion that different oxygen species are involved in the two reactions [6,7].

Another interesting feature in selective oxidation catalysis is the influence of water. Water is often added with the feed as a diluent to reduce the tendency of


coking of a catalyst. Less understood is the effect of water in the feed changing the surface hydroxyl concentration and the acidity and basicity of the catalyst. A high concentration of surface hydroxyl groups enhances the possibility of reaction of surface intermediates with hydroxyl groups. For example, in the oxidation of methane with N 2 0 over a silica-supported molybdena catalyst, methanol is produced at a much higher rate in the presence of water [8.9]. Similarly, in the oxidation of ethane with N20 over a supported Moo3 catalyst, the presence of water results in a higher yield of acetaldehyde primarily at the expense of ethene [10,11]. In general, the presence of water vapor enhances the selectivity for alcohols and saturated ketones in the oxidation of alkenes. For example, propene is first hydrated to 2-propanol which is then oxidatively dehydrogenated to acetone when propene is passed over a Sn-Mo-0 catalyst with water and oxygen [12]. In the absence of water and at a higher temperature, no acetone is formed. Instead, acrolein is produced with 45% selectivity [13]. One of the most interesting observations in this regard is the selective formation of acetic acid and acetaldehyde from butene in the presence of oxygen and water over a V-Ti-0 catalyst [ 141, whereas maleic anhydride is a major partial oxidation product over many vanadia catalysts when the feed does not contain water.

11.4 CHEMICAL FACTORS AFFECTING SELECTIVITY

Past effort by many research workers has provided significant understanding of heterogeneous catalytic selective oxidation processes. Some of the efforts have been to find correlations between activity and selectivity and properties of the catalysts. The following correlations have been proposed:

a) Presence of cation vacancies: Correlations have been found between the rates of production of selective oxidation products in the oxidation of butene and propene on schcelite catalysts derived from PbMo04 [15] and the presence of cation vacancies. Later work concluded that the cation vacancies associated with Bi ions are the important species [16].

b) Metal-oxygen bond strength: Since the discovery that lattice oxygen is incorporated into the product acrolein in propene oxidation, it has been proposed that an optimal M a bond strength is required for high selectivity in oxidation. More recently, correlations have been made between selectivitjj and the M-O bond strengths measured as the rates of change of the h a t s of reduction of oxides with the degree of reduction [17], or measured as a function of M a bond lengths [18].

c) Ability to form shear structures: The ability of a catalyst to form shear structures facilitates the oxidation and reduction cycle that a catalyst undergoes during reaction without the need for major structural changes [19].

d) Optimal density of active oxygen: The active site must not contain too many or too few active oxygen ions that can participate in the reaction. Too many would lead to excessive oxidation, and too few would lead to

176

an inactive catalyst [201. e) Acid-base properties of the oxide: Reactants and products of oxidation

reactions can be classified as Lewis acids or bases according to their ionization potentials. Whether a catalyst contains strong or weak Lewis acid or base sites will determine the strength of interaction of the reactants and products with the solid [21], which in turn determines whether a reactant can be readily adsorbed (and presumably activated) or a product can be readily desorbed.

f) Electron binding energy of lattice oxygen: A correlation has been observed between the electron binding energies of lattice oxygen in four catalysts as determined by XPS and their activities for selective oxidation [22]. This correlation is based on the concept that the electron binding energy can be used as a measure of the basicity of the lattice oxygen, and that for the same cation, a more basic lattice oxygen ion could more easily abstract an allylic H atom from an alkene. This last point will be discussed further later.

g) Crystallographic plane: Different surface planes have different atomic arrangements, which may show different activities and selectivities. For example, different selectivities have been observed on different crystallographic planes of MoOj in the oxidation of methanol and ethanol [23,24]. However, only a phenomenological explanation is presently available.

h) Presence of M=O bonds: It has been noted that terminal MLO bonds are present on many selective oxidation catalysts such as Moo3, molybdates [25,26], and VzOs [27]. Thus it has been suggested that their presence is important. However, not all selective catalysts contain such M=O bonds, whereas some nonselective catalysts do. This will be discussed further later.

desorption to determine which oxide adsorbs oxygen weakly, a correlation has been found that for simple binary oxides, those that are selective in oxidation do not have weakly adsorbed oxygen [28]. This observation is consistent with those which show that weakly adsorbed oxygen causes combustion of adsorbed hydrocarbons [29].

i) Absence of weakly adsorbed oxygen: Using temperature programmed

Further discussions of these correlations are found in the sections that describe individual reactions. The existence of so many correlations indicates the complexity of the problem. Because selective oxidation is a multistep process and there are different types of selective oxidation reactions, the important step that determines activity may well be different from the step that determines selectivity, and these steps may be different for different reactions. Even for the same reaction, the critical step may depend on the oxide and operating conditions. In other words, because of the diversity of this area, it is unlikely that one corrclation is sufficient to explain all processes under all conditions on all catalysts. To expand on this, consider the examples of the oxidation of propene to acrolein and of butcne to butadiene. Under industrial operating conditions, the rate-limiting step for both reactions is the breaking of the allylic C-H bond (first C-H bond


breaking). A successful correlation would be rate versus factors that affect the rate of breaking of this bond. At lower temperatures, however, the desorption of acrolein [30] or butadiene [31] are rate limiting. In this case, a successful correlation for the rate would be with factors that affect rates of desorption of these products, Since selectivity is determined by factors other than those that determine rate, these correlations would not be successful for selectivity.

That selective oxidation is a multi-step sequence of reactions is well recognized. The need for every step to proceed rapidly on a good catalyst is also known. Grasselli and coworkers have identified three major functions of a catalyst for the oxidation and ammoxidation of propene [6,32]: allylic hydrogen abstraction, oxygen insertion, and oxidation-reduction of the catalyst. Different cations important for each function have been identified. An active and selective catalyst must then contain cations for every function. We shall discuss this in detail later.

Approaching from a different point of view, Haber analyzed the question of selectivity by considering the types of oxygen available for the reaction [33]. He proposes that the reactivity of an oxygen depends on whether it is electrophilic or nuclcophilic. Electrophilic oxygen species, such as adsorbed 02- and 0- are very active. They attack hydrocarbon molecules at the regions of high electron densities. Saturated aldehydes are formed which readily undergo total oxidation. Nucleophilic oxygen species, such as lattice oxygen are less reactive and are suitable for partial oxidation. In a reaction that makes use of nucleophilic oxygen, activation of the hydrocarbon molecules is the rate-determining step, whereas in a reaction that uses electrophilic oxygen, adsorption of oxygen to form the electrophilic species is the rate-limiting step. In Haber's model, a selective oxidation catalyst should be capable of adsorbing and activating a hydrocarbon molecule for nucleophilic attack by oxygen. It should be able to insert nucleophilic lattice oxygen efficiently into the hydrocarbon molecule, and be able to replenish rapidly the lattice oxygen consumed. It should not generate electrophilic oxygen species.

Haber's treatment, while generally plausible, is not universal. For example, the statement that electrophilic 02- or 0- are not suitable for selective oxidation needs qualification. In the recent examples of selective oxidation of methane on Si02-supported M a 3 134-361 and V205 [37,38] catalysts, N20, which decomposes rcadily to adsorbed 0-, is an effective oxidant. In another example, the oxidation of ethylcne to ethylene oxide, the standard heat of reaction is -24.7 kcal/mole. Thus the heat of adsorption of the oxygen involved must be smaller than this value. In other words, it must be a relatively weakly adsorbed oxygen.

Another approach to understanding selective oxidation is to analyze each step in the reaction sequence individually, and consider the factors that affect each step. This approach makes it possible to understand why changes in some catalyst propcrtics would rcsult in the observed changes in the reaction characteristics when the ratc-limiting and the selectivity-determining steps are known.

To illustrate this approach the oxidation of butane to maleic anhydride is used as an cxamplc. A simplificd scheme for the sequential oxidation of butane is shown in Fig. 11-3. This reaction contains all the essential features of most

178

0

CO + C02 + H20

Figure 11-3 A simplified scheme for the oxidation of butane.

selective oxidation reactions. Step 1 involves the activation of an alkane. Step 2 involves the activation of an alkene. Steps 3 and 4 involve insertion of oxygen into the hydrocarbon. The chemical transformations involved in these steps have been analyzed, and the catalyst properties important in these steps have been postulated [391. A summary of the results are presented below.

a) Alkane activation (step 1). Because alkane molecules are very inert and interact only very weakly with most catalysts, it is believed that the transition state in the breaking of the first C-H bond in alkane activation is reactant-like. The energetics of the process favors the dissociation of the C-H bond in a manner similar to the production of free radicals. Furthermore, because of the strong C-H bonds in alkanes, this reaction step requires a very reactive weakly adsorbed oxygen species to be energetically favorable. Weakly adsorbed 02- and 0- species are some of the possible candidates. Once the first C-H bond is broken, a weakly adsorbed alkyl radical or a surface alkyl species would be formed, and breaking of the second C-H bond to form an alkene should proceed easily.

b) Alkene activation (step 2). Unlike alkanes, the C=C bonds of alkenes can interact strongly with the cations of a catalyst. The strong bonding makes possible charge delocalization between the adsorbate and the cation. Since the charge on the cation in the solid can be much better stabilized by electrostatic forces in the solid than charges on the adsorbate outside the solid, it becomes possible that the allylic C-H bond breaking in this step is achieved with charge transfer between the x-ally1 and the surface cation:

(1 1-1) RCH-C H X H 2 - - - - I - - - RCH=CHCH3 + Mn+02- --+ [

M 0 2 -

This step which is heterolytic dissociative adsorption of alkene is enhanced if (i) the cation can readily undergo reduction; (ii) the surface oxygen is a strong Brdnsted base; and (iii) h e cation is a soft acid.

Following the allylic C-H bond breaking, a second C-H bond must also be


broken to form diene. Breaking of the second C-H bond can be homolytic and does not need to be accompanied by charge separation. It can be enhanced by weakly adsorbed highly reactive oxygen species.

In both dehydrogenation of alkane to alkene and alkene to diene, the products must desorb rapidly. Since unsaturated hydrocarbons are soft basic ligands, desorption is facilitated by hard acid cations. For this step, the solid is preferably ionic, and the cation is in a high oxidation state.

At this point, it becomes apparent that different, and sometimes opposite, properties are desirable for different steps. Take for instance the desirable properties of a cation in the dehydrogenation of alkene to diene. It is desirable for the activation of an alkene to have a soft cation, but for the desorption of diene to have a hard cation. Therefore, to determine the role of a promoter or in seeking meaningful correlations, it is important to be able to identify the critical factor (i.e. catalyst property) in the rate-limiting step.

c) Oxygenate formation (steps 3 and 4). The formation of oxygenates from a diene or an alkene involves breaking C-H bonds, forming C-O bonds (i.e. oxygen insertion), and desorption of the oxygenates. In the oxygen insertion step, it is important that the appropriate number of oxygen atoms be incorporated into the molecules. Excess incorporation or incorporation at the wrong position leads to undesirable combustion. There are at least two important oxide properties that affect this step: geometric effect and metal-oxygen bond energy.

The geometric effect states that an ideal active site must not have too many or too few oxygen ions or atoms that can participate in the reaction [6,32,40]. Too many would lead to excessive oxidation, too few would lead to an inactive catalyst. This concept points to the importance of the crystal structure and the surface atomic structure.

The importance of metal-oxygen bond strength has long been recognized. It is generally believed that too weak bonds result in nondiscriminative C 4 bond formation which leads to combustion, and too strong bonds lead to unavailability of lattice oxygen that could participate in the reaction. Since the M-O bond strength usually increases with the degree of reduction of the oxide, one concept is that there is an optimal value for the rate of increase in the heat of reduction of an oxide with the extent of reduction near the steady state of the oxide [17].

Both the geometric and the M-O bond strength effects are short-ranged. An efficient way for a promoter to affect them is by the formation of a compound. The beneficial effects in the formation of bismuth molybdate, vanadium pyrophosphate, etc. could be in part due to this reason. A correlation mentioned carlier that krminal M=O groups are important may also be another manifestation of the geometric and M a bond strength effect.

The lattice oxygen that is used for C 4 bond formation must be replenished to maintain the catalyst at the optimal steady state. The ability of an oxide to form shcar structure facilitates this reduction-oxidation process [9]. However, this is not a necessary condition for selective oxidation catalysts, as some of them (e.g. silver metal, copper oxide) are not known to form shear planes.

After the desired oxygenates are formed, they must be desorbed. Many

180

Table 11-3 Desirable Surface Properties in Selective Oxidation Reactions. (From Ind. Eng. Chem. Prod. Res. Devel., 25, 171 (1986), copyright American Chemical Society).

Reaction Properties

Alkane activation:

C-H bond dissociation Highly reactive surface oxygen (weakly adsorbed oxygen and/or surface lattice defects).

Alkene activation:

C-H bond dissociation (i) Surface oxygen strong Brbnsted base (ii) Cation readily undergoes reduction (iii) Cation soft acid

Diene (or alkene) desorption Cation hard acids

Water desorption Cation soft acids

Oxygenate formation:

Oxygen insertion (i) Limited number of available oxygen (ii) d AH,/& too large for further reduction

Oxygenate desorption Cation soft acids

Prevent combustion: (i) Short residence time of surface intermediates (ii) Weak adsorption of desired product (iii) No weakly adsorbed oxygen (iv) No combustion site

oxygenates are adsorbed via the electron lone pair of oxygen. They are hard bases. Thus desorption is easier from soft than from hard cations.

d) Combustion reaction. This is an undesirable reaction and should be minimized. This can be achieved by (i) shortening h e surface residence time of surface intermediates, especially the desired products; (ii) eliminating weakly adsorbed oxygen species which are very reactive; (iii) eliminating active sites that only lead


to combustion products. Property (ii) has been discussed earlier. For (i), very reactive oxygen is desirable for activating the strong C-H bonds of alkanes, but is undesirable for other processes. (iii) is self-explanatory. A summary of the desirable properties of a catalyst for each of the steps described is presented in Table 11-3.

11.5 OXIDATION OF PROPENE TO ACROLEIN AND AMMOXIDATION TO ACRYLONITRILE

Catalytic partial oxidation of propene can result in a number of products such as acetone, propionaldehyde, and acrolein. In the presence of ammonia, ammoxidation to acrylonitrile is also possible. Among these reactions, the catalytic oxidation to acrolein:

CH3CHxH2 + 0 2 300-4500c) CH2=CHCH0 + H20 (11-2)

and the ammoxidation to acrylonitrile:

CH3CHKH2 + NH3 + 3/22? (11-3)

400-4600c+ CH2=CHCN + 3 H 2 0

are commercially very important and the most studied and understood. We shall discuss them in this section.

There have been successive generations of commercial catalysts for these reactions based on molybdates, cuprous oxide or antimonates. Bismuth molybdates, tin antimonates and uranium antimonates were the catalysts in the recent past. Work on tin antimonates has been recently reviewed [41]. The latest catalysts are the multicomponent molybdate catalysts, M.2+Mb3+Bi,Moy0, (I@+ = Ni, Co, Mg, Mn, and M3+ = Fe, Cr, Al, Ce) that are based on bismuth molybdate, multicomponent antimonate catalysts, M,MbF%SbyO,, and Ce-Te-Mo oxide. Experience has shown that a catalyst that is active and selective for propene ammoxidation to acrylonitrile is also excellent for oxidation to acrolein.

The selectivity for acrolein or acrylonitrile over the commercial catalysts are very high, usually in excess of 80% even at high conversions. In the oxidation reaction, the byproducts include carbon dioxide, carbon monoxide, acetaldehyde, formaldehyde, propionaldehyde, propionic acid, formic acid, and acetic acid.

Since most of the selective catalysts are based on molybdates, the published work has concentrated on the molybdate system. In particular bismuth molybdate, which exists as Bi2Mo06 (y phase), Bi2Mo3012 (a phase), and Bi2M0209 (p phase), has been heavily studied. In this system, the y phase is the most selective and active, while the a phase is inferior to the other two.

182

Reaction Mechanism Based on work by various workers, detailed mechanisms for both oxidation

and ammoxidation over bismuth molybdate have been proposed by Burrington, et al. 142,431. They are shown in Fig. 114.

In these mechanisms, the surface active site is a Bi-Mo pair site (I) that is composed of a Bi-O group responsible for allylic H abstraction, and a molybdenum dioxo group for oxygen insertion or diimido group for nitrogen insertion. The molybdenum cation is believed to be the propene adsorption site, although the discussions later would suggest that this assignment of sites is not yet universally accepted. The oxidation reaction proceeds via dissociative adsorption of propene to produce a x-ally1 species (III). Under reaction conditions, whether dissociation of propene occurs on adsorption (I + I11 directly) or propene is first adsorbed molecularly as a II; complex and allylic H abstraction follows (I + I1 + 111) is not clear. This step is followed by the formation of a C-0 bond (I11 + IV), and a second hydrogen abstraction in the form of a 1.4 shift in species IV to produce adsorbed acrolein and Mo-OH. Desorption of the product and reoxidation of the catalyst complete the cycle.

In the ammoxidation cycle, a surface molybdenum diimido species (VI) is formed by the reaction of &ox0 groups with ammonia (I -+ VI). Propene is activated in the same way as in the oxidation to acrolein. The reaction proceeds via dissociative adsorption of propene (VI + VIII), formation of a C-N bond (VIII -+ IX), and a second and a third hydrogen abstraction to form acrylonitrile. The reduced surface site is reoxidized, and then reconverted to the diimido species to complete the cycle. It has also been suggested that the diimido species is important at high propene pressure (10 P a ) , and a monoimido species is important at low propene pressure (4 P a ) [431.

In both oxidation and ammoxidation, the first step of the reaction is the dissociative adsorption of propene to form an adsorbed symmetrical x-ally1 species (I + I n and VI + VIII). The involvement of a symmetric x-ally1 species has been convincingly shown 6rst by Adams and Jennings on bismuth molybdate and cuprous oxide catalysts [44,45]. When propene labeled with deuterium at either end is oxidized, deuterium atoms are found only at the end carbons of acrolein, and the distribution of deuterium is the same regardless of which deuterated propene is used.

CH2=CH-CH2D or -- Bi-Mo-o CH@CH=CHD + CDHH=CH2 (11-4)

CD2=CH-CH3

Experiments using carbon labeling also substantiate this conclusion. It is found that the carbon labels at either end of propene is always equally distributed in the end carbons of acrolein. This has been confirmed over bismuth molybdates [46-48] as well as cuprous oxide [49], uranium antimonate [50], tin-antimony oxide [51], and supported rhodium, ruthenium [52] and gold [53] catalysts.

In principle, the symmetric distribution of labeled carbon or deuterium in

SEI ,E

CT

IVE

OX

IDA

TIO

N I

183

184

acrolein can be due to the abstraction of a hydrogen atom to form rc-ally1 or the addition of a hydrogen atom to form a 2-propyl species:

When the reaction is studied over a deuterated catalyst, the 2-propyl route would yield approximately 50% monodeuterated acrolein, while the rc-ally1 route would yield 0%. Experimentally, it is found that very little deuterium is incorporated in the acrolein. thus confirming the rc-ally1 route [46].

This rc-ally1 intermediate behaves like an adsorbed radical. Martin and Lunsford have detected with electron spin resonance the desorption of rc-ally1 radicals from a bismuth molybdate catalyst [54,55].

In the oxidation of para-substituted phenylpropene @-XC6&CH2CH~H2), the rates of oxidation of these compounds over bismuth molybdate at 320°C relative to the unsubstituted phenylpropene (X=H). rXhH are 1.55, 1.98 and 3.03 for X = CH3, C1, and W H 3 , respectively. These rates compare well with the rates of decomposition of psubstituted phenyl azoethanes (p-XC6&CH(CH3)-N=b- (CH3)CH@-XC6&)). These azocompounds decompose into allyl radicals. The good agreement suggests that adsorbed ally1 radicals are formed in the oxidation of phenylpropenes [56].

Other evidence is also consistent with this adsorbed n-ally1 radical intermediate, but the data may also be interpreted as involving a n-ally1 cation intermediate. The rates of oxidation of various alkenes depend on the type of allylic hydrogen. The relative rates over bismuth molybdate at 460°C for alkenes with a tertiary, a secondary and a primary allylic hydrogen are 75, 14, and 1 [57], which parallel the order of stability of the rc-ally1 cations or radicals, and the order of the C-H bond energies. Exposure of the catalyst to propene results in the reduction of Mo6+ to Mo5+ which can be followed by EPR [58] . Therefore, there is electron transfer from the surface intermediate to the cation. If the allylic hydrogen is abstracted from propene as a proton, the electron transfer would result in an adsorbed x-ally1 radical (see eq. 11-1). If the hydrogen is abstracted as a hydrogen atom, the electron transfer would result in a rc-ally1 cation.

The oxidation products of propene on bismuth molybdate have also been compared with those of n-ally1 radicals generated by the decomposition of azopropene [59], and n-ally1 cations generated by the decomposition of allyl iodide [60]. Both azopropene and allyl iodide yield selective oxidation products similar to propene, but azopropene produces more C 0 2 , acetaldehyde and benzene, and less acrolein than allyl iodide. While this evidence seems to support an allyic cation intermediate, the effect of iodide is not clear nor is the consequence of azopropene and ally1 iodide.

Above 400°C. the ratc-limiting step in the catalytic oxidation of propene over bismuth molybdate is the abstraction of the allylic hydrogen. This makes it


difficult to study by kinetics alone the steps subsequent to the first hydrogen abstraction, and there is no conclusive evidence as to whether the abstraction of the second hydrogen precedes or follows oxygen or nitrogen insertion.

In one report, it is claimed that the details of the next step depends on the catalyst [61]. On bismuth molybdate, the oxidation of (E)-propene-l-dl yields a product mixture that contains 1:l: 1 (E)-acrolein-3-dl, (Z)-acrolein-3-dl, and acrolein-dl .

(E)-propene-1-dl (Ei)-acrolein-3-dl (Z)-acrolein-3-dl acrolein-1-dl

This can be explained by rapid interconversion of x-ally1 and 0-ally1 intermediates, which equilibrates the (E) and (Z) spccies:

The ratio of these products is 1:1:1.6 over cuprous oxide. Thus the x-ally1 and the 0-ally1 intermediate do not equilibrate as rapidly on this catalyst [62].

It is not yet established whether the abstraction of the second H precedes or follows 0 or N insertion to the adsorbed sc-allyl. The H-D kinetic isotope effect in acrolein production on bismuth molybdate is consistent with the assumption that abstraction of the second hydrogen precedes oxygen insertion. This is shown by the results [44,45] of experiments which monitor the deuterium content in the various positions of acrolein (CH@CH=C(DJ-Q2 and C(D,H)O-CH<H2) during the oxidation of CH2=CH-CH2D and CHD=CH-CH3. The experimental results are compared with the calculated values which are based on the relative probability of a deuterium atom being abstracted relative to that of a hydrogen atom, assuming that the abstraction of the second hydrogen precedes oxygen insertion. The results agree well. The kinetic isotope effect, kH/kD, for the second hydrogen abstraction at 450°C is found to be 1.82, while kH/kD for the overall reaction of acrolein production is 2.4 at 320°C [63].

A different conclusion is obtained for the second hydrogen abstraction in the oxidation of allyl alcohol (CH2=CH-CD20H) [63]. Heterolytic dissociative adsorption of allyl alcohol on a bismuth molybdate catalyst produces 0-0-ally1 molybdate (species IV in Fig. 11-4):

186

This species is the same as the one that would be formed by oxygen insertion to a surface n-allyl. Thus the acrolein produced from allyl alcohol would represent the result of a mechanism in which oxygen insertion precedes abstraction of the second hydrogen. The interpretation of the results, however, is complicated by the isomerization of allyl alcohol which equilibrates the two ends of the molecule:

H H H2C=C-CD20H C- HOHzC-C=CDz (11-7)

The results of the oxidation and ammoxidation of a mixture of 1,l-d2-ally1 alcohol and pyridine (the latter is added to suppress the isomerization) have been compared with those of propene-l,l-dz. Although there are some differences, similarity between results with the two reactants is large enough to indicate that 0 or N insertion precedes the second H abstraction.

In addition to the surface reactions shown in Fig. 11-4, it has also been established that at 45OoC over bismuth molybdate, desorbed allyl radicals react with gas phase oxygen to form peroxide species, which undergo a homogeneous gas phase reaction with propene to form propene oxide [64,65]:

It is believed that these surface-initiated homogeneous reactions would become important at high propene to oxygen ratios and in reactors having large post- catalytic volumes. This hydroperoxide species, however, does not participate in the surface reaction on Bi-Mo-oxides, although it is a possible intermediate over USb301o [MI.

The detailed mechanisms for the production of the side products (products other than acrolein or acrylonitrile) are less clear. The same kinetic isotope effect, kH/kD. is observed in the formation of acrolein and carbon dioxide on Bi2M03012 and BizM006 [66]. This suggests that the same rate limiting-step applies to the formation of both products. There are indications based on isotope labeling experiments that carbon dioxide arises mostly from further oxidation of acrolein at the vinyl group [65,67], whereas formaldehyde is derived from the carbonyl group of acrolein [68].

On Bi3FeMo2012, however, the kH/kD ratio is smaller for C02 than for acrolein. Thus there may be an additional mechanism leading to combustion on this catalysts [66].

Kinetics Propene oxidation proceeds readily above 300°C. On bismuth molybdate


Table 11-4 Activation Energies for the Oxidation of Propene, 1-Butene, and the Reduction of Oxides by Propenea

Oxidation of redn by propene Catalyst propene kJ/mole 1-butene kJ/mole kJ/mole

Bi2Mo06 59 46 65 Bi2MozO9 63 46 67 Bi2(M004)3 71 46 67 Bi203 59b

Footnotes: a) From J. Peacock, et al.. J. Catal., 15, 398 (1969), copyright Academic Press. b, From W. Martin, and J. Lunsford, J. h e r . Chem. SOC., 103, 3728 (1981).

However, a value of 92 kJ/mole has been quoted in another report (M. White. and J. Hightower, J. Catal., 82, 185 (1983)).

and cobalt molybdate, the reaction is close to fist order in propene and zeroth order in oxygen. The ammoxidation reaction requires a higher temperature (over 400OC). For a given propene and oxygen concentration in the feed, the relative rates of production of acrylonitrile versus acrolein increases with the NH3/propene ratio [42,59].

The activation energies for the oxidation of propene for different Bi-Mo- oxides are shown in Table 11-4 [58]. Also shown in the table are the activation energies for the oxidation of 1-butene and the reduction of the oxides by propene. Since abstraction of the allylic hydrogen from propene is believed to be the rate- limiting step in the oxidation of propene, the similar activation energies for this reaction over Bi203 and Bi-Mo oxide of various B W o ratios suggest that Bi is the site for propene adsorption and formation of n-allyl. This suggestion is supported by the fact that the activation energy for the production of gas phase n-ally1 radicals from propene on Bi203 is 67 kJ/mole, which is close to the value for propene oxidation [69]. It is also supported by the observed constant activation energy for the oxidation of 1-butene, which is also believed to proceed via the formation of n-methylallyl, and for the reduction of the oxides by propene, assuming that its rate-limiting step is allylic hydrogen abstraction from propene.

Callahan, et al. have also determined the activation energy for the oxidation of propene [70], and reported values of 80-88 kJ/mole for various molybdates. These values are the same for ammoxidaton of propene over the same catalysts, which suggests that the rate-determining steps are identical in both reactions. The differences in the activation energies between these values and those in Table 11-4 have not been explained.

On the multicomponent catalyst (M,MbBi,Mo,O,) at about 440°C, the

188

Table 11-5 Kinetic Isotope Effect in the Oxidation of Propene over Bi-Mo Oxide. (From J. Caul., 87, 363 (1984), copyright Academic Press).

Reactant Relative Rate of Oxidation

1 .oo 0.85 0.98 0.55

difference in the activation energies between C02 formation and ammoxidation of propene is 39 kJ/mole, and C02 formation and selective oxidation to acrolein is 5 kJ/mole. This implies a 34 kJ/mole difference in the overall process for nitrogen versus oxygen insertion [42].

The activation energy for propene oxidation over cobalt molybdate ranges from 51 to 67 kJ/mole [71,72]. The activation energy for acrolein formation is 37 kl/mole. There are reports that the rate of propene reaction is retarded by the presence of the products acrolein and acrylic acid [73].

The kinetic isotope effect for propene oxidation over a Bi-Mo oxide catalyst has been determined by Adams and Jennings using C3H, and C3D6 [44,45]. In these studies, it is assumed that the first and second H abstraction have the same isotope effect. Then the value of kH/kD is deduced from the data. This value depends on the temperature. It is 2.3 k 0.4 at 365"C, and 1.8 f 0.3 at 475°C. These values are very close to the theoretical values for the dissociation of a C-H or versus a C-D bond. It is further found that the isotope effect is much larger if the D atom is located in the CH3 group than in the CH2 group. This is illustrated in Table 11-5.

Using [2,3,3,3-&]-propene, Krenze and Keulks have observed kH/kD values of 1.7 to 1.8 at 450°C and 1.5 to 2.2 at 350°C over Bi2M03012, Bi2M006, and Bi3FeMo2OI2 [66]. The close agreement of this value on Biz03 and Bi-Mo oxides again suggests that the abstraction of the allylic H to form an adsorbed x-ally1 species is the rate- determining step, that Bi ion is involved in this step, and the intermediate is a rc- allyl. The isotope effect on the activation energy in the formation of desorbed gas phase x-ally1 radicals from Bi2O3 has been determined over 365475°C to be 67 kJ/mole for C& and 74 kJ/mole for C3D6 [541.

Ammoxidation exhibits a similar kinetic isotope effect as oxidation when investigated using [3,3,3-d3]-propene. This shows that the two reactions have the same rate-limiting steps [44,45].

On Bi203, the value is 1.7 at 400°C [54].


Nature of Catalysts (Bismuth Molybdate) According to the mechanism in Fig. 11-4, the catalyst provides the function

for allylic hydrogen abstraction, C-O or C b J bond formation, and activation of gaseous oxygen. In addition, the catalyst undergoes redox cycles during the reaction. Grasselli and Bunington have summarized these requirements and assigned individual components in multicomponent oxides based on bismuth molybdate and other selective oxidation catalysts to these functions [59]. Their summary is shown in Table 11-6.

These assignments are very helpful towards understanding the role of the catalysts. Although the emphasis in these assignments is on the nature of the cations, it is understood that the oxide ions bonded to these cations are just as imponnat.

For bismuth rnolybdate, the oxide ion bonded to Bi is proposed to be responsible for allylic hydrogen abstraction. The similar kinetic parameters already mentioned earlier for this catalyst and bismuth oxide support this proposal. The x-ally1 species is proposed to adsorb on a Mo ion. This is supported by some but not all evidence. For example, it has been observed that propene oxidation on bismuth oxide yields hexadiene, which can be produced by coupling of two ally1 radicals [74,75]. Thus n-ally1 is produced and presumably can adsorb on Bi ions. lSO-labeled bismuth molybdate can be prepared such that l80 is concentrated either in the Mo layer or in the Bi202 layer. When propene is oxidized over this oxide, the oxygen label in acrolein is consistent with the view that propene is first oxidized by the Bi202 layer [76]. This implies that propene is adsorbed on a Bi ion.

On the other hand, unsupported Moo3 can oxidize propene to acrolein. Thus adsorption of x-ally1 on Mo ion is possible. Furthermore, as mentioned earlier, adsorption of propene on Bi-Mo-0 results in the formation of MoS+ ions [58 ] , which suggests that electron transfer from adsorbed x-ally1 to Mo ions occurs.

There are data which suggest that Bi ions in conjunction with cation vacancies are important. Sleight and coworkers [77] have shown that in the oxidation of propene and butene on PbMo04, the rate of production of selective oxidation products increases with the amount of cation vacancies introduced when Pb is substituted by Bi to produce Pbl-3xBi2x@xMo04. Later work by Brazdil, et al. shows that the important species are cation vacancies associated with Bi ions [16]. In the Pb0.84-3xB~.08Na,,08La2x@xMo04 system, the Bi content can be kept constant, and cation vacancies can be introduced by the incorporation of La. It is found that with increasing cation vacancy density, the rate of ammoxidation of propene and the acrylonitrile yield increase, while the selectivity remains constant. However, in the Pb-La-Mo-0 system that does not contain bismuth, the yield and selectivity for acrylonitrile are both low and independent of the density of cation vacancies. Thus cation vacancies are important only when they are associated with Bi ions.

Although the current view is that sites involving Bi are responsible for allylic hydrogen abstraction, there are data that suggest other possibilities. For example, when l-butene is oxidized over BiOCl and BiOBr, which have the same type of (Bi2022+)n layer structure as Bi2Mo06, butadiene is formed very selectively. The

190

Table 11-6 The Functions of Various Components in the Selective Oxidation Catalysts. (From Ind. Eng. Chem. Prod. Res. Devel., 23, 393 (1984), copyright American Chemical Society).

Function

Catalyst allylic H alkene chemisorption redox coupleb abstraction' and 0-insertion

Multicomponent Bi3+(5d106s26p0) Mo6+(4d05s0) Fe2+/Fe3+ bismuth molybdate

TezMo07, Te4+(4d1 5s2 5p0) Mo6+(4d0 5s') ~ e 3 + / ~ e 4 + (Te,Ce ,Mo)O,

Fe,S b,O, Sb3+(d1 O5s25po) S6+(5S05PO) Fe2+/Fe3+

USb3010 u5 +(5S26d07S0) s bs +( 5s05pO) u5+/u6+

Footnotes: a) The lattice oxide ions bridging these cations and Mo6+ or Sb5+ are for allylic

b, For activation of gaseous oxygen. H abstraction.

catalyst activity is not high, and no isomerization of butene occurs. In contrast, the h 2 M d 6 catalyst that has MOO^^-),, layers is active in isomerization of butene, but not active in selective oxidation. These results lead to the conclusion that butene adsorption and allylic hydrogen abstraction take place on the Mo layers, while oxidation takes place using the lattice oxygen of the Bi layer [78].

The formation of carbon-oxygen bond (0 insertion) is a critical step in the production of oxygenates. The mechanism in Fig. 11-4 shows that the surface lattice oxygen is the active species in this step. On bismuth molybdates, it is now well established that at the higher operating temperatures of about 45OoC, lattice oxygen accounts for nearly 100% of the oxygen incorporated into acrolein and carbon dioxide [66,79-811. When a mixture of propene and 1 8 0 2 are passed over a catalyst containing l 6 0 , the acrolein and C02 produced initially contain practically only l 6 0 , which can only come from the lattice. Exchange between the lattice oxygen and gaseous oxygen or acrolein can be independently determined to be too slow to account fo the results under similar conditions. It is further shown that in some catalysts, all of the lattice oxygen can participate in the reaction due to rapid diffusion of lattice oxygen.


The extent of lattice versus adsorbed oxygen as being the proximate origin of oxygen for the C-O bond formation at a lower temperature of about 350" is not unequivocally established. One report shows that all acrolein is formed from lattice oxygen [66], whereas another one shows that the contribution from adsorbed oxygen increases with decreasing temperature [81].

That lnttice oxygen is the major proximate source of oxygen in the formation of acrolein has now been demonstrated for Sn-Sb-0, U-Sb-0 [82,83]. and Sn- Mo-0 [84], in addition to Bi-Mo-0. This property is not universal. In the case of cuprous oxide, acrolein is formed with adsorbed oxygen [85]. There is also IR evidence that adsorbed propene reacts with 0 2 - on ZnO to form an adsorbed acrolein [86,87].

The formation of a C-O bond using lattice oxygen implies reduction of the cations in the oxide, which must be reoxidized for a steady state operation. Indeed, exposure of bismuth molybdate to propene causes the reduction of Mo6+ to MoS+ [88], which can be detected by EPR. Coupled with the earlier discussion that the B i 4 group is involved in allylic hydrogen abstraction, it is possible that during catalysis, the oxygen neighboring bismuth is removed in water formatkn, the lattice oxygen at the molybdenum center is removed in the formation of acrolein. Both oxygen species are replenished by the diffusion of lattice oxygen, and eventually the lattice is reoxidized by gaseous oxygen.

According to this picture, there must be an inlet and an outlet in the catalyst for oxygen. One outlet is the oxygen species in the inolybdooxo group (Mo=O). The inlet is a pair of oxygen vacancies (I$), perhaps in an ensemble with Bi ions, and the reoxidation proceeds as [89-911:

(Bi++)202- + O2 - (~i3+02-)~02- (11-9)

In addition to isotopic studies, the involvement of the lattice oxygen can also be demonstrated by monitoring the vibrational spectra of the oxide. Molybdenum oxide, molybdates and vanadates possess characteristic bands in the 800- lo00 cm-' region that can be assigned to terminal M=O vibrations (see Chapter 2). The frequencies of these bands depend on the oxygen isotope, and possibly on the metal oxidation state of the cation. For example, the shift in this band has been interpreted as due to the transformation of Mo6+=0 to Mo5+=0 on adsorption of butadienes [92]. Broadening and shifts of bands in this region of a bismuth molybdate catalyst have been observed after it is used in the oxidation of propene with lSO2 [93,941. From the magnitudes of the shifts, it was concluded that practically all of the lattice oxygen in Bi2M006 (y phase) and Bi2M0209 (p phase) are involved in the reaction, while much less is involved in Bi2M03012 (a phase) [93]. This is probably due to the much slower diffusion of lattice oxygen in the a phase than in the other two phases.

These results support the scheme in Fig. 11-4 that M H groups are involved in the reaction. That the presence of M H groups is important for high selectivity on molybdates has been pointed out by Trifib. It has been observed that the molybdates of Bi, Fe, Co and Mn, which show strong infrared absorption in the 920-970 cm-' region, are selective oxidation catalysts, while molybdates of

192

Ca, Pb, and Tl, which do not show strong absorption in that region (implying no M d bonds) are nonselective catalysts [96,97]. In addition, the higher activity of bismuth molybdate than Fe, Co or Mn molybdates is correlated with its slightly weaker Mo=O bond, as indicated by a lower vibrational frequency. However, the presence of M d groups is not a sufficient condition. Na and K molybdates have Mo=O groups, but are nonselective catalysts [981.

The selectivity of bismuth molybdate for acrolein depends little on the partial pressure of gaseous oxygen. The selectivity on Co, Fe(I1) and Mn molybdate, however, decreases with increasing oxygen partial pressure. This has been correlated with the octahedral coordination of Mo in Bi-Mo-0, and tetrahedral coordination in the other molybdates [99]. No direct correlation between the amount of Mo=O and the activity or selectivity of the catalyst has been reported. It has also been suggested that molybdates that contain isolated or comer-sharing Moo6 octahedra as in Mo17047 and BizMO06 [loO,lOl] are more active or selective catalysts than those with edge-sharing octahedra as in Bi2M03012. or those with layer structures as in Moo3 [1021.

In the scheme in Fig. 11-4, the active site is shown as a cluster of bismuth ion and molybdenum dioxo (I) or diimido (IV) species. Experimental results correlating X P S intensities of bismuth and molybdenum in Bi-Mo-0 with catalytic activity support the view that the active sites consist of bismuth-molybdenum pairs [103,104]. There is no direct evidence to support the presence of dioxo species. The involvement of diimido species is supported by data in ammoxidation that the acrylonitrile/acrolein ratio increases linearly with the [NH3]2/[propene] ratio [42]. which fits the model in which substitution of two Mo=O species by two MwNH occurs at the active site.

Other Molybhtes Cobalt molybdate CoMd4 is also a selective oxidation catalyst for the

production of acrolein. It has two polymorphic modifications: a high temperature (>4O0-45O0C) p and a low temperature a form [105]. Excess M a 3 is required for high selectivity for acrolein in the Co-Mo-0 system. It has been reported that excess Moo3 stabilizes the a modification [ 1061 whereas excess Co304 stabilizes the p form [107]. Although the matter has not been discussed, it is also possible that excess Moo3 is the active site for the reaction, as is observed in the Ni-Mo oxide system [108]. It has also been reported that with Ni-Mo oxides, only samples containing more Moo3 than needed for NiMo04 produce acrylic acid in propene oxidation [110].

When excess Moo3 is present in Co-Mo-0, acrolein can be produced from propene in the absence of gaseous oxygen [lo91 until an equivalent of one monolayer of lattice oxygen is consumed, The acrylic acid yield, however, decreases rapidly with increasing degrees of catalyst reduction, and the rate of reaction is slower in the absence than in the presence of gaseous oxygen.

The dependence of the activity and selectivity on the surface crystallographic orientation of Moo3 has been studied recently. Volta, et al. have studied the reaction over a series of Mo03/graphite catalysts prepared with different calcination times and temperatures [ 11 1,1121. These catalysts have different


distributions of exposed crystal faces. A linear correlation is observed between the yield ratio of acroleiqK02 and the ratio of (100)/(010) surface areas. It is concluded that acrolein is produced on the (100) face of Moo3, and C02 is produced on the (010) face. Other explanations of these data have been advanced [ 113,1141. In one report, it was concluded that acrolein was produced on the (010) face, and combustion, on the (100) and the (101) faces. This latter assignment is supported by the observation that the yield of acrolein in the oxidation of ally1 halides is proportional to the surface area of the (010) face of Moo3 [114]. Another example of crystal-face specificity has been reported for the multicomponent Mnl -x4xV2-2xM02x0,j catalyst [ 1151.

Effect of Modifiers to Molybdates Various modifiers have been used to further improve the activity and

selectivity of bismuth molybdates. These have led to the development of the multicomponent catalysts. However, detailed understanding of the chemical effects of the additives is lacking. In many cases, only observations of the effect are reported.

When Ce is substituted for some Bi in Bi2(M00~)~, high catalytic activities in propene ammoxidation are observed at Bi/Ce ratios that correspond to the maximum solubility of Bi in Ce(Mo04) and of Ce in Bi2(Mo04),. A high activity is also observed at a Ce concentration that correspnods to the point of equal solubility of Ce in Bi molybdate and Bi in Ce molybdate [116,117]. Substitution of Bi by La is less effective than by Ce [95]. This is interpreted as due to the fact that Ce undergoes redox (e- + Ce4+ + Ce3+) more readily than La or Bi, and more effectively facilitates the reoxidation and reconstruction of the Bi-Mo-0 phase.

Thcre are reports that addition of Fe203, CrzO3, CuO, Te02 or Se02 facilitates the formation of the a and p phases of bismuth molybdates during the preparation of the catalysts [118]. Addition of BiP04, F%MqO12, or Cr2M03012 to Bi2M03012 also increases the formation of the a phase [119].

The current industrial multicomponent catalyst, Me,F%B&M%PiK,O, where Me = Ni, Co, or Mg is essentially a Me-Bi-Fe molybdate promoted by P and K. In some recent publications, it is suggested that the primary components of this catalyst are femc molybdate and Bi3(Fe04)(M00~)~ [ 120,1211.

Alkali metals are promoters in the supported Moo3 system. At low alkali metal concentrations, the promoting effect decreases as Cs> Rb > K > Na > Li [92]. This trend is inversely related to the electronegativity of the alkali metal.

Addition of V [I221 or Sn [123] to Moo3 does not produce a selective catalyst for acrolein in propene oxidation. On the other hand, Te02-Mo03 is very active and selective both for propene oxidation to acrolein [124,125] and for ammoxidation [126]. Ce-Mo-Te oxide is also an active and selective ammoxidation catalyst. A catalyst prepared by coprecipitation of all the components consists of essentially a ternary (Ce,Mo,Te) oxide, a-Ce2Mo4OI5 and/or p- Ce2M03013 [127]. Ni-Mo-0 and Cu-Mo-0 have also been reported to have reasonable selectivities [128,129,130].

In the Co-Mo-0 system, a small amount of Te has been found to be an

194

effective promoter in the oxidation of both propene [128] and butene [131]. Te is found to be enriched at the surface [105]. Addition of Fe, Bi and V also increases the activity and selectivity in acrolein production [105]. In general the Co-Mo-0 system produces more acrylic acid than the Bi-Mo-0 system.

11.6 EFFECT OF WATER ON PROPENE OXIDATION

When water is added to a feed of propene and oxygen, 2-propanol is often formed either as the major product or as the initial product which is then dehydrogenated to acetone.

On a Mo03/Al203 catalyst [132], acetone is produced with 83% selectivity at 300°C. This contrasts sharply the situation for acrolein production where the proximate source of the oxygen atom in acrolein is a lattice oxide ion. If D20 is used, the deuterium atoms rapidly exchange with the hydrogen atoms of the end carbons of the reactant propene to form (DIH)3C-CH=C(D,H)2. The hydrogen at the center carbon does not undergo exchange. When deuterated propene is used, no kinetic isotope effect is observed if the deuterium atoms are at the end carbons, as would be expected if the exchange of these atoms with hydrogen atoms in water is rapid. A kH/kD ratio of 2.2 is observed if the deuterium atom is at the center carbon. This behavior of the kinetic isotope effect is opposite that observed in acrolein production.

It has been proposed that the reaction proceeds via a carbenium ion intermediate [132]:

If H21S0 is used, the l 8 0 atom is incorporated into acetone.

MoS+ + H20 -+ MoS+OH- + H+(ad) (11-10)

H+(ad) + C3H, + C3H7+(ad)

C3H7+(ad) + Mo5+OH- + Mo5+ + C3H70H

C3H70H + 1/202 + (CH3)ZCO + H2O

In this mechanism, a surface coordinatively unsaturated MoS+ is assumed to be the active site. However, there is no direct evidence for the presence of MoS+ in the catalyst.

Whether desorbed 2-propanol is a significant product is questionable at temperatures higher than about 30O0C when the equilibrium concentration of 2- propanol is small. Under these conditions, it is possible that acetone is formed from a surface alkoxide [ 1341:

H 3 q H p - h b H -H 0 I

H 0

C& + J F=== ? -+ (CH&CO + J (11-11) *


The mechanism in eq. (1 1-10) explains the data over a Sn-Mo oxide catalyst at 150°C [133]. On this catalyst, 2-propanol has been shown to be the only product at low conversions. Acetone begins to appear as the conversion increases. The reaction rate at high conversions is suppressed by methyl ethyl ketone.

In the presence of water, propene can also be oxidized to acetone on a VzOs catalyst at 200°C [135]. As the temperature increases, the acetone yield decreases, whereas the yields for acetic acid and C02 increase. Over a V-P oxide (VP about unity), selective production of acrylic acid is observed. Addition of Te suppressess the selectivity for combustion such that at 80% conversion, a combined selectivity to acrylic acid, acetic acid, and acrolein in excess of 60% is obtained. Decreasing the water partial pressure in the feed decreases the selectivity for acrylic acid and acetic acid, but increases that for acrolein [135].

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125. T. V. Andrushkevich. G.K. Boreskov. L.L. Kuznetsova, L.M. Plyasova, Y. N.

126. J. C. J. Bart, and N. Giordano, J. Cafal., 64, 356 (1980). 127. J. C. J. Bart, N. Giordano, Ind. Eng. Chem. Prod. Res. Dev., 23, 56 (1984). 128, Ph. Jaeger, and J.E. Germain. Bull. SOC. Chim. France, 11-12, I(1982) 407. 129. J. Haber. Kinet. Carol.. 21, 100 (1980). 130. C. Mazzochia, P. Centola, R. Del Crosso, G. Terzaghi, and I. Pasguon, Chim. I d . ,

131. P. Forzatti, P.L. Villa, D. Ercoli. G. Gasparini, and F. Trifiro,

132. N. Giordano. A. Vaghi, J. Bart, and A. Castellan. J. Catal., 38, 11 (1975). 133. Y. Takita, Y. Moro-oka, and A. Ozaki, J. Catal., 52, 95 (1978). 134. J. Buiten. J. Catal., 10, 188 (1969); 13, 373 (1969); 27, 232 (1972). 135. M. Ai. J . Catal., 101, 473 (1986).

73, 651 (1976).

Tyurin, and Y. M. Shchekochikhin, Kinet. Katal., 15,424 (1974).

55, 687 (1973).

Id. Eng. Chem. Prod. Res. and Devel., 16, 26 (1977).

Chapter 12

SELECTIVE OXIDATION CATALYSIS I1

12.1 SELECTIVE OXIDATION OF BUTENES

The selective oxidative dehydrogenation of butenes to butadiene (eq. 12-l), and the selective oxidation of butenes to maleic anhydride (eq. 12-2) are both industrially important reactions:

C4Hg + 3 0 0 - 4 0 0 0 c > c4H, + H20 (12-1)

C4Hg + 3 0 2 4 0 0 - 5 0 0 0 c ) C4H203 + 3H2O (1 2-2)

Other important reactions of butenes include the production of acetaldehyde and acetic acid catalyzed by promoted vanadium oxides supported on rutile Ti02 [l] and other vanadates, and the oxidation of 2-methylpropene (isobutene) to methacrolein (2-methylprop-l-en-3-al) or ammoxidation to methacry lonitrile (2- cyanopropene) catalyzed by molybdates and cuprous oxide:

(12-3) 0 2 C4Hg =@ CH3COOH, CH3CHO

( 12-4a)

(12-4b)

Selective oxidative dehydrogenation (eq. 12-1) can be carried out on many catalysts including molybdates. vanadates. ferrites, uranium-antimony oxide, tin-

200

SELECTIVE OXIDATION II a01

antimony oxide and other antimonates. Selectivities for butadiene in excess of 85% have been obtained using these catalysts. Maleic anhydride can be produced selectively on catalysts based on vanadium oxide and/or molybdenum oxide, especially those promoted by phosphorous. A selectivity for maleic anhydride of higher than 70% has been reported in a number of patents [2].

Reaction Mechanism i) Oxidative Dehydrogenation:

water involves stepwise abstraction of two hydrogen atoms: It is likely that the oxidative dehydrogenation of butene to butadiene and

(12-5)

This mechanism has been proposed by analogy to propene oxidation in which dissociative adsorption of propene to form adsorbed x-ally1 is well established (see Chapter 11). Adsorbed x-bonded butene (C4H8(ad)) and x-ally1 (C4H7(ad)) on ferrites have been observed using infrared spectroscopy at room temperature [3,4]. It has also been shown that on ferrites, adsorbed butadiene is produced below 200°C. and the desorption of butadiene is rate limiting [5,6]. However, data at ordinary catalytic temperatures of about 350°C are not available.

Butene isomerization accompanies dehydrogenation, but there are few studies on this reaction during oxidation. Isomerization by reversing the allylic hydrogen abstraction step (step I1 in eq. 12-5) is a possibility. On MgF%04 [8] and CoF%04 [9], it has been proposed that isomerization and dehydrogenation take place on separate sites [7] based on the fact that when a mixture of trans-2-C4D8 and trans-2-C4H8 is oxidized over the catalysts, there is no H/D isotope mixing found in butadiene. Such mixing is expected if step I1 is readily reversible on the oxidation sites. Isomerization also shows a smaller deuterium isotope effect than oxidation [8,9], but its implication on the mechanisms has not been explored.

ii) Oxygenate Formation:

postulated to be [lo]: The primary route for the oxidation of butene to maleic anhydride has been

(12-6)

(ad or g)

\ /-nd

202

Other possible intermediates that have been suggested include 45 - dihydrofuran and 2.5-dihydrofuran [11,121. Methyl vinyl ketone, acetic acid, acetaldehyde, and carbon oxides are among the common side products [13]. This mechanism is supported by the following results using V-P oxides (V/P 1. 1/1) as catalysts. When used as the reactant feed under similar reaction conditions. butene, butadiene, crotonaldehyde, and furan all yield substantial amounts of maleic anhydride [13,14]. The selectivity for maleic anhydrik increases in the sequence: butene c butadiene c furan. In the oxidation of 1-butene, substantial conversion begins at -220°C. The major product below 250°C is butadiene which becomes negligible at 280°C. As the temperature increases, the yield of maleic anhydride increases and reaches a maximum at 300-330°C. This supports butadiene as both an intermediate and a precursor for maleic anhydride. The CO and C02 production increases with increasing temperature continuously. Small amounts of furan, acetaldehyde, methyl vinyl ketone, and crotonaldehyde are also observed between 250-330°C [141.

The product distribution on the V-P oxide also depends on the oxygen partial pressure. At low butene/02 ratios or at low conversions, maleic anhydride and COX are the only detectable products. At high 0 2 conversions, other intermediate oxidation products appear, including butadiene [141. In the absence of gaseous oxygen, butadiene is the major product until the catalyst is too extensively reduced to be active [15,161.

Selective production of mdeic anhydride from butene, butadiene, and furan (with selectivity increasing in this order) has also been reported for NiMo04 containing 15% excess Moo3 [17,18]. The reaction sequence (12-6) probably applies to this catalyst.

Kinetics i) Oxidative dehydrogenation:

On most catalysts, the oxidative dehydrogenation of butene to butadiene shows a positive order in butene and zero order in oxygen. In addition, a negative order in butadiene, indicative of product inhibition, is often reported, especially at lower temperatures. The order in butene depends both on the temperature and the pressure. It usually decreases with increasing butene pressure or decreasing temperature. For example, on bismuth molybdate, the rate is first order in butene over the temperature range 343-500°C [19]. The rate depends on the butene isomer. At 460°C. the rate for 1-butene is several times greater than for cis-2- butene, which is greater than for trans-2-butene [19]. Below 4OO0C, butadiene inhibits the reaction. In contrast, water or carbon dioxide has little effect on the reaction. The same kinetics is observed on Fe-promoted Bi-Mo oxide [20].

Unlike bismuth molybdates or ferrites, the reaction on the scheelite Pbl -3xBi2x$x(Mo04) system shows a zeroth order in butene and a positive order in oxygen [211.

The formation of butadiene shows a deuterium isotope effect. The ratio of rate constants kH/?.D is 3.9 at 300°C and 2.6 at 400°C on MgFe204 (8), and 2.4 at 430°C for CoFe204 [9]. The large isotope effects indicate that the breaking of C-H (C-D) bonds is involved in the slow step of the reaction. This conclusion is

SELECTIVE OXIDATION LI 203

supported by the relative reaction rates of various substituted butenes [22]. On a- Fe203, the relative rates at 270°C are 1.3 : 1.0 : 0.9 : 0.7 for

F 7 C I

C=C-C<<, C=G<<, C=C<<, and C-C=C<, respectively. These relative rates parallel the expected relative strengths of allylic C-H bonds in these compounds, suggesting that breaking of these bonds to form allylic species is involved in the slow reaction step. When the reactivities are normalized to per hydrogen atom at the allylic position, they are found to be 1.4, and 12 for primary, secondary and tertiary allylic C-Hs.

The activation energy for the reaction of butene on the scheelite Pbl-3xBi2x@n(Mo04) system [21] is 147 kJ/mole, which is close to the 139 kJ/mole for propene oxidation. This value does not depend on the concentration of cation vacancies, although the selectivity and the rate both increase with increasing values of x. The rate of isomerization increases with increasing x, and the ratio of isomcrization to oxidation is constant for a given temperature.

The values of the activation energy depend strongly on the catalyst and the reaction conditions. For a- and y-Fez03, the values are about 105 and 63 kJ mol-', respectively, at about 300°C [22]. For Bi3(Fe04)(Mo04),, the value decreases from 209 kJ mol-l below 400°C to 75 kJ mo1-l above 400°C due to a change in the rate-limiting step from butadiene desorption at low temperature to surface reaction at high temperature [23]. A similar decrease from 160 to 40 kJ mole-' has also been reported on Fe-promoted bismuth molybdates, and on defect scheelite systems such as Pbo.s~Ce+os$o.04 Moo4 and [email protected] M00.2~04

for which the activation energy decreases from about 147 to about 63 kJ/mole above 400°C [21].

The reported activation energy for bismuth molybdate varies widely. It ranges from 109 [19] to 168 kJ/mole [24] at about 400°C. The value also decreases from high to low temperature, in contrast to the other oxides mentioned earlier. In one report [21], it decreases from 159 to 59 kJ/mole. One interpretation assumes that the rate-limiting step at low temperatures is the desorption of butadiene, and the rate-limiting step at high temperatures is allylic hydrogen abstraction. This interpretation is consistent with the results from butene and butadiene on bismuth molybdate [25]. When compared under identical conditions, the activation energies for 1-butene oxidation are the same for a, p and y-bismuth molybdates (see Table 11-4).

ii) Oxygenate Formation: The kinetics of the oxidation of butene to maleic anhydride on V-P-0 has

been reported [26,27]. At about 380°C and for a butene pressure of about 1 P a in air, the rate of disappearance of butene follows the expression [27]:

k€l kO 'H '0

k, '0 -t % kH 'H rate = (12-7)

Where kH and are constants, PH and PO are the partial pressures of butene and oxygen, and % is the average number of oxygen molecules required per hydrocarbon molecule oxidized. This rate expression suggests inhibition by reactants at high butene partial pressures. This has been observed [30]. Below 1% concentration of butene, the reaction rate follows a first order kinetics [13,27]. 1-Butene and 2-butene show similar behavior. The activation energy is around 84 kJ/mole [27]. Isomerization of butenes accompanies oxidation [28a].

Nature of Catalysts i) Catalysts for Oxidative Dehydrogenation:

Bismuth molybdates, which are selective for the oxidation of propene to acrolein, are also selective for oxidative dehydrogenation of butenes. These oxides have been discussed in detial in Chapter 11, and will not be repeated here. However, it is interesting and important to point out two differences between oxidative dehydrogenation of butene and oxidation of propene to acrolein on these catalysts. The first difference is that the lattice oxygen involved in the formation of water during butene oxidation is different from that incorporated into propene to form acrolein [28] (see section 11-3). The second difference is that, although acrolein is produced with high yield and selectivity from propene on bismuth molybdate, selective dehydrogenation is the dominant reaction for butene. Little furan, maleic anhydride, or oxygenates are produced.

On unsupported [29] and graphite-supported Moo3 catalysts [31], oxidative dehydrogenation of butene has been shown to be crystal-face specific. A linear correlation is observed between the ratio of the surface areas of the (100)/(010) faces and the ratio of butadiene to C02 production in the oxidation of 1-butene. It has been established that butadiene is mostly formed on the (100) plane, and COX on the basal (010) plane.

On many selective oxidative dehydrogenation catalysts, the reaction proceeds with the same initial activity and selectivity in the absence and presence of gaseous oxygen (see Table 11-2). This implies that the formation of water involves lattice oxygen, although there has been no study to determine the proximate origin of the oxygen atoms in the product water under steady state conditions in the presence of gaseous osxygen. It has been reported that the selectivity for butadiene is higher in the absence than in the presence of gaseous oxygen [541.

It has been observed that fenites of the spinel structure are more selective than those of the corundum structure, especially at higher temperatures (near 400°C) [7,32]. The difference is particularly evident when comparing a-Fe203 (corundum) with y-Fe203 (spinel) [22]. In one report, the much higher selectivity for butadiene on y-Fe203 has been explained by two factors: adsorbed n-ally1 or butadiene on y-Fe203 is much less susceptible to degradation by gaseous oxygen, and the activation energy for desorption of butadiene is lower on y-Fe203 [32]. This latter effect might be due to the fact that surface coordinatively unsaturated cations derived from the bulk tetrahedral Fe3+ ions which are present only in the spinel structure bind butadiene less strongly than those derived from the bulk octahedral Fe3+ ions. Indeed, a rough correlation has been reported between the

Fe-Te oxide does not fit this generalization.

SELECTIVE OXIDATION II

fraction of bulk tetrahedral Fe ions and the production rate of butadiene on the Cr substituted MgFe204 [33]. Another explanation has been advanced that the diffusion of lattice cation is faster in a spinel than in a corundum because of the presence of rows of unoccupied octahedral holes in a spinel. The rapid diffusion allows rapid reduction and reoxidation of the oxide during reaction [9].

When a femte catalyst is reduced, such as from Fez03 to Fe304, the oxide is no longer active for dehydrogenation, although it is still active for cis-trans isomerization and double bond migration. This has led to the suggestion that Fez+ is responsible for isomerization [34,35]. Reoxidation of Fe304 at low temperatures (< 300°C) produces y-F%03 which is a highly active and selective catalyst for oxidative dehydrogenation of butenes [34]. Probably because of this conversion, prereduction of iron oxide also increases significantly the selectivity in the dehydrogenation of ethylbenzene to styrene [36].

In addition to the crystal structure, the selectivity also depends on the crystallite size. On a series of a-Fe203 with a crystallite size ranging from about 2.5 nm to 60 nm, the selectivity for butadiene decreases from about 80% to about 55% [37]. The exact reason for this dependence is not fully understood. Mossbauer effect, X-ray, and magnetic characterization of the samples show only the presence of a-F%03. This rules out the possiblity that the effect is due to the presence of other bulk phases. One explanation for the effect is that the reaction is crystal-face specific, and different proportions of various crystal faces are present on crystallites of different sizes. Another explanation is that smaller crystallites supported on silica are more difficult to be reduced than larger crystallites [38,39] and the different redox properties result in different catalytic properties. It is interesting that on the same series of oxides, the activation energy in N20 decomposition is found to be lower on the sample that contain very small crystallites [40]. Since desorption of oxygen is the rate-limiting step in N20 decomposition, this observation may indicate different reactivites of adsorbed oxygen on different iron oxide samples.

The activities and selectivities of a series of iron-containing compounds have been correlated to their heats of adsorption of l-butene. On a-F%03, Fe304, Fe4Bi209, FeSb04, FeAs04 and FeP04 [41], it has been found that oxides that have high AH's of adsorption are nonselective, and combustion is the major reaction. Oxides that have low AH's are active for isomerization, and high selectivity for oxidative dehydrogenation are found for oxides of intermediate AH's (FeSbO, and FeAs04). For these two oxides, two forms of adsorption of 1-butene are observed: a fast and weak and a slow and strong adsorption.

V-Mg oxides have also been found to be selective for this reaction. It has been proposed that the active site is associatcd either with octahedrally coordinated bulk vanadium ions that are exposed on the surface [42] or with surface V043- complexes [43]. It has also been reported that the activity correlates wilh the intensity of an EPR signal assigned to V02+ [44]. Addition of Na' ions to a silica-supported vanadia catalyst increases the selectivity for butadiene substantially, but decreases the activity [45].

206

ii) Catalysts for Oxygenate Formation The V-P-0 catalyst has been studied extensively and the results have been

recently reviewed [2,46]. It is established that a selective catalyst must maintain a P/V ratio between 0.9 to 1.3. V20~/Si02 without phosphorous is slightly inferior to those with phosphorous in selectivity for maleic anhydride from butene, although it is totally nonselective for butane oxidation (unlike V-P-0) [47]. V2O5/riO2 shows high selectivities for acetic acid and acetaldehyde when used in the oxidation of butene in air and steam [l].

It has been proposed that in these catalysts, the ability of the V ion to undergo redox cycles is very important. When a catalyst is quenched from reaction conditions, V(1V) is found to be the major component, but V(II1) is also present when the catalyst is used under conditions selective for butadiene production (lower temperature, higher butene/02 ratio), while V(V) is present when the catalyst is selective for maleic anhydride production (higher temperature, lower butene/02 ratio). These results are shown in Fig. 12-1 [46]. The data on the valence state of vanadium were obtained by chemical analyses and they were confirmed by diffuse reflectance measurements of the catalysts. Based on these findings and the fact that the states of the catalysts are interconvertible, the scheme in Fig. 12-2 has been proposed. In this scheme, the V(II1)-V(1V) couple is involved in the oxidative dehydrogenation of butene to butadiene, and the V(1V)-V(V) couple is involved in the oxidation of adsorbed butadiene to maleic anhydride. Depending on the reaction conditions (temperature and butene/O2 ratio), the catalyst surface can be in one of these states: (i) a reduced state (V(II1) and V(1V)) in which the main product is butadiene; (ii) a state with contemporary presence of V(II1). V(IV), and V(V) in which the main product is maleic anhydride, and (iii) a strongly oxidized state (V(1V) and V(V)) in which the main products are carbon oxides. In addition to changing the product selectivity, the different reaction conditions and surface valence states also affect the catalyst life. For example, deactivation is much more rapid in higher butene/02 ratios and at higher temperatures [55].

The role of phosphorous has now been understood: it forms compounds with vanadium which facilitate the redox processes among the various oxidation states of vanadium, and help maintain the vanadium ions in the optimal distribution of valence states.

In various V-P oxide preparations, a number of phases have been identified, including a-VOPO,, P-VOPO,, (V0)[email protected], P-(V0)2P207, an amorphous V(1V) phase, a B-phase and more recently y-VOP04, and Y- (VO)~P~O~ [561. An interesting aspect of these phosphorous compounds is their ability to vary their degrees of condensation, which allows the solid to accommodate a change in the valence state of V [46]. In earlier studies, the resulting crystalline phase after calcination has been found to depend on the method of preparation. If the catalyst is prepared in an aqueous medium by reduction of V205, a mixture of a-VOP0, and P-(VO2)P20, (vanadyl pyrophosphate) is obtained [48-511. During butane oxidation, the P-vanadyl pyrophosphate phase is the predominant phase present at about 40O0C, and the a-VOPO, phase is the predominant phase at 420°C [14,521. If the catalyst is prepared in an organic medium by the reduction of V205 or

Selective V-P-0 catalysts contain P/V ratios near unity.

SELECTIVE OXIDATION II 207

v >

0 I

250 300 350

Temperature, C

Figure 12-1 Effect of reaction temperature on: a. The selectivity of main products (butadiene BD, maleic anhydride MA, and COX) in 1-butene oxidation; b. The valence state of vanadium. Catalyst V-P-0 with V/P = 0.99 containing 100% VOV) before use; feed 0.6 % 1-butene, 12 % 02. (Ind. Eng. Chem. Prod. Res. Devel., 24, 221 (1985). copyright American Chemical Society).

VOP04 with HCl or alcohol, it contains vanadyl pyrophosphate [15]. p-VOP04 is a slightly less selective catalyst than a-VOP04. Both p-VOP04 and (V02)P207 have been tested in the pure form to produce maleic anhydride, and are found to interconvert slowly during butene oxidation [53].

It has been proposed that the active catalyst consists of p-(VO)2P2O7 and a- VOP04 which can readily interconvert by reduction and oxidation under reaction conditions [52,57]. This interconversion is regarded as crucial in the reaction mechanism because it interconverts V(IV) and V(V). The active p-(VO)2P207 phase may be nonstoichiomeuic. At high temperatures under oxidizing

208

C4H

Figure 12-2 Catalytic cycle of butene oxidation on a V-P-0 catalyst.

conditions, a-VOp04 converts to p-VOP04, which cannot be easily reconverted to the pyrophosphate phase. This causes a loss of the redox ability of the catalyst and the low selectivity.

A more recent crystallographic study [56] of the V-P-0 system shows that vOHp04.#20 is the precursor prior to calcination. If VOHP04-#120 is obtained by reduction in an organic medium, a 6-VOPO, form is formed on calcination which successively converts to y and p-VOP04. Under nitrogen, y- VOP04 is reduced reversibly to y-(V0)2P207. Repeated cycles of oxidation- reduction slowly converts this y-VOP04/y-(VO)2P20, couple into the p-VOP04/p- (V0)2Pz07 couple.

Calcination of VOHP04.+ H20 obtained in an aqueous medium produces a- VOPo4 which exists in two forms: a1 and all. a1 is converted to p-(V0)2P20, by reduction at 76OoC, a11 is converted to y-(V0)2P207 when heated under nitrogen at 700°C.

Crystallographic considerations shows that the interconversion of y-VOP04 to y-(VO)2P207 requires very small displacements of atoms [56]. The reduction of p- VOW4 to form p-(VO)2P207 involves crystallographic shear planes, and would be more difficult. Thus, p-VOP04, once formed, makes the catalyst less selective.

p-(VO)2P207 contains double chains of edge-sharing VO6 octahedra linked by pyrophosphate (V2@4-) units. The axial V-O bond lengths within an octahedra are not equivalent. One of the V-O bonds can be viewed as a vanadyl ( V 4 ) group [59]. An IR band at 963 cm-' assigned to V=O has been observed on this compound. A V=O stretching band at 1001 cm-I has also been observed on p-VOP04 [60].


The presence of V-0 might be important. There is a correlation between the density of V - 0 groups and the activity in butene and butane oxidation on unsupported V2O5 catalysts [47]. The activity per V=Q depends on the method of preparation of V205. Interestingly, butene is oxidized to maleic anhydride with about 45% selectivity on the catalysts, but butane is oxidized only to carbon oxides.

In the NiMo04-Mo03 system, excess Moo3 is found to be essential for the selective production of maleic anhydride from butene at 480°C. Pure Moo3 or pure NiMoO, produces butadiene at about 20-25% selectivity but no maleic anhydride. NiMo04 containing excess Moo3, however, produces maleic anhydride with up to 35% selectivity [17,611. A similar synergistic effect is also observed for the selective oxidation of butadiene to maleic anhydride [18]. Spectroscopic investigations lead to the conclusion that the selective catalysts are Moo3 particles whose surfaces are decorated with NiMoO, particles that effectively block the combustion sites of Moo3. Furthermore, NiMo04, being effective in producing butadiene, also serves to supply butadiene to Mo03 which produces furan and maleic anhydride more selectively from butadiene than from butene.

The production of maleic anhydride on Moo3 is crystal-face specific [29,31]. A correlation has been found between the ratio of the (lOO)/(OlO) areas and the ratio of maleic anhydride to C02 production in the oxidation of 2-methylpropene.

Effect of Promoters There are claims that many promoters are effective, and summaries of these

are found in a number of review articles and monographs [2,7,61a,62]. Unfortunately, there is relatively little understanding of the promoting effects.

In the oxidative dehydrogenation of butenes on ferrites, the promoting effects of Mg, Cr, Co, Zn, Bi, Mo, P, Sb, and V have been studied. The addition of Cr, Mg, or Zn increases the resistance of the ferrites to reduction [35]. In addition, Co, Zn and Mg form spinel ferrites with Fe2O3, which is a more desirable structure than corundum for this reaction.

The promoting effect of Mo, V and P has been interpreted as due to changes in the acidity and basicity of the oxide [41,63]. This interpretation is based on the idea that butene and butadiene are both Lewis bases as indicated by their rather low ionization potentials. The Lewis acidity or basicity of a selective catalyst must not be too strong so that it is sufficiently acidic to adsorb butene, yet sufficiently basic to prevent strong adsorption of butadiene which would poison the active sites. Addition of Mo, V and P increases the acidity of iron oxide, which is presumably too basic. In the case of Sb and As, the promoting effect is a result of the formation of FeAs04 and FeSb04 [41].

The effect of the support in supported F%03 has been investigated [64]. Si02, Ti02, A1203 and their mixed oxides have been used as supports, and the activity per gram of catalyst varies widely over two orders of magnitude, but the selectivity for butadiene remains in the range 65-88%.

The V-P oxide catalyst deactivates by the loss of phosphorous during operation. This problem can be reduced by operating at a lower temperature or by

210

stabilizing the composition of the catalyst with the addition of alkali metals. Li has been claimed to be effective for this purpose [2]. The rate-limiting step in butene oxidation on this catalyst is believed to be C-H bond breaking. Thus promoters such as Fe, Cr, Mg, and Zn that have dehydrogenation activities have been reported to enhance the catalytic acitivity.

In a number of studies on vanadium oxide catalysts [65-681, it is reported that there is an optimal concentration of promoters beyond which the selectivity decreases rapidly. This may be due to possible changes in the crystalline phases and/or in the redox potential of the vanadium ions in the solid when too high concentrations of promoters are added. This interpretation is consistent with the fact that the effects of some promoters depend on the feed composition, especially the hydrocarbon to oxygen (or air) ratio and the conversion. Perhaps different gas phase compositions are needed to obtain the optimal distribution of oxidation states of vanadium in the presence of different promoters.

Alkali promotion has also been reported for vanadium oxide. In one report, sodium addition significantly increases the selectivity for butadiene on a V205/SiOz catalyst but decreases the selectivity to maleic anhydnde [69].

12.2 SELECTIVE OXIDATION OF BUTANE

In all studies reported, butane is less reactive than butene and higher reaction temperatures are required. It is reasonable to expect that the oxidation of butane proceeds via dehydrogenation to form butene, which is then reacted to form butadiene, maleic anhydnde, or other oxidation products as discussed in the last scction. It is important to note, however, that because butane is less reactive and requires a higher temperature for reaction, some catalysts such as ferrites that are highly selective for butene oxidation are not as selective for butane oxidation.

The most often studied reactions of butane are the oxidative dehydrogenation to butenes and butadiene (eq. 12-8) and oxidation to maleic anhydride (eq. 12-9):

Nature of Catalysts A number of oxides have been reported to be quite selective in oxidative

dehydrogenation. Among these are NiO, Moo3, and c0304 [70], and mixed oxides of Co-Mo, Ni-Mo, Mg-Ni [71], and Mg-V [72]. A selectivity for dehydrogenation in excess of 60% at a conversion of 20% or more has been obtained with some of these mixed oixdes.

Among the mixed oxides, Mg-Mo oxide and Mg-V oxide have been studied quite extensively. High selectivities (60+%) for dehydrogenation (butenes and


butadiene) on Mg-Mo oxide have been observed both in the presence [71] and in the absence of gaseous oxygen in the feed [73]. Analyses of the selective catalysts show that they are mixtures of MgMo04, MgO, and an unidentified phase.

The selectivity for oxidative dehydrogenation on a Mg-V oxide is as high as 80% at low conversions [74]. It decreases to about 50% when the conversion increases to 40% [72]. Detailed analyses of the catalyst show that the active phase is magnesium orthovanadate (Mg3(V04)2). This conclusion is confirmed by the similarity in the catalytic behaviors of a stoichiomeuic Mg orthovanadate and a mixed Mg-V oxide. Interestingly, neither Mg pyrovanadate (Mg2VzG) nor Mg metavanadate (MgV206) is selective for this reaction [74]. On the other hand, other orthovanadates including LuV04 and SmV04 are about as selective as Mg3(V04)2 [751.

The differences among the various vanadates have led to the suggestion that the bonding of oxygen in the oxide lattice is a factor that determines selectivity. The oxygen ions in an orthovanadate are bridged between a vanadium ion and another cation such as Mg, Lu, or Sm. If this other cation is not easily reducible, the bridging oxygen ion cannot be readily removed from the lattice, and the catalyst is selective. In a metavanadate or a pyrovanadate, there are oxygen ions bridging between two vanadaium ions. Since V5+ ions are easily reducible, the bridging oxygen ions are readily removable, and the catalysts show high combustion activities.

The selective oxidation of butane to maleic anhydride is quite demanding and only a few catalysts have been reported to have high selectivities. Molybdenum- phosphorous oxide and V-P-0 are among the few. Of these two, V-P-0 has been very extensively studied as described in section 12.1. Selectivities as high as 78% for maleic anhydnde at low conversions of butane and 65% at high conversions have been reported in the patents [2].

Although the active phases in V-P-0 in butane oxidation are believed to be the same as those in butene oxidation, some differences have been reported between the behavior of this catalyst in the two reactions. It has been found that although p-VOP04 transforms into (V0)2P207 during butene oxidation, the transformation is much slower during butane oxidation [53]. The areal activities of V-P-0 catalysts prepared either by aqueous reduction of V2O5 with H3P04 or by reduction with H3P04 in an organic medium are the same in butene oxidation, but are different in butane oxidation [76]. This observation may indicate different rate-limiting steps in the two reactions.

The participation of lattice oxygen in a V-P-0 catalyst in butane oxidation has becn demonstrated. It has been found that only the first one to five layers of lattice oxygen are involved [14,15]. This is very different from the behavior of bismuth molybdates in propene oxidation (see Chapter 11). It has been porposed that the active surface of V-P-0 is a layer of VOP04, and scrambling of lattice oxygen occurs rapidly, creating a pool of five exchangeable oxygen atoms per surface V ion [ 5 8 ] . When one of the five oxygen atoms is consumed in the reaction, the vacancy is filled by reoxidation by gaseous oxygen. The replenished oxygen atom scrambles rapidly with the remaining four oxygen atoms in the surface layer. This picture is consistent with the observation that in a pulse experiment, 90% of the

212

reaction products contain oxygen atoms already present in the catalyst, and 10% contains oxygen from gaseous oxygen.

Finally, if the catalyst is first treated with D20, nearly 2f3 of the product maleic anhydride contains two deuterium atoms in a pulse experiment. Thus the reaction proceeds with exchange of H atoms in the surface intermediate (such as a x-allyl) and surface OD groups.

Mechanism The mechanism and the kinetics of butane oxidation to maleic anhydride on a

V-P-0 catalyst have been studied. By studying the deuterium isotope effect at different positions of a butane molecule on the p-(VO)2P2@ catalyst, it has been demonstrated that the first step of the reaction is the production of adsorbed butyl species [58]. At 400°C. a sizable isotope effect, kH/kD of 2.1 to 2.2 is observed for C-H (C-D) bonds at the secondary carbons of a butane molecule, but there is almost no isotope effect at the two primary carbons. The k H / k D values of 2.1 to 2.2 are close to the theoretical maximum value of 2.27 for a homolytic dissociation of a C-H bond based on the difference in zero point energy of the C-H and C-D bonds, which is 4.4 kJ/mole. Thus the dissociation of a methylene C-H bond of butane to form an adsorbed sec-butyl species is the rate-limiting step in the reaction.

It is likely that the formation of an adsorbed butyl species is followed by a second C-H bond breaking which results in an adsorbed butene molecule. The butene molecule then reacts in the same manner as described in the last section. An in situ FTIR study identified maleic acid on the catalyst together with some unidentified alkenic species, in addition to butane, adsorbed maleic anhydride, and adsorbed carbon dioxide [1201. Little mechanistic information is provided by the kinetics [58,77-791.

There is no information on the subsequent steps.

12.3 SELECTIVE OXIDATION OF METHANE

The selective oxidation of methane is one of the most challenging processes because the C-H bond in methane is the strongest among all hydrocarbons, and because the selective oxidation products are generally more reactive than methane such that high yields of products are difficult to obtain.

This reaction has received considerable attention recently because of the abundant reserve of natural gas in the world, and two recent review articles have appeared on this subject [80,81]. In the review by Pitchai and Klier [81], different methods to oxidize methane are mentioned. These methods include partial oxidation to methanol or formaldehyde, oxidative dimerization to acetylene, ethene, and ethane, oxidative methylation, and others. The first two processes, oxidation of methane to methanol or formaldehyde and oxidative dimerization will be discussed here.

Formatwn of Methanol or Formaldehyde A large number of oxides can act as catalysts for the oxidation of methane to

methanol or formaldehyde using oxygen [80,81]. Oxides of zinc, lead, nickel,

SELECTIVE OXIDATION 11 213

chromium, copper, iron, silver, cobalt, molybdenum, manganese, vanadium as well as mixed oxides containing these have been claimed to be active and selective. In general, however, the conversions are low, thus the yields of methanol and formaldehyde are low. Among these oxide, vanadium oxide, and especially molybdenum oxide are the more thoroughly investigated.

The first extensive report on the use of Moo3 and mixed oxides of Moo3 with Fe2O3, ZnO, U02, or V 0 2 is in a patent by Dowden and Walker [82]. In a highly oxygen-lean feed of about 97% CH, and 3% oxygen, and at a low methane conversion of a few percent, a selectivity of methanol of up to 75% is achieved with Mo03-U02. 3MoO3-F~O3 shows a methanol selectivity of 6396, and a formaldehyde selectivity of 8%. A later patent claims CuO-Moo3 to be about as effective as the other molybdates [83].

Selective oxidation to methanol and formaldehyde at a much higher conversion is reported by Iwamoto [84] and by Lunsford and coworkers on silica- supported Moo3 catalysts using N20 as the oxidizing agent [85]. In the latter report, CH30H and formaldehyde are produced with 55.1% and 23.9% selectivity, respectively, at a conversion of 12.5%. The feed is 72 torr CH,, 277 torr N20 and 266 torr H20, the reaction temperature is 580°C and the catalyst is 1.7 weight % Mo03/SiOz. Water is found to be essential for the high selectivity to the formation of partial oxidation products. Higher partial pressures of water increase the selectivity. At low water partial pressures, the activity increases with increasing water partial pressure, but becomes constant for P H ~ ~ > 75 tom. Later work show that below 5OO0C, methanol is formed almost exclusively [86.87]. Above 5OO0C, HCHO selectivity continues to increase and HCHO becomes the major partial oxidation product above about 550°C. Up to this temperature, no CO or C02 is detected even at a methane conversion of 1.9%. At higher temperatures, CO formation becomes considerable. At 580°C. the selectivity to partial oxidation products increases with increasing PcH4/P~20 ratio in the feed, and reaches 100% when the ratio is 1.2. Although the conversion to methanol increases as PHz0 increases from 0 to 260 torr, the conversion to HCHO stops increasing for P H ~ ~ larger than 40 torr, and the conversion to CO decreases slightly. The activity of the catalyst also increases with PHZo up to 50 tom. beyond which there is no further effect. The steam reforming of methane and the water-gas shift reactions are negligible under these reaction conditions.

Selective production of formaldehyde is also observed using oxygen as the oxidant over a M a 3 catalyst supported on silica-alumina [81]. Formaldehyde, CO and C02 are the only products reported. No methanol is formed. The selectivity for HCHO increases with increasing CH4/02 ratio, reaching a value of 46% at 600°C and a CH,/02 ratio of 17. As with N20 as the oxidant, addition of water to the feed is beneficial. The selectivity for HCHO increases with the amount of water added, although the CH, conversion decreases. At 6OO0C, a HCHO selectivity of over 90% is achieved with water in the feed and a C&/O2 ratio of 9.

From these reports, a number of interesting differences is observed between using 0 2 and using N 2 0 as the oxidant. Methanol is not produced using O2 but is produced using N20. With increasing temperature, the selectivity for the partial

214

oxidation products increases using 0 2 , but decreases using N2O. The CH,J02 ratio required to obtain high selectivity to partial oxidation is much higher than the CH,/N20 ratio.

The need to have water in the reaction mixture for high selectivity could limit the long term stability of the molybdenum oxide catalysts. Water reacts with Moo3 to form a volatile molybdenum hydroxide species 1811:

This could lead to loss of molybdenum under reaction conditions. At 540°C. the production of CH30H in CH, oxidation with N 2 0 on Moo3/

Si02 is roughly first order in CH, and H20 and zeroth order in N 2 0 [87]. The production of HCHO is zeroth order with respect to all reactants. At 580°C. the overall conversion of CH, is zeroth order in CHq and first order in N20 [86]. For some reasons, the rate constant is lower for P N ~ O < 80 torr than for higher PNz0. Between 480 and 520°C the activation energy for methanol production is 170 kJ/mole. Above 520". subsequent oxidation of methanol becomes rapid, making it difficult to estimate the true activation energy. The activation energy for HCHO production is 344 kJ/mole from 480 to 540°C. and 168 kJ/mole from 540 to 590°C [87]. The activation energy for the disappearance of methane is 176 kJ/mole from

The following sequence of reactions has been proposed for HCHO production 550 to 594OC [86].

[ 861 :

MO' + N20

MoV'O- + CH, + MoV1 OH- + CH3

MoV'02- + C H ? + MoVOCH3-

MoVOCH3- + H 2 0

MoV'OH- + MoVOH- + MoV + MoV'02- + H2O

MoV'O- + N2

Mo'OH- + CH30H

and for the combustion reaction, the overall reaction is:

8MoV' + 4 0 2 - + C&

2MoV + N20 - 2MoV' + N2 + 02-

8 MoV + 2 H20 + C02

(12- 1 1)

(12-12)

(12- 13)

(12- 14)

(12-15)

(12-16)

(12-17)

Bands in the C-H stretching region assigned to methoxy (peaks at 2955,2928, and 2857 cm-') are observed by IR when a reduced MoV'/Si02 catalyst, reoxidized by N20 , is exposed to CH, at 25°C. The production of methyl radicals on the catalyst on exposing CH, to a reduced catalyst that is reoxidized by N20 has also bccn demonstrated by EPR [86].

The active site is postulated to consist of oxomolybdenum clusters such as the

SELECTIVE OXIDATION I1 215

following species (4 is an anion vacancy):

This species could be the one on the right hand side of eq. (12-15). Oxide ions of silica could make up some of the oxide ions in this species.

Selective oxidation of methane by N20 can also be carried out using a V205/Si02 catalyst [88,89]. In fact, with catalysts of comparable loadings and at comparable conversions, a supported V205 catalyst produces methanol and formaldehyde selectively at 460-500°C. a temperature lower than that required for Mo03/Si02 [881.

The selectivities for CH30H and HCHO decrease with increasing temperature. They also decrease with increasing contact time due to secondary reactions of these molecules to form carbon oxides. At 500°C. replacing N20 with O2 reduces the selectivities for partial oxidation products from nearly 100% to less than 20%, and combustion to carbon oxides becomes the dominant reaction. Addition of a small amount of O2 to N 2 0 enhances the conversion with some loss in selectivity.

At 500°C with a small amount of oxygen in the feed stream, the kinetic orders are found to be:

rate (CH30H) a [CHJ'/2[N20]1/2 [H20]1/2 (1 2- 18)

rate (CH20) a [CI-LJ1/2 [N20]1/3 (12-19)

rate (CO) a [CIl,]1/2 [N20]' [H20]-' ( 12-20)

Over the temperature range 460-5Oo0C, the activation energies for the formation of HCHO, CH30H, and CO are 92, 168, and 197 kJ/mole, respectively.

It should be mentioned that in addition to the direct observation using EPR of the presence of methyl radicals at low temperatures on these catalysts, the importance of this species in this reaction is inferred from the fact that the addition of dimethyl ether, which decomposes to produce CH3 and CH30 species, decreases the required reaction temperature substantially [80]. Similarly other hydrocarbons such as ethane can also be used as initiators [81].

Oxidative Dimerization Many oxides possess some catalytic activities for oxidative dimerization of

methane [90-931. Using a high CIl,/02 ratio and at low conversions, many rare earth oxides show good selectivities for C2 products (primarily ethane and ethene). Carbon selectivities for Cz better than 80% are obtained, and Sm2O3 and Dy203 show selectivities better than 90%. Exceptions to this are Ce02 and TbO,. Among the other oxides, the basic oxides Bi2O3 and CaO also show good selectivities. On the other hand, acidic oxides such as silica-alumina, hydrogen-Y zeolite, and

216

ZSM-5 zeolite are nonselective. Thus far, only a few transition metal oxides are selective for this reaction. Based on data reported, it appears that only a few oxides that can undergo oxidation-reduction readily are selective.

Addition of alkali metals sometimes improves the activity and selectivity of a catalyst. Addition of Li to MgO has been shown to be very effective. A C2 selectivity of 89% has been obtained at a CH, conversion of 5% [91]. Li-ZnO is also quite active and selective, and a C2 selectivity of over 65% has been obtained WI .

This oxidative dimerization reaction has been studied in detail on the Li- promoted MgO and L a 2 0 3 [91,93], and interesting similarities and differences are detected between the two systems. On Li-promoted MgO, the products are ethane, ethene, and C02. Only small amounts of CO, methanol, and formaldehyde are detected. For 02/CH, ratios less than unity, a lower ratio leads to a higher C2 selectivity but a lower methane conversion at 720°C. C21-& is the major C2 product at 620"C, but C2H4 becomes dominant at 720°C. The total C2 selectivity increases with increasing temperatures up to 7OO0C, beyond which it decreases. The apparent activation energy for the reaction is 230 kJ/mole over 560 to 660°C.

Methyl radicals are detected downstream from the catalyst bed of Li-MgO. Thus ethane is formed by the coupling of two methyl radicals in the gas phase, Ethene is then formed by the oxidative dehydrogenation of ethane. These steps are:

c2H6 + W a d ) + C2Hs(ad) + OH-(ad) (12-22)

.C2HS(ad) + 02-(s) --+ C2H50-(ad) + e-(lattice) (12-23)

CzHsO-(ad) 4 C2&(g) + OH-(ad) (12-24)

The rate-limiting step is likely the catalytic production of methyl radicals. C02 is formed from methane at low conversions by the oxidation of surface methoxy groups and gas phase peroxide CH3O2. species. This peroxide species is formed by the reaction of a gaseous methyl radical and an oxygen molecule, and its equilibrium concentration decreases with increasing temperature. This latter effect is a major reason why very high temperatures (700°C or higher) are required for substantial yields of C2 products.

On Li-MgO, the active site is postulated to be a Li'O-center that is produced when Li is incorporated into the MgO matrix. Methane reacts at this center to form methyl radicals:

Li+O- + CH, --+ Li'OH- + C H 3 ( 12-25)

The active site is regenerated by the desorption of water to produce Li+ @ (where @ denotes an anion vacancy), which then reacts with oxygen to form the active site. The presence of this Li'O- center is detected by EPR. In addition,


LiC03 is also present in the catalyst, and might contribute to the reaction. On La2O3, C2 products are also produced via the gas phase coupling of

methyl radicals. C02 is formed by the oxidation of these radicals. But at high methane conversions, combustion of C2& is extensive and accounts for the low selectivity to C2 products. Under oxygen-lean conditions, increasing temperature decreases the COX selectivity, and increases Cz& selectivity, while C2& selectivity remains roughly constant.

The active site on La203 is probably different from that on Li-MgO. In particular, no 0- species has been detected with EPR on La203. Instead, a signal that can be assigned to 02- is detected [93]. Experiments at temperatures below 200°C show that 02- is much less reactive toward alkanes than 0-. At the reaction temperature of about 70O0C, this species may become active, or it may acquire an electron from the lattice to become two 0- species which are active. Without more evidence, the active site has been tentatively assumed to be some form of sorbed oxygen, such as 02-.

Selective production of C2 products has also been reported in reactions (noncatalytic) of methane with oxides [95,96]. Si02-supported manganese oxide, In203, Ge02 , PbO, and Sb2O3 show selectivities higher than 70% when methane is passed over them in the absence of gaseous oxygen at about 700°C. The oxides are reduced in the process, and the conversion of methane drops rapidly as the reduction procecds. Reoxidation of the reduced oxides regenerates the initial activity. Ethane and ethene are the major products, and a few percent of other higher hydrocarbons including benzene and toluene are detected. If the reaction is carried out under catalytic conditions using a methane/oxygen mixture, the selectivities for C2 products are somewhat lower.

From the change in product distribution as a function of contact time, it is concluded that ethane is the primary product on supported manganese oxide. Ethcne, higher hydrocarbons, and COX are all secondary products. One pathway to form higher hydrocarbons is methylation of alkenes, which has been demonstrated to occur under reaction conditions by passing a mixture of methane and alkenes over the oxide. Evolution of methyl radicals into the gas phase has also been detected.

Silica-supported manganese oxide is among the best oxides reported for this reaction. Mn7SiOI2 has been detected as a component on this oxide. After use, MnSi03 is formed. It has also been found that silica is a better support than alumina for manganese oxide because of higher selectivity for C2 products. Addition of sodium phosphate further improves the performance.

Of all the process reported, this cyclic noncatalytic oxidation-reduction process appears to be the most promising for commercialization. The process has now been successfully demonstrated in a pilot-plant scale test by ARCO.

12.4 SELECTIVE OXIDATION OF METHANOL

The selective oxidation of methanol to formaldehyde can be carried out effectively on many catalysts based on vanadium oxide and molybdenum oxide.

218

As in other oxidation reactions, the activity and selectivity depend on the nature of the support and the presence of other components that form compounds with V or Mo oxides.

The commercial process is carried out at 200-300°C over an iron molybdate catalyst containing excess Moo3. For example, the commercial Harshaw catalysts is a 3:l mixture of Mo03 and Fe2(Mo04)3. The process runs at 400°C under methanol-lean conditions with methanol conversion of 95-99% and selectivities of

On a M a 3 or Fe2(Mo04), catalyst at 200-300°C, pulse reaction studies show that dimethyl ether, formaldehyde and water are the primary reaction products [98]. In a flow reactor, the selectivity to various products typically depends on the conversion in the manner shown in Fig. 12-3 [99]. At low conversions, formaldehyde is the main product, followed by dimethoxymethane and dimethyl ether. With increasing conversion, the selectivity for formaldehyde increases, that for dimethoxymethane drops rapidly, while that for dimethyl ether drops slowly. Methyl formate is formed as the major secondary product at intermediate conversions. Selectivity for CO is low except at high conversions, whereas C02 is never a major product. At higher temperatures and lower conversions, the selectivities for CO and HCHO both increase while those for all other products decrease.

Addition of a small amount of water increases the selectivity for HCHO, but suppresses the activity [99,10O]. In the absence of water, the reaction rate at low partial pressures shows a small positive order in oxygen, and an order near 0.5 in methanol [99]. In one paper [loll, the order in both components drops to zero above 180 torr, but an order of 0.5 has been reported at 318 torr in another paper [loo]. Over 162-251°C. the activation energy is 70 M/mole on F%(MoO~)~, 75-88 kJ/mole on M a 3 , and 79 Id/mole on the mixed Fe-Mo oxide [99,100].

Between 220 and 300°C over MOO3, Fe2(M0O4), and their mixture, the ratio of rate constants for CH30D and CH30H varies from 0.83 to O.%, while the ratio for CD30D and CH30H varies from 0.15 to 0.37 [102]. Thus deuteration of the methyl group shows a much larger kinetic isotope effect than the 0 - H group. This indicates that the breaking of the C-H bond is rate limiting. This conclusion is supported by temperature programmed desorption results which show that the desorption temperatures of formaldehyde, methanol and water are shifted to slightly higher temperatures when CD30H is adsorbed instead of CH30H [l03]. The slower rate for CH30D than CH30H is attributed to the slower dchydroxylation rate of surface OD groups than OH groups, which results in a slower rate of regeneration of active sites. CH30H and CH30D give similar products, but CD30D gives much more dimethyl ether and less formaldehyde than CH30H, consistent with the expectation that the formation of formaldehyde involves C-H bond breaking as a rate-limiting step.

It has been shown by Mossbauer spectroscopy that the Fe ions in FQ(MoO~)~ and Fe2(M004)3.M003 can be reduced by methanol at catalytic temperatures [104]. This suggests the possibility that the catalyst undergoes oxidation and reduction during reaction and that lattice oxide ions are involved. Using a feed that contains *02, no 80-labeled water is found initially. The amount of H2'*0

91-94% [97].

SELECTIVE OXIDATION 11

'-O I 219

Fraction of Methanol Converted

Figure 12-3 Product distribution in methanol oxidation on a Fe-Mo-0 catalyst as a function of methanol conversion. Products are CO, methylformate MF, dimethyl ether DME, dimethoxymethane DMM, and formaldehyde F. (From Proc. 4th Intern. Cong. Chem. Uses Molydenum, p.411 (1982), copyright Climax Molydenum).

increases with time, and l80 is detected in the catalyst after use. These results support the possibility that lattice oxygen participates in the reaction, although oxygen-exchange between water and lattice oxygen would lead to the same observations.)

A proposed reaction mechanism on Mo03 is shown in Fig. 12-4 [105]. In this mechanism, methanol is adsorbed dissociatively on a dioxomolybdenum species. Abstraction of a hydrogen atom from the methyl group of an adsorbed methoxy is accomplished by an adjacent dioxomolybdenum species. Desorption of formaldehyde and water, and reoxidation by gaseous oxygen complete the cycle. The product water is formed initially with the oxygen atom from the surface.

The reaction of adsorbed methanol has been studied with IR spectroscopy. One report claims that dissociatively adsorbed methanol is observed at 100 K on Moo3 and ferric molybdate [106]. Another claims that methanol initially chemisorbs on M a 3 molecularly and, with increasing coverage, reacts in a bimolecular process to form chemisorbed methoxy and water [103]. Whether the catalyst is reduced is the process has not been addressed. Methoxy is also formed by exposing an oxidized M a 3 (pretreated with 0 2 at 250°C) to methanol at 100°C [ 106,1071. Temperature programmed desorption of adsorbed CD30H yields CD30D as one of the products [103]. This species must be formed by the recombination of a surface methoxy and a deuterium atom from surface OD group that is formed when a C-D bond is dissociated.

N 2 0 can also be used as the oxidant instead of 0 2 on Md3/Si02 . However, the resulting activity and selectivity for formaldehyde are lower. In fact, the selectivity decreases rather sharply with increase in temperature beyond 200°C

220

o+Mo:'OH 4 %

Figure 12-4 Mechanism for methanol oxidation to formaldehyde on MOO,. (From J. Catal.. 92. 127 (1985), copyright Academic Press).

[loo]. Inspite of these differences, the orders of the reaction in O2 or N20 are the same. These observations have led to the conclusion that adsorbed 0- species is not involved in the selective oxidation reaction.

Nature of Catalysts The commercial Fe-Mo oxide catalyst consists of Fe2(M00~)~ and Moo3.

The M a 3 is required to maintain high activity and selectivity. It has been claimed that a catalyst with a Fe/Mo ratio of about 1.7 shows the maximum activity [108,109]. It has been suggested in the past that the activity is due to the Fe2(M004)3 phase and Moo3 serves as a structural promoter to increase the surface area and as a source to replenish the Mo lost during operation [104,110,111]. Recent papers report little difference in the kinetic order dependence on methanol and oxygen, the areal activities, the product distributions, and the activation energy among samples of pure Moo3. pure F%(M0O4),, and M o O ~ / F ~ ~ ( M O O ~ ) ~ of comparable surface areas [99,104]. Thus, it is possible that the excess MOO, is also an active phase.

FeMo04, F ~ P M O ~ ~ O ~ ~ and Fe4(SiMo12040)3 show activities and selectivities comparable to Moo3, Fe2(Mo04)3, and their mixtures. Bismuth molybdates and aluminum molybdates show good selectivities but lower activities. Other molybdates such as chromium molybdate, cobalt-nickel molybdate, Co-P-Mo-0 and Ru-P-Mo-0 show poorer

Other molybdate catalysts have been tried.

SELECTIVE OXIDATION II 22 1

selectivities [991. The initial oxidation state of Mo is not very important. It has been found that

at steady state, M a 2 , Mo205, and Moo3 are all active and highly selective for formaldehyde. However, Mo2O5 appears to be somewhat more active. X P S investigations of catalysts used in a pulse reaction study show a rough correlation between the amount of formaldehyde formed and the intensity of the Mov signal. Since the catalyst undergoes a redox cycle during reaction, it is suggested that Mo(IV), Mo(V), and Mo(V1) are involved in the reaction cycle [loo], although the mechanism described in Fig. 124 uses only Mo(V) and Mo(V1).

The reaction on Moo3 is crystal-face specific. Under ultra high vacuum conditions, an ordered (010) surface of Moo3 does not adsorb methanol. Methanol is adsorbed much more readily after this surface is reduced and disordered by ion bombardment [ 1031. Presumably, ion bombardment generates a higher density of surface Mo ions that are extensively coordinatively unsaturated, reduced Mo ions, or other high energy defect structures.

Such structure sensitivity has also been demonstrated on crystalline samples of Moo3 [112,113]. The selectivity for oxidation to formaldehyde or methylal versus dehydration to dimethyl ether depends on the ratio of the crystal faces exposed:

+CH30H - 2 v CH20 CH2(mH3)2

CH30H (12-26) +CH+ ( c H ~ ) ~ o -H 2 0

Fig. 12-5 shows this dependence. The selectivity for formaldehyde versus methylal varies linearly with the ratio of the basal (010) plane to the side (100) plane, and the selectivity for ether versus oxidation to formaldehyde is proportional to the ratio of the apical (001) and (101) planes to the basal plane. These results are opposite to what would be expected from results mentioned earlier that an ordered (010) surface does not adsorb methanol. This is probably due to the fact that in these catalytic studies, the (010) faces of the crystals used have high densities of defects such as steps, which provide active sites for the reaction. The very different pressures of methanol used in these two types of studies may also contribute. Structure sensitivity is also observed in ethanol oxidation [112,114] and in the reaction of methanol in the absence of oxygen [106,112]. In the latter case, the dependence of activity on the crystal faces is somewhat different from that in the presence of oxygen.

The results of these studies suggest that the basal (010) plane of Moo3 is very selective for dehydrogenation, and it is much more active in the presence than in the absence of oxygen. The active site is probably associated with the terminal Mo=O groups. In the absence of 0 2 , the C-H bond breaking occurs on this surface, but water is not formed. The hydrogen accumulates on the surface to form bronze H , M d 3 . The apical (001) and (101) or (701) surfaces are active in dehydrogenation and dehydration both in the presence and in the absence of

222

100)

%(001) and (101)/%(010)

Figure 12-5 Dependence of selectivity in methanol oxidation on Moo3 catalyst. a. Ratios of formaldehyde and methylal selectivities versus area ratios of basal (010) to side (100) faces of M o Q ; b. Ratios of dimethyl ether and formaldehyde selectivities versus area ratios of apical (001) and (101) to basal (010) faces. (From J. Catal., 72. 375 (1981).

. , * . 3 * n \


oxygen. The (100) face catalyzes dehydrogenation in the presence of gaseous oxygen when MoSO groups exist. In the absence of oxygen, this face catalyzes dehydration where the active sites are proposed to be the exposed Mo cations which act as Lewis acid sites.

In addition to the dependence on the crystal face orientation, the selectivity in methanol oxidation also depends on the polymorphic crystalline form of Mo03 [115]. The orthorhombic form is less active and catalyzes mainly the oxidation to methylal. The hexagonal form is more active, and the main reaction is dehydration to dimethyl ether.

In addition to molybdenum oxides, vanadium oxides are also selective in the production of formaldehyde. Under optimal conditions, 87% yield of formaldehyde is obtained with an unsupported V2O5, and a Ce02-supported V2O5 shows a nearly 100% selectivity at low conversions [116]. On other supports, the selectivity decreases in the order: a-A1203 (83%) > Ti02 (rutile) = zirconia > Ti02 (analase) [116,117]. In general, the selectivity decreases rapidly with increasing temperature because of increase in conversion.

When VzOs is used in a solid solution with other oxides (A1203, Sn02, or Ti02), the selectivity for HCHO is very low at low vanadium concentrations. At high concentrations, the selectivity increases up to almost 100% [118]. V2O5 supported on A1203 or TiOz shows poor selectivity. On V2OS/Al2O3, dehydration dominates; on V205/ri02, combustion dominates [ 1191.

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Chapter 13

CATALYTIC REACTIONS BETWEEN

HYDROGEN AND CARBON OXIDES

13.1 INTRODUCTION

CO and C02 can be hydrogenated on transition metal oxides to alcohols and hydrocarbons. Since the oxides of the Group VIII transition metals have lower hydrogenation activities, they may show higher selectivities for higher hydrocarbons and alkenes. Transition metal oxides also have a lower tendency to dissociate CO than their metals. Thus, they often produce alcohols more selectively. In addition, some transition metal oxides produce branched products quite selectively, especially 2-methylpropane, 2-methylpropene, and 2- methylpropanol. However, in some of the systems that have been extensively studied, such as Cu-Zn oxide, rhodium or palladium oxide, it has been controversial as to whether the active components are the metal or the oxide.

The water-gas shift reaction, which is the reaction between carbon monoxide and water to form hydrogen and carbon dioxide, also involves hydrogen and carbon oxides. This reaction is readily reversible, and the reverse reaction, known as reverse shift reaction, is between hydrogen and carbon dioxide.

Various aspects of earlier work on these reactions catalyzed by oxides have been rcviewed [l]. For example, it has been discussed in these earlier reviews that catalysts based on ZnO, such as Zn-Cr oxide are effective in the hydrogenation of CO to methanol, which can also be catalyzed by copper catalysts. Higher alcohols can be produced by promoting these catalysts with alkali ions. On the other hand, Z r O 2 and Tho2 are effective for the production of branched products. Since these earlier reviews, new catalysts have been discovered and significant progress in understanding the reaction mechanisms has been made. The discussions here will emphasize the more recent developments.

221

228

13.2 OTHER CU-BASED CATALYSTS

ALCOHOL SYNTHESIS ON COPPER-ZINC OXIDE AND

Methanol Synthesis Copper-zinc oxides are very active and selective catalysts for the production

of methanol from a feed mixture containing CO, C 0 2 and H2 [2-51. Highly selective catalysts can be made with a C a n ratio between about 3/7 to 7/3. Often the catalysts are suppported, and A 1 2 0 3 is the support in the commercial catalyst. At present, it is controversial as to whether the active phase is highly dispersed copper metal supported on ZnO or Cu(1) ion in a ZnO matrix. This will be discussed further later.

Highly active and selective catalysts to produce methanol from a CO and H2 mixture have been reported using other copper compounds. In particular, rare earth-copper intermetallics appear to be quite effective 121. For these catalysts, the active phase is believed to be metallic copper.

There has been a large volume of work reported on the study of this reaction, and a very comprehensive review has been published [2]. The readers should refer to this review for detailed discussions regarding this reaction.

Methanol can be produced by the hydrogenation of CO or C02:

CO + 2H2 + CH30H (13-1)

CO2 + 3 H2 + CH30H + H2O (1 3-2)

Industrially a mixture of CO, C02 and H2 is used as the feed. The high activity of Cu-Zn oxide has made it possible to synthesize methanol at a lower temperature (c30O0C) and a lower pressure (c10 MPa) than those used for the older generation of catalysts such as Zn-Cr oxide.

Kinetics and Mechanism of Methanol Synthesis Typical industrial reaction conditions for a copper-zinc oxide catalyst are 200

to 275°C and 5 to 10 MPa total pressure. A typical feed consists of a mixture of CO (8 to lo%), C02 ( 5 to 6%), and the balance being Hz. At this temperature, the methanol synthesis catalyst is also an active catalyst for the water-gas shift reaction:

CO + H20 .I- C02 + H2 (13-3)

The rate of methanol production depends on the partial pressures of the various components. On Cu-Zn oxide, the rate from a C02/H2 mixture is faster than from a C0/H2 mixture [6-81. In a mixture of CO, C02, and H2, the ratc depends both on the CO/C02 ratio and the conversion. At low conversions, the rate increases with increasing C02 content [6]. At high conversions, the rate first increases rapidly with the addition of a small amount of C02 to a HdCO feed, but decreases on further addition beyond an optimal concentration [9,10]. This behavior is shown in Fig. 13-1. A small amount of water in the feed also enhances

CO HYDROGENATION 229

2 0 30 m a

20 c z W 0

W a 1 0

a

I I 0 10 20 30

PERCENT C o p IN SYNGAS, C O + c o 1 = 3 O

Figure 13-1 Effect of COz partial pressure on the rate of methanol production on a Cu-Zn oxide catalyst. (From J. Catal.. 74, 343 (1982), copyright Academic Press).

the rate of methanol production from a mixture of Hz and CO (Fig. 13-2) [ll], although high concentrations actually suppresses the rate [6,11].

The behavior shown in Fig. 13-1 can be interpreted possibly by two factors. The first is that methanol is formed faster by the hydrogenation of COz than by CO. Thus addition of C02 enhances the rate at low C02 partial pressures. The apparent suppression of the rate at high C02 partial pressures is probably due to limitation by the approach to chemical equilibrium. However, it may also be due to poisoning of the catalyst by carbonate formation, oxidation of the catalyst, or the presence of a high partial pressure of water due to the reverse water-gas shift reaction and the hydrogenation of COz.

The behavior shown in Fig. 13-2 can be similarly explained. In the absence of water, CO hydrogenation is the only reaction possible. Addition of a small amount of water results in the formation of C02 by the water-gas shift reaction, and the rate of methanol production is enhanced. Excess water results in suppression of the rate because water competes effectively for the active sites.

That C02 is the main source of carbon for methanol has been confirmed by isotope labeling experiments, first by Rozovskii and coworkers [12-141, and later by others [15-171. These results can be summarized as follows. If the carbon atom of a carbon dioxide molecule in a feed is labeled, and the feed is passed over a

230

0 20 40 60 80

Water Addition, mol/h X lo3

Figure 13-2 The dependence of methanol synthesis rate on the quantity of water added to the binary H&O = 70/30 vol% synthesis gas with GHSV = 6120 l(STP)/kg cat/hr at 7.5 MPa and temperature of 1: 190DC, 2: 215"C, 3: 22SoC, and 4: 235°C. Catalyst weight = 2.45 g. (From F'roc. 8th Intern. Congr. Catal., II. 47 (1984), copyright Elsevier Scientific h b l . ) .

Cu-Zn oxide catalyst under conditions that are close to the industrial ones, the carbon label of the methanol molecule is identical to that of carbon dioxide when thc conversion is low (high space velocities). As the conversion increases (lower space velocities), the carbon label begins to appear in carbon monoxide, and the fractions of carbon dioxide and methanol in the reactor outlet that contain the label decrease. These results are shown in Fig. 13-3 [15,16].

Since water is formed in the experiments shown in Fig. 13-3 by reactions (13-2) and (13-3), these results suggest that the presence of a small amount of water has little effect on the main source of carbon for methanol. However, there are indications that a large amount of water suppresses the rate of hydrogenation of C02 such that the contribution from CO hydrogenation becomes significant [17].

The effect of C02 on the rate of methanol synthesis from CO differs on other catalysts. On CuCr02, addition of a small amount of C02 to a CO and H2 feed actually decreases the rate of production of methanol by about 25%, probably due to competition by C02 for the adsorption site of H2. Further addition does not show any more effect [18]. On a Cu/Th02 catalyst, addition of C02 to the C0/H2 feed

CO HYDROGENATION 23 1

a 0

m 5 0.8

CT \ \ \ \ . ,exit CO

\ \

% - - inlet CO -I

'1 - - - - 0 1'0 2b

Space Velocity (h-lX

Figure 13-3 Variation of specific radioactivity of the products with space velocity in the reaction of CO:C02:H2 = 10:10:80 over Cu/Zn/Al/oxide. Total pressure = 5 MPa, 25O"C, and the feed contained 14C02. (From Appl. Catal., 30, 333 1987). copyright Elsevier Scientific hbl . ) .

has no effect on the rate of methanol production, although it prevents the deactivation of the catalyst that occurs at low HJCO ratios [19]. On 21-0~. methanol is produced as a minor product from a HJCO feed only if a small amount of water is also present [20]. In this case, it is believed that water is needed to form surface hydroxyl groups which react with CO to form surface formate species, which are believed to be intermediates in the formation of methanol on this oxide.

The detailed mechanism of methanol formation on a Cu-Zn oxide catalyst has been investigated by a number of workers, but no agreements have been reached. An earlier infrared study has reported the observation of two bands at 2270 and 2661 cm-' which have been assigned to a formyl species [21]. In a later in situ study, these two bands were not observed. Instead, a surface formate, which can be hydrogenated to adsorbed formaldehyde and then to adsorbed methoxy was observed [22]. This surface formate could be formed by the hydrogenation of a surface carbonate. A carbonyl species with a vibrational frequency of 2090 cm-' was also observed which was assigned to CO adsorbed on copper. A surface formate adsorbed on copper metal has also been identified in the reaction of COz with H2 [231.

The presence of surface formate has also been confirmed by chemical trapping with dimethyl sulfate [7]. The presence of surface formyl and formaldehyde species are inferred from results of chemical trapping with diethylamine, methyldiethylamine [ 111, and isopropylamine [24]. In these experiments, the chemical trapping agent is injected into the feed stream of CO and H2 under reaction conditions for methanol synthesis. The extra products are then

analyzed so as to deduce from them the nature of the adsorbed species that have reacted with the trapping agent, For example, in the study using isopropylamine as the trapping agent, methylisopropylamine is formed. Since condensation between methanol and the amine is slow under the reaction condition, it is proposed that methylisopropylamine is formed by the reaction of isopropylamine with adsorbed formaldehyde:

H R1 R2N C h H

I M

-H 11. R1 R2NCH20H

+ M

H20 11-H20

R 1 R2NCH3

+

H2 H2O ____) (13-4)

+

M

The result of this experiment therefore supports the presence of surface formyl and/or adsorbed formaldehyde species, which is formed by the reaction of adsorbed CO and adsorbed hydrogen.

Based on these results, a likely mechanism for the formation of methanol from C02 is as follows:

H +H 2

H I * + 4 + COz(ad) + * + -*-OH - H > C 4 H & (13-5)

*

+ H CH3 I +H -- 0 -4 CH30H J

Nature of Copper-Zinc Oxide Catalyst in Methanol Synthesis There are two proposals for the nature of the active component in copper-zinc

oxide. Thcy arc Cu(1) ions stabilized in a matrix of ZnO, and highly dispersed copper metal on a ZnO support. In the former model, the Cu(1) ions are proposed to be sites for the adsorption of CO, whereas the ZnO matrix is for the adsorption of H2. This model was first advanced by Klier [3,25]. The data that support this


model include the following: a) Optical diffuse reflectance results show that the active catalysts possess a

new absorption band that is not characteristic of ZnO, copper metal, or cupric oxide.

b) Electron microscopic and surface analyses show intimate contact between copper and ZnO.

c) The active catalysts contain a copper phase that is amorphous to X-ray diffraction.

d) The methanol synthesis activity has been reported to be directly correlated to the amount of strongly adsorbed CO [3], and only Cu(1) species are known to adsorb CO strongly [26]. In one in-situ infrared spectroscopic study of the catalyst using CO and H2, a carbonyl species at 2130 cm-I was observed whose intensity was linearly proportional to the methanol synthesis activity [27]. In another study, a band at 2093 cm-I was observed that appears to be involved in the synthesis reaction [lo]. These band frequencies suggest adsorption on an oxidized form of copper.

e) XPS studies of an active catalyst show that Cu(1) is stabilized by ZnO [28]. f ) CuCr02, which contains Cu(I), is active in methanol synthesis [18,29].

In this catalyst, the activity could be correlated to the amount of Cu(1) detected by X P S .

In the second model, copper metal is believed to be the active component. Zinc oxide serves to stabilize the copper metal in a state of high dispersion, thereby maintaining a large surface area of copper. Under high pressures, the reactants for methanol synthesis COz and H2 adsorb on copper. Participation of zinc oxide in the activation of H2 is also possible. The evidence that support this model include: a) The methanol synthesis activity has been found to be linearly proportional

to the surface area of copper [15,16,30], which, for example, can be determined by the stoichiomeuic decomposition of nitrous oxide.

b) Cu/Si02 is also an active and selective catalyst for methanol synthesis [27,31]. The activity is comparable to Cu-Zn oxide [27], but it decreases with time on stream due to sintering of copper [31]. Similarly, highly active Cu/A1203, Cu/MgO, Cu/MnO and C o h o 2 catalysts have been reported [15,32]. It is not known that Cu(1) can be stabilized by these supports.

It is unlikely that the question of the active site can be resolved readily because at present it is quite difficult to characterize a catalyst under reaction conditions. Since the state of the catalyst, especially the oxidation state of copper, responds rapidly to the gas composition, there will be uncertainties in the applicability of the data from various characterization techniques to the working catalyst.

Effect of Promoters on Cu-Zn Oxide The performance of a Cu-Zn oxide catalyst can be improved by the addition

of promoters. Aluminum is a structural promoter that significantly increases the thermal stability of the catalyst. Many other promoters have been mentioned, including boron, chromium, manganese, rare earth, and some alkaline metal ions

234

ethanol

600

E 6 t methyl /- formate

I

7

Nominal Cs Conc. mole %

Figure 13-4 Yields of methanol, methyl formate, and ethanol as a function of Cs loading on a Cu/Zn/O catalyst. Reaction conditions: Hz/CO = 2.333, P = 7.6 MPa, GHSV = 6120 l(STP)/kg cat-h, 523 K. (From J. Chem. SOC. Chem. Commun.. 193 (1986). copyright Royal society of Chemistry).

[2 ] , but there is little understanding of their functions. Potassium, barium, and cesium ions are among the alkali and alkali earth

promoters. The effect of the addition of Cs has been studied in detail [33,34]. It has been observed that the rate of methanol production by the hydrogenation of CO is greatly increased by adding a small amount of Cs to the catalyst. This is shown in Fig. 13-4. Addition of a large amount of Cs, however, blocks the catalyst surface and causes decline of the activity.

It is proposed that Cs participates in the synthesis reaction in the following manner [34]:

(13-6)


In addition to methanol, Cs promotion also enhances the production of C2 compounds such as ethanol and methylformate (Fig. 13-4). as well as higher alcohols [33,34].

Whether CO or methanol is the carbon source for ethanol and methylformate produced on a Cs-Cu-Zn oxide has been investigated by isotopic labeling experiments [34]. By using a mixture of 13CH30H and 12C0/H2 as the reactant, it is found that the 13C label in methylformate resides almost exclusively in the methyl group. In ethanol and propanol, the carbon atoms are about equally labeled.

These labeling data can be explained by the following. Methylformate is formed by CO insertion into the surface-oxygen bond of a methoxide, which might be bound to a Cs ion:

CH3 co Y H 3 -H20 I CSOH + CH3OH e==k 0' -+ w=O (13-7)

+H20 Cs@ @c s

+HzO r-- CSOH + HCOOCH3 -H20 --2

The nearly equal extent of labeling of the carbon atoms in ethanol and propanol suggests that CO insertion is not involved in their formation. It is proposed that they are formed by coupling of alcohol molecules. For example, the overall reaction for the formation of ethanol is:

2CH30H CH3CH20H + H2O (13-8)

which might proceed via an aldehydic intermediate:

(13-9)

Higher alcohols are then produced by analogous coupling reactions. Promotion of a Cu-Zn oxide catalyst by both Co and A1 results in a catalyst

that produces almost as much ethanol as methanol at 563 K and 6 MPa pressure using a feed of HJCO ratio of two [35]. Higher alcohols are also produced the amounts of which decrease with increasing chain length.

236

Unlike Cs, promotion by potassium enhances the production of branched alcohols [36]. This will be described in the section on isosynthesis.

13.3 G R O W VIII METALMETAL OXIDE CATALYSTS

A number of group VIII metal oxides have been reported to be active in the hydrogenation of carbon oxides. Rather extensive work has been performed on systems based on Pd, Rh, and Fe. These will be discussed here. However, in all of these cases, it is not yet settled whether the active component in the catalyst is the metal or the metal oxide.

Methanol Production on Palladium Catalysts It is well known that Pd catalyzes CO hydrogenation to methane. However,

the production of methanol on Pd has also been known for a long time [37]. Recently, it has been found that methanol can be produced with high selectivity [38,39]. Since then, it has been found that the methanol selectivity depends on the support used. For example, methanol is produced almost exclusively on Pd supported on La2O3. On other supports, the selectivity for methanol decreases as: La203 > Si02(II) > ZrOz > ZnO - MgO > Ti02 > A1203 2 Si02(I), and the product on Pd/Si02(I) is practically exclusively methane [40]. It is interesting that the selectivities are quite different on the two different sources of silica, Si02(I) and Si02(II). This difference for different SO2 supports has later been confirmed [41]. In general, basic oxide supports favor methanol production. Different promoters, such as Na [39,42], Mg, and L a [43] also enhance the selectivity for methanol.

The reasons behind the dependence of the selectivity for methanol versus methane on the nature of the promoter or the support are not well understood. One proposal is that methanol is formed via a formate intermediate. The role of the alkali ion promoter is to stabilize the formation of the formate [42], which has been detected by infrared spectroscopy on a Na-promoted Pd catalyst. By isotopically labeling this formate species and following its reaction, its participation in methanol production is established. However, this proposal has not been supportcd by other evidence. In one study, chemical trapping experiments on a Mg-promoted catalyst did not detect any surface formate. Instead, a good correlation between the activity for methanol production and the concentration of a surface formyl species was obtained [44]. In another study, a surface formate species was detected only on a Pd/La203 catalyst but not on a Pd/SiO2 catalyst, although the kinetics of methanol synthesis on both catalysts were very similar P51.

Another factor that might determine methanol selectivity is the morphology of the Pd crystalliks. It has been suggested that a Pd(100) surface in a Pd/La203 catalyst is more active than a Pd(ll1) surface, and for a fixed morphology, Pd/La203 is more active than Pd/SiO [461

Yet another proposal is that Pd ion is involved in methanol production. It has been found that the rate of methanol production can be correlated to the amount

3 . -


of Pd extractable by acetylacetone on the promoted catalysts [43]. Since acetylacetone extracts Pd' ions, this correlation strongly suggests that on these promoted catalysts, methanol formation is associated with an oxidized Pd center. The role of a promoter or a suitable support is to stabilize the Pd ion in a reducing atmosphere. Indeed, Pd' has been detected by EPR on MgO and La203-supported catalysts [471.

This last proposal is also supported by the effect of oxygen treatment and isotope labeling results [42]. When Pd is supported on 2 3 0 2 or TiO2, the catalyst pretreated with oxygen at 723 K followed by a low temperature reduction results in a higher methanol production rate than one without the oxygen pretreatment. The effect is removed by a high temperature reduction. It is suggested that Zr02 and Ti02 can stabilize a certain form of surface oxygen at or near Pd that results in the active site for methanol production. The effect of oxygen treatment, however, is absent if the catalyst is promoted with Na, because the special active site is already stabilized by the alkali ion promoter.

When C1*O is used as a feed over a Na-promoted Pd catalyst, the methanol produced initially is CH3160H [42]. The amount of such l60 product is much smaller on catalysts without Na promotion. Thus there is a surface oxygen species present on catalyts active in methanol production that is stabilized by the alkali ion promoter.

+

Rhodium Catalysts Some catalysts based on rhodium oxide have been found to hydrogenate CO

to produce hydrocarbons, methanol and higher oxygenates. The product distribution depends on the reaction conditions, the presence of promoters, and the pretreatment of the catalyst 1481.

Table 13-1 lists some examples where the production of oxygenates has been reported. Among the first rhodium oxide-based catalyst studied was a surface- oxidized rhodium foil [49a]. At a pressure of 0.6 m a , and a HJCO ratio of 3 to 1/3, up to about 10% oxygenates were produced, which are primarily methanol, ethanol, and acetaldehyde. Decreasing the HJCO ratio or the reaction temperature shifted the product distribution towards higher molecular weight hydrocarbons, and higher ethylene to ethane ratios. Lowering the temperature also increased the oxygenate yield, while decreasing the HJCO ratio favored the production of acetaldehyde at the expense of methanol and ethanol. About the same product distribution was obtained using a H2/C02 feed instead of a HJCO feed [49b], but the reaction rate was higher.

In comparison. the products from Rh metal that has not been preoxidized contain no oxygenates but a larger fraction of methane. Furthermore, the activity of the surface-oxidized rhodium, as measured by the methane turnover number, (that is the rate of methane production pcr exposed rhodium metal atom before oxidation), is about ten timcs that of rhodium [49a,b]. The use of a H2/C02 feed on Rh metal instead of HJCO results in a higher C b selectivity at the expense of longcr chain hydrocarbons [49b].

Rho2, Rh203, and Rh203.5H20 have been studied as catalysts [50]. It has been found that under thc reaction conditions of 300°C, 0.6 MPa of 1:l CO/H2,

238

Table 13-1 Product Distribution in CO Hydrogenation over Rhodium Oxides

Products. wt %

Catalyst 'PC H2/CO' CH, C2H, C,H, C,+ CH,OH C,H,OH MeCHO Oxyge- ref.b nates

oxidized 300 3/1 Rh foil 300 1/3 and 250 3/1 Rh(ll1) 250 1/3

heavily 300 3/1 oxid. Rh( 11 1)

R$O,. 300 1/1 5Hz0

LaRhO, 300 1/1 225 1/1

8 0 4 7 7 - 2 i 56 15 3 22 - 3 1

60 5 6 17 - 12 i 35 10 2 48 - 10 1

70 4 6 10 - 10 i

29 22(C2) 25 3 tr 16 11

... 16 9(C,+) 17 14 31 111

12 6(C,+) 40 13 23 111 ...

Footnotes: a) 0.6 MPa total pressure. b, References:

i) D.G. Castner, R.L. Blackadar, and G.A. Somorjai. J. Catal., 66, 257 (1980). ii) P.R. Watson, and G.A. Somorjai, J. Catal.. 72. 347 (1981). iii) P.R. Watson, and G.A. Somorjai. J. Catal., 74, 282 (1982).

only Rh203.5H20 resists complete reduction to the metal. The selectivity for methane on this catalyst is much lower than on a Rh metal or oxidized Rh metal (see Table 13-1). Almost all Cz+ hydrocarbons are alkenes, and a significant quantity of oxygenates is produced. In addition to the products shown, a small amount of propanal and butanal are produced. At the same temperature and total pressure, the C b yield increases with an increasing HJCO ratio, but the product distribution changes little with total pressure. Increasing the temperature increases the CH, selectivity.

Addition of C 2 b to the feed enhances the production of C~HSCHO. This can be explained by the general scheme for the formation of aldehyde via carbonylation of surface hydrocarbon species:

+ c o + (z+ l )H M-C,H, ---+ M - C O - C , H , A M + C,H,+,CHO (13-10)


Addition of C2H4 to the feed increases the surface concentration of M-C2& (or M- C2H5) which, upon carbonylation and hydrogenation, yields propanal.

Results using LaRh03 as the catalyst follow the trend described thus far. At 225 and 30O0C, 0.6 MPa and using a 1/1 C0/H2 mixture, about 80% of the products were oxygenates, of which ovTr half were acetaldehyde and ethanol. Among the hydrocarbons, methane was the dominant product. The higher hydrocarbons were mostly alkenes. In comparison, LaRh03 supported on a- A1203 produced methanol very selectively [2]. The surface condition of W h o 3 during reaction is not known. In one study, surface analyses by AES showed that the LaRh03 surface contained some Rh metal, but most of the rhodium was in the +1 state [51]. In another study, surface analyses by XPS showed the presence of Rh metal and Rh(II1). No Rh(1) was detected [52].

The higher rates of oxygenate and alkene production on rhodium oxides can be explained by the lower hydrogenation activity and CO dissociation ability of the oxides than the mctal. On going from heavily oxidized metal to rhodium oxide and LaRh03, the degree of oxidation and the stability of the oxide increases. (This has been shown by temperature programmed reduction studies [52]). Concurrently one sees an increase in the selectivity for higher hydrocarbons, alkenes, and oxygenates. The increase in oxygenate production is also a result of the higher concentration of adsorbed molecular CO on the oxides that leads to a greater rate of CO insertion into a surface MXH, species to produce Cz+ oxygenates, and to more direct hydrogenation of CO to methanol.

It should be noted that the discussions above do not explain the higher activity due to surface oxidation, especially in methane production. This phenomenon is not yet satisfactorily explained.

The role of oxidized Rh appears to be also important in other supported Rh catalyst. On a Rh/MgO catalyst that has a high selectivity for methanol and other oxygenates, a Rh(I1) EPR signal has been detected [47]. Isotope labeling experiments show that adsorbed oxygen and/or surface lattice oxygen on a Na- promoted Rh/Si02 is incorporated into the product acetaldehyde, and that the C 4 bond is formed by CO insertion into a surface CH, species [421.

However, the participation of rhodium oxide has not been universally accepted. Recently, it has been reported that the product distributions are similar on a Rh/Si02, reduced Rh2O3, and prereduced or unprereduced LaRh03 catalyst. This has led to the suggestion that Rh metal is the active component [52].

Iron Catalyst Recently, it has been suggested that iron oxide is active in CO hydrogenation

because the catalytic activity of an iron foil is increased ten times when the foil is preoxidized in dry oxygen before exposure to a C0/H2 mixture [531. This high activity decreases with time as the surface oxide is reduced. Subsequently, it was reported that the pseudo-steady state CO hydrogcnation activity of an unsupported catalyst whose precursor is a-Fez03 is higher, if the iron oxide has not been prereduced, than if it is prercduced to a-Fe [54]. The bulk of the unprereduced catalyst is converted to Fe304 at pseudo-steady state. However, there is no information on the surface conditions. The a-Fe catalyst is converted to iron

240

carbide during reaction, and the used catalyst contains more surface graphitic carbon than the used unprereduced catalyst. These results show that bulk Fe304 is at least as active in CO hydrogenation as bulk metallic iron or iron carbide.

Most of the observations above have been confinned in another study [55]. In particular, the pseudo-steady state bulk compound of a a-Fe203 catalyst is confinned to be Fe304. Some X-iron carbide is also present. If the a-Fe203 is prereduced to Fe before reaction, the pseudo-steady state catalyst contains x- carbide and small amounts of €'-carbide and a-Fe. However, the product distributions over these two catalysts are the same. This has led to the conclusion that neither F@03 nor Fe304 are active for hydrocarbon synthesis. It is proposed that the superior catalytic acitivity of the unprereduced a-FezO3 catalyst is due to the formation of very small crystallites of a-Fe and X-iron carbide under reaction conditions, which are the active phases.

13.4 ISOSYNTHESIS REACTION

It has been found that on some basic oxides, especially Zr02 and Tho2 [56- 601, and some rare earth oxides such as La203 and Dy2O3 [61], branched alcohols and hydrocarbons are produced quite efficiently, particularly 2-methylpropane, 2- methylpropene, and 2-methylpropanol. Table 13-2 shows some sample data for Zr02 [58]. At the low conversions shown, alcohols, ether, and hydrocarbons are the major products, and substantial amounts of 2-methylpropane and 2- methylpropene are produced. These two branch-chain molecules are among the major hydrocarbon products also at much higher conversions of CO, although then the predominant reaction producct is COz, which is formed by the water-gas shift reaction [W]. The reaction that leads to the production of these branched products is called the isosynthesis reaction. In addition to the oxides mentioned above, some alkali-promoted ZnO [56] and potassium-promoted Cu-Zn oxide [36] also produce branched products.

The formation of alcohols and hydrocarbons in the isosynthesis reaction are closely related. For example, on Th02, branched alcohols are produced in substantial yields below about 375°C. At higher temperatures (425-450°C), branched C4 hydrocarbons are formed. Above 50O0C, methane, ethane, and propane are the principal products [62].

A mechanism for chain growth and chain branching in the isosynthesis reaction on Zr02 is shown in Fig. 13-5 [58,63]. It is based on analogous compounds and reaction mechanisms involving coordination complexes. An important intermediate in this mechanism is the bound aldehyde (species I). The lowest member of the bound aldehyde is the bound formaldehyde (R=H), which can be formed by hydrogenation of adsorbed CO or surface formate. The hydrogenation of adsorbed CO may take place by insertion of CO into a metal- hydrogen bond. A model for this step is the reaction of CO with [(Cp)2ZrHCIl, to give [(Cp)2ZrC1]2(p-CH20) [ a ] . This latter compound has a bridging formaldehyde moiety in which the oxygen atom of the formaldehyde bridges the two metal atoms and the mcthylene fragment is bound to one Zr and the 0 atom.


Table 13-2 Product Distributions in CO Hydrogenation on Zr02 (425°C. 3.5 MPa total pressure, CO/H2 = l/l). (From J. Catal., 109, 284 (1988), copyright Academic Press).

CO conversion % Product Selectivity (mole %)

methane methanol dimethy lether ethane and ethene propane and propene linear C4's branched C4's. linear C5's branched C5'sb

C4 alkenes/C4 alkanes (mole ratio)

catalyst 1

0.6

42.8 4.1 16.8 11.8 3 .O 4.6

13.1 0.5 2.3

16.4

catalyst 2

0.4

41.9 5.9

10.4 11.1 4.7 5.9 16.2 0.4 3.5

20.2

Foomotes: a) The ratio of 2-methylpropene/branched C4's is greater than 0.93. b, Monomethylated butanes only.

C2 compounds are formed by insertion of CO into a bound formaldehyde to form a cyclic acyl species (I -+ TI). Hydrogenation of species II (I1 -+ IV) would eventually lead to a linear alcohol that is one carbon longer.

The acyl species @I) has an isomeric structure (HI), which permits 1,2-H shift to form an enediolate (I11 -+ V). The formation of enediolate by a similar process has been suggested in complexes of thorium, uranium, and hafnium [65,66]. Chain branching results when the enediolate is hydrogenated to form a bound diol (V -+ VI) which dissociates to form an enolate and a surface hydroxyl (VI 4 VII). The latter step is similar to what would be involved in the catalytic dehydration of alcohols. Species VII has a resonance structure VIII which is a q3-enolate. Addition of H to VIII leads to a bound aldehyde structure OX). Insertion of CO into this species like the step I -+ I1 complctcs the cycle for chain branching.

The validity of this mechanism, especically the importance of the bound aldehyde structure, has been tested by monitoring the incorporation of propanal and acetone into the isosynthesis products. It has been found that on ZrOz [58], addition of propanal to a feed of CO and H2 increases the yields of 1-butene, 2- mcthylpropene, and 2-methyl- 1-propanol. By using 13C-labeled propanal, it is confirmcd that these products can indeed be formed from propanal. Similarly, when acetone is added, the yields of I-butene, 2-methylpropene, and 2-methyl-l-

242

[I

n

1 I

N

In N

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u \/o

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\n

1

u II

k5

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I-

u

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N

/b

s

o

S

:T m In

(Y

I

u \

/

u

II

I

0 N

I 0,

'L

- N

2

/

rJt x

II 0

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n

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Y

In

I

u-

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=/

\L

\u /N

II

Y

In

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F\o

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0

1


CH3 I

z; j o

CH2 I CH

\

( 1 )

*

- H L - + H

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I CH20H

+ H 1 7

- H

1 CH 3

I )o

* I CH3-CH

I CH

Z r ( X I 1 1

I,

CH3 I 0 I Zr

( X I )

+

Figure 13-6 Aldol condensation mechanism in the formation of branched products on Z Q . (From the same source as Figure 13-5. copyright Academic Press).

butene are increased. The enhanced yields of these products support the participation of surface

species of the bound aldehyde structure (I and IX). A bound propanal would produced 1-butene by the reaction sequence I + I1 + I V 4 butanol, which is dehydrated to 1-butene. The same reaction sequence for a bound acetone would produce 2-methylpropanol, which on dehydration would result in 2-methylpropene. The production of 1-butene from acetone is interesting. It can be accounted for if the steps leading from species II to IX are all rapidly reversible.

It has also been shown that when 2-propanol is added to the reaction feed, there is little enhancement in the yields of the C4 products [%I. Thus 2-propanol cannot be converted easily into an intermediate for chain growth or chain branching. This rules out the possibility that acetone is first hydrogenated to 2- propanol before i t is converted to the C4 products.

The enhanced production of 2-methylpropene or 2-methylpropanol from propanal, or 2-methyl-1 -butene from acetone cannot be explained by this mechanism. However, these are important products. In the experiments where propanal was added, it was observed that the yields of 2-methyl-1-propanol and 2- methylpropanal were both increased substantially in a parallel fashion. It was

244

concluded that these oxygenates were the primary products. They were subsequently converted to 2-methylpropene by dehydration and hydrogenation [%I. The enhanced yields of these products suggest the presence of a parallel mechanism, which is also suggested by other evidence. If the mechanism in Fig. 13-5 is the only process on the catalyst, it would predict a statistical distribution of products with respect to the number of carbon atoms in the molecules. In particular, the amount of each product would decrease in a proportional manner as the carbon number increases, (that is, it would follow the Anderson-Schultz-Flory distribution). The data in Table 13-2 shows that this is not the case. The amounts of C4 products are substantially more than those of C3 and C5.

To account for these observations, an additional mechanism has been proposed. This mechanism is shown in Fig. 13-5. It involves alkylation of an q3- enolate by a nearby methoxide group at the P-carbon. This is a condensation reaction, and it is assumed that during synthesis from CO and H2, this reaction is faster for the C2 species than for the C3 or C4 species, and has no equivalent in the C1 reactions [63].

The participation of this condensation reaction has been confirmed. It has been observed that when '3C-melhanol is added to the CO/H2 feed, 13C is incorporated into 2-methylpropene, 2- and 3-methyl-1 -butene, and 2-methyl-2- butene. Interestingly, no 13C incorporation is observed in the C4 or C5 linear products [ S S ] . Since it has been shown that methanol adsorbs as methoxide on the catalyst [67], the incorporation of 13C into the branched products supports the condensation mechanism.

This mechanism explains the production of 2-methylpropanol and 2- methylpropene from propanal. Condensation of a methoxide with a bound propanal results in these products. The production of 2-methyl-1-butene from added acetone can be understood also. Condensation of a methoxide with a bound acetone would produce 2-butanol or 2-butanone. CO insertion to a bound 2- butanone followed by the reaction sequence leading from I to IV would produce 2- methyl-1-butene.

13.5 WATER-GAS SHIFT REACTION

The water-gas shift reaction is an industrial process to generate hydrogen from steam and carbon monoxide:

CO + H20 + C02 + H2

The reaction is carried out between 1 to 3 MPa. The commercial catalysts are of two types. An iron-chromium oxide catalyst, sometimes promoted by other oxides such as A1203, or supported on Si02, is used at a higher temperature of 600- 725 K. A newer catalyst, based on copper-zinc oxide, promoted by Cr, Co, A!, and Ni, is used at a lower temperature of about 500 K. The advantage of the copper-zinc oxide catalyst over the iron-chromium oxide catalyst is that the catalyst is much more aclive and can be used at lower temperatures where the


equilibrium is more favorable to the formation of H2. The disadvantages are sensitivity to poisoning by sulfur and chlorine, and shorter catalyst life.

The copper-zinc oxide catalyst used for this reaction is essentially the same as the one used in methanol synthesis. In fact, during methanol synthesis, the water- gas shift reaction proceeds at a comparable rate. Thus the readers are referred to section 13.2 for discussions of this catalyst. When this catalyst is used for water- gas shift, methanol production becomes thermodynamically unfavorable because of the high water and low hydrogen partial pressures.

It is believed that the mechanism of the reaction over copper-zinc oxide involves a surface formate that is formed by the insertion of CO into a surface hydroxyl group:

CO(g) + CO(ad) (13-1 la)

H20(g) + OH-(ad) + H+(ad) (13-llb)

CO(ad) + OH-(ad) + HCOO-(ad) ( 13- 1 1 C)

HCOO-(ad) _L C02(ad) + H-(ad) (13- 1 Id)

A correlation between the ability to form a surface formate and the water-gas shift activity of a catalyst has been observed [68]. On Cu-Zn oxide, this mechanism has been discussed for methanol synthesis (similar to reaction 13-6). A surface formate has been identified by chemical trapping [7] and by infrared spectroscopy for the reaction of H2 + C02 [22,23]. Further support comes from the fact that the rate for formic acid decomposition is comparable to that for water- gas shift reaction [69].

In spite of the fact that a surface formate is probably involved as an intermediate in both the water-gas shift and the methanol synthesis reaction, there are no indications that these two reactions are necessarily correlated. In fact, the relative activities for the two reactions on different catalysts could be quite different.

The iron-chromium oxide is a mixture of Fe304 and Cr203 under reaction conditions. It is generally agreed that the addition of Cr greatly helps retain the surface area of the catalyst. There is no evidence that chromium chemically affects the iron oxide. The mechanism with which the Cr promoter functions is not firmly established. It has been suggested that Cr2O3 forms a protective layer around the a-Fe203 particles (which has the same corundum structure as Cr2O3) during calcination, thereby prevcnting particle growth of the iron oxide [70]. Indeed, determination of the surface composition of an Fe203-Cr203 solid solution has suggested the formation of a surface compound which might serve as the protective laycr [71]. It has been shown on another system of Fe304 supported on

246

Si02 that the formation of a surface shell of iron-silicon oxide on a Fe304 particle stabilizes the particle from sintering during oxidation-reduction [72].

There are two proposed mechanisms for the iron-chromium oxide catalyst. One mechanism involves a surface formate as shown in eq. 13-11 173,741. The other is a redox mechanism in which a surface oxygen species is involved as is shown below:

(13-12)

slow p L - M +CO2

s low 0, p -2 - M + H 2

In this mechanism, the active site is proposed to be an anion-cation pair

O P site ( \M ). CO first adsorbs weakly on the coordinatively unsaturated metal cation of this site. This weakly adsorbcd molecule then reacts with surface oxygen in the site to form a bidcntate carbonate, which decomposes into C 0 2 leaving

an anion vacancy in the site (M ). The anion vacancy is the site for the adsorption of water. The oxidized surfacc site is regcncratcd by thc desorption of H2 [751.

Support for this redox mechanism (also called rcgencrative mcchanism) has been provided by chemisorption expcriments undcr catalytic conditions which show that: (1) H2 adsorbs dissociatively and H 2 0 adsorbs associatively, and the isotherm for concurrent adsorption of bolh molcculcs follow the form for competitive adsorption [75]; (2) whcn a Fe3O4 catalyst is exposed to a mixture of CO and C 0 2 , thc adsorption isotherms of thcse molecules also follow the form for competitive adsorption. Furthermorc, thc saturation coverage of CO corrcsponds directly to thc coverage of anion vacancics, while changes in the C02 saturation coverage correlate with changcs in the covcragc of the reducible oxygcn spccics on thc surface 1761; (3) by following the CO-C02 isotopic exchange reaction using 13C-labeling 1771, it is concludcd that CO and C02 arc indistinguishable once adsorbcd on thc catalyst as is prcdictcd by rcaction (13-12); (4) the rates of production of CO from thc rcaction of C02 with a catalyst surfacc, of C02 froin CO as mcasurcd h y thc chcinisorption cxpcrirncnts, and of 1-12 from H 2 0 and H2O

/O


0 50 100 150 200

Fe 0 Particle Size, nm 3 4

Figure 13-7 The F q 0 4 particle size dependence of the water-gas shift activity of magnetite. Open circles represent unsupported Fe304 (except that the smallest particle-size sample is a commercial catalyst) and solid circles represent silica-supported catalysts. (From J. Catal., 76, 93 (1982). copyright Academic Press).

from H2 are comparable to the rate of th water-gas shift reaction under similar conditions [78]; and (5) it is observed that the rate of the wixr-gas shift reaction corresponds to the rates at which H20 oxidizes and CO reduces the surface of Fe304 [79,80].

It has been further proposed that the active site consists of octahedrally coordinated Fe ions of magnetite that are exposed on the surface. The proposal is based on the following observations [75]. It is found that, the rate of the water-gas shift reaction per surface Fe ion (the amount of exposed surface Fe ions is dctcrmined by NO adsorption) is relatively ilJdependent of the crystallite size of unsupported Fe304 (see Fig. 13-7) 1811. Hovever, the rate per surface Fe ion decreases sharply with decreasing crystallite size on Si02-supported magnetite catalysts. On the other ha. neither unsupported nor supported catalysts show crystallite-six dependence when the number of active sites is determined by CO/C02 adsorption (which measures the amount of removable oxygm on the catalyst surface). I t is further shown that Si ions substitute into the magnetite

248

lattice at the surface of the crystallites in the Si02-supported catalysts. The Si ion substitution occurs at the tetrahedral sites [72].

The data in Fig. 13-7 can be explained as follows. The catalytically active site consists of a surface-exposed Fe cation in an octahedral site and the surrounding oxygen anions. Then the apparent crystallite-size effect on the supported samples is due to the following. In a Si02-supported catalyst, Si4+ substitution in the magnetite surface causes all of the Fe ions to be in the octahedral sites as well as a decrease in the electron density at these ions. The latter effect perturbs the balance between Fez+ and Fe3+ at these sites and reduces their ability to undergo oxidation and reduction, as is required in the regenerative mechanism. This effect is manifested in a decrease in the amount of reducible oxygen on the surface as determined by CO/CO2 chemisorption, and a corresponding decrease in catalytic activity. Since the Si subsitution is confined to the surface, its effect is larger for smaller crystallites, and affects more strongly the more weakly adsorbed gases such as CO and C02. On the other hand, it does not affect the adsorption of more strongly adsorbed gases such as NO. This might be due to surface reconstruction induced by NO adsorption on cations that are inaccessible to more weakly adsorbed gases [81,82].

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33. J. Nunan, K. Klier, C.-W. Young, P. B. Himelfarb, and R. G. Herman,

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76, 1 1 2 (1982).

Chapter 14

PHOTO-ASSISTED SURFACE PROCESSES

14.1 INTRODUCTION

A general picture of the bulk electronic band structure of the valence electrons of transition metal oxides may be described as follows: the 0 2p orbitals overlap to form a filled valence band, and the s and d orbitals of the cations overlap to form the conduction band. The conduction band is at a higher energy than the valence band and separated from it for most transition metal oxides at their highest oxidation states. The separation between the two bands is the band gap (Fig. 14- la). At thermal equilibrium, the Fermi energy (that is, the electrochemical energy) of the valence electrons in the solid is between the conduction band energy and the valence band energy. The conduction band is populated with only a few electrons that are thermally excited into the band from the valence band. This excitation also leaves holes (unoccupied electron orbitals) in the valence band. The density of electrons in the conduction band can be increased by incorporating donor ions in the solid that donate valence electrons to the conduction band, and the density of holes in the valence band can be increased by acceptor ions that accept electrons from the valence band. The presence of these donor or acceptor ions also causes a corresponding adjustment of the Fermi energy such that at thermal equilibrium, the density of electrons in the conduction band, n, and of holes in the valence band, p, are given by:

where E,, Ef. E, are the lowest energy of the conduction band, the Fermi energy, and the highest energy of the valence band, rcspectively. N, and N, are constant. N, is related to the effective mass and mobility of an electron in the conduction band, and N, is related to the effective mass and mobili, y of a hole in the valence

252

PHOTO-ASSISTED PROCESSES

1 /Ill/ conduction band

253

I- - - - - - - - band gap

valence band

Figure 14-1 Energy level diagram of a semiconducting oxide. a. In vacuum; b. In contact with a solution containing a A-/A redox couple; c. As in b but under illumination. The photo-generated hole migrates to the surface to accept an electron from A-; the photo-generated electron migrates to the bulk of the solid and to a counter-electrode.

band. k is the Boltzmann constant, and T is the temperature. This simple picture applies to the bulk of an ideal oxide. As discusscd in

Chapter 3, truncation of the solid to form surfaces might generate surface

254

electronic states. In addition, the position of the Fermi energy may be different under different situations such as when the oxide is in contact with another medium, when there are adsorbates on the oxide surface, or when the surface is partly reduced. These situations could result in a surface region of the solid that is very different from the bulk, and a potential gradient is developed between the surface and the bulk. Associated with this potential gradient is a distribution of densities of electrons and holes (Fig. 14-lb). Absorption of light of energy larger than the band gap energy by the solid also changes these populations. Depending on the solid, the light intensity and its wavelength, photo-excitation could increase the population of electrons and holes by thousands or millions times. When these photo-generated holes and electrons migrate to the surface, light-induced surface chemistry becomes noticeable.

A photo-enhanced surface electron transfer process is schematically illustrated in Fig. 14-1. The energy level diagram of an oxide neglecting any surface effects is shown in Fig. 14-la. When this sample is placed in water which contains a redox couple A-/A, electron transfer between the solid and the solution occurs until the electrochemical potential of the valence electrons in the solid equals that in the solution. If the redox couple A-/A is the only species that can exchange electrons with the solid, the electrochemical potential of the solution is defined by the standard redox potential of this couple and the relative concentrations of A- and A. Fig. 14-lb shows the situation where before the solid contacts the solution, the Fermi energy of the solid is higher than the electrochemical energy of the elcctrons in the solution, and after contact, electrons flow out of the solid to the solution so as to establish equilibrium (that is, equalization of electrochemical potcntial) across the interface. The outflow of electrons from the solid results in a net positive charge near the surface in the solid, which eventually sets up a potcntial barrier to stop the electron flow. Equilibrium is then established. The region of the solid near the surface that has a net charge is called the depletion layer.

When the solid absorbs a photon of light of an energy larger than the band gap, an electron is excited from the valence band to the conduction band (Fig. 14- lc). Because of the upward-bending of the electron bands near the surface. the photo-generated hole tends to migrate to the surface, while the electron tends to migrate to the bulk. This results in an increased hole concentration at the surface, which facilitates surface reactions involving electron transfer to the solid, such as oxidation reactions. This situation persists as long as the solid is illuminated and there is a mechanism to remove the electron that has migrated to the bulk. Flow of electrons from the oxide via an external circuit to a counter electrode is one way to remove them.

If there are no mechanisms to remove the electrons from the bulk, continuous illumination will result in accumulation of elcctrons and holes near the surface, to the extent that the bending of the electron bands is removed. The photo-generated elcctrons become available at the surface in significant concentrations, which facilitate surface reactions that involve transfer of elcctrons from the solid, such as reduction reactions.

Recently thcrc has becn a strong intercst to replace the external circuit that

PHOTO-ASSISTED PROCESSES

A A-

D-

3

P t Ti O2

255

Figure 14-2 Schematic diagram showing a Ti02 particle deposited with Pt. Absorption of a photon generates an electron-hole pair. a. The hole migrates to the Ti02 surface to accept an electron from a donor species, and the electron migrates to thz Pt surface to be donated to an acceptor species. b. The energy level diagram of this system under illumination.

connccts the counter electrode and the irradiated oxide by direct contact of the two elcctrodes. Platinum deposited on TiOz is the most popular system that has been investigated. In this system, the photo-generated holes are consumed in the surface reaction at Ti02. The photo-generated electrons migrate to platinum where they are consumed in a reaction at the platinum surface. The complete oxidation- reduction cycle can be accomplished on one composite particle. This is schematically shown in Fig. 14-2.

It is seen from this description that the photo-effect is significant only when a large concentration of photo-gcnerated holes and/or electrons are present at the oxide surface. Since most of the light is absorbed by the region of an oxide up to a

256

few hundred nanometers from the surface, these holes and electrons must have sufficiently long lifetimes to reach the surface. Therefore, reduction of an oxide and the presence of lattice defects and impurities which shorten the lifetime of these holes and electrons will reduce the photo-effect. The lifetime also depends on the effective masses of these carriers, which depend on the band structure of the solid. For example, photo-enhanced processes are readily observed on ZnO and Ti02, but are much weaker on V205.

It is also apparent that light of an energy larger than the band gap energy is needed to produce photo-effects. It is possible to generate photoelectrons with light of a lower energy if there are discrete energy levels (commonly called traps) in the band gap. Absorption of a photon may excite an electron from the trap to the conduction band. Similarly, it is possible to generate photo-holes by exciting electrons from the valence band into the unfilled inter-band gap traps. However, these processes are normally not efficient. The concentration of these traps are limited so that the possible densities of photo-generated holes and electrons are also limited. These traps usually act as electron-hole recombination centers that shorten the lifetime of the photo-generated carriers.

The description for the scheme shown in Fig. 14-1 is for an n-type semiconducting oxide. An analogous scheme can be readily drawn for a p-type semiconducting oxide, whose Femi energy is much closer to the valence band than n-type semiconductors. Because of the position of the Fermi energy, it is common that electrons flow into a p-type oxide from an external redox couple. This results in a bending of the electron bands in a direction opposite to those shown in the Fig.

A side effect that occurs on illumination is heating of the oxide. A majority of electrons and holes recombine nonradiatively. The energy released heats up the oxide. In many experiments, light sources are used that give a broad spectral distribution. Some light of energy less than the band gap energy is absorbed by the phonon mode of the oxide and heats up the solid. The extent of heating depends on the oxide and the light source used. A temperature rise to 60°C from room temperature has been reported [19].

Thus the magnitude of the photo-effect depends on the oxide.

14-1.

14.2 PHOTO-ASSISTED ADSORPTION AND DESORPTION

As described in the last scction, absorption of light of energy larger than thc band gap energy increases the surface concentrations of electrons and holes of a semiconducting oxide significantly and enhances rate processes such as adsorption and dcsorption that involve electron transfer.

Photo-assisted adsorption and desorption of oxygen is the most studied process in this category [ 1,23:

Adsorption: e- + O2 + 02-(ad) (14-3)

Desorption: h+ + 02-(ad) --+ 02(g) (14-4)

PHOTO-ASSISTED PROCESSES 257

Both adsorption and desorption can be enhance by irradiation. Which process dominates depends on the experimental conditions and the nature and pretreatment of the oxide. It is commonly found that the relative rates of these two processes depend on the pressure of oxygen. Band gap illumination of ZnO at room temperature results in photo-assisted adsorption of oxygen at low pressures. The effect decreases with increasing pressure, and eventually becomes photo- assisted desorption at high pressures [3]. However, the effect of pressure is opposite at 400°C: photo-assisted desorption at low pressures, and adsorption at high pressures. The effect is not well understood.

A similar pressure dependence has been reported on Ti02 at 500°C [41. At low pressures of oxygen, irradiation enhances desorption; at high pressures, it enhances adsorption.

Detection of photo-assisted adsorption can be made by following changes in the conductivity of the sample. Since such adsorption changes the concentration of charged species on the surface and consumes photoelectrons or holes, it changes the conductivity of the sample. It has been found that the surface returns only slowly to the equilibrium state after the light is turned off, thus the enhanced adsorption due to irradiation must be retained for a long time, and these species are practically irreversibly adsorbed. Their desorption can be enhanced by heating the sample [51.

The nature of the adsorbed species has been deduced from the dependence of photoconductivity 0 with pressure. Up to lo2 Pa of oxygen, 0 shows a -1/2 power dependence on the pressure on ZnO and ZrOz [ 10.1 11. There are indications that the order changes to -1 at higher pressures on ZnO [12]. On anatase TiOz, the order dependence changes from -1/2 to -1 as the pressure increases [lo]. An order of -1 has been reported on a single crystal rutile TiOz sample [13]. Furthermore, cr increases linearly with light intensity.

The following processes have been proposed to explain the order dependence [lo]:

02(ad) + e- + 02-(ad) (14-7)

O(ad) + e- + 0-(ad) (14-8)

The order in oxygen changes from -1/2 to -1 when thc dominant reactions change from (14-6) and (14-8) to (14-5) and (14-7). Other processes have been excluded. The reaction 02-(ad) + e- + 2 0-(ad) would lead to only a -1 order in oxygen, and the react. m: 02(ad) + 2 e- + 022-(ad) + 2 0-(ad) would lead to a half order dependence on light intensity.

The surface condition is important. It has been shown that the amount of oxygen photo-adsorbed on Ti02 increases with the amount of hydroxyl groups on the surface [6]. A similar observation has been reportcd that dehydroxylation of

258

the surface reduces photo-adsorption of oxygen on TiO2, which can be restored by rehydrating the surface [7,14]. It is believed that hydroxyl groups are important in trapping photo-generated holes, making photo-electrons available for chemisorption:

OH- + h+ -+ .OH(ad) (14-9)

In this case, the fate of the hydroxyl radical is unclear. In the presence of hydrocarbon molecules, the radical may participate in oxidation reactions, as will be discussed later.

In addition to photo-enhanced adsorption or desorption of oxygen, photolysis of the lattice resulting in evolution of lattice oxygen has been reported on ZnO [8,9]. On the other hand, photo-decomposition of Ti02 has not been observed

The photo-response of a sample depends on its pretreatment, especially if the pretreatment results in a slightly reduced or fully oxidized sample. Reduction usually leads to n-type semiconductivity, and the extent of reduction determines the density of conduction electrons and the lifetime of photocarriers. For some oxides, severe oxidation may lead to excess lattice oxygen and p-type conductivity. Coupled with the effect of surface hydroxylation, it is not surprising that there are conflicting reports in the literature on the effect of illumination.

Doping of an oxide changes its semiconducting properties and response to irradiation. Addition of Li to ZnO enhances photo-adsorption of oxygen, while addition of Ga or A1 reduces it [3,4].

Ti02 and ZnO are among those that demonstrate large photo-effects. Under the same conditions where these two oxides show photo-adsorption behavior, adsorption of oxygen on V205 does not show any response to light, and W 0 3 shows only a weak response [lo]. This behavior parallels their photo-enhanced catalytic oxidation activities which will be discussed next.

[91.

14.3 PHOTOCATALYSIS: GAS PHASE REACTIONS

A large number of gas phase catalytic reactions have been found to be greatly enhanced by illuminating an oxide with light of band gap energy. These reactions include:

oxidation of hydrocarbons oxidation of alcohols oxygen isotope exchange oxidation of CO, NH3 H2-D2 exchange

In general, these reactions proceed readily at elcvated temperatures in the dark. Thus most of the studies of light enhancement have bcen conducted near room temperature. Product selectivities are generally different between thermal and light-assisted reactions.


Table 14-1 Partial Oxidation Products in the Gas Phase Photocatalytic Oxidation of Hydrocarbons at Low Conversions.

Hydrocarbon Oxide Partial Oxidation F’roducts Major Minor

ethane propane

propene

2-methylpropane 2-methylpropene 2,2-dimethyl-

propane butane

2-methylbutane

2,2-dimethylbutane

2-methyl- 1 -butene 3-methyl- 1-butene 3-methyl-2-butene pentane

2-methy lpentane

toluene p-alkyltoluene

Ti02 Ti02 M a 3 wo3

wo3

Ti02

Ti02 Ti02 Ti02

Ti02

Ti02

Ti02

Ti02 Ti02 Ti02 Ti02

Ti02

TiOz Ti02

ethanal acetone acetone acetone ethanal, acrolein

ethanal, acrolein acetone acetone acetone

butanone

acetone

3,3-dimethylbutan- 2-one

butanone, ethanal acetone acetone, ethanal pentan-2-one, pentan-3-one 2-methylpentan- 3-one, 4-methylpenta.n- 2-one, acetone phenol substituted benzaldehyde

propanal, ethanal propanal

propene oxide, acetone, propanol

2,2-dimethylpropanal

butanal, propanal, ethanal butanone, ethanal, 3-methylbutanone butanone, acetone, 3,3-dimethylbutanal, propanal, ethanal acrolein, acetone

pentanal, butanal, propanal, ethanal 2-methylpentanal,

propanal, ethanal

Oxidation of Hydrocarbons Oxidation of hydrocarbons is among the most studied reactions. In

particular, the oxidation of alkanes have been studied heavily. Table 14-1 lists thc partial oxidation products for a number of alkanes. It is clear from the data using

260

Table 14-2 Selectivities for Ketone and Aldehydes for Different Alkanes. (From Farad. Disc. Chem. Soc., 58, 185 (1974). copyright Royal Society of Chemistry).

Alkanes % Selectivity (excluding C o d

Ketones Aldehydes

n-alkanes 56 20 alkanes possessing one or two 61 19

neoalkanes 43 14 tertiary carbon atoms

TiOz as the catalyst that photocatalytic oxidation produces a wide range of products. In addition, combustion is usually a significant reaction, especially at conversions higher than a few percent. However, there are some notable exceptions. The oxidation of 2-methylpropane to acetone on Ti02 proceeds with up to 75% selectivity [15], and the oxidation of 3-methyl-1-butene to acetone proceeds with 87% selectivity [16].

In general, ketones and aldehydes are the partial oxidation products for the hydrocarbons shown in Table 14-1, as well as for hexane, heptane, octane [17] and p-alkyltoluenes [52]. It should be noted that cracking of the carbon chain to products of lower carbon numbers is a common process. No dehydrogenation products (alkenes, dienes, etc.) have been reported as significant, although they have been suggested as reaction intermediates. The production of ketones is more preferred than aldehydes at low conversions. Teichner and coworkers have calculated the average selectivities for a large number of hydrocarbons, and grouped these values according to their structures [17]. These values are shown in Table 14-2. Apparently neoalkanes are more susceptible to combustion than other alkanes. Combustion is also the only reaction reported for butenes [15].

The reactivity also depends on the nature of the alkane. In particular, the reactivity of various types of carbon atoms on Ti02 follows the sequence:

The reactivity do not seem to follow the ionization potential of the hydrocarbons

Although Ti02 has been the most extensively studied, and it shows the highest quantum efficicncy, it is not as sclcctive as W03, Moo3 or some nontransition metal oxides such as antimony oxide and tin oxide [ 18,191.

In the formation of ketones and aldchydes on Ti02, it has been proposed that the first step of the reaction is the formation of an alcohol. If a primary alcohol is formed, it may dehydrogenate to an aldchyde. Dehydrogcnation of a secondary alcohol would form a ketone. A primary, secondary, or tertiary alcohol may also

~ 7 1 .

PHOTO-ASSISTED PROCESSES 26 1

dehydrate to form an alkene, which could then be oxidatively cleaved to shorter chain ketones and aldehydes [161. These reactions are illustrated below for 2- methylbutane:

CH3 I CH3 I CH3

I CH3CH2CHzCH3 0 CH3CHCH2CH2OH 5 CH3CHCH2CHO

(14- 10) -H2

-H20 1 CH 3 I CH 3 I

CH3CHCH=CH2 CH3CHCHO + COz (14-11) H2d

CH 3 I CH3

I 0 ( 14- 12) CH3 I

CH3CHCH2CH3 0 CH3CHCHCH3 CH3CHCCH3 I

OH 6 -H20 J

CH 3 I

CH3CHCH=CH2 ::& CH3CHCHO + COz (14-13a) 2

+ CH 3 I 0 CH3C=CHCH3 CH3COCH3 + C02 (14-13b)

2

OH

HZ CH3COCH3 + CH3CH0

This mechanism has been substantiated by studying the photo-oxidation of the appropriate alcohols and alkenes. Oxidation of 3-methyl-1-butanol yields 52% 3- methylbutanal, 20% 2-methylpropanal and acetone: 3-methyl-2-butanol yields 39% ethanal, 34% acetone, 17% 3-methyl-2-butanone, and 10% 2-methylpropanal; and 2-methyl-2-butanol yields 50% acetone and 42% ethanal. Oxidation of 3- methyl- I-butene yields 87% acetone and small amounts of 2-methylpropanal, acrolein and ethanal; and 2-methyl-2-butene yields 35% acetone, 29% ethanal, and some unidentified products. These results support the mechanism illustrated above.

262

Another mechanism for the production of alcohols and ketones involves a peroxy intermediate. This is illustrated for the oxidation of propane [20]:

C3Hs + O-(ad) + 0 2 + C3H700. + OH-(ad) (14- 15)

The peroxy species then further reacts to form an aldehyde. In the photocatalytic oxidation of propane on silica-supported Mo03, it is

determined by varying the residence time that acetone, propanal, and propanols are primary products. As the residence time increases, the secondary products C02, CH3CH0 and methanol increase. The production rates of all primary products show a nearly first order dependence on light intensity, while the secondary products show an order dependence that is substantially larger than unity [19]. Since acetone appears as a primary product, its production cannot be via an alcohol intermediate as suggested earlier. The following mechanism has been proposed for Moo3:

C3Hs + 0-(ad) -+ C3H70+ + e-(lattice) (14- 16)

n-C3H70. + O2 -+ C2HsCHO + H02- (14- 17)

n-C3H70. + C3Hs -+ n-C3H70H + C3H7 (14-18)

i-C3H70. + O2 + (CH3)2CO + HO2. (14- 19)

i-C3H70- + C3H8 -+ (CH3)2CHOH + C3H7 (14-20)

Since a C-H bond is stronger than a C&H bond, it is not clear how efficient are reactions (14-18) and (14-20) for the production of propanols. It is possible that there are other routes to produce these alcohols.

Although most oxidation reactions produce a broad distribution of partial oxidation products, the oxidation of 2-methylpropane (isobutane) on anatase Ti02 produces almost exclusively acetone [21]. The following accounts for up to 95% of the reaction of 2-methylpropane:

CH3 I

CH3-CHXH3 + 5/2 0 2 + CH3COCH3 + C02 + 2 H20 (14-21)

Since the activity in the photocatalytic oxidation of hydrocarbons is proportional to the light intensity, and the reaction only proceeds in the presence of oxygen, it has been assumed that photo-adsorbed oxygen and photo-generated holes are both important. The detailed roles of these species, however, have not been unequivocally identified. The photo-adsorbed 0-(ad) species is often assumed to be the species that activates alkanes by the abstraction of a hydrogen atom [21]:

CnHzn+2 + 0-(ad) -+ .C,,H2"+, + OH-(ad) (14-22)


or 0-(ad) + h+ -+ O(ad) (14-23)

An alternate activation mechanism involves surface OH species that are formed when a surface hydroxyl group captures a hole:

OH-(ad) + h+ + .OH(ad) (14-25)

or 02-(ad) + .OH(ad) + 0-(ad) + 02H(ad) (14-27)

Reactions (14-27) and (14-28) have been suggested because the reaction requires gaseous oxygen and does not take place with N 2 0 as the oxidant [17]. This has been taken to indicate that 02-(ad) is necessary. The participation of OH groups, however, may not be important on TiO2. It has been shown that a highly dehydroxylated anatase sample photo-oxidizes 2-methylpropane as actively as a partially dehydroxylated sample [22].

On the other hand, using isotopically labeled oxygen, it has been observed on silica-supported vanadia that lattice oxygen participates in the photo-oxidation of propene to acetaldehyde [32]. Since the photo-activity of vanadia is relatively poor, and the participation of lattice oxygen in thermal oxidation reactions on vanadia is well established, it is not clear if this observation applies to titania or other oxides. The fact that lattice oxygen does participate in CO oxidation on titania suggests that this may be important in some cases [33].

The participation of photo-adsorbed oxygen is also supported by the fact that the activity in photocatalytic oxidation over various oxides can be broadly correlated with their photo-adsorption properties [10,18]. Thus TiOz and Zr02, which are very active in photo-assisted adsorption of oxygen, are also very active in photocatalytic oxidation of propene. V205, which does not demonstrate photo- assisted adsorption of oxygen, is quite inactive. The different selectivities on different oxides for partial oxidation products seem to parallel the trend in thermal oxidation. However, these correlations are at best qualitative.

Changes in oxide morphology after reaction have been reported. On silica- supported molybdena, isolated tetrahedrally coordinated Mo6+ ions crystallizes to Moo3 crystallites after being used in the photocatalytic oxidation of cyclohexane 1231. Similar crystallization has also been reported for silica-supported V2O5 [24]. It is interesting that V2O5 is inactive in the oxidation of propene, but V205/Si02 is active in the oxidation of cyclohexane.

Oxidation of Alcohols Photocatalytic oxidation of alcohols at room temperature in the presence of

oxygen proceeds readily on a number of oxides. For example, the oxidation of 2-

264

propanol on rutile Ti02 results almost exclusively in acetone [14,25]. A small amount of acetaldehyde is also produced which may be formed from the photo- oxidation of propene which is a dehydration product of 2-propanol. Continual exposure of the reaction mixture to illuminated Ti02 results in further oxidation of acetone to formic acid, and eventually C02 and water. However, the rate of oxidation of acetone is slow [171.

Other selective oxidation of alcohols have been reported. On Ti02 butanol is oxidized to butanal, and 2-butanol to butanone [26], although there is a report that acetaldehyde is also a significant product from 2-butanol [30]. Thus oxidative cleavage of the carbon chain may occur. In some cases, this is the predominant reaction. For example, 2-methyl-2-propanol is oxidized to acetone quite selectively [26], 3-methyl-2-butanol is oxidized to ethanal and acetone, and 2- methyl-2-butanol to acetone and ethanal. 3-Methylbutanol is oxidized primarily to 3-methylbutanal but there is a substantial production of 2-methylpropanal as well [16]. From the product distributions in the reactions of the appropriate alkenes, it has been suggested that oxidation of alcohol is the first reaction step for primary alcohols. For secondary and tertiary alcohols, the first step is dehydration to the corresponding alkenes, which are then oxidized by the photo-generated adsorbed oxygen 1271.

Photo-enhanced oxidative dehydrogenation of 2-butanol and 2-propanol to the corresponding ketones have also been observed on MnOz and ZnO [30,31]. The photo-activity declines with time. Illumination has no effect on V205. Cr2O3, Fe2O3. Co3O4. NiO, or CuO.

Photo-oxidation of methanol at 403 K on anatase Ti02 results in methyl formate up to about 60% conversion [281. It is proposed that an adsorbed methoxy is photo-oxidized to an adsorbed formate, which reacts with another methoxy to form methyl formate. The reaction rate at room temperture is very low, although formaldehyde is readily oxidized [22]. When the titania is impregnated with Moo3, the quantum efficiency decreases with increasing molybdena content to about 10% of the molybdena-free value when there is about a monolayer equivalent of molybdena. The rate of production of methyl formate is low, and the major product is dimethoxymethane. It is proposed that an adsorbed methoxy is first oxidized to an adsorbed formaldehyde which condenses with adsorbed methoxy groups to form dimethoxymethane [28]. The photo- oxidation of methanol to methyl formate has also been reported for ZnO [29].

Oxygen Isotope Exchange

over Ti02 and ZnO. Both the homophasic exchange reaction: Light-enhance oxygen isotope exchange reaction has been studied in detail

and the hcterophasic exchange reaction:

(14-29)

l8O2(g) + l6O(lattice) 4 l60l8O(g) + ls0(lattice) (14-30)


proceed readily at room temperature under band gap illumination. Exposure of oxygen to an illuminated Ti02 at 77 K results in the appearance

of an EPR signal assigned to 03-(ad). This species is thermally unstable and decomposes to adsorbed 0- and O2 on warming to 137 K, but its intensity can be correlated with the catalytic activity for homophasic exchange of oxygen over the sample at 137 K [34]. Thus this species has been proposed as a reaction intermediate in the homophasic exchange reaction which proceeds as:

l802(g) + l60-(ad) + 03Jad) -----) l60l8O(g) + l80-(ad) (14-31)

There is evidence that isotopic scrambling between adsorbed 0 2 (or 0 2 3

species does not contribute to this reaction. On a Ti02 sample previously exposed to 1802 and 1602, it has been found that equilibrium distribution of isotopic oxygen species in the gas phase can be obtained rapidly when the sample is illuminated. However, the isotopic composition of the adsorbed oxygen is very different from the gas phase composition at the end of the reaction [35]. In fact, it is similar to the composition of adsorbed oxygen from a sample that has not been used for the reaction. This has been found by thermally desorbing the adsorbed oxygen after reaction. Thus adsorbed O2 species are not involved in the homophasic exchange reaction.

Monatomic oxygen species are also involved in the heterophasic isotope exchange reaction on Ti02. When anatase Ti02 is exposed to l S 0 2 under illumination, 160180 first appears in the gas phase followed by 1602 in a manner characteristic of stepwise reactions: 1802 * l60l8O + 1602. Thus the reaction proceeds as [36]:

1802(g) + 160(ad) + l6OI8O(g) + lS0(ad) (14-32a)

160180(g) + 160(ad) + 1602(g) + I80(ad) (14-32b)

Both homophasic and heterophasic reactions occur simultaneously on an illuminated oxide [37,38]. The homophasic reaction proceeds faster than the hctcrophasic reaction on ZnO pretreated with oxygen or prereduced. This contrasts the thermal reaction for which the homophasic reaction is greatly suppressed by pretreating the oxide with oxygen. The extent of surface hydroxylation may be important also. The photocatalytic activity of a partially dehydroxylated TiOz is higher than a fully hydroxylated sample [21].

Other photo-oxidation reactions compete with the photocatalytic oxygen isotope exchange reaction. In the presence of 2-methylpropane, oxygen isotope exchange is completely suppressed on Ti02 until the oxidation of the hydrocarbon is complete [36]. On ZnO, CO oxidation suppresses the oxygen isotope exchange reaction [33,39]. These results suggest that these reactions either proceed on the same sites or involve a common intermediate. Indeed, on TiO2, the activity per unit area for oxygen isotope exchange correlates with that for 2-methylpropane oxidation [36], and the participation of 0-(ad) in both reactions has been suggested.

266

Other Reactions Many other photocatalytic reactions have been observed [2,40]. Oxidation

of CO [41] and NH3 [42], H2-D2 exchange [43], decomposition of NO [26] and N 2 0 [ a ] , and reduction of N2 [45] have been reported. In the oxidation of NH3 over Ti02, N2 and N20 are the products observed in addition to H20. From conductivity measurements, it appears that the species involved in this reaction, NH3, N2 and N20 do not compete for photo-generated electrons with adsorbed oxygen. The decomposition of NO produces N2 as the product at low NO pressures, and N 2 0 at high NO pressures. Adsorbed oxygen is the other product at room temperature.

The photo-enhanced H2-D2 exchange reaction has been analyzed in detail [2]. It is assumed to proceed by heterolytic dissociative adsorption of hydrogen.

Participation of lattice oxygen in CO oxidation has been demonstrated for Ti02 [33]. Using a mixture of CO and the photo-enhanced production of C1602 is much larger than C1601s0 at the beginning of the reaction.

Photo-reduction of N2 to form NH3 has been successfully demonstrated by passing N2 saturated with water vapor over illuminated W 0 3 [45]. No N2H4 is formed. If the tungsten oxide is slightly reduced before the reaction, it is reoxidized during the reaction. The reverse reaction, decomposition of ammonia, is also photocatalyzed. Photo-reduction of N2 also occurs on illuminated anatase TiO2. Ammonia and traces of N2H4 are formed [461. Rutile Ti02 doped with Fe is quite effective in this reaction.

14.4 PHOTOCATALYSIS: LIQUID PHASE REACTIONS

In gas phase photocatalytic reactions, charged reaction intermediates must remain adsorbed on the surface. On the other hand, charged species can readily exist in the liquid phase. Thus it becomes possible to remove an electron from an anion in the solution to form a neutral species by a photoelectrode, or inject an electron into a neutral species to form an ion. The resulting species could then further react in solution.

Oxidation Reactions The oxidation of oxalate and formate ions to C 0 2 by an illuminated ZnO in an

aqueous suspension has long been reported [471. Hydrogen peroxide is a byproduct:

C2Od2- + 0 2 + 2 H2O ---+ 2 C02 + 2 OH- + H202 (14-33)

HCOO- + 0 2 + H20 - C02 + OH- + H202 (14-34)

Under similar conditions, 2-propanol is oxidized to acetone:

(CH3)ZCHOH + 0 2 (CHj)ZCO + H202 (14-35)


The oxidative decarboxylation of acetate ions can be effected with illuminated mile TiOz. In the absence of oxygen, ethane is a product which is formed by coupling of two methyl radicals [48]:

CH3COz- + h+ - CH3 + COz (14-36)

2 C H 3 - Cz& (14-37)

The production of ethane is greatly suppressed by the presence of oxygen. Presumably the methyl radicals react to form methanol and formaldehyde [49].

Photo-assisted oxidation has also been demonstrated at a TiO, electrode for hydroquinone, p-aminophenol, I-, Br-, C1-, Fez+, Ce3+, and CN- ions. Many of these species are oxidized at potentials negative of their standard redox potential P O I .

It is interesting to note that in the photo-assisted decarboxylation of a dicarboxylate on TiOz, the reaction stops at the monoacid form. Perhaps this is due to the displacement of the monoacid from the surface by the diacid [Sl]:

C02H T i O z , hu + COz (14-38)

O C 0 2 H -- C H 3 C N 0- COzH

Photo-assisted oxidation of hydrocarbons has been studied rather extensively [511. In an aqueous suspension of TiOz, substituted toluene reacts to form both oxidation and coupling products [55,56]:

(14-39)

R

This contrasts sharply the relatively selective production of substituted benzaldehyde in the gas phase reaction [52]. Cyclohexene is likewise oxidized to a mixture of cyclohexenol and cyclohexenone [51].

Oxidative cleavage of 1 ,l-diphenylethylene to benzophenone can be achieved efficiently with TiOz [53]:

(14-40)

268

Substituted naphthalene is photo-oxidized with cleavage on one of the benzene rings. For example [51]:

(14-41)

Amines can also be photo-oxidized. On ZnO, aniline is oxidized to azobenzene, and toluidines are oxidized to the corresponding azo compounds. The relative rates depend on the substituent, and the order meta > ortho > para- toluidine > aniline has been observed [54].

The mechanism of liquid phase photo-assisted oxidation of hydrocarbons with TiO, has been investigated in detail [51]. Several observations support the involvement of a radical cation intermediate:

RH + h+ ---+ RH+ (14-42)

Stilbene radical cation is observed spectroscopically after a Ti02 colloidal suspension in acetonitrile containing trans-stilbene has been flashed with a nitrogen laser [55]. Efficient photo-conversion of diphenylelhylene is achieved only in solvents of high dielectric constants. This is consistent with the involvement of a highly polar transition state. When para-substituted diphenylethylene is oxidized with Ti02, a linear Hammett plot with a negative slope is obtained between the relative rate of reaction and the substitution parameter 1561. The relative rates of reaction of other substituted diphenylcthylene molecules also show reduced reactivity with decreased K- electron density of the double bond. Thus these relative rates reflect the stability of the one-electron oxidation product, consistent with the proposed involvement of radical cations.

Subsequent reactions of the radical cations may involve capture by triplet molecular oxygen. They may also involve attack by superoxide. In either case, the reactions are not photo-assisted. If the hydrocarbon molecule is an alkene, dioxetane is then formed which would be cleaved to a ketone or an aldehyde [51].

The oxidation of water has been studied extensively. Earlier work has been on the formation of hydrogen peroxide 12,471. The reaction mechanism can be either:

H20 + 0-(ad) --+ HOO-(ad) + H(ad) (14-43)

HOO-(ad) + H+(ad) + H202 (14-44)

or H20 + 0-(ad) + OH-(ad) + OH(ad) (14-45)

OH-(ad) + OH(ad) ---+ Hz02 + e-(lattice) (14-46)


Adsorbed electron acceptors retard the reaction. Thus adsorbed 0, [57] or adsorbed bicarbonate [58] slow the reaction.

Recently, the interest has been shifted to the photo-assisted electrochemical oxidation of water to oxygen. It has been discovered that band gap illumination of a Ti02 anode substantially reduces the applied potential required for the evolution of oxygen, as shown in Fig. 14-3 [59]. Many other oxides have been subsequently shown to demonstrate a similar effect [60,61,68,69]. They include W03 [62], n- Fez03 [63,64], and mixed oxides such as SrTi03 [65], KTa03 [66] and others. In the absence of other more readily reducible species, proton is reduced to hydrogen at the cathode when a sufficiently high potential is applied. Thus the net reaction is the decomposition of water:

2 H20 + 2 H2 + 0 2 (14-47)

With the more efficient photo-anodes such as those listed above, the external potential needed to achieve this reaction is substantially less than the thermodynamic value of 1.23 V when the photo-anode is illuminated. Therefore, such a system serves to convert light energy into chemical energy. However, only the large band gap oxides such as SrTi03 and KTa03 can sustain the reaction without any applied external potential. In fact, it appears that the magnitude of the applied potential is roughly inversely correlated with the width of the band gap [611.

The oxidation reaction at the anode may involve the photo-production of OH radicals which form H202 [671:

Hydrogen peroxide decomposes to water and oxygen in a basic solution. The reduced titania is reoxidized by water to regenerate the hydroxylated surface. In fact, if deuterated water is used in the reoxidation, D, is evolved 1671:

Reduction Reactions Photo-assisted reduction of metal complexes to metallic particles deposited

on the oxide can be achieved in a number of systems. The reduction is accompanied by oxidation of another species. For example, copper ions can be reduced to copper metal deposited on illuminated Ti02 or W03 particles in water and oxygen is produced [701.

270

cu I I E 0

+ c 1.0 2 L 3 0 0 .- U 2 0 a

I V(SCE) I -0.5 1

Figure 14-3 Anodic current as a function of voltage (versus SCE) at a semiconducting anode in a pH 4.7 electrolyte. The potential for the evolution of oxygen from the oxidation of water is lowered by illumination. a. Anode illuminated with strong light; b. with weak light; c. No illumination. (After Nature, 238, 37 (1972), copyright Mcmillan Publisher).

T i 0 2 , hu Cu2+ + H20 -- CU + 1/2 0 2 + 2 H+ (14-50)

In the presence of acetate, acetate is oxidized to C02:

Cu2+ + 2CH3COO- Cu + C2& + 2C02 (14-51)

Copper can also be deposited on KTa03 and SrTi03 [71]. The deposition of Pt from a solution of hexachloroplatinic acid and carbonate onto irradiated Ti02 has also been achieved. The carbonate is converted to C02 [72]. Other platinum complexes can likewise be used [73]. The reduction of chloroplatinic ion is accompanied by a release of proton that has been quantitatively detected. In the reduction of dinitrodiaminoplatinum, N20 and N2 are evolved. Transmission electron microscopy shows that platinum is initially deposited as very small particles distributed over the Ti02 particles. Deposition of other metals, including Pd, Ag, Rh. Au, and Ir on Ti02 and on other oxides including ZnO, Nb2OS and Tho2 have been reported [73, and references therein]. The rate of deposition depends on the strength of adsorption of the metal ion and the quantum efficiency for reduction. For the reduction of Ag+- the adsorption of the ion on ZnO is weaker than on Ti02, but the quantum efficiency for reduction is higher. The relative rates of deposition on the two oxides also depend on the concentration of the ions in the solution [74].

There are reports on the photo-assisted reduction of water by oxides. p-Type

PHOTO-ASSISTED PROCESSES 27 1

Fez03 has been successfully used with n-type F e 0 3 to decompose water. Water is reduced to H2 at the p-type iron oxide [75]. p-LuRh03 has also been reported as effective in this reduction [76].

14.5 PHOTOCATALYSIS BY METALMETAL OXIDE AND OXIDE/OXIDE COMPOSITES

It can be Seen from the examples described above that some oxide particles or electrodes are quite efficient in effecting photo-assisted electron transfer. For an n-type semiconducting oxide, this has resulted in enhancement in many oxidation reactions. However, for an effective photocatalyst, every step of the catalytic cycle must proceed efficiently. This requirement may not be met by many oxides. For example, photo-oxidation on Ti02 is commonly observed; but photo-reduction on this oxide is much less known. Thus electron transfer out of Ti02 is more difficult than into the oxide. Therefore, attempts have been made to deposit metal or another oxide on a semiconducting oxide that will facilitate electron transfer in the difficult direction. Noble metals and RuO2 have been used as deposits for this purpose.

Sometimes a composite of n- and p-type semiconducting oxides are used when a single oxide does not generate sufficient photoelectrochemical driving force to complete a desired reaction. Then a composite of two appropriate semiconducting oxides could be used to supply a combined photoelectrochemical driving force. Successful decomposition of water using visible light has been achieved using such devices.

Such composite systems have been used in the oxidation of ethanol. Ethanol is oxidized in a water mixture by an irradiated Wi02 to acetaldehyde and acetic acid. Hydrogen is also evolved [77]:

CH3CH20H hu, CH3CH0 + H2

CH3CHO + H20 -!% CH3COOH + H2

(14-52)

(14-53)

After prolonged irradiation, methane and C02 are observed due to the decarboxylation reaction of acetic acid [78].

CH3COOH C& + C02 (14-54)

A side reaction may also occur via the reaction of acetaldehyde with OH radical, which eventually leads to combustion:

CH3CHO + 3 H2O 5 HzO + 2 C02 (14-55)

Photo-assisted decarboxylation of acids takes place efficiently on Wanatase Ti02 in the absence of oxygen. Wanatase is more effective in this reaction than Wrutile, possibly because anatase has a slightly larger band gap than rutile. The

272

Table 14-3 Products in the Photo-assisted Decarboxylation of Acids on m i 0 2

Acid Major Products Minor Products

acetic cH4, co2 C2&. H2 propionic C2&. co2 c2H4, H2 n-butyric propane, C02 H2 n-valeric n-butane, C02 H2 pivalic isobutane. C02 isobutene. H2 adamantane- 1 - adamantane, C02

adipic butane, C02 valeric acid carbox ylic

products of the reaction are listed in Table 14-3 [78,8S]. These products can be accounted for by a general reaction scheme:

h+ + RC02- Re + CO2 (14-56)

e- + RCOOH - H(ad) + RCOO-

e- + R. + RCOOH - RH + RCOO-

2 H(ad) - H2

R. + H(ad) - RH

It is of significance to note that H2 is a product in these reactions because Pt is an excellent catalyst for the recombination of H atoms. Pd or Ru02 can be used instead of R, but the efficiency in H2 evolution follows the order: W i 0 2 2 Pd/Ti02 > RuOJI’iO2 >> Ti02 1791.

Complete decarboxylation of adipic acid, HOOC(CH3)4COOH, to C4H1 0 can also be achieved with illuminated m i 0 2 [85]. Oxidation of alkanes, including hexane, cyclohexane, heptane, nonane, and decane in an aqueous solution results in practically complete combustion [86].

The photo-decomposition of water has been studied extensively. Pt/SrTi03 is an efficient catalyst [80], more efficient than Wrutile Ti02. This is because SrTi03 has a smaller electron affinity and therefore a larger band bending at the surface [81]. For the Pt/TiO2 system, anatase is more efficient than mile. A small amount of Nb205 doping further enhances the efficiency [82]. Platinked


KTa03 also decomposes water on illumination [711. Oxide/oxide composites have been used for this reaction. A NiO/SrTi03

composite, prepared by impregnation with Ni(NO3)2 followed by reduction and reoxidation at 500°C decomposes water effectively upon irradiation. Nearly stoichiometric amounts of H2 and O2 are produced [83]. The reduction step is necessary to produce metallic Ni to form a NiO/Ni/SrTi03 interface in which Ni serves as an ohmic contact. The activity in NaOH is higher than in water, perhaps because hydrogen peroxide decomposes more efficiently in basic than neutral solutions.

Another system reported to be successful is a n-Fe203/pF%03 composite [75,84]. p-Fe203 is made by Mg doping. The actual material in the composite is a mixture of MgO, p-Fe203 and MgFe204. The activity of the system declines slowly with time. It can be regenerated by oxygen treatment of the oxide.

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84. 3207 (1980).

INDEX

Acetone, adsorption, 66 Acid sites, formation, 74, 75-80 Acidity, 72 -, determination, 80, 83, 87 -, effect of pretreatment, 82 -, selective oxidation, 175 -, strength, 74. 78 Acrolein, from propene. 181 Acrylonitrile, from propene, 18 1 Adsorbed oxygen, 110 -, desorption, 11 1 -, detection, 11 1

-, photo-adsorption, 256 -, photo-desorption, 256 -, reactivity, 116, 118 -, selective oxidation, 176 -, superoxide, 114

Adsorption, 61 -, abstractive, 62 -I and cus, 57 -, heterolytic dissociative, 3, 35, 57, 61 -, homolytic dissociative, 61 -, molecular, 57, 61 -, pretreatment effect, 63 -, reductive, 62 AES, 4 Aerogel, preparation, 124

-, IR, 114

-, types, 110

Alcohol, adsorption, 66 -, decomposition. 146 -, photo-oxidation, 261, 263 Aldol condensation, in isosynthesis, 243 Alkane, activation, 178 -, deuterium exchange, 161 -, photo-oxidation, 259 -, reaction with adsorbed oxygen, 116 Alkene, activation, 178 -. adsorption, 64 -, deuterium exchange, 161 -, hydrogenation, 156 -, reaction with adsorbed oxygen, 118 Alkyne, adsorption, 65 Alumina, isoelecmc point. 85 Ammonia, adsorption, 81 -, photo-oxidation, 266 Anion vacancies, surface, 55-57 Anisotropy constant, 43 Antiferromagnetism, 35 --, crystallite-size effect, 39 -, temperature effect, 39 Antimonoy-tin oxide, propene oxidation,

181

Band gap. 252 Bi-Fe-Mo-0, propene oxidation, 186 Bismuth molybdate, 189, 204 -, butene oxidation, 173, 204

277

278

-, propene oxidation, 173, 181. 182.

Brdnsted acid, 72 -, formation, 75 Butane oxidation, 210 Butadiene, from butane, 210 -, from butene, 200 Butene, adsorption, 65 .-, dehydrogenation, 201, 203 -, hydrogenation, 159 -, oxidation, 88, 189, 200

185, 187, 188

Carbon-oxygen bond, formation, 179 Carbanion, in isomerization, 143 Carbenium ion, in isomerization, 142 Carbonate, adsorbed, 62 Carboxylate, adsorbed, 62 Catalytic processes, 2 Chromia (see chromium oxide) Chromium-iron oxide, 209, 244 Chromium oxide, adsorbed oxygen, 116 -, butene hydrogenation, 160 -, butene isomerization. 141, 144 -, CO adsorption, 58, 61 -, dehydroxylation, 55, 58 -, deuterium exchange, 161 -, methanol decomposition, 149 -, NO reduction, 165 -, preparation, 126 -, propene hydrogenation, 159 CO, adsorption, 58, 67 -, copper catalysts, 228, 230, 232 -, hydrogenation, 228 -, in C02 hydrogenation, 228 -, iron catalyst, 239 -, Pd catalyst, 236 -, Rh catalyst, 237 Cobalt molybdate. propene oxidation,

Cobalt oxide, ethene hydrogenation, 158 -, isomerization, 143 Conduction band, 252 Contracting sphere model, 95 COZ, hydrogenation, 228 Coordinative unsaturation, 53 .-, conjugate base, 72

188, 192

-, effect on structure, 15 -, formation, 53. 55 -. Lewis acid, 72 Copper oxide, methanol decomposition,

-, NO reduction, 162 -, propanol decomposition, 151 -, reduction, 98 Copper-thorium oxide, 230 Copper-zinc oxide, 228, 232, 245 Corundum structure, 10 Cubic close-packing, 8 Curie temperature, 39 Cuprous oxide, propene oxidation, 182.

Cus (see coordinative unsaturation)

149

185

Decarboxylation, 267, 271 Decomposition, methanol, 146 -, water, 269, 272 Decoration effect. 102 Dehydrogenation, butane, 210 -, butene. 200 Dehydroxylation, 54 -, C r 2 4 , 58 -, zno, 58 Diamagnetism, 36 Dimerization, 215 Dissociative adsorption, 3, 35

Electronic structure, bulk, 252 -, surface, 46 ELS, 4 EPR, 4 Ethane, from carbon oxides, 235 -, from methane, 215 Ethanol, from carbon oxides, 235, 238 -. photo-oxidation, 271 Ethene. from carbon oxides, 235 -, from methane, 215 -, hydrogenation, 156 EXAFS. 4

Fez& (see iron oxide) Fermi energy, 252 Femmagnetism. 38

INDEX 279

Ferrites (see iron oxide) Formaldehyde, from methane, 212 -, from methanol, 217 Formate, 62, 147, 245 m, 4

Hexagonal close-packing, 8 HREELS, 4 Hydrogerb adsorption, 63 -, reaction. adsorbed oxygen, 11 8 Hydrogen-deuterium scrambling, 161 Hydrogenation, 155 Hydroperoxide, in propene oxidation, 186 Hydroxyl, Brdnsted acid, 72 -, in photo-assisted processes, 257

Impregnation, 130 Indicators, 87 Ion-exchange, in preparation, 129, 132 Ionic charge, effect on structure, 15 Ionicity, determination, 28 -, Pauling's, 33 -, Phillip's, 33, 34 -, values, 30 IR, 4 Iron-molybdate, methanol oxidation, 218,

Iron oxide, alcohol decomposition, 152 -, anisotropy constant, 43 -, butene dehydrogenation, 201, 203, 204 -, CO hydrogenation, 239 -, lattice dilation, 41 -, magnetic hyperfine field, 42 -, magnetic properties, 39 -, Morin transition, 41, 42 -, oxygen desorption, 11 1 -, preparation, 126, 131 -, quadrupole splitting, 43, 46 -, reduction, 98 -, structure, 10 -, water decomposition, 273 -, water-gas shift, 246 Isoelectric point. 83 Isomerization, akenes, 140 -, carbenium ion, 142 -, effect of reduction, 141

220

-, metathesis, 141 Isosynthesis, 240 ISS, 4

Lanthanum oxide, methane coupling, 217 Lattice oxygen, in selective oxidation,

170, 177, 190, 211 -, in photo-oxidation, 263 Lattice self-potential, 29, 32 LEED, 4 Lewis acid, 72 -, formation, 75

Magnetic properties, 36, 38 Magnetic hyperfine field, 4 1 4 3 Magnesium oxide, methane coupling, 216 Maleic anhydride, from butane, 210 -. from butene, 200 Manganese oxide, methane coupling, 217 Metathesis, 136 -, degenerative, 137 -, isomerization, 141 -. mechanism, 137 -. productive, 137 -, substituent effect, 138 -, support effect, 140 Methane, dimerization, 215 -, oxidation, 212 Methanol, decomposition, 146 -, from carbon oxides, 228, 231 -, from methane, 212 -, oxidation, 217 -, photo-oxidation, 264 Methyl formate, from carbon oxides, 235 Methyl radical, 116. 216 Mixed oxides, acidity, 75 -, alkane reaction, 118 -, reduction, 100 Molybdenum-nickel oxide, butene

oxidation, 202, 209 Molybdenum oxide, adsorbed oxygen, 11 6 -, alcohol adsorption, 66 -, alkane reaction, 117, 118 -, alkene isomerization, 141. 143 -, butene oxidation, 204 -, ethanol decomposition, 152

-, isoelectric point, 85 -, methane oxidation, 213 -, methanol oxidation, 218, 221 -, metathesis, 138, 139 -, photo-oxidation 260 -, preparation, 124, 131 -, propene oxidation, 189, 192. 194 -, reduction, 96 -, shear structure, 12 -, structure, 10, 22 Molybdenum-tin oxide, propene oxidation,

173 Morin transition temperature. 41 Multicomponent oxidation catalyst, 187

Ne61 temperature, 38 Nickel oxide, acetone adsorption, 66 -, reduction, 95 -, structure, 6, 24 -, support effect, 99 -, surface phonon. 48 -, water decomposition, 273 Niobium oxide, ethanol decomposition,

-, reduction, 98 Nitric oxide, adsorption, 67 -, reduction, 162 Nitrous oxide. decomposition, 153 NMR. 4 Nonpolar surface, 15 Nucleation model, 93

152

Oxidation, photo-assisted, 259 -, selective (see selective oxidation) Oxide, properties, 3 -, structure, 6 Oxide catalysts, 2 Oxygen, adsorbed (see adsorbed oxygen) -, isotope exchange, 264 --, on chromia, 58, 11 6 -, on iron oxide, 11 1 -, on zinc oxide, 114, 116 -, photo-adsorption, 256 -, photo-desorption, 256

Palladium, CO hydrogenation, 236

Paramagnetism, 38 Perovskite. structure. 12 Phonon, 48 -, and ionicity, 35 Photo-assisted processes, 256 -. adsorption, 256 -, desorption. 256 -, oxidation, 259, 263 -, reduction. 269 Photoelectrons, 254 Photo-holes, 254

-, radicals, 186 Polar surface, 15 Preparation, 121 -, complexation, 132 -, coprecipitation, 133 -, crystallographic phase, 124 -, freeze-drying. 132 -, impregnation, 130 -, ionexchange, 129. 132 -, morphology, 125 -, pH effect, 130 -, special techniques, 130 -, spraydrying, 132 Propene. adsorption, 64 -, ammoxidation, 181 -, hydrogenation, 158 -, metathesis, 137 Propene oxidation, 173. 181 -, kinetics, 185, 186 -, mechanism, 182 -, promoters, 193 -, water effect, 194 Promoters, butene oxidation, 209 -, COX hydrogenation, 233 -, propene oxidation, 193 Propanol decomposition, 149 Pyridine. adsorption, 8 1

n-allyl, 64. 178, 182

Quadruple splitting, 43, 46

Reconstruction, surfaces, 15 Reduced oxide, effect on metal, 102 -, properties, 100 Reduction, 91

28 1

-, bronze formation, 100 -, kinetics, 93 -, mechanism, 96 -, mixed oxides, 100 -, support effect, 98 -, thermodynamics, 92 Relaxation, surfaces, 18 Rhodium, CO hydrogenation, 237 Rhodium oxide, support effect in

reduction of, 99 Rocksalt structure, 6 Rutile structure, 8, 10

Scheelites, in oxidation, 189, 203 Selective oxidation, 169 -, oxygen source, 173 -, redox mechanism, 170 -, selectivity in, 175 -, types, 170 -, water effect. 174 SEM, 4 Shear structure, 12 -, in selective oxidation, 175, 179 Solid solutions, surface composition, 28 Spinel structure, 8, 10 STEM, 4 Strontium titanate. surface phonon, 49 Structure, supported oxide. 21 Structure sensitivity, alcohol

decomposition, 148 Supported oxide, preparation, 129 -, structure, 21 Surface, anion vacancies, 55-57 --, Composition, 27 -, electronic structure, 46 -, polar and nonpolar, 15 -. reconstruction, 15 -, relaxation, 18 -, structure, 15 Surface vibration, 48

TEM, 4 Thorium oxide, isosynthesis, 240 Til03, surface structure, 55 Titanium oxide, acidity, 75 --, adsorbed oxygen, 114

-, alcohol, decomposition, 152 -, -, photo-oxidation, 261, 264 -, alkane, photo-oxidation, 260 -. ammonia oxidation, 266 -, decarboxylation, 267, 271 -, methanol decomposition, 148 -, 0, isotope exchange, 265 -, photo-adsorption. 257 -. photo-desorption. 257 -, photo-oxidation, 267 -, photo-reduction, 269 -, preparation, 124, 131 -, propanol decomposition, 151 -, reduced. decoration effect, 102

-, reduction, 96 -, structure, 10 -, surface electronic structure, 47 -. surface structure. 18, 55 -. water decomposition, 269, 272 -, water oxidation, 268 Titration, acidity, 87 TPD, 4 TPR, 4 Tungsten oxide, alkane, photo-oxidation,

-, alkene isomerization, 143 -, reduction, support effect, 98 -, supported, structure, 24

_ - , , reactivity, 101

260

UPS, 4 uv-vis. 4

Vacancies, in selective oxidation, 175, 189 -, surface, 55-57 Valence band, 252 Vanadates, butane oxidation, 21 1 -. butene oxidation, 205 Vanadia (see vanadium oxide) Vanadium oxide, isoelectric point. 85 -, methane oxidation, 215 -. methanol oxidation, 223 -, NO reduction, 162 .-, propanol decomposition. 151 -, propene oxidation, 195 -, reduction, support effect, 99

282

- structure, 10 - supported, structure, 21

- butane oxidation, 21 1 - butene oxidation, 202, 203 ~ propene oxidation, 195

V-P-0, 206

Water, adsorption, 57 - in COX hydrogenation, 230 - in selective oxidation, 174 - photodecomposition, 269 - photo-oxidation, 269 Water-gas shift, 244 WO, (see tungsten oxide) Wurtzite structure, 8

XANES. 4 XPS, 4

Zinc oxide, acetone adsorption, 66, 150 -, acidity, 75 -, adsorbed oxygen, 114, 11 6 -, alcohol, adsorption, 66 -, -, decomposition, 152 -, alkane reaction, 116 -, alkene isomerization. 144 -, butene, adsorption, 65 -, -, hydrogenation, 160 -, CO adsorption, 60, 67 -, dehydroxylation, 58 -, deuterium exchange, 161 -, ethene hydrogenation, 158 -, hydrogen adsorption, 63 -, methanol decomposition, 146 -, O2 isotope exchange, 265 --, photo-adsorption, 257 -, photo-desorption, 257 -, photo-oxidation, 266, 268 -, preparation, 125 -, propanol decomposition, 149 -, propene adsorption, 64 -, propene hydrogenation, 158 --, surface electronic smcture, 47 -, surface phonon, 48 -, surface structure, 18

zirconium oxide, isosynthesis. 240 -. preparation, 124

283

STUDIES IN SURFACE SCIENCE AND CATALYSIS

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Preparation of Catalysts I. Scientific Bases for the Preparation of Heterogeneous Catalysts. Proceedings of the First International Symposium, Brussels, October 14- 17, 1975 edited by B. Delmon, P.A. Jacobs and G. Poncelet The Control of the Reactivity of Solids. A Critical Survey of the Factors that Influence the Reactivity of Solids, with Special Emphasis on the Control of the Chemical Processes in Relation to Practical Applications by V.V. Boldyrev, M. Bulens and B. Delmon Preparation of Catalysts II. Scientific Bases for the Preparation of Heterogeneous Catalysts. Proceedings of the Second International Symposium, Louvain-la-Neuve, September 4-7, 1978 edited by B. Delmon, P. Grange, P. Jacobs and G. Poncelet Growth and Properties of Metal Clusters. Applications t o Catalysis and the Photographic Process. Proceedings of the 32nd International Meeting of the Societe de Chimie Physique, Villeurbanne, September 24-28, 1979 edited by J. Bourdon Catalysis by Zeolites. Proceedings of an International Symposium, Ecully (Lyon), September 9- 1 1, 1980 edited by B. Imelik, C. Naccache, V. Ben Taarit, J.C. Vedrine, G. Coudurier and H . Praliaud Catalyst Deactivation. Proceedings of an International Symposium, Antwerp, October 13-15, 1980 edited by B. Delmon and G.F. Froment New Horizons in Catalysis. Proceedings of the 7th International Congress on Catalysis, Tokyo, June 30-July 4, 1980. Parts A and B edited by T. Seiyama and K. Tanabe Catalysis by Supported Complexes by Vu.1. Yermakov, B.N. Kuznetsov and V.A. Zakharov Physics of Solid Surfaces. Proceedings of a Symposium, Bechyhe, September 29- October 3, 1980 edited by M. LazniEka Adsorption at the Gas-Solid and Liquid-Solid Interface. Proceedings of an International Symposium, Aix-en-Provence, September 21-23, 198 1 edited by J. Rouquerol and K.S.W. Sing Metal-Support and Metal-Additive Effects in Catalysis. Proceedings of an International Symposium, Ecully (Lyon), September 14-1 6, 1982 edited by B. Imelik, C. Naccache, G. Coudurier, H . Praliaud, P. Meriaudeau, P. Gallezot, G.A. Martin and J.C. Vedrine Metal Microstructures in Zeolites. Preparation - Properties - Applications. Proceedings of a Workshop, Bremen, September 22-24, 1982 edited by P.A. Jacobs, N.I. Jaeger, P. JirG and G. Schulz-Ekloff Adsorption on Metal Surfaces. An Integrated Approach edited by J. Benard Vibrations at Surfaces. Proceedings of the Third International Conference, Asilomar, CA, September 1-4, 1982 edited by C.R. Brundle and H . Morawitz

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