fundamental study of noble metal and transition metal …

The Pennsylvania State University

The Graduate School

Department of Chemistry

FUNDAMENTAL STUDY OF NOBLE METAL AND TRANSITION METAL

OXIDE CLUSTERS IN THE PRESENCE OF CO:

INSIGHT INTO MECHANISMS OF HETEROGENEOUS CATALYSIS

A Thesis in

Chemistry

by

Nelly Moore Reilly

© 2007 Nelly Moore Reilly

Submitted in Partial Fulfillment of the Requirements

for the Degree of

Doctor of Philosophy

May 2007

The thesis of Nelly Moore Reilly was reviewed and approved* by the following:

A. Welford Castleman, Jr. Eberly Distinguished Chair in Science Evan Pugh Professor of Chemistry and Physics Thesis Advisor Chair of Committee

James B. Anderson Evan Pugh Professor of Chemisty and Physics

Nicholas Winograd Evan Pugh Professor of Chemistry

Robert J. Santoro George L. Guillet Professor of Mechanical Engineering

Ayusman Sen Professor of Chemistry Head of the Department of Chemistry

*Signatures are on file in the Graduate School

iii

ABSTRACT

The conversion of carbon monoxide (CO) to carbon dioxide (CO2) is one of the

most important catalytic reactions. Fundamental studies aimed at providing information

to aid in the design of more efficient and selective catalysts for the abatement of this

harmful environmental pollutant were conducted. Gas phase metal oxide cluster studies

are a valuable complementary technique to surface methods for elucidating the

mechanistic details and active sites of catalytic systems. To this end, gas phase studies

were utilized to uncover possible species responsible for the increased activity and

selectivity of noble metal and transition metal nanoparticle catalysts. A guided ion beam

mass spectrometer was employed to study the dissociation patterns and interactions with

CO of mass selected clusters. This provides information into specific reaction pathways

based on cluster size, ionic charge and oxidation state, stoichiometry and coordinative

saturation.

There has been recent interest in the activity of gold nanoparticles toward the

oxidation of CO and studies herein have focused on the effects of charge state in

elucidating the mechanism. Most reactions of gold oxide cations underwent oxygen

replacement by CO, however evidence of other reaction pathways including CO

oxidation for both Au4O2+ and Au4O3

+ species were also observed. Compared to anionic

gold oxide clusters of the same stoichiometry, CO oxidation was previously found to be

the dominant reaction channel. Anionic clusters were governed by a structure-reactivity

relationship in which a peripheral oxygen atom was the most efficient reaction center for

effecting CO oxidation. Density functional calculations found molecularly bound oxygen

iv

in the ground state structure of cationic clusters in which oxygen atom replacement by

CO was observed in experiments. High barriers to O-O bond dissociation suggest the

role of ionic charge may be an important factor governing the reaction. A mechanism

based on a triplet-singlet spin state transition was proposed that may provide a pathway

with lower barriers for O2 dissociation in order to explain the experimentally observed

reaction products.

Next, studies involving mixed metal oxide clusters were undertaken to probe to

effects of catalyst-dopant interactions and gain a deeper understanding of the principles

governing the oxidation of CO. Comparison of the reactions of pure metal oxides with

CO to experiments of bimetallic oxides can aid in determining the influence that each

individual metal exhibits and the changes that occur in reactivity when another metal is

present. Studies involving silver-gold dimer oxide clusters were conducted and provide

insight into the role of charge transfer between different metals. AgAuO1,2+ species

showed similar reaction products as pure gold oxide cationic clusters, while AgAuO3,4+

reacted similar to pure silver oxide clusters with the same stoichiometry. Only AgAuO2-

exhibited oxygen atom transfer which was different from reactions of either pure metal

oxide. This provides experimental evidence of density functional calculations that

previously predicted the reaction of AgAu- with O2 and CO promotes CO oxidation.

Studies concerning 3d transition metals were conducted to investigate their

activity as common catalytic support materials and their ability toward direct oxidation of

CO. Structural elucidation was conducted through CID experiments and density

functional calculations were performed for iron oxide ionic clusters. Reactions of iron

v

oxide cations with CO were compared to anionic clusters. Both ionic charge states were

shown promote CO oxidation; however, attachment of CO onto the metal cluster was

only observed for cationic clusters. Calculated energy profiles of the reaction

demonstrated that the oxidation efficiency was governed by the strength of oxygen

binding to the iron atom. The results provide important insight into the nature of the

active site for condensed phase catalysis, shedding light on the role of charge state for CO

oxidation in the presence of iron oxide clusters.

The behavior of several other 3d transition metal oxide clusters, cobalt, nickel and

copper, were undertaken to examine possible periodic trends. The dissociation patterns

and reactivity in the presence of CO for these metal oxides were similar to iron oxides for

each respective charge state. The most active stoichiometry for CO oxidation was MxOy+

where y=x for cationic clusters and MxOy- where y=x+1 for anionic clusters with the

exception of Cu3O3-. The same number of Cu atoms and oxygen atoms was a particularly

active stoichiometry. This may be a consequence of the electron density for copper,

which possesses a full orbital shell. Comparison of cations and anions showed that

cations were more efficient in their particular reaction channels requiring less CO

reactant gas to observe reaction products.

vi

TABLE OF CONTENTS

LIST OF FIGURES ................................................................................................. ix

LIST OF TABLES...................................................................................................xiii

ACKNOWLEDGEMENTS......................................................................................xiv

Chapter 1 Introduction .............................................................................................1

1.1 The Oxidation of Carbon Monoxide ............................................................1 1.2 Clusters as Models.......................................................................................2 1.3 Insight into Alternative Catalytic Material ...................................................4

1.3.1 Gold Oxides ......................................................................................5 1.3.2 Bimetallic Oxides ..............................................................................5 1.3.3 Transition Metal Oxides ....................................................................6

1.4 Synopsis......................................................................................................7 References ........................................................................................................9

Chapter 2 Experimental Setup ..................................................................................14

2.1 Introduction.................................................................................................14 2.2 Vacuum system ...........................................................................................14 2.3 Laser Vaporization Source...........................................................................17 2.4 Guided Ion Beam Setup...............................................................................20 2.5 Detection and Data Acquisition ...................................................................23 References ........................................................................................................25

Chapter 3 Comprehensive Study of the Reactivity of Gold Oxide Cluster Cations (AuxOy

+; x=1-4, y=1-5) in the Presence of Carbon Monoxide............................26

3.1 Introduction.................................................................................................26 3.2 Experimental Methods.................................................................................28 3.3 Results ........................................................................................................28 3.4 Discussion...................................................................................................32

3.4.1 CO Replacement Reactions ...............................................................32 3.4.2 Au Fragmentation/Replacement.........................................................42 3.4.3 Oxidation...........................................................................................44

3.5 Comparison to Anionic Gold Oxides ...........................................................45 3.6 Conclusions.................................................................................................46 References ........................................................................................................47

Chapter 4 Reactivity of Ag2Oy+,- and Bimetallic AgAuOy

+,- (y=1-4) Ionic Clusters with CO ............................................................................................................50

vii

4.1 Introduction.................................................................................................50 4.2 Experimental ...............................................................................................52 4.3 Results ........................................................................................................52

4.3.1 Pure Silver Oxide Ions.......................................................................53 4.3.2 Bimetallic Ag-Au Oxide Ions ............................................................56

4.4 Discussion...................................................................................................62 4.5 Conclusion ..................................................................................................64 References ........................................................................................................65

Chapter 5 Influence of Charge State on the Reaction of FeO3+,- with Carbon

Monoxide..........................................................................................................67

5.1 Introduction.................................................................................................67 5.2 Methods ......................................................................................................68 5.3 Results and Discussion ................................................................................70 5.4 Conclusion ..................................................................................................79 References ........................................................................................................80

Chapter 6 Experimental and Theoretical Study of the Structure and Reactivity of Fe1-2O≤6¯ Clusters with CO................................................................................82

6.1 Introduction.................................................................................................82 6.2 Methods ......................................................................................................84 6.3 Results and Discussion ................................................................................86

6.3.1 CID Fragmentation............................................................................90 6.3.2 Reactions with CO and N2 .................................................................98

6.4 Conclusion ..................................................................................................104 References ........................................................................................................105

Chapter 7 Experimental and Theoretical Study of the Structure and Reactivity of Fe1-2O1-5

+ Clusters with CO...............................................................................107

7.1 Introduction.................................................................................................107 7.2 Methods ......................................................................................................109 7.3 Results and Discussion ................................................................................111

7.3.1 CID Studies .......................................................................................114 7.3.2 Reactivity Studies..............................................................................118

7.4 Conclusion ..................................................................................................126 References ........................................................................................................127

Chapter 8 Gas Phase Study of Cobalt, Nickel, and Copper Oxide Cluster Ions: Dissociation Patterns and Trends in Reactivity with CO ....................................129

8.1 Introduction.................................................................................................129 8.2 Experimental ...............................................................................................132

viii

8.3 Results and Discussion ................................................................................133 8.3.1 Collision Induced Dissociation Studies ..............................................133 8.3.2 Reactivity Studies..............................................................................145

8.4 Comparison to Bulk.....................................................................................156 8.5 Conclusion ..................................................................................................157 References ........................................................................................................159

Chapter 9 General Conclusions ................................................................................163

9.1 Major Findings ............................................................................................164 9.2 Future Studies .............................................................................................167 References ........................................................................................................169

Appendix Kinetic Analysis of the Reaction between (V2O5)1,2+ and Ethylene ..........170

A.1 Abstract ......................................................................................................170 A.2 Introduction................................................................................................171 A.3 Experimental and Computational Methods .................................................174 A.4 Results and Discussion ...............................................................................175 A.5 Conclusion .................................................................................................192 References ........................................................................................................195

ix

LIST OF FIGURES

Figure 2-1: Guided Ion Beam Mass Spectrometer employed in gas phase metal oxide studies. ....................................................................................................15

Figure 2-2: Schematic setup of the dual rod LaVa source which includes the beam splitting mirror and conical nozzle. ..........................................................19

Figure 3-1: A typical mass distribution for cationic gold oxide clusters. The first gold oxide in each series is labeled with the addition of one oxygen atom to subsequent peaks...............................................................................................29

Figure 3-2: Predicted lowest energy structures for AuxOy+ where x=1-2, y=1-4.

Source: Figure courtesy of the group of Professor V. Bonačić-Koutecký at the Humboldt University in Berlin.....................................................................33

Figure 3-3: Branching ratios for a) AuO+, b) Au2O+, and c) Au3O

+ which all show a large O atom replacement by CO reaction channel.................................36

Figure 3-4: Logarithm of the intensity ratio of ln[Ir/Io] versus CO reactant concentration for a) AuO+, b) Au2O

+, and c) Au3O+...........................................37

Figure 3-5: Branching ratio for AuO3+ containing a peripheral and molecular

oxygen unit in the lowest energy structure calculation. ......................................39

Figure 3-6: Branching ratio for AuO2+ which shows a selective pathway for O

atom replacement by CO. ..................................................................................39

Figure 3-7: Branching ratios for a) Au4O2+ and b) Au4O3

+ and their respective products with increasing carbon monoxide reactant gas.....................................43

Figure 4-1: Branching ratios for a) Au2O2+, b) AgAuO2

+, c) Ag2O2+, and d)

AgAuO2- showing the products in the reaction with CO reactant gas. ................55

Figure 4-2: a) Typical anionic silver-gold oxide mass distribution and b) cationic silver-gold oxide mass distribution obtained from dual rod LaVa source. ..........57

Figure 4-3: Branching ratios for a) Au2O+ and b) AgAuO+ with increasing CO

pressure.............................................................................................................59

Figure 4-4: Plots of ln[Ir/Io] versus CO reactant concentration for a) Au2O+ and b)

AgAuO+ ............................................................................................................60

Figure 5-1: Branching ratios for a) FeO3+ and b) FeO3

- with increasing CO pressure.............................................................................................................71

x

Figure 5-2: Ground state geometries for O2, CO, CO2, and FeOy+,- clusters.

Source: Figure courtesy of the group of Professor S. N. Khanna at VCU. ..........73

Figure 5-3: a) Graph of the energy required to remove an O or O2 from the FeOy+

cluster. b) Graph of the energy required to remove an O, O2, O- and O2

- from the FeOy

- cluster. Source: Figure courtesy of the group of Professor S. N. Khanna at VCU.................................................................................................75

Figure 5-4: Profile of the reaction energy for FeO3+ with CO including the

corresponding change in energy, ΔE, for each reaction step. Source: Figure courtesy of the group of Professor S. N. Khanna at VCU...................................76

Figure 5-5: Profile of the reaction energy for FeO3- with CO including the


Figure 6-1: A typical mass distribution produced for iron oxide anionic clusters. The first iron oxide in each series with subsequent peaks in the series having one additional oxygen atom...............................................................................87

Figure 6-2: The ground state geometries of O2, CO, FeO1-3- and FexOy

- clusters. Source: Figure courtesy of the group of Professor S. N. Khanna at VCU. ..........89

Figure 6-3: Graphs of the fragmentation energies for atomic and molecular oxygen loss in a) FeOy

- and b) Fe2Oy- clusters. c) Graph of the fragmentation

energy for neutral metal and metal oxide loss from Fe2Oy- clusters. ...................93

Figure 6-4: The change in energy (ΔE) for the dissociation pathway of parent cluster, Fe2O4

- producing FeO3- and FeO2

-. Source: Figure courtesy of the group of Professor S. N. Khanna at VCU. .........................................................95

Figure 6-5: The change in energy (ΔE) for the dissociation pathway of parent cluster Fe2O3


-. Source: Figure courtesy of the group of Professor S. N. Khanna at VCU. .........................................................97

Figure 6-6: The branching ratios for a) FeO2- and b) Fe2O3

- with increasing CO reactant gas pressure..........................................................................................100

Figure 6-7: The change in energy (ΔE) for the reaction pathway of Fe2O3- with

CO. Source: Figure courtesy of the group of Professor S. N. Khanna at VCU....101

Figure 6-8: The branching ratios for a) FeO4- and b) Fe2O6

- with increasing CO pressure, both clusters show molecular oxygen loss...........................................103

xi

Figure 7-1: Typical mass distribution for iron oxide cation clusters produced when employing a) a 27 mm conical nozzle and b) a 51 mm conical nozzle at the exit of the source. ........................................................................................112

Figure 7-2: The ground state geometries for a) FexOy+ clusters and b) Fe(CO)1-2

+ and FexOyCO+ clusters. Source: Figure courtesy of the group of Professor S. N. Khanna at VCU….. .......................................................................................116


+ and b) Fe2Oy+ clusters. c) Graph of the fragmentation

energy for neutral metal and metal oxide loss from Fe2Oy+ clusters. ..................119

Figure 7-4: Branching ratios for a) FeO+, b) FeO2+, c) Fe2O

+, and d) Fe2O2+

reacted with CO.. ..............................................................................................123

Figure 7-5: Branching ratio for Fe2O3+ reacted with CO. .........................................123



Figure 7-7: Branching ratio for FeO4+ reacted with CO showing at higher

pressures two CO molecules become attached to bare iron. ...............................125

Figure 8-1: Typical anionic mass distributions for a) Nickel oxides, b) Cobalt oxides, and c) Copper oxides. ............................................................................134

Figure 8-2: Typical cationic mass distributions for a) Nickel oxides and b) Cobalt oxides................................................................................................................135

Figure 8-3: Branching ratios for a) Co2O3-, b) Ni2O3

-, and c) Cu3O3- reacted with

carbon monoxide. ..............................................................................................147

Figure 8-4: Branching ratios for a) CoO2+ and b) NiO3

+ reacted with carbon monoxide.. ........................................................................................................148

Figure 8-5: Branching ratios for a) NiO2- and b) NiO2

+ reacted with carbon monoxide.. ........................................................................................................151

Figure 1-1: General scheme of the energetic profile and structural transformation of the mechanism for the reaction between (V2O5)n=1,2

+ and C2H4. Source: Figure courtesy of the group of Professor V. Bonačić-Koutecký at the Humboldt Unisversity in Berlin………………………………………………....173

Figure 1-2: Snapshots of the ab initio MD trajectory calculated using DFT methods for the reaction of a) V2O5

+ with C2H4 and b) V4O10+ with C2H4

xii

leading to the formation of acetaldehyde. Source: Figure courtesy of the group of Professor V. Bonačić-Koutecký at the Humboldt Unisversity in Berlin. ...............................................................................................................177

Figure 1-3: Branching ratios for the energetic analysis of V2O5+ with ethylene

conducted under the following pressures: a) 0.1 mTorr, b) 0.2 mTorr, c) 0.3 mTorr, d) 0.4 mTorr, e) 0.5 mTorr and f) 0.6 mTorr of ethylene........................178


conducted under the following pressures: a) 0.1 mTorr, b) 0.2 mTorr, c) 0.3 mTorr, d) 0.4 mTorr, e) 0.5 mTorr and f) 0.6 mTorr of ethylene. .......................179

Figure 1-5: a) Ratio of product ion intensity, Ip, to total ion intensity, I0, versus the reaction cell gas pressure for reactions of V2O5

+ with ethylene at ECM= 0 eV and b) at ECM= 2.0 eV. .................................................................................183

Figure 1-6: a) The rate constant versus energy for V2O5+ with ethylene calculated

from cross section data. b) The rate constant versus energy for V4O10+ with

ethylene calculated from cross section data........................................................183

Figure 1-7: a) Logarithm of the intensity ratio [V2O5+/V2O5

++ΣIp] versus the concentration of C2H4 calculated as an ideal gas at ECM=0 eV and b) at ECM= 2 eV. .................................................................................................................185

Figure 1-8: a) The rate constant versus center-of-mass energy for the reaction of V2O5

+ with ethylene calculated from the velocity of the ions and reaction time and b) The rate constant versus center-of-mass energy for the reaction of V4O10

+ with ethylene calculated from the velocity of the ions and reaction time...................................................................................................................185

Figure 1-9: a) Energy dependence of the total reaction cross sections for V2O5+

with ethylene for the theoretical model (solid line) and experimental measurements (diamonds). b) Energy dependence of the total reaction cross sections for V4O10

+ with ethylene. .....................................................................193

Figure 1-10: Dependence of the reaction cross section (σ(E)) on the energy for different values of the barrier for hydrogen transfer: 1.45 eV (black line), 3.75 eV (blue line), 3.85 eV (green line), 3.88 eV (red line). Source: Figure courtesy of the group of Professor V. Bonačić-Koutecký at the Humboldt Unisversity in Berlin. ........................................................................................193

xiii

LIST OF TABLES

Table 3-1: Reaction products for AuxOy+ with CO reactant gas and nitrogen

collision gas. .....................................................................................................31

Table 4-1: Product list for the dimer of pure silver, pure gold, and mixed Ag-Au oxide clusters in the presence of CO..................................................................54

Table 6-1: The CID fragmentation channels, calculated dissociation energies (DE), and reaction products with CO and N2 for Fe1,2O2-6

- clusters. ...................92

Table 7-1: Results of CID studies showing the fragmentation channels for selected Fe1-2Oy

+ clusters...................................................................................117

Table 7-2: Reaction products for the reaction between selected iron oxide cluster cations and carbon monoxide…………………………………………………. ..120

Table 8-1: CID fragmentation channels for selected MxOy- (M=Co, Ni, Cu)

anions................................................................................................................137

Table 8-2: CID fragmentation channels for selected MxOy+ (M=Co, Ni) cations. ....138

Table 8-3: Reaction products for selected MxOy- (M=Ni, Co, Cu) anions with CO

and N2. ..............................................................................................................140

Table 8-4: Reaction products for MxOy+ (M=Co, Ni) cations with CO and N2.........142

xiv

ACKNOWLEDGEMENTS

I would like to thank my thesis advisor Professor A. Welford Castleman, Jr. for

giving me the opportunity to grow as a scientist and a person. I would also like to thank

past and present members of the Castleman group for providing support, knowledge, and

helpful discussions, especially my lab partners Dina Justes and Grant Johnson. Also, I

thank Richard Bell for help in troubleshooting the instrument and Michele Kimble for

providing her knowledge and friendship. I also acknowledge and thank the members of

my thesis committee, Professor James B. Anderson, Professor Nicholas Winograd, and

Professor Robert J. Santoro.

I would like to thank my collaborators, Christian Bürgel, Roland Mitrić, and

Professor Vlasta Bonačić-Koutecký at the Humboldt University in Berlin, Germany and

J. Ulises Reveles and Professor Shiv N. Khanna at Virginia Commonwealth University in

Richmond, Virginia. Their theoretical calculations and insightful discussions have

contributed to a greater understanding of the experimental studies herein.

Most importantly I want to express my gratitude to my family who have always

supported and encouraged me. Finally, I dedicate this thesis to my husband, Brian Reilly.

He has stood by all my decisions and kept me going through the rough times.

Chapter 1

Introduction

1.1 The Oxidation of Carbon Monoxide

Carbon monoxide (CO) is a harmful atmospheric pollutant produced by

incomplete combustion of fossil fuels. The conversion of CO to CO2 is one of the most

important catalytic reactions pertaining to environmental and industrial applications. For

example, the ability to remove byproducts such as CO in fuel cell systems through the

water-gas shift (WGS) reaction at low temperatures and at fast reaction rates would aid in

the development of lower cost fuel alternatives (1).

From a fundamental point of view, the energy gain in going from CO to CO2 is

5.85 eV while the energy required to break the O-O bond in O2 is 5.23 eV. Therefore

based on the pure energetics of the reaction, the conversion of CO to CO2 should occur

spontaneously in the presence of O2. It has been suggested that the rate limiting step in

this process is the activation of O2 because the breaking of the O-O bond involves a large

energetic barrier (2). Catalysts that can lower the reaction barrier and hence promote the

reaction are needed to assist the conversion. The basic idea of the catalyst is then to

weaken/break the O-O bond through charge transfer and/or bonding, thus facilitating the

reaction. The ability to tailor the design of more energy efficient catalysts for the

oxidation of CO however remains elusive and fundamental studies on the active sites and

reaction mechanisms would aid in creating more selective catalysts.

2

1.2 Clusters as Models

Studies on the activity and selectivity of heterogeneous catalysts are highly

inconsistent depending on the chosen preparation method (3-5) and molecular level

insight into the active sites and catalytic mechanisms is difficult employing surface

science techniques alone. Gas phase studies are a complementary method which offer the

opportunity to study isolated potential catalytic active sites in the absence of solvent

effects and surface inhomogenieties that often complicate condensed phase catalytic

research (6-7). To this end, there has been a growing acceptance of the use of gas phase

metal oxide clusters to model the reactions that occur over metal oxide surfaces.

Muetterties and Witko et al. have suggested that a metal oxide surface may be viewed as

a collection of clusters and isolated reaction sites (8-9). In addition, Somorjai has

described a collection of atoms to be “cluster-like” (10), further establishing the value

that cluster studies may have in elucidating molecular level mechanisms.

Numerous correlations between the products of gas phase experiments and

heterogeneous catalysis have been uncovered (9-12). In our own laboratory, significant

correlations between gas phase experiments and analogous condensed phase reactions

have been obtained. In one study, vanadium and niobium oxide cluster cations with

methanol produced the dehydrogenation of methanol yielding neutral formaldehyde (13).

A major product produced over condensed phase vanadia surfaces is formaldehyde (14).

Another study involving a comprehensive experimental and theoretical evaluation of

vanadium oxides observed that only (V2O5)1,2+ clusters transfer an oxygen atom to

ethylene reactant gas, producing acetaldehyde (12, 15). Acetaldehyde has been produced

during ethylene oxidation over vanadia surfaces with V2O5 sites as the proposed active

3

sites (16). In addition, differences between vanadium oxide clusters and niobium and

tantalum oxides have been elucidated in reactions with n-Butane (17-18). Similar to

analogous bulk catalytic systems (19), V2O5+ transfers an oxygen atom to the

hydrocarbon, while Nb2O5+ and Ta2O5

+ exhibit C-C cracking as the dominant reaction

pathway.

Beyond our own work (12-13, 15, 17-18, 20), additional support for the

correlation between gas phase and condensed phase studies is being realized. Zamaraev

et al. studied the oxidation of methanol on MoxOy+ clusters and observed similar

reactions occurring over the corresponding molybdenum-oxygen sites in homogeneous

and heterogeneous catalysts (11). Gas phase studies conducted by Plattner and

coworkers found C-H activation on cationic iridium (III) complexes (21-22). Their

studies were used to design a more efficient and reactive solution phase catalyst

containing Ir (23). Wallace and Whetten and Wöste and coworkers have combined gas

phase experiments and theoretical calculations to elucidate information on the full

catalytic cycle of CO oxidation with proposed reaction steps (2, 24-25). Schwarz et al.

has also demonstrated a full gas phase catalytic cycle for the oxidation of NO in the

presence of platinum oxide cation clusters (26). The ability to probe the active sites and

reaction intermediates of complex catalytic reactions through gas phase ion-molecule

studies has aided many researchers in unraveling the related mechanisms (26-29).

Furthermore, Armentrout and coworkers have revealed ion-molecule reaction kinetics

yielding information on the thermochemistry, structure, and bond energy of transition

metal clusters (30).

4

Investigating the nature of catalytic active sites is of utmost importance for

understanding the reaction mechanisms and designing better catalytic systems. Gas

phase studies are especially amenable to probing ionic charge state and electron density,

two factors that have been proposed to assist in activating reactions (31-34). Studies by

Grzybowska-Świerkosz have focused on the role of ionic centers that arise on certain

sites in vanadium catalysts and the properties of these charged sites (35). It has been

pointed out by Davis that further investigation of ionic sites will expand our

understanding of active sites in catalytic systems (36). Therefore the above mentioned

findings illustrate the significance of gas phase ionic cluster studies in elucidating such

effects as cluster size, stoichiometry, charge and oxidation state, and composition on

reactivity. These studies provide a useful means to gain important insight into the basic

processes involved in heterogeneous catalysis.

1.3 Insight into Alternative Catalytic Material

The most common catalytic materials employed in the oxidation of CO presently

are Pdn and Ptn nanoparticles; however these materials are expensive and require high

operating temperatures (1, 37-38). New efforts are being put forth to find suitable

alternative materials with good active lifetimes that are selective toward the oxidation of

CO.

5

1.3.1 Gold Oxides

Bulk gold has long been known to be catalytically inert (39-40). However,

findings by Haruta and coworkers demonstrated that nanosized gold particles deposited

onto select metal oxides exhibit high activities for the oxidation of carbon monoxide (39,

41-42). These studies also established that the catalytic activity was highly dependant on

the type of metal oxide support and the gold particle size.

Past studies have focused on elucidating the structures of charged gold through

ion mobility measurements and theoretical methods (43-47). The adsorption properties

(48-51) and reactivity of O2 and CO with gold clusters have also been extensively

explored to uncover potential active sites and mechanistic details in the oxidation of CO

(52-54). Several researchers suggest that charge transfer between gold and adsorbed O2

allows for elongation and activation of the O-O bond (2, 55). However, there is no

consensus on the active charge state for the catalytic conversion, thus further

investigation is warranted.

1.3.2 Bimetallic Oxides

The introduction of another metal into a catalytic system can greatly impact the

structure and electronic properties of the catalyst (56-57). In a study by Metiu and

coworkers on AgmAun clusters, replacing active sites with either Au or Ag resulted in a

change in the bond strength (58). This change in strength was directly related to the

LUMO energy; the lower the LUMO energy, the stronger the bond energy became. The

ability to tune the properties by changing the number of atoms of each component in

6

mixed clusters may be particularly valuable for designing new catalytic materials with

enhanced activity. There is a great disparity in the information available for pure metal

oxide systems compared to bimetallic systems and further experiments on bimetallic

reactivity may aid in elucidating the active sites and reaction mechanisms responsible for

increased catalytic activity.

1.3.3 Transition Metal Oxides

An alternate approach to overcome the O2 rate limiting step is to use systems that

contain O atoms without the O-O bonds. Transition metal oxides with oxygen atoms

bound only to the transition metal without O-O bonding offer an attractive possibility.

Indeed, recent experiments on iron oxide nanoparticles show that it may be possible to

convert CO to CO2 at low temperatures in the presence or absence of O2 (59-62).

Theoretical studies on small clusters containing less than 10 atoms reveal that the ease in

structural rearrangement at small sizes almost eliminates the reaction barriers (63-64).

This is provided that changes in the spin multiplicity are minimal and offer a spin

allowed process. Employing small size clusters therefore may offer the possibility of

spontaneous conversion of CO to CO2.

Since charge transfer and electron density are predicted to play a role in the active

sites of catalysts, it is valuable to gain information about different transition metal

systems each with their own structural and electronic properties. Studies involving

several 3d transition metal clusters, iron, nickel, cobalt, and copper, will provide insight

into the role of additional d electrons on the reactivity toward CO oxidation.

7

Comprehensive studies on transition metals have revealed periodic trends in reaction

energy and the role of d orbitals in bonding (65). The influence of different electron

affinities for each metal composition may provide insight into how regions of electron

deficiency and enhancement affect catalytic activity. Individual studies have probed the

structures, bond dissociation and adsorption energies, and reactivity of iron (66), nickel

(67), cobalt (68), and copper (69). However, there remain many unanswered questions

involving the catalytic efficiency of these metals.

1.4 Synopsis

The following thesis contains an investigation of ionic gold oxides, gold-silver

bimetallic oxides, and several transition metal oxide species toward elucidating the

reactivity for alternative catalytic materials in the conversion of CO to CO2. The

technique employed in the experimental studies is described in Chapter 2. Reaction

pathways for cationic gold oxides are discussed in Chapter 3 and compared to reactions

of anionic clusters with similar stoichiometry. An emphasis on the oxygen atom

replacement pathway and a proposed mechanism is presented where the role of charge

state is an important factor in this observed reaction product. The reactivity of bimetallic

gold-silver dimer oxides are examined in Chapter 4. The ability to tune the properties of

clusters by changing the composition is revealed based on similarities and differences in

the reaction products observed. The structure of ionic iron oxide clusters and charge state

influences in the reaction pathways are explored in Chapters 5-7. Reactivity and collision

induced dissociation studies involving cobalt, nickel, and copper oxide clusters are

8

investigated in Chapter 8. The periodic trends in the first row 3d transition metals are

discussed. Chapter 9 presents the major finding and draws conclusions from the

experiments presented in the previous chapters. Finally, in addition to CO reactivity

studies discussed throughout this thesis, a study describing the reaction between

(V2O5)n=1-2+ and ethylene is presented in the appendix.

9

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13

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Chapter 2

Experimental Setup

2.1 Introduction

The experiments discussed herein were conducted on a guided ion beam (GIB)

mass spectrometer. The GIB mass spectrometer is capable of studying the reactivity of

ion-molecule reactions and collision induced dissociation (CID) of mass selected cluster

ions. The design of this apparatus has been described in detail previously (1-2).

Figure 2-1 shows a schematic of the GIB with the different chambers and individual

components labeled. The instrument consists of a laser vaporization source, quadrupole

mass selector and analyzer, octopole reaction cell, a vacuum system, and a detection and

data collection system. Below is a discussion of the individual components and the

conditions with which the transition metal and noble metal oxide experiments were

conducted.

2.2 Vacuum system

The GIB apparatus has three regions, the source, reaction, and detection, each

with its own vacuum pumps in order to maintain differential pumping. The source region

is pumped by a 6” diffusion pump (Varian VHS-6) with a pumping speed of 3000 L/s and

backed by a Sargent-Welch mechanical pump (model 1397) which pumps 500 L/min.

15

Figure 2-1: Guided Ion Beam Mass Spectrometer employed in gas phase metal oxide studies.

Gas Inlet

OctopoleReaction Cell

LaserVaporization

Source

Deflectors CEM

Diffusion Pump Diffusion Pump Diffusion Pump

Quadrupole 2Quadrupole 1

DiffusionPump

Conversion dynode

MKS Baratron

Source

Reaction

Detection

Skimmer

Gas Inlet

OctopoleReaction Cell

LaserVaporization

Source

Deflectors CEM

Diffusion Pump Diffusion Pump Diffusion Pump

Quadrupole 2Quadrupole 1

DiffusionPump

Conversion dynode

MKS Baratron

Source

Reaction

Detection

Skimmer

16

This mechanical pump has a switching valve that allows it to rough-out the mixing tank

or chamber as needed. The source region operates at ~5x10-5 Torr and is monitored by a

thermocouple gauge (Fil-tech, model G-100-P). The reaction chamber region is

evacuated by a 6” diffusion pump (Varian VHS-6) and backed with a Sargent-Welch

mechanical pump (model 1397), similar to the source region. The operating pressure

inside the reaction chamber is ~6x10-7 Torr monitored using an ionization gauge and ion

gauge controller (Granville Phillips, model 270). The skimmer serves to separate the

source and reaction chamber effectively creating two differentially pumped chambers.

The detection chamber uses both a 10” diffusion pump (Varian HS-10) with an

effective pumping speed of 4200 L/s and a 6” diffusion pump (Varian NRC-6) having a

pumping speed of 3000 L/s. The 10” pump has a freon cooled baffle and is backed by a

Sargent-Welch mechanical pump, while the 6” is backed by a direct drive mechanical

pump (Varian SD-700) having a pumping speed of 765 L/min. The operating pressure

inside the detection chamber is ~6x10-7 Torr and is monitored by an ionization gauge and

ion gauge controller (Granville Phillips, model 270).

All the diffusion pumps have a manual gate valve to separate the pump from the

chamber. An in-house built safety system monitors the pressure and temperature of the

diffusion pumps. If the vacuum pressure increases to above 10-4 the safety will switch off

the main power. Also, if the pressure of the cooling water decreases below the set

threshold (20 psi) a McDonnell flow switch (FS-1), installed in-line with the cooling

lines, will switch off the main power.

17

All the mechanical pumps and the freon cooling baffle are mounted on spring

isolators (Kinetics Noise Control, FDS 1-35 & FDS-70). This is to dampen the vibrations

of these pumps and allow for stability of the components in the chamber.

2.3 Laser Vaporization Source

The guided ion beam apparatus is coupled to a laser vaporization source (LaVa)

employed in the formation of the ionic clusters and based on the design of Smalley et al.

(3-4). Utilizing a LaVa source, metal oxide clusters of varying stoichiometries can be

produced by ablating a rotating and translating metal rod with the second harmonic (532

nm) of a Nd:YAG laser (Spectra-Physics, INDI-50, 20 Hz). The laser is focused by a 20

cm focal length lens onto a 1/8” metal rod which is continuously rotated at a rate of 1-1.5

revolutions per minute to ensure a fresh metal ablation surface each time the laser is

pulsed. The laser fluence is adjusted between 5 to 20 mJ/pulse depending on the metal

and size of clusters being studied. Generally lower laser power is used for gold because

it is a soft metal, whereas silver is run at higher laser power. Also, it is generally

recognized that lower energy pulses produce larger metal oxide ions and higher energy

pulses generate smaller metal oxides due to the higher energy atomizing the metal.

A pulse valve (General Valve, Series 9, 28 V) is employed to deliver a carrier gas

composed of a mixture of high purity oxygen (99.999%, MG Industries) and ultra high

purity helium (99.9999%, MG Industries). The ratio of oxygen to helium is dependent on

the metal cluster ion being studied and ranges from 1% for iron anionic clusters to 80%

for anionic gold clusters. The pulse valve is operated by a home built pulse valve driver

18

that delivers a pulse of mixed gas into the source region with a fixed backing pressure,

~5.5 to 6.5 atm. The timing of the pulse valve and laser pulse is controlled by a home-

built 20 Hz pattern generator. The pulse valve delay ranges from 20 to 150 μs while the

laser pulse delays range between 300 to 450 μs. It has generally been found that larger

mass clusters, like gold, use a longer gas delay time and shorter laser delay to maximize

the overlap between the pulse nozzle being open (240 μs) and when the laser pulse hits

the rod.

As the mixture of gas is pulsed it is expanded in a waiting room before it passes

over the ablated rod and interacts with the vaporized metal plasma composed of metal

atoms and ions. The helium in the carrier gas serves to induce clustering and thermalize

the particles through collisions. A conical expansion nozzle (shown in Figure 2-2 for the

dual rod source) placed at the exit of the source region aids in the clustering process by

allowing more third-body collisions. Two different nozzles were employed depending on

the metal being studied. For anionic gold oxides, a 51 millimeter long conical nozzle

with a 2 mm inner diameter and a 12 mm conical expansion cone of 30 degree internal

angle was used, while cationic iron oxides had a wider cluster distribution using a 27

millimeter long nozzle composed of a 15 mm long channel with a 4 mm diameter and 12

mm conical expansion cone with a 30 degree total internal angle. When the clusters exit

the nozzle they rapidly expand into the high vacuum region (~10-7 Torr) and undergo

supersonic expansion, which cools the clusters internally to near thermal energy. The

clusters then travel 25 cm in a field free region and are collimated through a 3mm

skimmer having a 50 degree total internal angle.

19

Figure 2-2: Schematic setup of the dual rod LaVa source which includes the beam splitting mirror and conical nozzle.

Pulse nozzle Conical

Expansion

Nozzle

2 rods

Pulse nozzle Conical

Expansion

Nozzle

2 rods

20

A LaVa source able to hold two metal rods and ablate them using one laser was

built based on the design of Wagner et al. (5) in order to study the effects of bimetallic

clusters. An expanded view of the dual rod LaVa source is shown in Figure 2-2. A half

circular mirror was used to split the laser beam from a single laser into two beams able to

ablate two separate rotating and translating 1/4” metal rods. All other aspects of this dual

rod source are similar to the single metal LaVa source describe herein. This source was

only employed for the mixed Ag-Au oxide experiments described in Chapter 4.

2.4 Guided Ion Beam Setup

The mass spectrometer used in the present studies consists of three

radiofrequency components, a mass selecting quadrupole, an octopole reaction cell, and a

mass analyzing quadrupole, shown schematically in Figure 2-1. Each of the components

is aligned with the beam of a HeNe laser which creates the path the ion beam travels

through the instrument.

The molecular ion beam created by the LaVa source is further collimated by a set

of electrostatic lenses and two pairs of ion deflectors before being directed into the first

quadrupole. All the lenses and deflectors here and further down the setup are powered by

205B-01 Bertan power supplies (0-1 kV, 0-30 mA). Each lens is individually set to

optimize the ion intensity and kept at the lowest possible voltages in order to avoid

transferring energy to the ions.

Both quadrupoles used in the experiments are the same dimensions having four

cylindrical stainless steel rods, 9.5 mm in diameter and 20 cm in length. Each rod is

21

symmetrically aligned around a 0.8 cm diameter circle giving an optimal quadrupolar

field for the ions to pass through. The quadrupoles are each mounted in a vented

stainless steel housing that is also used to attach the electrostatic lenses. An external

oscillator is used to produce the rf voltage in order to avoid phase shifts between the two

quadrupoles. A radio frequency of 880 kHz is generated and this signal is amplified by

separate 300 W power supplies (Extrel, model QC-150). These power supplies are

connected through RG-8 coaxial cables connected to feedthroughs on the outside of the

chamber. In addition to the rf component, a dc potential is applied to opposite pairs of

rods and controlled by a mass programmer (Extranuclear Laboratories, model 091-3).

The detectable mass range for the quadrupoles employed in these experiments is 0-4000

amu, which is determined by the rf signal and radius of the rods.

The first quadrupole is used to mass select the cluster ion of interest. By applying

a dc potential to the opposite pairs of rods the quadrupole can act as a mass filter to only

transmit ions of specific a mass/charge (m/e) ratio (6-8). This dc potential in combination

with the rf voltage and angular frequency of the oscillating potential determine which

ions have stable trajectories through the quadrupole. Once the ion is selected in the first

quadrupole it is directed into the octopole reaction cell by a second set of electrostatic

lenses.

The guided ion beam apparatus functions similar to the triple quadrupole, initially

introduced by Teloy and Gerlich (9). However, it employs an octopole reaction cell

instead of a quadrupole. The octopole ion guide, based on the design of Armentrout et al.

(10) using the approach of Denison (11) for calculating the rf potential of non-ideal

multipoles, has eight circular stainless steel rods, each 40.6 cm long and 3.57 mm in

22

diameter. The rods are set in a circle with a radius of 5.18 mm. The rf and dc potentials

are supplied by a home-built rf generator based on the design of Anderson (12). The

generator is powered by an adjustable 0 to 1000 V, 100 mA power supply (Bertan, model

915-1 P). The dc float voltage is set to zero in reactivity studies where the ions

experience near thermal energies; however, in CID experiments the voltage on the

octopole rods may be increased. This voltage is supplied by a bipolar 0-1 kV, 0-15 mA

power supply (Bertan, model 230-01F).

There are several advantages to employing an octopole in these experiments. The

octopole has a higher trapping efficiency, allowing 4π collection of ionic products, and

does not impart significant off axis energy that would scatter the ions and decrease signal.

The ion collection efficiency depends on the cell geometry and direction of scattered

products and is determined using eq 2-1 for the effective radial potential energy (Ueff) (9,

12),

s

n

oo

oeff U

r

r

rm

VqnrU +⎥

⎦

⎤⎢⎣

⎡=

−22

22

222

4)(

ω. (2-1)

The number of poles is n, q is the charge, m is the mass of the ion, Vo is the rf potential on

the rods, r is the radial distance from the multipole axis, and ro is the inner circle radius of

the rods.

The ion-molecule reactions occur in a stainless steel reaction cell, 11.4 cm long

with a 5 cm diameter, within the octopole rods. The calculated effective path length of

the cell was determined to be 12.9 cm using the trapezoidal pressure fall-off

approximation (2, 13). The neutral reactant gas pressure, carbon monoxide in the present

studies, is monitored by a Bertan capacitance manometer (MKS) controlled by a signal

23

conditioner (model 270). Pressures range between 0-10 mTorr for cationic and 0-20

mTorr for anionic cluster studies.

The reaction products generated in the octopole are then focused by a third set of

electrostatic lenses into the second quadrupole. The second quadrupole, employed as a

mass analyzer, is operated in rf mode only in order to separate ions with differing m/e

ratios. The products are then directed into the channel electron multiplier.

In addition to reactivity studies, CID studies were conducted to aid in the cluster’s

structural determination and bond energy strength. In these experiments the dc float

voltage on the octopole rods was increased to raise the kinetic energy of the ions in the

reaction cell. Teflon screws were used to hold the octopole and second quadrupole in

their housing to effectively isolate them from ground so that voltages could be applied.

An inert gas, xenon, was added to the reaction cell under single collision conditions,

~0.09 mTorr. The mass analyzing quadrupole was then floated by a dc potential. In

order to attract the product ions, the potential of the quadrupole needed to be greater than

the birth potential of the ions, ie. the potential of the octopole. The third set of lenses,

extraction lenses, was increased to attract the ions also. By increasing the pole bias and

voltage on the lenses, signal during these studies greatly diminished, therefore only metal

oxides with significant signal initially were studied.

2.5 Detection and Data Acquisition

The mass analyzed product ions are then directed into a channel electron

multiplier, CEM, (Detector Technology, Inc., model 402A-H) by a single electrostatic

24

lens. The CEM has a conversion dynode which is used to attract the ions with a high

voltage, ±5000 V. The channeltron operates at the same polarity for positive and

negative ions, in which secondary electrons and secondary positive ions are detected,

respectively. In the case of cations, the impact of the cation with the dynode creates a

cascade of secondary electrons that are then accelerated by the gradient electrostatic field

(14). The CEM operates with a voltage drop from approximately -2300 V on the front to

ground on the rear of the channeltron. This decrease directs the secondary electrons

through the detector generating a large gain of approximately 107 to 108. Similarly for

anions, a +5000 V dynode attracts the negative ions, creates a cascade of secondary

positive ions which are then accelerated to the -2300 V to ground drop across the

channeltron. The CEM is operated in pulse counting mode to distinguish spurious noise

from input signal.

The raw signal is filtered and sent through a 50 MHz preamplifier-discriminator

pad (MIT, model F-100T). The ion pulses are counted by a linear ratemeter (Mech-

Tronics, model 777) and also sent to a multi-channel scalar (MCS) card

(Tennelec/Nucleus, Inc., MCS-II) installed in a PC. Before entering the MCS card, the

signal is sent through a home-built time-resolved mass gate that filters random noise and

fast ions. The data is converted into mass spectra and analyzed by Grams/32 software

package.

25

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4. Dietz, T. G.; Duncan, M. A.; Powers, D. E.; Smalley, R. E. J. Chem. Phys. 1981, 74, 6511.

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10. Ervin, K. M.; Armentrout, P. B. J. Chem. Phys. 1985, 83, 166.

11. Denison, D. R. J. Vac. Sci. Technol. 1971, 8, 266.

12. Landau, L. D.; Lifshitz, E. M. Mechanics, 3rd Ed. Oxford: New York 1976, 3, 53.

13. Rosion, S.; Rabi, I. I. Phys. Rev. 1935, 48, 373.; Amdur, I.; Jordon, J. E. Adv. Chem. Phys. 1966, 10, 29.

14. Strobel, H. A.; Heineman, W. R. Chemical Instrumentation: A Systematic Approach John Wiley and Sons: New York, 1989, 805.

Chapter 3

Comprehensive Study of the Reactivity of Gold Oxide Cluster Cations (AuxOy+;

x=1-4, y=1-5) in the Presence of Carbon Monoxide

3.1 Introduction

Nano-scale gold particles are shown to be catalytically active species useful for

the oxidation of carbon monoxide to carbon dioxide (1) at low temperatures (2-3).

However, the reaction mechanism and intermediates of active gold species are still under

much debate. Controversy exists over the role of charge state in gold nanoparticles and

most past studies have focused on the anionic gold species (4-5). For anionic gold, the

existence of surface oxygen vacancies in support material enables the transfer of

electrons from the support to the gold (6). Therefore, it has been suggested that anionic

gold is the active site for weakening/breaking molecular oxygen and allowing for CO

oxidation to proceed (4, 7-8).

However, size-selected cationic Au deposition studies showed clusters as small as

Au3 promote the oxidation of CO (9). Numerous studies have provided evidence of both

metallic and cationic gold species present in working catalysts (10-16). Corma and

coworkers observed both cationic and metallic gold played a role in CO oxidation and

their findings established that the cationic species were the catalytic active sites (17).

They proposed Au3+ was an active site for the oxidation of CO, leaving both Au+ and Au0

reduced species. While some studies advocate the influence of Au+ sites toward

27

oxidation activity (18-19), others report that CO on Au+ sites are detrimental to the

oxidation of CO (20).

Fierro-Gonzalez and Gates prepared a gold catalyst and employed XANES

spectroscopy to show the reduction of Au3+ to Au+ in the oxidation of CO, with no

metallic Au being produced (13). Kinetic studies of this catalyst containing only cationic

gold exhibited a decrease in the reaction rate for CO oxidation versus catalysts containing

both cationic and metallic species. In other studies, it was proposed that an electron is

transferred from the metal to the support; this was represented by an FTIR absorbance

spectra band corresponding to CO absorbed to Auδ+ (21). These positive gold sites were

found to form strong Au-CO bonds which consequently decreased the oxidation activity.

Therefore, the influence of charge density and charge transfer between the gold and metal

oxide support plays an important role in catalytic activity (22-26); however, there is little

agreement on the role positive gold sites play in the activity of CO oxidation.

Previously, anionic gold monomer through trimer oxide cluster experiments

conducted in our laboratory found structure-reactivity relationships for CO oxidation and

CO association reactions (27-30). A comparison of their reaction pathways is included in

the discussion herein. This chapter addresses mass selected gold ions reacted with CO,

adding to prior investigations utilizing a flow tube reactor (31). The results offer

important insight into reaction mechanisms responsible for driving condensed phase

catalysis, with particular emphasis on the role of charge state.

28

3.2 Experimental Methods

Gas-phase cluster studies were performed using a guided ion beam mass

spectrometer coupled to a laser vaporization source, detailed in Chapter 2. Briefly, the

second harmonic of a Nd:YAG laser was used to ablate a rotating and translating gold

rod (SurePure Chemetals, 99.99% purity). At a predetermined time, oxygen seeded in

helium (~20%) was pulsed over the rod, forming hot dense plasma. A 51 millimeter

conical expansion nozzle with a 2 mm channel diameter was placed at the exit of the

source to allow for more collisions and clustering. The clusters underwent supersonic

expansion in a field free region before being passed through a 3 mm skimmer which

created a molecular beam. The cooled clusters were then focused by a set of electrostatic

lenses and deflectors into the first quadrupole. Each cluster species was individually

selected in the first quadrupole and passed through a second set of lenses into the

octopole reaction cell where they were decelerated to thermal energy. Reactant gas,

ranging from 0-10 mTorr, was added to the reaction cell and the pressure was monitored

by a MKS baratron. The products were directed through a third set of lenses to a second

quadrupole where they were mass analyzed and finally detected using a channel electron

multiplier.

3.3 Results

The laser vaporization source employed in these studies allows for the

atomization of oxygen and creates gold oxide clusters with atomic as well as molecular

oxygen.

29

Figure 3-1: A typical mass distribution for cationic gold oxide clusters. The first gold oxide in each series is labeled with the addition of one oxygen atom to subsequent peaks.

150 300 450 600 750 900

Au+ A

u 2+

AuO

2+

Au 3

+

Au 2

O3+

Au 4

O2+

Au 3

O4+

Au 4

O4+

Mass (amu)

Ion

Inte

nsity

(ar

b. u

nits

)

150 300 450 600 750 900

Au+ A

u 2+

AuO

2+

Au 3

+

Au 2

O3+

Au 4

O2+

Au 3

O4+

Au 4

O4+

Mass (amu)

Ion

Inte

nsity

(ar

b. u

nits

)

30

Figure 3-1 shows the typical cluster mass distribution at near thermal energy for gold

monomer through tetramer oxide species. The present reactivity studies include an

investigation of AuO1-5+, Au2O1-5

+, Au3O1-5+, and Au4O2-5

+ species. Au4O+ was the only

atomic O species not produced (see Figure 3-1).

Table 3-1 lists the products of each cluster species in descending order of

intensity reacted with either carbon monoxide or nitrogen. For reaction verification each

cluster was exposed to nitrogen gas under the same experimental conditions of pressure

and energy as CO reactant gas. Any products obtained with nitrogen as the collision gas,

are assumed to represent fragmentation products. These studies aid in identifying

different structural elements and verifying whether atomic or molecular oxygen was

adsorbed to the gold cluster. Oxygen rich clusters, AuO4+, AuO5

+, Au2O2+, Au2O3

+,

Au2O4+, Au2O5

+, Au3O2+, Au3O3

+, Au3O4+, Au3O5

+, Au4O2+, Au4O4

+, and Au4O5+, with

molecular O2 units dissociated an oxygen molecule at higher N2 pressures. Clusters with

more than five oxygen atoms were assumed to contain weakly bound O2 units that

fragmented to form more stable, less oxygen rich clusters. An oxygen atom was lost in

reactions with N2 for the following cluster species: AuO2+, AuO4

+, AuO5+, Au2O5

+, and

Au3O+.

Notice from Table 3-1 that most of the products of reactions with carbon

monoxide contain CO attached to the metal oxide cation. Adsorption of CO onto anionic

gold oxide clusters was not observed for monomer or dimer species and only Au3OCO-

species was observed during previous anionic studies with CO (27-29).

31

Table 3-1: Reaction products for AuxOy+ with CO reactant gas and nitrogen collision gas.

AuxOy+

(x, y) Products with CO

Products with N2

AuxOy+

(x,y) Products with CO

Products with N2

(1,1) AuCO+ none (3,1) Au3CO+ none

(1,2) AuOCO+ none (3,2) Au3+

Au3CO+ Au3

+

(1,3) AuOCO+ Au(CO)2

+ AuCO+

AuO+a

none (3,3) Au3O+

Au3OCO+

Au2OCO+ Au3CO+

Au3O2CO+a

Au3O+

(1,4) AuO2CO+

AuCO+ AuOCO+

AuO2+ (3,4) Au3CO+

Au3+

Au3OCO+

Au3O2+

Au3+

(1,5) AuO3CO+ AuO4

+ AuO2CO+

AuO3+a

AuO3+

AuO4+

(3,5) Au3O3+

Au3CO+ Au3O2CO+

Au3O+a

Au3O3+

Au3O+

(2,1) Au2CO+

Au2+a

none (4,2) Au3CO+ Au4

+ Au4OCO+

Au4O+

Au3+a

Au4+

(2,2) Au2CO+

Au2+

Au2+ (4,3)

Au4O

+

Au3OCO+

Au4O2CO+ Au4O2

+ Au3

+a

Au4O+

(2,3) Au2CO+

AuOCO+ Au2O

+ Au(CO)2

+

Au2O2CO+a

Au2O+ (4,4) Au4CO+

Au3CO+

Au4O+

Au4O3CO+ Au4OCO+a

Au4O2+

Au4+

(2,4) Au2CO+ Au2O2+ (4,5) Au4O3

+ Au4O2CO+

Au4O2+

Au3+a

Au3CO+a

Au4O3+

(2,5) Au2OCO+ Au2O3

+ Au2O4

+

Au2O3+

aDenotes a minor product channel whose relative intensity is less than 1% and not shown in the branching ratios for clarity.

32

3.4 Discussion

For anionic gold, structure-reactivity relationships were developed for the three

different reaction centers identified, including a peripheral oxygen, bridging oxygen, and

molecular oxygen, in order to show the reactions with CO were largely structure driven.

Structural calculations for small cationic gold monomer and dimer oxides conducted by

the Bonačić-Koutecký research group found either molecular oxygen or bridging oxygen

bound to gold in the calculated lowest energy structure (see Figure 3-2). As shown in

Figure 3-2, AuO+ was the only cluster with a peripheral oxygen atom, while AuO3+

contained both peripheral and molecular oxygen. Au2O+ and Au2O3

+ both have a

bridging oxygen atom and even numbered oxygen species, AunO2,4+, have structures with

molecular oxygen attached. The following sections are presented according to the

reactions that the clusters undergo, including oxygen replacement by CO, Au

fragmentation, and CO oxidation, with structural and energetic implications to the

reaction pathways being addressed. According to the observed reactions, the role of

charge density may be an important driving force for cationic reactions.

3.4.1 CO Replacement Reactions

The most common reaction product observed for cationic gold oxide clusters was

the replacement of I) an oxygen atom, or II) molecular O2 group by CO.

I) Oxygen Atom Replacement

It is energetically demanding to break the strong O-O bond in gold oxides that

contain an O2 subunit based on the structures of Figure 3-2.

33

Figure 3-2: Predicted lowest energy structures for AuxOy

+ where x=1-2, y=1-4. The superscripts represent the spin multiplicity. Source: Figure courtesy of the group of Professor V. Bonačić-Koutecký at the Humboldt University in Berlin.

34

Therefore, according to the three structural features identified in anionic studies (27), it

was expected that O atom replacement only occur in clusters that had a peripheral oxygen

atom.

Peripheral Oxygen: The only cluster shown to have oxygen bound to a peripheral

site on cationic gold was AuO+. Figure 3-3(a) shows the branching ratio for AuO+ which

displays a selective reaction channel for O atom replacement by CO. Atomic oxygen

replacement on AuO+ easily occurs as the lone pair of CO is drawn to the positive charge

of the gold atom. This reaction channel was therefore energetically possible and very

selective. It should be noted that CO oxidation was also an energetically feasible channel

since the positive charge on gold would bind oxygen weaker than anionic gold centers.

However, oxidation was not observed for this cluster experimentally.

The CO replacement reaction rate, assuming pseudo-first order kinetics, was

calculated based on the decrease in parent intensity versus the reactant concentration,

shown in eq 3-1.

[ ]tRkI

I

o

r =⎥⎦

⎤⎢⎣

⎡ln (3-1)

Here, Ir and Io is the parent ion intensity with and without the addition of CO reactant

gas, respectively. The pressure of the CO reactant molecule, [R], is calculated as a

concentration, and t is the time the ion spends interacting in the octopole reaction cell.

Figure 3-4 shows the plot of ln[Ir/Io] versus the concentration of the reactant ion, in which

the slope is equal to –kt. In order to calculate the time the ion spends in the reaction cell,

the velocity of the ions is calculated. Since no kinetic energy is added into the reaction

35

cell, the velocity depends only on the supersonic expansion velocity. The supersonic

expansion velocity is given by eq 3-2,

2/1

2

11 ⎥

⎥⎥⎥

⎦

⎤

⎢⎢⎢⎢

⎣

⎡

⎟⎠⎞

⎜⎝⎛ −+

=T

Ts

Mm

RTMv

γγ

,

(3-2)

where MT is the mach number, m is the mass of the ion, and γ is the specific heat capacity

ratio (32). In this case, the heat capacity was calculated for atoms because the carrier gas

was mostly comprised of helium being pulsed through the nozzle. The velocity is equal

to the time divided by the length of the reaction cell, 0.129 cm. Finally, the slope divided

by the time reveals a rate constant of 7.6x10-12 ± 0.5 cm3s-1 for AuO+ as shown in Figure

3-4(a). This reaction rate was calculated from the average of several experiments

collected on different days.

For AuO3+ species, oxygen is bound both peripherally and molecularly. The

reaction products, Au(CO)2+ and AuCO+, lose both of these oxygen subunits as shown in

the branching ratio of Figure 3-5. The binding energy of O2 is approximately 0.6 eV and

the binding of the O atom is much stronger at 1.4 eV (32). The overall energy required to

remove both O2 and O is therefore 2 eV. This pathway is slightly energetically possible

because the binding of CO onto the gold cluster is approximately 2.2 eV. Thus the

energetic pathway of this reaction is feasible because the gain in binding CO is greater

than the energy needed to break both bonds. It can be seen in Figure 3-5 that the product

ion, AuOCO+, is more intense than the other two since only one O2 bond is broken to

form this product.

36

Figure 3-3: Branching ratios for a) AuO+, b) Au2O

+, and c) Au3O+ which all show a

large O atom replacement by CO reaction channel.

0

0.25

0.5

0.75

1

0 2 4 6 8 10

CO (mTorr)

AuO+

AuCO+

AuO+

AuCO+

0

0.25

0.5

0.75

1

0 2 4 6 8 10

CO (mTorr)

Au3O+

Au3CO+

0

0.25

0.5

0.75

1

0 2 4 6 8 10

CO (mTorr)

Au3O+

Au3CO+

Au3O+

Au3CO+

Au3O+

Au3CO+

a)

b)

c)

Rel

ativ

e In

tens

ity

Rel

ativ

e In

tens

ity

Rel

ativ

e In

tens

ity

0

0.25

0.5

0.75

1

0 2 4 6 8 10

CO (mTorr)

Au2O+

Au2CO+

Au2O+

Au2CO+

b)

37

Figure 3-4: Logarithm of the intensity ratio [Ir/Io] versus CO reactant concentrationcalculated as an ideal gas at a) AuO+, b) Au2O

+, and c) Au3O+.

y = -2.26x10-15xR2 = 0.826

-1.2

-1.0

-0.8

-0.6

-0.4

-0.2

0.00.0 0.5 1.0 1.5 2.0 2.5 3.0 3.5

Conc [CO] 1014 mc/cc

ln[I

r/Io]

b)

y = -4.48x10-15x

R2 = 0.993

-1.6

-1.4-1.2

-1.0

-0.8

-0.6

-0.4

-0.2

0.00.0 0.5 1.0 1.5 2.0 2.5 3.0 3.5


ln[I

r/Io]

c)

y = -4.23x10-15xR2= 0.994

-1.4

-1.2

-1.0

-0.8

-0.6

-0.4

-0.2

0.00.0 0.5 1.0 1.5 2.0 2.5 3.0 3.5


ln[I

r/Io]

a)

y = -2.26x10-15xR2 = 0.826

-1.2

-1.0

-0.8

-0.6

-0.4

-0.2

0.00.0 0.5 1.0 1.5 2.0 2.5 3.0 3.5


ln[I

r/Io]

b)

y = -2.26x10-15xR2 = 0.826

-1.2

-1.0

-0.8

-0.6

-0.4

-0.2

0.00.0 0.5 1.0 1.5 2.0 2.5 3.0 3.5


ln[I

r/Io]

y = -2.26x10-15xR2 = 0.826

-1.2

-1.0

-0.8

-0.6

-0.4

-0.2

0.00.0 0.5 1.0 1.5 2.0 2.5 3.0 3.5


ln[I

r/Io]

b)

y = -4.48x10-15x

R2 = 0.993

-1.6

-1.4-1.2

-1.0

-0.8

-0.6

-0.4

-0.2

0.00.0 0.5 1.0 1.5 2.0 2.5 3.0 3.5


ln[I

r/Io]

c)

y = -4.48x10-15x

R2 = 0.993

-1.6

-1.4-1.2

-1.0

-0.8

-0.6

-0.4

-0.2

0.00.0 0.5 1.0 1.5 2.0 2.5 3.0 3.5


ln[I

r/Io]

c)

y = -4.23x10-15xR2= 0.994

-1.4

-1.2

-1.0

-0.8

-0.6

-0.4

-0.2

0.00.0 0.5 1.0 1.5 2.0 2.5 3.0 3.5


ln[I

r/Io]

a)

y = -4.23x10-15xR2= 0.994

-1.4

-1.2

-1.0

-0.8

-0.6

-0.4

-0.2

0.00.0 0.5 1.0 1.5 2.0 2.5 3.0 3.5


ln[I

r/Io]

a)

38

Bridging Oxygen: For Au2O+ and Au3O

+, shown in Figure 3-2, the oxygen is

bound at a bridging site between two gold atoms. The barriers for removing bridging

oxygen are high and it was expected that no reaction would occur based on the study of

anionic gold clusters containing bridging oxygen (27). A selective reaction channel for O

atom replacement was observed as shown in the branching ratios of Figure 3-3(b-c). This

reaction channel was very selective and a transfer of charges may play a large role in the

observed products. It is possible that oxygen is only bound electrostatically or that

charge density may lower the barriers for O atom removal through creating a strong Au-C

bond. A comparison of the rate of O atom replacement by CO with peripheral oxygen

versus bridging oxygen reveals a slight decrease in the reaction rate. The plots in Figure

3-4 for Au2O+ and Au3O

+ show a kinetic rate of 3.6x10-12 ± 1.6 cm3s-1 and 2.8x10-12 ± 3.1

cm3s-1, respectively, for replacement of the bridging O atom. Therefore the calculated

rate constants are decreased by a factor of 2 compared to AuO+ which contains a

peripheral oxygen atom. However, the large standard deviation in the rate constant

indicates that the kinetic rate is relatively unaffected and the reaction proceeds.

Molecular Oxygen: The cluster species that undergo atomic oxygen replacement

by CO and also contain adsorbed molecular oxygen are AuO2+, AuO4

+, and Au3O4+.

These species are unique because they seemingly have the ability to break the strong O-O

bond in O2 subunits attached to gold. The branching ratio for AuO2+ is shown in Figure

3-6, and is presented below as a test case. Consider the reaction energies of the following

equations for neutral species:

CO + O2 CO2 + O, (3-3)

Au-O2 + CO Au-CO2 + O. (3-4)

39

Figure 3-5: Branching ratio for AuO3

+ containing a peripheral and molecular oxygen unit in the lowest energy structure calculation. Notice that CO is attached to all the reaction products observed.

Figure 3-6: Branching ratio for AuO2

+ which shows a selective pathway for O atom replacement by CO.

0

0.25

0.5

0.75

1

0 2 4 6 8 10

CO (mTorr)

AuO3+

AuOCO+

Au(CO)2+

AuCO+

Rel

ativ

e In

tens

ity

0

0.25

0.5

0.75

1

0 2 4 6 8 10

CO (mTorr)

AuO3+

AuOCO+

Au(CO)2+

AuCO+

0

0.25

0.5

0.75

1

0 2 4 6 8 10

CO (mTorr)

AuO3+

AuOCO+

Au(CO)2+

AuCO+

Rel

ativ

e In

tens

ity

0

0.25

0.5

0.75

1

0 4 8 12

CO (mTorr)

AuO2+

AuCO2+

0

0.25

0.5

0.75

1

0 4 8 12

CO (mTorr)

AuO2+

AuCO2+

Rel

ativ

e In

tens

ity

40

Equations 3-3 and 3-4 are exothermic by 0.4 eV and 0.6 eV, respectively. Since Au-O2

and CO form AuCO2 in an exothermic reaction, the reaction may proceed if any energetic

barriers are overcome. For eq 3-3 there is a 2.3 eV barrier without gold present and a 4

eV barrier if gold is attached (eq 3-4). The positive charge density may act in lowering

the barriers of this reaction by triplet-singlet spin state flipping (33). This process may

account for the breaking of the strong O-O bond by removing the large barrier to CO2

formation and replacing it with two smaller barriers that are surmountable (less than 0.2

eV). In order for spin state flipping to occur, both singlet and triplet spin states must

cross or lie close in energy on the potential energy surface. A transition to the singlet spin

state would increase the reactivity of the species and account for our experimental

findings. It is interesting to note recent theoretical studies by Vach et al. found mild

molecular collisions were able to induce triplet to singlet oxygen transitions explained by

a nonadiabatic “ladder climbing” mechanism (34). It is also important to point out that

no dimer clusters exhibited the O-O bond breaking. This may be due to the fact that

dimer species possess doublet ground states which are harder to induce transitions.

Other researchers have observed intermediate species that correspond to products

observed herein. Studies by Bernhardt and coworkers describe a proposed catalytic cycle

for anionic gold dimers in the presence of O2 and CO and reveal that Au2CO2- is a short

lived weakly bound species (35). Although this species is not observed in the anionic

case, for cations the strong Au+-C bonds formed may allow for a more stable

intermediate species to be observed in experiments. Additionally, a photoelectron spectra

of anionic, neutral, and cationic silver dimers found a temperature dependent

intermediate of the cationic silver dimer can dissociate O2 and bind atomic oxygen (36).

41

Finally, Kung and coworkers observed an IR peak which they assigned to an OC-Auδ+-

Oδ- species upon CO adsorption and O2 exposure (3). The appearance of products with O

atom loss further substantiates the observed mechanisms in the present studies as

intermediate species. Further calculations are being conducted by our collaborators in the

Bonaćič-Koutecký research group to explore the possible energetic pathways.

II) Molecular Oxygen Replacement

The dominant reaction channel for oxygen rich gold clusters is O2 replacement by

CO. This reaction pathway is observed for the following cluster species: AuO3+, AuO4

+,

AuO5+, Au2O2

+, Au2O3+, Au2O4

+, Au2O5+, Au3O2

+, Au3O3+, Au3O4

+, Au3O5+, Au4O3

+,

Au4O4+, Au4O5

+. By comparing collisional products observed in studies with N2 with CO

products (Table 3-1), this reaction channel was identified. CO association was also a

possible reaction pathway in the case of oxygen rich clusters, AuO4+ and AuO5

+, since O2

was lost under nitrogen. However, the percentage of O2 loss observed with nitrogen was

lower than products produced with CO reactant gas. This is evidence that the chemical

reaction, O2 replacement by CO, occurs. As described previously the bond energy for O2

on Au is ~0.6 eV which is overcome from the energy released in binding CO to the

cluster.

A maximum of two CO molecules bound to gold, as demonstrated in Figure 3-5

for AuO3+, was observed. The only products identified that associate two CO molecules

were Au(CO)2+ from AuO2

+, AuO3+, Au2O

+, and Au2O4+ cluster species and Au3O(CO)2

+

produced from Au4O5+. Both AuCO+ and Au(CO)2

+ were found to be more stable

species than Au and CO separately as shown in density functional theory (DFT)

calculations (37-38). Products with a maximum of two CO molecules bound to the gold

42

clusters were also observed in flow tube reactor experiments (31). Previous flow tube

studies verified the initial appearance of AuCO+, Au(CO)2+, AuO2CO+, Au2CO+, and

Au2OCO+ species and with additional CO reactant gas each of these species decreased or

disappeared (31). While there was no observance of the addition of Au(CO)3+ or species

with increasing numbers of CO attached, these species with CO attached may be key

reaction intermediates. The presence of both oxygen and carbon monoxide on a gold

cluster is important in CO oxidation for changing the electronic structure of the cluster

and allowing for the activation of bonds (39).

The clusters most selective for replacement of an oxygen atom or molecule by CO

however, may not be good candidates in the design of more active catalysts. When CO

binds to a Au cluster, it may displace the oxygen necessary to participate in the CO

oxidation. Studies have shown that too much CO can poison active gold sites (6, 9). If

CO binds to the cluster first, it may block O2 activation. It is interesting to note that on

free gold anions, O2 and CO has been found to cooperatively co-adsorb to the cluster with

each preferring different sites (40). From the observed CO attachment products and the

strength of Au+-CO bonds, these studies may show that O2 and CO prefer the same site to

adsorb, thus creating a competitive adsorption and decreasing the oxidation channel.

3.4.2 Au Fragmentation/Replacement

The fragmentation of a Au atom and replacement by CO was also observed in

select larger cluster species including Au4O2+ and Au4O3

+ shown in Figure 3-7.

43

Figure 3-7: Branching ratios for a) Au4O2+ and b) Au4O3

+ and their respective products with increasing carbon monoxide reactant gas.

0

0.25

0.5

0.75

1

0 2 4 6 8 10

CO (mTorr)

Au4O3+

Au4O+

Au3OCO+

Au4O2CO+

Au4O2+R

elat

ive

Inte

nsit

y

0

0.25

0.5

0.75

1

0 2 4 6 8 10

CO (mTorr)

Au4O3+

Au4O+

Au3OCO+

Au4O2CO+

Au4O2+R

elat

ive

Inte

nsit

yR

elat

ive

Inte

nsit

y

0

0.25

0.5

0.75

1

0 3 6 9 12CO (mTorr)

Au4O2+

Au4+

Au3CO+

Au4OCO+

Au4O+R

elat

ive

Inte

nsit

y

0

0.25

0.5

0.75

1

0 3 6 9 12CO (mTorr)

Au4O2+

Au4+

Au3CO+

Au4OCO+

Au4O+

)

)

a)

b)

44

There was dissociation of Au, AuO, AuO2, AuO3, AuO4, and AuO5 from the selected

cluster, which was not directly observed in our experiments. Assignments were based on

the selected cluster minus the detected ionic product. For example, Figure 3-7 shows

Au3OCO+ as a product from Au4O3+ suggesting fragmentation of neutral AuO2. No

fragmentation products were observed with nitrogen experiments under similar

experimental conditions of pressure and energy; therefore, CO was responsible for

activating the fragmentation of a Au-O or Au-Au bond in the cluster. It is reasonable for

Au to fragment from the selected metal cluster since Au-C bonds are stronger than Au-

Au bonds (33). Therefore these reaction products show that the weaker bond (Au-Au)

was broken and replaced by the stronger bond (Au-C), further evidence that Au-N was

not a strong enough bond to fragment the cluster.

3.4.3 Oxidation

Au4O2+ and Au4O3

+ were identified to be slightly active toward the oxidation of

CO; however, this was not the dominant reaction pathway as shown in the branching

ratios of Figure 3-7. Calculations on larger cationic gold clusters were not conducted;

therefore, the structures for these two species are not known. However, according to the

monomer and dimer structures in Figure 3-2, Au4O2+ and Au4O3

+ are expected to contain

bridging and molecular oxygen groups.

Heiz, Landman and coworkers showed anionic gold clusters were very size

dependent and that Au8 deposited on MgO was the smallest size cluster to catalyze the

oxidation of CO (41). Here, at least four gold atoms are necessary in order to observe the

45

oxidation reaction with a positive charge. These studies show the importance of size on

reactivity and the ability of charge to redistribute on larger clusters of gold to possibly aid

in the oxidation of CO.

3.5 Comparison to Anionic Gold Oxides

As discussed throughout this chapter, there were remarkable differences between

small anionic and cationic gold oxide clusters and the reaction pathways observed in the

presence of CO. For anionic monomer and dimer gold oxide species, AuO-, AuO3-,

Au2O-, Au2O3

-, and Au2O4-, CO oxidation was found to occur (27, 29); whereas, Au-CO

bonds were prevalent for cationic clusters. Both AuCO+ and Au(CO)2+ were calculated

by DFT to be more stable species than Au and CO separately (37-38) due to shorter CO

bond lengths (42). Gold carbonyl cations are generally easier to form compared to anions

because of the ability to form strong Au+-CO bonds (5, 43). The lone pair electrons on

CO bind to positive gold which has deficient electron density allowing for σ donation

from the sσ* orbital of CO into the metal σ orbital (44). Carbon monoxide saturation

studies of bare cationic gold clusters showed that uptake of CO was dependent on the

cluster size, geometry and available binding sites while anionic cluster uptake of CO was

governed by different mechanisms (38). Weaker CO bonds are formed with anionic gold

due repulsion between the filled metal s orbital and CO sσ* orbital (44). This difference

in the behavior of anionic and cationic gold is an example of how charge density affects

the reactivity of small gold clusters.

46

Cations also were more efficient in their respective reaction pathways compared

to gold anionic clusters. As the kinetic analysis of oxygen replacement by CO shows a

rate constant on the order of 10-12 cm3s-1, the anionic gold oxides that underwent CO

oxidation were on the order of 10-13 cm3s-1 (27).

3.6 Conclusions

This chapter provides support for the condensed phase catalyst study of Gates et

al., which showed that cationic gold alone was able to oxidize CO to CO2 (13). Gas phase

studies of cationic gold clusters found both Au4O2+ and Au4O3

+ species to be the only

active stoichiometries in the oxidation of CO to CO2. Although this reaction channel was

identified, it was not a very selective pathway. Competition between other pathways

including replacement by CO reduced the efficiency of the oxidation channel. Further

calculations need to be conducted and a study involving larger clusters would help

elucidate trends and establish a relationship for the oxygen atom transfer pathway. Small

cationic gold oxide clusters were found to be less active and selective in the oxidation

mechanism than their anionic counterparts.

These experiments focused on the oxygen replacement by CO reaction pathway.

Findings show O2 molecules were fragmented to form products with a stoichiometry of

AuxOy-1CO+. Although the energetic pathways of this O atom replacement channel are

not clear, a possible mechanism to account for the experimental findings was proposed.

Therefore the charge state and the strength of CO binding have a great effect on the

observed reaction pathways.

47

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7. Yoon, B.; Häkkinen, H.; Landman, U. J. Phys. Chem. A 2003, 107, 4066.

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10. Guzman, J.; Gates, B. C. J. Am. Chem. Soc. 2004, 126, 2672.

11. Bond, G. C.; Thompson, D. T. Gold Bulletin 2000, 33, 41.

12. Guzman, J.; Gates, B. C. J. Phys. Chem. B 2002, 106, 7659.

13. Fierro-Gonzalez, J. C.; Gates, B. C. J. Phys. Chem. B 2004, 108, 12529.

14. Fu, Q.; Deng, W.; Saltsburg, H.; Flytzani-Stephanopoulos, M. Appl. Catal. B 2005, 56, 57.

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Kung, M. C.; Kung, H. H. Appl. Catal. A 2003, 243, 15.

16. Costello, C. K.; Kung, M. C.; Oh, H.-S.; Wang, Y.; Kung, H. H. Appl. Catal. A 2002, 232, 159.

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17. Guzman, J.; Carrettin, S.; Corma, A. J. Am. Chem. Soc. 2005, 127, 3286.

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N. Catal. Today 2002, 72, 157.

20. Dekkers, M. A. P.; Lippits, M. J.; Nieuwenhuys, B. E. Catal. Lett. 1998, 56, 195.

21. Boccuzzi, F.; Cerrato, G.; Pinna, F.; Strukel, G. J. Phys. Chem. B 1998, 102, 5733.

22. Kung, H. H.; Kung, M. C.; Costello, C. K. J. Catal. 2003, 216, 425.

23. Haruta, M.; Daté, M. Appl. Catal. A 2001, 222, 427.

24. Oh, H. S.; Costello, C. K.; Cheung, C.; Kung, H. H.; Kung, M. C. Stud. Surf. Sci. Catal. 2001, 139, 375.

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Mitrić, R.; Bonaćič-Koutecký, V. J. Chem. Phys. 2006, 125, 204311.

28. Kimble, M. L.; Bürgel, C.; Moore, N. A.; Mitrić, R.; Bonaćič-Koutecký, V.; Castleman, A. W., Jr. “Is O2 Dissociation Sufficient to Effect the Oxidation of CO on Anionic Gold Oxide Clusters?” in prep.

29. Kimble, M. L.; Castleman, A. W., Jr.; Mitrić, R.; Bürgel, C.; Bonaćič-Koutecký,

V. J. Am. Chem. Soc. 2004, 126, 2526.

30. Kimble, M. L. Ph. D. Thesis The Pennsylvania State University, 2005.

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44. Bernhardt, T. M. Int. J. Mass Spec. 2005, 243, 1.

Chapter 4

Reactivity of Ag2Oy+,- and Bimetallic AgAuOy

+,- (y=1-4) Ionic Clusters with CO

4.1 Introduction

In contrast to the widely investigated structural, electronic, and reactivity

properties of pure gold (1-3) and silver clusters (4-7), there is considerably less

information available for bimetallic silver-gold clusters (8-12). Calculations by Bonačić-

Koutecký and coworkers illustrated the role of charge transfer on catalytic activity and

the differences in s and d electrons for bonding in bimetallic systems compared with pure

gold or silver (13). Due to the larger relativistic effects of gold, where the 6s and 5d

orbitals lie close in energy, there is significant contribution from the d electrons. For

silver clusters, the s and d shells are energetically well separated and findings for

bimetallic Ag-Au clusters exhibit structures similar to pure silver (14-15).

It has been theoretically predicted that the interaction of AgAu- with O2 and CO

would promote the oxidation of CO (13-14). Charge transfer from Ag to Au would

create negatively charged gold subunits (13). Negatively charged gold was expected to

be the active site in the oxidation of CO. Indeed, previous experiments of pure gold

anions found the oxidation of CO to CO2 on Au2O-, Au2O3

- and Au2O4- (16-17).

However, positively charged Ag species may provide an active site that enhances CO

oxidation through O-O bond activation which has been observed in previous pure silver

studies (5, 18). Wang and coworkers found a synergistic effect between silver and gold

51

that resulted in higher catalytic activity with silver proposed to be a key factor in oxygen

activation (19). Although reaction schemes for the oxidation of CO are proposed, no

agreement has been reached. In fact, experiments by Bernhardt et al. have reported no

reactive behavior towards CO with Au2-, Ag2

- or AgAu- (4).

Kappes et al. have studied cationic AgmAun+ clusters with CO and observed

cluster carbonyls of the form AgmAunCO+ (11). They also revealed that CO binds to the

gold atoms, with binding energies ranging from 0.77 to 1.09 eV. Calculations of surface

Ag(100) find CO adsorbs with lower energies ranging from 0.12 to 0.19 eV (7).

Structural calculations of larger clusters, AgmAun0,+,- where n+m ≥ 3, revealed gold atoms

to be embedded on Ag in exposed positions and also found hetero-bonding more

prevelant than homo-nuclear bonding in order to enable the withdraw of charge from Ag

atoms (13).

Through experiments and theory some information on the properties of mixed

clusters has been achieved; however, the active sites and mechanisms for the reactivity of

silver-gold clusters are still not fully understood. Comparison of the reactions of pure

metal oxides with CO to experiments of bimetallic oxides can aid in determining the

influence that each individual metal exhibits and the changes that occur in reactivity

when another metal is present. This chapter focuses on elucidating the role of charge

state and composition of Ag-Au oxide clusters in the presence of CO.

52

4.2 Experimental

Experiments were conducted utilizing a guided ion beam mass spectrometer

explained in detail in Chapter 2. Briefly, the clusters were formed using the modified

laser vaporization source which employed the second harmonic of a Nd:YAG laser to

ablate two rotating and translating metal rods (gold, silver, SurePure Chemetals). A half

circular mirror was used to split the laser beam from a single laser into two beams able to

ablate two separate metal rods. Oxygen, approximately 20% for cations and 50% for

anion clusters, seeded in helium was pulsed over the two rods. After exiting the source

through a 51 millimeter conical expansion nozzle, the clusters were cooled by supersonic

expansion in a field free region. The clusters were then passed through a 3 mm skimmer

and directed by a set of electrostatic lenses and deflectors into the first quadrupole for

mass selection of the reactant species. The cluster of interest was focused through a

second set of lenses into the octopole reaction cell. CO reactant gas, ranging in pressure

from 0-10 mTorr for cationic and 0-20 mTorr for anionic clusters, was added to the

reaction cell. The pressure was monitored using a MKS baratron. The remaining

reactant and product clusters were then focused by a third set of lenses into a second

quadrupole for mass analysis and detected by a channel electron multiplier.

4.3 Results

Experiments of mixed Ag-Au dimer oxide clusters reveal similarities and

differences in reaction products compared to the dimer of either pure silver or gold oxide

clusters. Pure silver ionic clusters are presented first followed by bimetallic Ag-Au oxide

53

clusters. For completeness, gold oxide reactions were compared here to pure silver and

bimetallic clusters in the discussion that follows. Pure gold oxide anions were reported

previously (16-17, 20) and detailed information on cationic clusters was discussed in

Chapter 3.

4.3.1 Pure Silver Oxide Ions

The reaction products of pure silver dimers are shown in Table 4-1. The dominant

reaction pathway for cationic silver dimer oxides was oxygen atom replacement by CO.

The product stoichiometry of AgxOy-1CO+ was observed for Ag2O2+, Ag2O3

+ and Ag2O4+.

The following equation, eq 4-1, shows that at multiple collision conditions, the oxidation

of CO may occur on silver cations through cooperative effects.

Ag2O2+ + CO Ag2CO2

+ + CO2 (4-1)

Thereby, two CO molecules may be able to dissociate the strong O-O bond as a result of

multiple collisions and excess CO molecules. This accounts for the first reaction product

observed in Figure 4-1. The following equation, eq 4-2, describes a further reaction

channel that results from the product Ag2CO2+ shown in eq 4-1, as CO2 desorbs from the

silver cation.

Ag2CO2+ Ag2

+ + CO2 (4-2)

This desorption reaction may be possible as the Ag2CO2+ molecule has time to relax and

rearrange. This also gives rise to the other product, Ag2+, detected in the branching ratio.

The fact that two products were observed in the spectra signifies that this reaction may

occur sequentially.

54

Table 4-1: Product list for the dimer of pure silver, pure gold, and mixed Ag-Au oxide clusters in the presence of CO.

Cations Anions

Ag2Oy+ Au2Oy

+ AgAuOy+ Ag2Oy

- Au2Oy- AgAuOy

-

y=1 AgCO+ Au2CO+ AgAuCO+

Ag2+

y=2 Ag2 + Au2

+ AgAu + y=2 No rxn No rxn AgAuO-

Ag2CO2+ Au2CO+ AgAu-

y=3 Ag2O2CO+ Au2CO+ AgAuO2CO+ y=3 AgO- Au2O2- AgAuO-

Ag2O+ Au2O

+ AgAuCO+ AgAu-

AuOCO+ AgAuO+

Au(CO)2+ AgO2

+

y=4 Ag2O3CO+ Au2CO+ AgAuO3CO+ y=4 Au2O3- AgAuO2

-

Ag2CO+ AgAuO2+

Ag2O(CO)2+

Ag2O2(CO)2+

55

Figure 4-1: Branching ratios for a) Au2O2

+, b) AgAuO2+, c) Ag2O2

+, and d) AgAuO2-

showing the products in the reaction with CO reactant gas.

0

0.25

0.5

0.75

1

0 2 4 6 8 10

CO (mTorr)

Au2O2+

Au2CO+

Au2+

0

0.25

0.5

0.75

1

0 2 4 6 8 10

CO (mTorr)

AgAuO2+

AgAu+

0.96

0.98

1

0 7 14 21 28

CO (mTorr)

0

0.01

0.02

AgAuO2-

AgAu-

AgAuO-

0

0.25

0.5

0.75

1

0 2 4 6 8 10

CO (mTorr)

Ag2O2+

Ag2CO2+

Ag2+

Rel

ativ

e In

tens

ity

a)

b)

c)

Rel

ativ

e In

tens

ity

Rel

ativ

e In

tens

ityR

elat

ive

Inte

nsit

y (P

aren

t ion

) Relative Intensity (Product ions)

d)

0

0.25

0.5

0.75

1

0 2 4 6 8 10

CO (mTorr)

Au2O2+

Au2CO+

Au2+

0

0.25

0.5

0.75

1

0 2 4 6 8 10

CO (mTorr)

AgAuO2+

AgAu+

0.96

0.98

1

0 7 14 21 28

CO (mTorr)

0

0.01

0.02

AgAuO2-

AgAu-

AgAuO-

0

0.25

0.5

0.75

1

0 2 4 6 8 10

CO (mTorr)

Ag2O2+

Ag2CO2+

Ag2+

Rel

ativ

e In

tens

ity

a)

b)

c)

Rel

ativ

e In

tens

ity

Rel

ativ

e In

tens

ityR

elat

ive

Inte

nsit

y (P

aren

t ion

) Relative Intensity (Product ions)

d)

56

This sequential mechanism is in contrast to the proposed singlet to triplet transition

discussed in Chapter 3 for AuO2+. First, the doublet ground spin state hinders the

proposed spin transition and no intermediate was observed as in the present case for

silver dimers.

Oxygen atom transfer was observed for Ag2O+ species reacted with CO, while no

oxygen atom was lost with N2. The positive charge on silver may weaken the Ag-O bond

to allow for easy formation of CO2. However, this was a very minor channel. The

dominant channel for Ag2O+ was AgO replacement by CO in which the weaker Ag-Ag

bond was broken forming a stronger Ag-C bond.

Finally, no reaction was detected for Ag2O2- and Ag2O3

- cluster species showed

O2 loss.

4.3.2 Bimetallic Ag-Au Oxide Ions

Typical mass distributions for anionic and cationic clusters presented in Figure 4-

2 show pure silver, pure gold, and mixed silver-gold oxides in the spectrum. The spectra

demonstrate the ability to form mixed dimer cluster species, specifically for clusters of

AgAuOx-,+ where x=1-5, employing the dual rod source described in Chapter 2. The laser

vaporization source allowed for dissociated and molecular oxygen to adsorb onto the

clusters. This enabled the study of Ag-Au oxide clusters with oxygen atoms bound.

Previous calculations have been conducted on the interaction of Ag-Au clusters and

molecular oxygen, but have not focused on dissociated oxygen (14).

57

Figure 4-2: a) Typical anionic silver-gold oxide mass distribution and b) cationic silver-gold oxide mass distribution obtained from dual rod LaVa source.

50 150 250 350 450

150 200 250 300 350 400

AgO

4-

AuO

-AgO

6-

AuO

2-A

uO3-

Ag 2O

3-A

uO4-

Ag 2O

5-

Ag 2O

6-

AuA

gO-

Ag 2O

7-

AuA

gO2-

AuA

gO3-

Ag 2O

8-

AuA

gO4-

AuA

gO5-

AgO

3+

AgO

+

Ag+

AgO

2+

AgO

4+A

u+

AuO

+

AuA

g+

Ag 2

O+

AuA

gO+

AuA

gO3+

AuO

2+

AuA

gO4+

Au 2+

Ag 2O

2+

Ag 2O

3+A

uO3+

Ag 2O

4+

AuA

gO5+

Mass (amu)

Arb

itrar

y In

tens

ity

Arb

itrar

y In

tens

ity

a)

b)

50 150 250 350 450

150 200 250 300 350 400

AgO

4-

AuO

-AgO

6-

AuO

2-A

uO3-

Ag 2O

3-A

uO4-

Ag 2O

5-

Ag 2O

6-

AuA

gO-

Ag 2O

7-

AuA

gO2-

AuA

gO3-

Ag 2O

8-

AuA

gO4-

AuA

gO5-

AgO

3+

AgO

+

Ag+

AgO

2+

AgO

4+A

u+

AuO

+

AuA

g+

Ag 2

O+

AuA

gO+

AuA

gO3+

AuO

2+

AuA

gO4+

Au 2+

Ag 2O

2+

Ag 2O

3+A

uO3+

Ag 2O

4+

AuA

gO5+

Mass (amu)

Arb

itrar

y In

tens

ity

AgO

+

Ag+

AgO

2+

AgO

4+A

u+

AuO

+

AuA

g+

Ag 2

O+

AuA

gO+

AuA

gO3+

AuO

2+

AuA

gO4+

Au 2+

Ag 2O

2+

Ag 2O

3+A

uO3+

Ag 2O

4+

AuA

gO5+

Mass (amu)

Arb

itrar

y In

tens

ity

Arb

itrar

y In

tens

ity

a)

b)

58

The present experiments therefore provide a report of bimetallic Ag-Au adsorbed with O

and O2 and the interactions in the presence of CO. Table 4-1 also lists the products for

mixed silver-gold cationic and anionic dimer clusters reacted with CO. Below, each

ionic dimer cluster is presented according to the amount of oxygen attached to the metal

ion.

AgAuO+,-: For a predissociated oxygen atom attached to cationic Ag-Au, the

bimetallic species behaved similar to the gold dimer oxide. Both Au2O+ and AgAuO+

underwent O atom replacement by a CO molecule. It can be seen in Figure 4-3 the

branching ratio of Au2O+ was more reactive toward O atom replacement compared to the

mixed AgAuO+. A kinetic analysis of the replacement rate constant value reveals the

presence of silver decreased the efficiency of the reaction channel. Figure 4-4 shows

plots of ln[Ir/Io] versus the CO reactant concentration in which the slope is equal to –kt.

Solving for the kinetic rate constant showed Au2O+ has a rate of 3.6x10-12 ± 1.6 cm3s-1

while AgAuO+ is only 1.5x10-12 ± 0.2 cm3s-1. This decrease in the rate may be a

consequence of charge transfer from silver to gold, creating more electron density around

gold and a weaker CO bond.

None of the anionic clusters with one oxygen atom were formed in our laser

ablation source and therefore are not reported.

AgAuO2+,-: The most studied interaction in the literature is of AgAu with O2. The

present studies found that the AgAu+ cation behaved similar to both pure gold and pure

silver with an O2 loss reaction channel shown in Figure 4-1. It is interesting to note that

both pure gold and pure silver also underwent other reaction pathways which were more

dominant than O2 loss. Au2O2+, besides losing O2, also formed Au2CO+.

59

Figure 4-3: Branching ratios for a) Au2O

+ and b) AgAuO+ with increasing CO pressure.

0

0.25

0.5

0.75

1

0 2 4 6 8 10

CO (mTorr)

Au2O+

Au2CO+

0

0.25

0.5

0.75

1

0 2 4 6 8CO (mTorr)

AgAuO+

AgAuCO+

Rel

ativ

e In

tens

ity

Rel

ativ

e In

tens

ity

a)

b)

0

0.25

0.5

0.75

1

0 2 4 6 8 10

CO (mTorr)

Au2O+

Au2CO+

0

0.25

0.5

0.75

1

0 2 4 6 8CO (mTorr)

AgAuO+

AgAuCO+

Rel

ativ

e In

tens

ity

Rel

ativ

e In

tens

ity

a)

b)

60

Figure 4-4: Plots of ln[Ir/Io] versus CO reactant concentration for a) Au2O

+ and b) AgAuO+

y = -9.73x10-16xR

2= 0.874

-0.3

-0.25

-0.2

-0.15

-0.1

-0.05

0.00.0 0.5 1.0 1.5 2.0 2.5 3.0


ln[I

r/Io]

y = -2.26x10-15xR2 = 0.826

-1.2

-1.0

-0.8

-0.6

-0.4

-0.2

0.00.0 0.5 1.0 1.5 2.0 2.5 3.0 3.5


ln[I

r/Io]

a)

b)

y = -9.73x10-16xR

2= 0.874

-0.3

-0.25

-0.2

-0.15

-0.1

-0.05

0.00.0 0.5 1.0 1.5 2.0 2.5 3.0


ln[I

r/Io]

y = -2.26x10-15xR2 = 0.826

-1.2

-1.0

-0.8

-0.6

-0.4

-0.2

0.00.0 0.5 1.0 1.5 2.0 2.5 3.0 3.5


ln[I

r/Io]

a)

y = -2.26x10-15xR2 = 0.826

-1.2

-1.0

-0.8

-0.6

-0.4

-0.2

0.00.0 0.5 1.0 1.5 2.0 2.5 3.0 3.5


ln[I

r/Io]

a)

b)

61

In the case of pure silver, an oxygen atom replacement product was also detected

(Ag2CO2+), as described above in the previous section. Therefore the bimetallic cluster

may aid in enhancing the selectivity one of the reaction channels, O2 loss.

For anionic clusters, Figure 4-1 shows the branching ratio of AgAuO2- which

revealed an oxygen atom transfer pathway and also O2 loss channel. It is difficult

determine whether this is a sequential reaction mechanism, in which AgAuO- transfers its

first oxygen to CO, followed by its second oxygen atom, shown below in eqs 4-3 and 4-4.

AgAuO2- + CO AgAuO- + CO2 (4-3)

AgAuO- + CO AgAu- + CO2 (4-4)

The small intensity of the products and the fact that this reaction occurs very slowly

complicates the ability to definitively assign the mechanism. The behavior of the mixed

metal is different from pure gold in which no reaction is observed and pure silver which

showed no reaction was detected at the highest pressures.

AgAuO3+,-: As the Ag-Au cluster becomes more oxygen rich, there is a shift in the

behavior toward pure silver oxide reactivity. For cationic Ag2O3+ and AgAuO3

+ clusters,

the major product was Ag2O2CO+ and AgAuO2CO+, respectively. This corresponds to an

oxygen atom replacement by a CO molecule and could be an intermediate species for CO

oxidation to CO2. These two clusters also show Ag2O+ and AuAgO+ for silver and

bimetallic species which suggests dissociation of a CO2 molecule for the metal center.

The bimetallic Ag-Au cluster also had a reaction product corresponding to neutral AuO

loss. The AgO2+ product species may arise because the Ag-Au bond is weaker in energy

then the Au-C bond. The dissociation of a metal atom has been observed with larger

62

mixed clusters, Ag1-3Au1-3+, by Neumaier et al. (11). They showed that the CO binding

energies were greater than the activation energy for losing a metal atom.

In the case of anions, mixed AgAuO3- behaved similar to Ag2O3

- with both

species showing loss of molecular oxygen. On the other hand, pure gold oxide produced

oxygen atom transfer to CO.

AgAuO4+,-: Both cationic Ag2O4

+ and AgAuO4+ clusters underwent O atom

replacement by a CO molecule. For the silver-gold cluster, AgAuO2+ was also observed.

In the case of anionic AgAuO4- an O2 loss is the only reaction product detected.

This does not correspond with either pure metal, as silver signal was too low to obtain

and the gold anion dimer underwent oxygen atom transfer.

4.4 Discussion

For both AgAuO+ and AgAuO2+, the bimetallic species exhibit products that

correspond to the reactivity of pure gold dimer oxide. There is evidence, however, that

the reactivity of the mixed species was affected by the presence of silver. For the

reaction of AgAuO2+, there was only one reaction product similar to that detected in pure

gold or silver oxide reactions with CO. Silver may have influenced the gold toward a

selective pathway for O2 loss. This supports the existence of structural and electronic

changes for catalytic systems that contain two different metals and the ability to tune the

properties of a catalyst based on composition.

In the case of AgAuO3+ and AgAuO4

+ when more oxygen is surrounding the

cluster, the bimetallic species produced similar products to pure silver dimer oxides. This

63

may be the result of charge redistribution for highly coordinated silver. It has been

shown that as more oxygen binds to silver, silver induces an O2 bond elongation which

results in activation (4, 21). This would account for the increased dominance by silver as

O-O bonds become weaker and more active.

The major reaction channel observed for AgAuO3-4+ was O atom replacement by

CO. This was an important reaction pathway exhibited by Ag2O2-4+ and select cationic

gold monomer oxide clusters. This pathway for gold oxides was proposed to be the result

of charge density effects which allowed for a spin state transition. However, as outlined

above for silver oxides, the mechanism for silver and bimetallic oxide clusters may differ

from gold oxides. The appearance of other products in the branching ratio suggests that O

atom replacement products may be intermediate species in the oxidation of CO. Other

studies have shown the existence of AgxOy-1CO+ and AuOCO+ intermediate species (5,

18, 22). It is also possible that multiple CO molecules induce cooperative effects to

promote the observed reactions. It has been shown in gold clusters by Whetten and

coworkers that when multiple CO molecules adsorb onto the cluster, changes in the

electronic structure occur (23).

Since the Ag2O4- products with CO are unknown, it is hypothesized that the

pattern in mixed cluster behavior would be similar to the silver behavior. In the case of

AgAuO3+ and AgAuO3

- both reaction products were analogous to silver reactions. It is

logical that since AgAuO4+ corresponds to silver products, that AgAuO4

- would also

likely follow this trend.

64

4.5 Conclusion

The bimetallic Ag-Au oxide studies show that oxygen deficient cationic clusters

behave similarly to pure gold dimers while clusters that are more oxygen rich have

comparable reaction products to pure silver dimers. For anionic clusters, there was less

connection between the pure and mixed clusters. Trends in the behavior have been

hypothesized; however, most often the Ag-Au species behaved differently then either

pure dimer oxides.

Collaboration with the Bonaćič-Koutecký research group to determine possible

structure-reactivity relationships that may be present for several cluster species is

ongoing. The study of larger mixed clusters with multiple gold and silver atoms would

shed light on the effects of size and stoichiometry for the observed reaction pathways.

65

References



3. Gilb, S.; Weis, P.; Furche, F.; Ahlrichs, R.; Kappes, M. J. Chem. Phys. 2002, 116,

4094.

4. Bernhardt, T. M.; Socaciu-Siebert, L. D.; Hagen, J.; Wöste, L. Appl. Catal. A 2005, 291, 170.

5. Socaciu, L. D.; Hagen, J.; Heiz, U.; Bernhardt, T. M.; Leisner, T.; Wöste, L.

Chem. Phys. Lett. 2001, 340, 282.

6. Weis, P.; Bierweiler, T.; Gilb, S.; Kappes, M. Chem. Phys. Lett. 2002, 355, 355.

7. Qin, C.; Sremaniak, L. S.; Whitten, J. L. J. Phys. Chem. B 2006, 110, 11272.

8. Negishi, Y.; Nakamura, Y.; Nakajima, A.; Kaya, K. J. Chem. Phys. 2001, 115, 3657.

9. Mitrić, R.; Hartmann, M.; Stanca, B.; Bonaćič-Koutecký, V.; Fantucci, P. J. Phys.

Chem. A 2001, 105, 8892.

10. Bauschlicher, C. W., Jr.; Langhoff, S. R.; Partridge, H. J. Chem. Phys. 1989, 91, 2412.

11. Neumaier, M.; Weigend, F.; Hampe, O.; Kappes, M. M. J. Chem. Phys. 2006,

125, 104308.

12. Cottancin, E.; Lermé, J.; Gaudry, M. Pellarin, M; Vialle, J. L.; Broyer, M. Phys. Rev. B 2000, 62, 5179.

13. Bonačić-Koutecký, V.; Burda, J.; Mitrić, R.; Ge, M.; Zampella, G.; Fantucci, P. J.

Chem. Phys. 2002, 117, 3120.

14. Mitrić, R.; Bürgel, C.; Burda, J.; Bonaćič-Koutecký, V.; Fantucci, P. Eur. Phys. J. D 2003, 24, 41.

66

15. Lee, H. M.; Ge, M.; Sahu, B. R.; Tarakeshwar, P.; Kim, K. S. J. Phys. Chem. B 2003, 107, 9994.



17. Kimble, M. L.; Castleman, A. W., Jr.; Mitrić, R.; Bürgel, C.; Bonaćič-Koutecký, V. J. Am. Chem. Soc. 2004, 126, 2526.

18. Socaciu-Siebert, L. D.; Hagen, J.; Le Roux, J.; Popolan, D.; Vaida, M.; Vajda, Š.;

Bernhardt, T. M.; Wöste, L. Phys. Chem. Chem. Phys. 2005, 7, 2706.

19. Wang, A. Q.; Chang, C. M.; Mou, C. Y. J. Phys. Chem. B 2005, 109, 18860.; Liu, J. H.; Wang, A. Q.; Lin, H. P.; Mou, C. Y. J. Phys. Chem. B 2005, 109, 40.; Wang, A. Q.; Liu, J. H.; Lin, S. D.; Lin, T. S.; Mou, C. Y. J. Catal. 2005, 233, 186.

20. Kimble, M. L.; Bürgel, C.; Moore, N. A.; Mitrić, R.; Bonaćič-Koutecký, V.;

Castleman, A. W., Jr. “Is O2 Dissociation Sufficient to Effect the Oxidation of CO on Anionic Gold Oxide Clusters?” in prep.

21. Hagen, J.; Socaciu, L. D.; Le Roux, J.; Popolan, D.; Bernhardt, T. M.; Wöste, L.;

Mitrić, R.; Noack, H.; Bonaćič-Koutecký, V. J. Am. Chem. Soc. 2004, 126, 3442.

22. Henao, J. D.; Caputo, T.; Yang, J. H.; Kung, M. C; Kung, H. H. J. Phys. Chem. B 2006, 110, 8689.


Chapter 5

Influence of Charge State on the Reaction of FeO3+,- with Carbon Monoxide

5.1 Introduction

Stimulated by recent theoretical findings regarding the activity of iron oxide

clusters for CO oxidation (1-2), a systematic experimental and theoretical investigation of

FexOy+,- clusters with CO was undertaken. Recent experiments on iron oxide

nanoparticles show that it may be possible to convert CO to CO2 at low temperatures in

the presence or absence of O2 (3-6). Transition metal oxides with oxygen atoms bound

only to the transition metal without O-O bonding offer an attractive possibility to

overcome the need for a catalyst to weaken/break the O-O bond through charge transfer

and/or bonding. In the past, Pdn and Ptn nanoparticles, and more recently charge Aun

clusters, have attracted attention (7). In comparison to precious metal based catalysts

used in effecting the oxidation of CO, iron oxide is a much more affordable and readily

available material.

The basic issue addressed herein is whether it is possible to use the stored oxygen

in highly oxidized and charged metal clusters to convert CO to CO2. Note that such a

possibility avoids the use of external O2 and hence may allow CO conversion in oxygen

free environments. The ratio of oxygen atoms to metal atoms in bulk metal oxides is

generally less than two. It is, however, possible to generate metal-oxygen clusters with

oxygen to metal ratios much higher than in the bulk (8). Further, unlike bulk, small

(Reproduced in part with permission from Chem. Phys. Lett. 2007, 435, 295-300. Copyright 2007 Elseveir.)

68

clusters offer the possibility of different ionic charge states. Numerous studies involving

charged iron and iron oxide clusters have been able to provide fundamental information

corresponding to bulk phase iron processes (8-12).

In this chapter, FeO3 is presented as the test case. Studies of both the cationic and

anionic clusters show that charge has a dramatic effect on the oxidation behavior. As

opposed to the case of Aun clusters where the findings have been that anions are more

effective in converting CO (7), here FeO3+ cations are far more efficient than the anions.

Also, the reaction pathways for cations are different from those for anions. In most cases,

the reaction proceeds without barriers and hence may allow CO oxidation at all

temperatures.

5.2 Methods

Experiments were conducted using a guided ion beam mass spectrometer

explained in detail in Chapter 2, and described here briefly. The clusters were formed by

laser vaporization employing the second harmonic of a Nd:YAG laser to ablate a rotating

and translating iron rod (PVD Materials Corp., 99.95% purity). Oxygen (1%) seeded in

helium was pulsed over the iron rod into the hot metal plasma where the clusters form.

After exiting the source through a 27 millimeter conical expansion nozzle, the clusters

were cooled by supersonic expansion in a field free region. The clusters were then

passed through a 3 mm skimmer and directed by a set of electrostatic lenses and

deflectors into the first quadrupole for mass selection of the reactant species. The cluster

of interest was focused through a second set of lenses into the octopole reaction cell. CO

69

reactant gas, ranging in pressure from 0-10 mTorr for cationic and 0-20 mTorr for

anionic clusters, was added to the reaction cell. The pressure was monitored by a MKS

baratron. The remaining reactant and product clusters were then focused by a third set of

lenses into a second quadrupole for mass analysis and detected by a channel electron

multiplier.

Theoretical studies were carried out by the Khanna research group using a first

principles electronic structure scheme employing a density functional formalism (13).

The exchange correlation effects were incorporated through a generalized gradient

approximation (GGA) functional as proposed by Perdew, Burke and Ernzerhof (14). The

electronic structure was determined using a linear combination of atomic orbitals

molecular orbital approach. The wave function for the cluster was constructed by a linear

combination of Gaussian type orbitals centered at the atomic positions in the cluster. The

actual calculations were performed using the deMon2k software (15). In this

implementation, an auxiliary function set is used for the variational fitting of the

Coulomb potential (16-17). The numerical integration of the exchange-correlation

energy and potential were then performed on an adaptive grid (18). In the present

studies, we employed the double zeta valence polarized (DZVP) basis sets (19) for C and

O and the Wachters-F basis set (20-22) for Fe. The GEN-A2 auxiliary function set for C

and O and the GEN-A2* auxiliary function set for Fe were used. To determine the

ground state, the configuration space was sampled by starting from several initial

configurations and optimizing the geometry by moving atoms in the direction of forces

until they dropped below a threshold value. Since transition metal atoms are marked by

non-zero spin multiplicities, the ground state determination included investigation over

70

several spin multiplicities. The geometries were optimized without any symmetry

constraint using delocalized internal coordinates with the rational function optimization

(RFO) and the Broyden, Fletcher, Goldfarb and Shanno (BFGS) update in the deMon2K

program (23).

5.3 Results and Discussion

Investigation into the reaction products for both charge states includes a look at

the observed experimental products, cluster structures, and kinetics involved in specific

reaction pathways. The branching ratios for FeO3+,- in the presence of increasing CO

pressure are shown in Figure 5-1. This figure shows the intensity of the parent ion as

well as all the product ions. Notice that the maximum CO pressure for the anion species

is twice that for the cation species. The experimental results reveal three important

findings. (1) The FeO3+ is far more reactive than the FeO3

-. Even though the maximum

pressure in the cation reaction is half the amount used in the anion reaction, the intensity

of the products is far higher for the cations. (2) For FeO3+, the reaction proceeds to the

final stages forming FeO2+, FeO+, and Fe+ as the products. For anions, only FeO2

- is

observed. (3) Only one reaction product with association of CO is observed, namely

FeOCO+. The intensity of this product is initially higher than that of FeO2+ indicating that

it is produced in the initial stages of reaction.

71

0.5

0.75

1

0 2.5 5 7.5 10CO (mTorr)

0

0.05

0.1

0.5

0.75

1

0 5 10 15 20

CO (mTorr)

0

0.05

0.1

FeO3+

FeOCO+

FeO+

Fe+

FeO2+

Rel

ativ

e In

tens

ity

(par

ent

ion)

Relative Intensity (product ions)

Rel

ativ

e In

tens

ity

(par

ent

ion)

Relative Intensity (product ion)

FeO3-

FeO2-

0.5

0.75

1

0 2.5 5 7.5 10CO (mTorr)

0

0.05

0.1

0.5

0.75

1

0 5 10 15 20

CO (mTorr)

0

0.05

0.1

FeO3+

FeOCO+

FeO+

Fe+

FeO2+

Rel

ativ

e In

tens

ity

(par

ent

ion)

Relative Intensity (product ions)

Rel

ativ

e In

tens

ity

(par

ent

ion)

Relative Intensity (product ion)

FeO3-

FeO2-

)

)

Figure 5-1: Branching ratios for a) FeO3

+ and b) FeO3- with increasing CO pressure. The

relative selected ion intensity is plotted on the left-hand axis and the relative product intensities on the right-hand axis. Note that the pressure used in the reactions of cationic FeO3

+ is half the amount used in reactions of anionic FeO3- to acquire reaction products

of the same order of magnitude.

a)

b)

72

Further evidence for this comes from the observation that its intensity is higher than that

of FeO2+ suggesting that it is not the product of association of CO to previously formed

FeO+ but that it is produced directly from the reaction of FeO3+. Theoretical studies on

the ground state geometries of FeOy+,- and the reaction pathways were carried out to shed

light on some of these findings.

Figure 5-2 shows the ground state geometries of FeOy+,- and FeOyCO+,- clusters

containing up to three oxygen atoms. For the FeOy+, clusters, note that in the lowest

energy state, there are no O-O bonds and that all the oxygen atoms are bound to the metal

atom. Our studies on larger sizes of metal clusters indicate that O-O bonds do form at

higher coverage. These will be presented in Chapters 6 and 7. A Mulliken population

analysis of the overall charge density was carried out to identify the location of missing

charge in cation clusters, extra charge in anion species, and any other charge transfers.

The calculated Mulliken charges are marked below the individual atoms. Irrespective of

the charge state, Figure 5-2 shows that there is a charge transfer from iron sites to the

oxygen atoms indicating that the FeO bond is stabilized by such transfers. The change in

the charge on oxygen, therefore, leads to a change in the strength of the Fe-O bond. To

show this more quantitatively, the energy required to remove an O atom and an O2

molecule from the various clusters was calculated. (It should be noted that for some

anionic clusters, it takes less energy to remove O-. In this work, however, the primary

focus is on neutral O and O2.) The relevant energy values are shown in Figure 5-3. Note

that the energy required to remove a single O atom from cationic clusters is significantly

lower than the energy required for removing an O atom from the corresponding anion of

the same size.

73

Figure 5-2: Ground state geometries for O2, CO, CO2, and FeOy

+,- clusters. The bond lengths are given in Angstroms and superscripts indicate the spin multiplicity. The Mulliken charges are marked below each atom. Source: Figure courtesy of the group of Professor S. N. Khanna at VCU.

74

Even more puzzling is the trend in the energy to remove an O2 molecule. For cations, at a

given size, the energy to remove O2 is far less than to remove O atoms. On the other

hand, it takes approximately the same energy to remove an O2 or an O atom in FeO3-.

These trends are rooted in the fact that the oxides are stabilized by a charge transfer from

Fe to oxygen (see Figure 5-2). The addition of an extra electron facilitates the O-

formation and hence leads to more strongly bound species.

In order to understand the observed products more clearly, the changes in energy

and the various reaction products are depicted in the energy diagram shown in Figure 5-4.

Starting from FeO3+, a CO molecule was placed around the cluster in various

configurations and the system was allowed to relax to determine the ground state

geometry. What was interesting is that the CO molecule approached an O atom leading to

the formation of CO2. As Figure 5-2 shows, the ground state of FeO3CO+ corresponds to

a CO2 molecule bound to FeO2+. To further examine any reaction barrier for the

formation of CO2, the reaction path for the entire process was calculated. Starting from

well separated FeO3+ and CO, the total energy as a function of Fe-C separation was

calculated. For each separation, the position of Fe and that of C were fixed and the

remaining atoms were allowed to relax without any symmetry constraint.

The CO molecule approached an O atom of the FeO3+ cluster, without any barrier,

resulting in the formation of the complex. A profile of the reaction intermediates and

their energies are given in Figure 5-4. The heat of formation is sufficient to drive the

reaction product FeO3CO+ to two possible channels. The first one involves the loss of

CO2 while the other corresponds to the loss of an O2 leaving behind the FeOCO+

complex.

75

Figure 5-3: a) Graph of the energy required to remove an O or O2 from the FeOy+ cluster.

b) Graph of the energy required to remove an O, O2, O- and O2

- from the FeOy- cluster.

Source: Figure courtesy of the group of Professor S. N. Khanna at VCU.

b)

a)

76

ΔE (eV) 2FeO3

+ + 1CO 2FeO3CO+ -4.29 2FeO3CO+ 2FeO2

+ + 1CO2 1.64 2FeO2

+ + 1CO 6FeO2CO+ -3.98 6FeO2CO+ 6FeO+ + 1CO2 1.31 6FeO+ + 1CO 4FeOCO+ -2.88 4FeOCO+ 4Fe+ + 1CO2 1.10 2FeO3CO+ 4FeOCO+ + 3O2 2.66


corresponding change in energy, ΔE, for each reaction step. The superscripts indicate spin multiplicity. Source: Figure courtesy of the group of Professor S. N. Khanna at VCU.

77

Indeed, experimentally we observe both these products as they lie close in energy and are

lower in energy than the initial reactants.

Starting from FeO2+, an additional CO was allowed to approach the cluster from

various directions. As shown in Figure 5-2, the most favorable path corresponds to a CO

approaching one of the O atoms, forming a CO2 molecule, and resulting in the FeOCO2+

complex. A detailed analysis of the reaction path again showed no reaction barriers. The

heat of formation for this process is 3.98 eV. One can again consider two possibilities,

namely (1) the loss of CO2 leading to FeO+ or (2) the loss of O leading to FeOCO+.

However, our calculations show that the energy required to remove an O atom is 4.73 eV.

Hence, the logical choice is the loss of the CO2 molecule which requires only 1.25 eV.

This is further evidence that the FeOCO+ species observed in Figure 5-1 arises from the

earlier pathway and not this latter channel. It is important to note that the ground state of

FeO+ has a spin multiplicity of 6 while FeO2+ has a spin multiplicity of 2. Hence the

formation of the ground state of FeO+ requires a change in total spin multiplicity.

Starting from the ground state of FeO+, a CO was again allowed to approach the

molecule. The ground state indicates the formation of the CO2 molecule. Again, the

detailed reaction path indicates that there is no barrier, leading finally, to the appearance

of Fe+ as found in the experiments.

Figure 5-5 shows the energetic pathways for the corresponding processes for the

anions. As mentioned before, the removal of an O atom from the anions requires much

more energy than from cations. Hence the energy gain in the formation of CO2 is not as

large as for cations. This is clearly seen in Figure 5-5 that shows that the heats of

formations are significantly reduced.

78

ΔE (eV) 4FeO3

- + 1CO 2FeO3CO- -1.96 2FeO3CO- 4FeO2

- + 1CO2 1.44 4FeO2

- + 1CO 4FeO2CO- -2.47 4FeO2CO+ 4FeO- + 1CO2 2.19 4FeO- + 1CO 4FeOCO- -2.35 4FeOCO- 4Fe- + 1CO2 2.22 4FeOCO- + 1CO 2FeOCOCO- -2.28 2FeOCOCO- 4FeCO- + 1CO2 2.58

Figure 5-5: Profile of the reaction energy for FeO3- with CO including the corresponding

change in energy, ΔE, for each reaction step. The superscripts indicate spin multiplicity. Source: Figure courtesy of the group of Professor S. N. Khanna at VCU.

79

Starting from FeO3-, our studies indicate that the approaching CO molecule does not

encounter a barrier. In this case, however, the heat of reaction is only 1.96 eV as opposed

to 4.29 eV for FeO3+. Consequently, the reaction channel involving the loss of an O2

molecule that was feasible for the cation, is energetically prohibited. The only pathway is

the formation of FeO2-. This is in agreement with the experimental findings that do not

detect any FeOCO- product. We also found that while the oxidation of FeO2- and FeO-

does not involve any appreciable barrier, the heats of formation are consistently low.

Combined with the low concentration of FeO2-, this may account for the fact that we do

not see any other products in significant amounts in our experiments and that one

probably needs to go to much higher amounts of CO to see subsequent products.

5.4 Conclusion

It is demonstrated that the charge state of a FeO3 cluster has a strong effect on its

ability to oxidize CO to CO2 using stored oxygen. An analysis of the reaction pathways

provides evidence that the observed dependence on the charge state is driven by the

binding energy of oxygen to the metal atom. As the clusters are stabilized by the charge

transfer from metal to oxygen, more energy is required to remove oxygen from anionic

clusters. For cations, the positive charge on the iron atom weakens the Fe-O bonds that

are formed and enables higher reactivity. We therefore provide an economical alternative

for the direct oxidation of CO using highly oxidized iron clusters thus eliminating the

need for costly precious metals to effect the oxidation of CO.

80

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6. PalDey, S.; Gedevanishvili, S.; Zhang, W.; Rasouli, F. Appl. Catal. B 2005, 56, 241.


8. Schröder, D.; Jackson, P.; Schwarz, H. Eur. J. Inorg. Chem. 2000, 2000, 1171.

9. Wang, L.-S.; Wu, H.; Desai, S. R. Phys. Rev. Lett. 1996, 76, 4853.

10. Schröder, D.; Schwarz, H.; Clemmer, D.; Chen, Y.; Armentrout, P. B.; Baranov, V. I.; Böhme, D. K. Inter. J. Mass Spec: Ion Proc. 1997, 161, 175.

11. Schultz, R. H.; Crellin, K. C.; Armentrout, P. B. J. Am. Chem. Soc. 1991, 113, 8590.

12. Armentrout, P. B.; Koizumi, H.; MacKenna, M. J. Phys. Chem. A 2005, 109,

11365.

13. Kohn, W.; Sham, L. J. Phys. Rev. 1965, 140, A1133.

14. Perdew, J. P.; Burke, K.; Ernzerhof, M. Phys. Rev. Lett. 1996, 77, 3865.

15. Köster, A. M.; Calaminici, P.; Casida, M.; Flores-Moreno, R.; Geudtner, G.; Goursot, A.; Heine, T.; Ipatov, A.; Janetzko, F.; Patchkovskii, S.; Reveles, J. U.; Vela, A.; Salahub, D. R. deMon2k V. 1.08, The international deMon developers community, 2005, http://www.deMon-software.com.

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16. Dunlap, B. I.; Connolly, J. W. D.; Sabin, J. R. J. Chem. Phys. 1979, 71, 4993.

17. Mintmire, J. W.; Dunlap, B. I. Phys. Rev. A 1982, 25, 88.

18. Köster, A. M.; Flores-Moreno, R.; Reveles, J. U. J. Chem. Phys. 2004, 121, 681.

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Chapter 6

Experimental and Theoretical Study of the Structure and Reactivity of Fe1-2O≤6¯

Clusters with CO

6.1 Introduction

Transition metal catalysts, specifically those composed of iron nanoparticles, have

been employed in industrial and biological processes as well as pollution abatement

applications (1). Lin et al. observed nano-sized iron oxides to be highly active for CO

oxidation at low temperatures (2). In addition to other studies probing the oxidation of

carbon monoxide with iron (3), iron oxides showed activity for the oxidation of methane

(4) and various hydrocarbons (5). Iron oxides are also promising because they exhibit

good catalytic lifetimes and resistance to high concentrations of moisture and CO2, which

often poison catalysts (6). Therefore, iron oxides may become economical alternatives to

costly precious metal catalysts.

Studies involving iron and iron oxide have been conducted previously employing

theoretical calculations (7-9) and photoelectron spectroscopy experiments (10-12). FeO,

Fe2O3 and Fe3O4 are stable stoichiometries in the condensed phase, thus they are the most

studied species (10, 13). It was suggested by Li and coworkers that Fe2O3 was the most

active species for CO oxidation, showing that Fe2O3 behaved as a catalyst in the presence

of O2 as well as a reagent in the absence of oxygen, oxidizing CO to CO2 (13).

Sequential reduction of Fe2O3 in the absence of O2 produced Fe3O4, FeO, and Fe species.

(Reproduced in part with permission from J. Phys. Chem. A 2007, Accepted. Copyright 2007 American Chemical Society.)

83

Previous calculations of the reaction of neutral Fe2O3 with CO have proposed a

mechanism with a viable energetic pathway to produce CO2 (14-16). Through

cooperative effects, the adsorption of CO onto the cluster was predicted to weaken a Fe-

O bond making this oxygen available to a second CO for easy formation of CO2.

Iron oxide layers are commonly employed in catalytic supports and have shown

active participation in the reactions (6, 17). Iron is unique and more active than other

supports because it is easily reduced and provides anion vacancy formation due to a

highly disordered structure (18, 19). Anion vacancies are important in the catalytic

process because they stabilize the supported clusters and transfer charge density.

Schubert et al. studied several different transition metal supports and found that iron

oxides were able to adsorb large amounts of oxygen (17). In addition to the adsorbed

oxygen participating in the reaction, they proposed a mechanism that also allowed for the

oxide support to supply the oxygen for the reaction. The ability of iron oxide to function

both ways greatly enhanced the oxidation activity.

This chapter presents a joint experimental and theoretical study that aids in

uncovering species with increased activity and selectivity for the oxidation of CO to CO2.

The results provide important insight into the nature of the active sites responsible for

condensed phase catalysis, shedding light on the role of cluster size, charge and oxidation

state, composition, and stoichiometry toward CO oxidation in the presence of iron oxide.

This chapter focuses on the anionic iron oxide behavior, and Chapter 7 addresses charge

state and electron density effects on the structure and bonding of small cationic iron oxide

clusters.

84

6.2 Methods

Gas-phase cluster studies were performed using a guided ion beam mass

spectrometer coupled to a laser vaporization source, explained previously (20) and in

Chapter 2. Briefly, the second harmonic of a Nd:YAG laser was used to ablate a rotating

and translating iron rod (PVD Materials Corp., 99.95% purity). At a predetermined time,

oxygen seeded in helium (~1%) was pulsed over the rod, forming dense, hot plasma. A

27 millimeter (mm) conical expansion nozzle was placed at the exit of the source

allowing for more third body collisions and aiding in larger cluster formation. The

nozzle was composed of a 15 mm long channel with a 4 mm diameter and 12 mm conical

expansion cone with a 30° total internal angle. The clusters then underwent supersonic

expansion in a field free region before passing through a 3 mm skimmer which created a

molecular beam. The cooled clusters were then focused by a set of electrostatic lenses

and deflectors into the first quadrupole. Each cluster species was individually selected in

the first quadrupole and directed through a second set of lenses into the octopole reaction

cell. Carbon monoxide reactant gas, ranging from 0-20 mTorr, was added to the reaction

cell and the pressure was monitored by a MKS baratron. The products were focused by a

third set of lenses to a second quadrupole where they were mass analyzed and finally

detected using a channel electron multiplier. Studies were also conducted with nitrogen

in the reaction cell under the same conditions of pressure and energy for verification of a

chemical reaction with carbon monoxide. Since both CO and N2 are of the same nominal

mass, experiments with N2 aided in identifying collisional fragmentation. Collision

induced dissociation (CID) experiments were conducted to study the fragmentation

85

patterns of the iron oxide clusters. In these experiments, inert xenon gas was introduced

into the reaction cell under single collision conditions (0.09 mTorr) while the kinetic

energy of the ions in the octopole reaction cell was slowly raised from 0 to 40 eV

laboratory frame energy. Experiments with slightly higher collision pressures, 0.2 mTorr

of Xe, were conducted to find sequential fragmentation patterns under multiple collision

conditions.

The theoretical studies were carried out using two different numerical schemes

that were developed within a density functional formalism by the Khanna research group

(21). The exchange and correlation effects were incorporated through the generalized

gradient approximation (GGA) via a functional proposed by of Perdew, Burke and

Ernzerhof (22). The electronic structure was determined using a linear combination of

atomic orbitals molecular orbital approach. The wave function for the cluster was

constructed by a linear combination of Gaussian type orbitals centered at the atomic

positions in the cluster. The actual calculations employed two different numerical

programs. Most calculations were carried using the Naval Research Laboratory

Molecular Orbital Library (NRLMOL) set of codes developed by Pederson and

coworkers (23-25). For these calculations, a 5s, 4p and 3d basis set for the C and O

atoms and 7s, 5p and 4d basis for the Fe atom was employed (25). In each case, the basis

set was supplemented by a diffuse Gaussian. Supplementary calculations were carried

out using the deMon2K software (26) in order to eliminate any uncertainties associated

with the choice of basis set or the numerical procedure. In these studies, a gradient

corrected density functional (22) and the double zeta valence polarized (DZVP) basis sets

(27) for C and O and the Wachters-F basis set (28) for Fe was employed. The GEN-A2

86

auxiliary function set for C and O and the GEN-A2* auxiliary function set for Fe were

used. For each cluster structure, the configuration space was sampled by starting from

several initial configurations. Then the geometry was optimized by moving atoms in the

direction of forces until they dropped below a threshold value. Since transition metal

atoms are marked by non-zero spin multiplicities, the calculations included optimizing

the spin multiplicities of each cluster.


The anionic iron oxide cluster distribution with both dissociated and molecular

oxygen adsorbed at near thermal energy is shown in Figure 6-1. Expansion gas mixtures

ranging from 1%-20% oxygen were investigated, with the lowest percentage of oxygen in

helium providing the broadest cluster distribution. Various conditions at the exit of the

source were also examined and experiments conducted with no conical expansion nozzle

produced only monomer iron oxide species. A conical nozzle placed at the end of the

source region before the field free expansion region allowed for more collisions and thus

better iron oxide clustering. A 22 mm nozzle was used because the longer nozzle (55

mm) produced species that were more oxygen rich possessing large numbers of

molecularly bound O2 units. We did not generate any stoichiometric FexOy- clusters with

x=y having the same number of iron and oxygen atoms. It is interesting to note that

Bernstein and coworkers produced neutral iron oxide clusters with oxygen deficient

stoichiometries by varying the oxygen concentration in the system (29).

87

Figure 6-1: A typical mass distribution produced for iron oxide anionic clusters. The first iron oxide in each series with subsequent peaks in the series having one additional oxygen atom.

0 100 200 300 400 500 600

Mass (amu)

Ion

Inte

nsit

y (A

rb. U

nits

)

FeO

3-

Fe2O

3-

Fe3O

5-

Fe4O

7-

Fe5O

9-

Fe6O

10-

Fe7O

12-

0 100 200 300 400 500 600

Mass (amu)

Ion

Inte

nsit

y (A

rb. U

nits

)

FeO

3-

Fe2O

3-

Fe3O

5-

Fe4O

7-

Fe5O

9-

Fe6O

10-

Fe7O

12-

88

They made oxygen deficient iron clusters employing a 0.1% oxygen mixture. However,

in the present source this concentration produced a narrower mass distribution of iron

oxide clusters than the 1% oxygen mixture. The following studies include FeO2-4- and

Fe2O3-6- anionic species, mostly with an oxygen-rich metal to oxygen ratio of x/y <1.

Structures. Figure 6-2 shows the calculated optimized lowest energy structures for

the FeO1-4- and Fe2O2-6

- clusters. For iron oxide clusters containing a single Fe atom,

oxygen atoms bind directly to the metal with no molecular oxygen units. The maximum

coordination number of Fe was four in the tetrahedral form of FeO4-. For FeO5

- (not

shown), the fifth oxygen attaches to an oxygen atom forming an O2 unit. In the case of

FeO6- (not shown), an FeO4

- unit with a O2 molecule is bound to an oxygen atom at a

distance of 3.1 Å. A higher energy isomer for FeO4- possessing a molecular oxygen

subunit was 0.78 eV above ground state energy. The FeO bond lengths for FeO1-4- are

relatively similar. The spin multiplicity changes from quartet in FeO1-3- species to

doublet for FeO4-.

In the iron oxide clusters containing two Fe atoms, a basic ring structure

composed of Fe2O2- with oxygen bridging each iron atom is formed. Fe2O3

- and Fe2O4-

have their third and fourth oxygen atoms attached to iron outside the ring. The addition

of oxygen atoms to the cluster shortens the FeO bond length from 1.85 Å to 1.79Å within

the ring and 1.65 Å to 1.62 Å outside. Maximum coordination was obtained for the

Fe2O6- cluster with two extra oxygen atoms bound directly to each Fe atom. A higher

energy isomer of Fe2O6- with an O2 subunit was found 1.23 eV above the ground state.

An anti-ferromagnetic spin coupling was found for Fe2O2-, Fe2O3

-, Fe2O4- and Fe2O5

-

cluster species.

89

Figure 6-2: The ground state geometries of O2, CO, FeO1-3

- and FexOy- clusters. The

bond lengths are given in Angstroms and the superscripts indicate the spin multiplicity. The arrows indicate the spin polarization at the Fe atoms for the Fe2Oy

- clusters. The Mulliken charges are marked below each atom. Source: Figure courtesy of the group of Professor S. N. Khanna at VCU.

90

However, the Fe sites in Fe2O6- were found to be coupled ferromagnetically. Similar

progressions of the exchange coupling with oxidation have been reported for the case of

Cr2On clusters (30).

6.3.1 CID Fragmentation

Collision induced dissociation studies were undertaken with inert xenon gas under

single (0.09 mTorr) and multiple (0.2 mTorr) collision conditions and the results are

listed in Table 6-1. The fragmentation of selected clusters and their corresponding neutral

loss aids in elucidating the structures of charged clusters and obtaining the general order

of bond strength. Trends have been established between the calculated dissociation

energies (DE) and the order of experimental fragmentation products, which are discussed

below.

For FeOy- clusters containing a single Fe atom, the observed fragmentation

channels with Xe correspond to the loss of an O atom or O2 molecule. For FeO2- and

FeO3-, the only dissociation observed was loss of an atomic O. For FeO4

-, the loss of O2

is the lowest energy pathway. At higher kinetic energies, FeO4- lost an O atom. More

interesting are the fragmentation channels for Fe2Oy- clusters containing two Fe atoms. In

the case of Fe2O3-, the product channel that is observed at the lowest collision energy is

loss of an O atom. In contrast, the product channel that is observed at the lowest collision

energy for Fe2O4- is loss of FeO. For Fe2O5

-, the lowest energy channel involves the loss

of a FeO2 and in Fe2O6-, the loss of an O2 molecule is the lowest channel. In each case,

91

other dissociation channels are observed as the kinetic energy is increased. The basic

question is then whether these pathways mirror the energetics of fragmentation.

To answer this question, the energies required to remove an O, O2, O-, and O2

-

from FeOy- and Fe2Oy

- clusters and the energies for fragmentation pathways leading to

the production of Fe, FeO, FeO2, and FeO3 in the case of Fe2Oy- clusters were calculated.

The results of these investigations are plotted in Figure 6-3. The calculated dissociation

energies for various reaction channels corresponding to the observed experimental

products are recorded in Table 6-1. The following is a discussion of the trends in the

theoretical results reported in Figure 6-3.

The current experiments detect anionic species, therefore the loss of neutral O and

O2 from the mass selected cluster can be experimentally observed. For FeOy- clusters,

one notices that the energy to remove an O atom or an O2 molecule decreases with

increasing oxygen coverage. As an example, while it takes 5.99 eV to remove an O atom

from FeO2- only 3.86 eV is needed to remove an O atom from FeO4

-. This decrease

makes the removal of O2 a favorable channel at higher coverage particularly since the

removed O atom can gain energy by combining to form O2 molecules. Indeed, for FeO4-,

the energy to remove an O2 molecule is noticeably less than to remove an O atom. While

for FeO5-, it will take only 0.95 eV to remove an O2 molecule. This is generally

consistent with the observed fragmentation patterns in Table 6-1 which shows the loss of

O as the preferred pathway for FeO2- and FeO3

- while the lowest energy pathway for

FeO4- fragmentation is the loss of an O2 molecule.

For Fe2Oy- clusters, the fragmentation pathways include the production of Fe,

FeO, FeO2 and possibly FeO3 fragments.

92

Table 6-1: The collision induced dissociation fragmentation channels, calculated dissociation energies (DE), and reaction products with CO and N2 for Fe1,2O2-6

- clusters.

FexOy-

(x,y) Products with Xea

Neutral(s) Lostb

DE (eV)

Products with CO

Products with N2

1,2 1,1 0,1 6.03 1,1 No products

1,3 1,2 0,1 5.8 1,2 No products

1,4 1,2 1,3

0,2 0,1

3.47 3.87

1,2

1,3c 1,2c

2,3 2,2 1,2 1,3

0,1 1,1 1,0

6.10 3.70 3.41

2,2 No products

2,4 1,3 1,2 2,3

1,1 1,2 0,1

3.94 4.68 6.03

2,3

1,3 No

products

2,5 1,3 2,4 2,3

1,2 0,1 0,2c

3.44 4.56 4.39

2,4

1,3 No

products

2,6 2,4 2,5 1,3

0,2d

0,1 1,3c

2.90 4.55 3.22

2,4

2,5c 2,4 ≈ 2,5c

aFragmentation channels are shown in order of observation with increasing collision energy. bNeutral loss is assigned based on the difference between the selected cluster and fragment ion formed. cDescribes channels taking place under multiple collision conditions. dDissociation occurs at near thermal energies.

93


- and b) Fe2Oy- clusters. c) Graph of the fragmentation energy for neutral metal

and metal oxide loss from Fe2Oy- clusters.

0

1

2

3

4

5

6

7

1 2 3 4 5 6

# of Oxygen atoms

DE

(eV

)

2

3

4

5

6

7

8

1 2 3 4 5 6# of Oxygen atoms

DE

(eV

)

2

3

4

5

6

7

8

9


DE

(eV

)

Fe2Oy- Fe2Oy-1

- + OFe2Oy

- Fe2Oy-2- + O2

Fe2Oy- Fe2Oy-1 + O-

Fe2Oy- Fe2Oy-2 + O2

-

Fe2Oy- Fe2Oy

- + FeFe2Oy

- Fe2Oy-1- + FeO

Fe2Oy- Fe2Oy-2

- + FeO2Fe2Oy

- Fe2Oy-3- + FeO3

FeOy- FeOy-1

- + OFeOy- FeOy-2

- + O2FeOy- FeOy-1 + O-

FeOy- FeOy-2 + O2-

a)

b)

c)

0

1

2

3

4

5

6

7

1 2 3 4 5 6

# of Oxygen atoms

DE

(eV

)

2

3

4

5

6

7

8


DE

(eV

)

2

3

4

5

6

7

8

9


DE

(eV

)

Fe2Oy- Fe2Oy-1

- + OFe2Oy

- Fe2Oy-2- + O2

Fe2Oy- Fe2Oy-1 + O-

Fe2Oy- Fe2Oy-2 + O2

-

Fe2Oy- Fe2Oy

- + FeFe2Oy

- Fe2Oy-1- + FeO

Fe2Oy- Fe2Oy-2

- + FeO2Fe2Oy

- Fe2Oy-3- + FeO3

FeOy- FeOy-1

- + OFeOy- FeOy-2


FeOy- FeOy-2 + O2-

FeOy- FeOy-1

- + OFeOy- FeOy-2


FeOy- FeOy-2 + O2-

a)

b)

c)

94

These dissociation energies are shown graphically in Figure 6-3. The dissociation energy

values associated with experimental products are recorded in Table 6-1. It is interesting

to note that the loss of O2 reaches an energetic minimum for Fe2O6- as the energy needed

to break two Fe-O bonds is likely to be overcome by the exothermic formation of O2

releasing 6.21 eV of energy. Oxygen recombination was also observed in CID studies

previously conducted in our laboratory with V2O5+ (31). The structure of V2O5

+

contained only atomic oxygen bonds to the metal (32); however, dissociation of O2 was

observed.

Fragmentation of Fe-Fe bonds occurred in the dimer clusters, mostly at higher

energies or under multiple collision conditions. The core structure of the dimer clusters

after fragmentation was either FeO3- or FeO2

-, which represent the stable building blocks

of the larger iron oxide clusters. Fragmentation of neutral FeO from the Fe2O4- cluster

requires 3.94 eV of overall energy as shown in Figure 6-4. The dissociation of Fe2O4-

begins with the breaking of the bond between an Fe atom and a bridging O atom, which

changes its coordination to become a terminal O atom, requiring 0.81 eV. Breaking the

Fe-Fe bond and internal rotation of an FeO2 subunit then leads to a chain structure and

requires an additional 0.29 eV. From this open chain structure it takes an additional 2.84

eV to lose neutral FeO and 3.59 eV to lose FeO2. The energy for neutral FeO2

fragmentation from Fe2O5- reaches a low of 3.44 eV. Figure 6-3 shows that more energy

is needed to fragment neutral FeO2 from Fe2O4- than neutral FeO; however, this order is

reversed in the case of Fe2O5-. Our experiments verify these theoretical results as

dissociation of FeO from Fe2O4- is observed at lower energy while fragmentation of FeO2

required additional energy.

95

Figure 6-4: The change in energy (ΔE) for the dissociation pathway of parent cluster, Fe2O4


-. The superscripts indicate the spin multiplicity. Source: Figure courtesy of the group of Professor S. N. Khanna at VCU.

96

These two clusters are unique because they form the stable fragment FeO3- prior to

oxygen loss. This indicates that the Fe-O bonds for Fe2O4- and Fe2O5

- are very strong,

and indeed they require 6.03 eV and 4.56 eV to break the Fe-O bond, respectively.

In the case of Fe2O3-, loss of an oxygen atom is the first fragment observed. This

fragment does not follow the pattern of dissociation energies found in Table 6-1. Figure

6-5 shows the change in energy for the dissociation pathways for neutral Fe and FeO loss.

In both pathways multiple bond breaks are required for neutral Fe and FeO loss. In the

first step a bridging Fe-O bond is broken and the oxygen either bonds to the Fe atom that

is less coordinated requiring 0.59 eV or stays on the iron atom with an oxygen atom

already attached requiring 0.90 eV of energy. At this stage in the reaction path, spin is

conserved in the doublet state for the later neutral FeO dissociation pathway. From the

[2Fe2O3-] transition state, the bonds are relaxed, and followed by Fe-Fe dissociation to

form a linear chain of alternating Fe-O bonds. Finally, neutral FeO is released with an

additional 3.08 eV of energy required. The other pathway for neutral Fe dissociation

also requires a spin change and bond rearrangements. Subsequent neutral Fe loss from

the cluster requires an additional 2.14 eV of energy. The bond breaking and

rearrangement of the atoms along the pathway causes these neutral fragments to occur

after O atom loss.

While a comparison of the two pathways may indicate that the loss of Fe is

energetically more favorable, the pathway for the loss of FeO requires more stable

intermediate steps. We believe that this accounts for the FeO fragment observed before

Fe loss.

97

Figure 6-5: The change in energy (ΔE) for the dissociation pathway of parent cluster Fe2O3


-. The superscripts indicate the spin multiplicity. Source: Figure courtesy of the group of Professor S. N. Khanna at VCU.

98

Further, the numerous bond breaking steps and rearrangement of the atoms along the

pathways may make these FeO and Fe loss mechanisms kinetically unfavorable

compared to the energetically unfavorable O atom loss channel. In both cases (loss of

FeO and loss of Fe), the need to undergo a spin transition from a doublet state to an octet

intermediate state is expected to slow the reaction considerably (33).

It is, therefore, possible that spin changes slow down the rate of dissociation and

FeO fragmentation pathways occur after oxygen atom dissociation from the cluster.

Additionally the calculated dissociation pathways for Fe2O5- and Fe2O6

- can be explained

with the same reasoning. For Fe2O6-, the fragmentation order observed with CID differs

from the calculated dissociation energies. For the same reasons discussed above for

Fe2O3-, FeO3 loss occurs after O atom loss. Multiple bonds fragmenting in FeO3 requires

more time than one Fe-O bond break and therefore is observed after O atom loss.

6.3.2 Reactions with CO and N2

The reaction products for each cluster species are recorded in descending order of

relative product intensity in Table 6-1. Theoretical studies indicate that the atomization

energy of CO and CO2 are 11.63 and 17.97 eV respectively. The formation of CO2 is

therefore energetically feasible in cases where it takes less than 6.34 eV to remove an O

atom from the cluster. However, the presence of reaction barriers can prevent the

formation of CO2. This is the reason that CO2 formation does not occur in the presence

of O2 even though the binding energy of O2 is only 6.20 eV. The energy calculated to

remove an O from the cluster species, FeO2-, FeO3

-, FeO4-, Fe2O3

-, Fe2O4-, and Fe2O5

- is

99

less than 6.34 eV and they are all active toward oxygen atom transfer to CO. This

process occurs as a dominant reaction channel for most anionic species. Since neutral

species are not detected, the presence of FexOy-1- products suggests that oxygen is

transferred from iron oxide clusters to CO to produce neutral CO2.

Figure 6-6 shows the branching ratios for FeO2- and Fe2O3

- in the presence of

increasing CO pressure. These cluster stoichiometries are the most efficient iron oxide

anions to affect the CO oxidation reaction channel. Notice that both clusters are

composed of one more oxygen atom than the number of iron atoms. The energy

calculated to remove an oxygen atom from these two clusters is approximately 6 eV.

Therefore the total energy needed is more than can be supplied through thermal collisions

alone. This is supported by the fact that no oxygen atom loss products are observed in

nitrogen studies conducted under the same experimental conditions (see Table 6-1).

Confirmation of the chemical reactivity of iron oxide anionic clusters with CO, therefore

can be seen by comparing the products with nitrogen.

The reaction paths for FeO2-, FeO3

-, FeO4-, Fe2O3

- and Fe2O4- with CO were

calculated in order to understand the different reactivity they displayed. For the most

active clusters, FeO2- and Fe2O3

-, the reaction was found to proceed without barriers and

follow a spin allowed path. The reaction path for Fe2O3- is shown in Figure 6-7. The first

CO can attach to the Fe site or approach the O atom to form CO2. Our studies indicate

that the more stable configuration corresponds to CO attached to the Fe site. This is

consistent with the electron donating behavior of CO and the partial positive charge

present on the Fe site as revealed by the Mulliken population calculation.

100

Figure 6-6: The branching ratios for a) FeO2

- and b) Fe2O3- with increasing CO reactant

gas pressure. Notice that both clusters are selective for the oxygen atom transfer reaction pathway.

0.8

0.9

1

0 5 10 15 20CO (mTorr)

0

0.1

0.2

0.8

0.9

1

0 5 10 15 20

CO (mTorr)

0

0.1

0.2

FeO2-

FeO-

Rel

ativ

e In

tens

ity

(par

ent)

Relative Intensity (product)

a)

Fe2O3-

Fe2O2-

b)

Rel

ativ

e In

tens

ity

(par

ent)

Relative Intensity (product)0.8

0.9

1

0 5 10 15 20CO (mTorr)

0

0.1

0.2

0.8

0.9

1

0 5 10 15 20

CO (mTorr)

0

0.1

0.2

FeO2-

FeO-

Rel

ativ

e In

tens

ity

(par

ent)


a)

Fe2O3-

Fe2O2-

b)

Rel

ativ

e In

tens

ity

(par

ent)


0.8

0.9

1

0 5 10 15 20

CO (mTorr)

0

0.1

0.2

FeO2-

FeO-

Rel

ativ

e In

tens

ity

(par

ent)


a)

Fe2O3-

Fe2O2-

b)

Rel

ativ

e In

tens

ity

(par

ent)


101

Figure 6-7: The change in energy (ΔE) for the reaction pathway of Fe2O3

- with CO. The superscripts indicate the spin multiplicity. Source: Figure courtesy of the group of Professor S. N. Khanna at VCU.

102

The subsequent CO then attaches to the O site forming CO2 as shown in Figure 6-7.

Previous studies on neutral Fe2O3 also obtained a similar reaction path (8).

On the other hand, for FeO3- and FeO4

-, the reactions with CO involve a change of

spin which may slow the reactions. For the Fe2O4- the reaction proceeds without barriers

and follows a spin allowed path. However, in the initial step of the reaction there is a

gain in energy of only 1.54 eV for the absorption of the CO molecule compared with a

gain of 4.20 eV for the absorption of two CO molecules in the Fe2O3- cluster. This is

because one of the Fe sites in Fe2O3- is coordinated to only two O atoms and the first CO

binds to this metal site. The difference in absorption energy explains the difference in

reactivity observed in the two clusters.

It is important to note that the reaction pathways involve intermediate species and

consequently involve structural rearrangements. For small clusters, the reduced size is

amenable to these rearrangements and consequently, the reaction barriers are far less than

for the case of bulk surfaces where the extended geometry is less flexible towards

structural changes.

Another reaction channel observed was the loss of molecular oxygen from FeO4-

and Fe2O6-. Branching ratios of these clusters are shown in Figure 6-8 with O2 loss as the

dominant reaction channel followed by minor O atom loss. In the case of FeO4-, it is

assumed that at the high CO pressures in which FeO2- is detected as a product, the loss of

molecular oxygen can be attributed to collision energy that promotes the cluster to the

structure containing an intact molecular oxygen unit. The same will apply for the

formation of the higher energy isomer of Fe2O6- with a molecular oxygen unit. Both

Fe2O6- and FeO4

- showed O2 loss in the presence of nitrogen as seen in Table 6-1.

103

Figure 6-8: The branching ratios for a) FeO4

- and b) Fe2O6- with increasing CO pressure,

both clusters show molecular oxygen loss.

0.8

0.9

1

0 5 10 15 20CO (mTorr)

0

0.1

0.2

0.8

0.9

1

0 7 14 21CO (mTorr)

0

0.1

0.2

Fe2O6-

Fe2O4-

Fe2O5-

FeO4-

FeO2-

FeO3-

a)

b)

Rel

ativ

e In

tens

ity

(par

ent) R

elative Intensity (products)

Rel

ativ

e In

tens

ity

(par

ent)

Relative Intensity (products)

0.8

0.9

1

0 5 10 15 20CO (mTorr)

0

0.1

0.2

0.8

0.9

1

0 7 14 21CO (mTorr)

0

0.1

0.2

Fe2O6-

Fe2O4-

Fe2O5-

FeO4-

FeO2-

FeO3-

a)

b)

Rel

ativ

e In

tens

ity

(par

ent) R

elative Intensity (products)

Rel

ativ

e In

tens

ity

(par

ent)

Relative Intensity (products)

104

These products were observed at higher nitrogen pressure under conditions where

multiple collisions are expected.

6.4 Conclusion

It is shown that the nature of fragments produced in collision induced dissociation

is governed not only by the overall energetics but also by the nature and multiplicity of

the intermediate motifs marking the fragmentation. For FeOy- clusters, the present studies

show that while the loss of atomic O is the favored channel at low energies, the loss of O2

is the dominant pathway at higher coverage. In the case of Fe2Oy- clusters, the lowest

fragmentation pathways involve the loss of Fe or FeO units except for Fe2O6-, where the

lowest energy pathway is the loss of an O2 molecule. Small anionic iron oxide clusters

are shown to enable the oxidation of CO at near thermal energies. The most active and

selective iron oxides were composed of one more oxygen atom then iron atom. The

increased reactivity is attributed to the following: (1) the energy required to remove an O

atom in these clusters is less than the gain in energy to form CO2 making the oxidation

thermodynamically feasible and (2) the reduced size (compared to bulk species) allows

structural rearrangements that eliminate the high reaction barriers and make the oxidation

kinetically possible.

105

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2. Lin, H-Y.; Chen, Y-W.; Wang, W-J. J. Nanopart. Res. 2005, 7, 249.


4. Tanaka, S.; Nakagawa, K.; Kanezaki, E.; Katoh, M.; Murai, K.-I.; Moriga, T.;

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106

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107

Chapter 7

Experimental and Theoretical Study of the Structure and Reactivity of Fe1-2O1-5+

Clusters with CO

7.1 Introduction

It is particularly important to understand the influence of ionic charge state on the

oxidation of CO with iron oxides. There is much debate over the role that charge transfer

plays on the active sites in supported condensed phase catalysts and a better

understanding of the mechanism governing the reaction may eventually lead to the design

of more economical and efficient catalysts. Previous studies of iron oxide layers

supported on gold have uncovered both Fe2O3 and FeO species present, with Fe3+ and

Fe2+ oxidation states, respectively, in the oxidation of carbon monoxide (1). PalDey and

coworkers employed XRD to show that mixed catalysts containing Fe2O3 exhibited

defect sites in the form of cationic vacancies (2). These vacancies, which constitute

regions of charge deficiency, may be important as active sites for CO oxidation.

Schröder et al. have employed three gas-phase techniques, ion cyclotron

resonance (ICR), guided ion beam (GIB), and selected ion flow tube (SIFT) mass

spectroscopy, to probe the reaction of FeO+ toward the oxidation of hydrocarbons (3).

Their results provide evidence for the importance of gas-phase studies which show

similar reactivity between all three methods and provide further understanding of the

molecular level mechanisms involved in condensed phase oxidation reactions. Studies by

108

Armentrout and coworkers have investigated the thermochemistry and bond energies of

gas-phase iron cations with CO and CO2 (4, 5), while Schwarz et al. have classified

dissociation patterns of small iron oxide cations (6). Ionic gas-phase studies are becoming

an acceptable means for elucidating the fundamental mechanisms and specific reaction

steps of catalytic reactions.

Previous theoretical calculations have predicted the thermodynamics and reaction

steps that occur over iron oxide catalysts. Findings by Yumura et al. for the oxidation of

formaldehyde with FeO+ showed that complete oxidation to CO2 was exothermic and

involved a carbon monoxide intermediate complex, H2O-Fe+-CO (7). Other theoretical

studies of the structure and reactivity of small iron oxide clusters have predicted weakly

bound oxygen units to be active in effecting the oxidation of CO to CO2 (8, 9). A

mechanism for the simultaneous oxidation of CO and reduction of NO utilizing iron

oxide species with regeneration of active sites was proposed (9, 10). This shows the

potential importance of iron oxide species in the abatement of harmful pollutants.

This chapter presents a joint experimental and theoretical study of the role of

charge state on the structure and reactivity of monomer and dimer iron oxide cluster

cations. Most previous studies of iron oxides have focused specifically on neutral FeO

and Fe2O3 as these are the most common bulk phase stoichiometries. Herein, studies of

cationic clusters of various size and stoichiometry were explored in order to uncover the

reaction pathways and associated thermodynamics.

109

7.2 Methods

Gas-phase iron oxide cluster studies were carried out employing a guided ion

beam mass spectrometer coupled to a laser vaporization source (LAVA), presented in

detail in Chapter 2. Briefly, the second harmonic of a Nd:YAG laser was used to ablate a

rotating and translating iron rod (PVD Materials Corp., 99.95% purity) ensuring a fresh

surface was ablated with each laser pulse. At a predetermined time, oxygen seeded in

helium (~1%) was pulsed over the rod, forming dense, hot plasma. A conical expansion

nozzle was placed at the exit of the source to allow for more collisions and aid in cluster

formation. Two different nozzle were used, one having a total length of 27 millimeters

with a 4 mm diameter, 15 mm long channel with a 30° total internal angle, 12 mm long

expansion cone. The other nozzle was 51 millimeters in total length with a 2 mm inner

diameter and similar cone dimensions as the other nozzle. The clusters were then allowed

to pass into a field free region where they underwent supersonic expansion before being

collimated through a 3 mm skimmer which created a molecular beam of iron oxide

clusters. The cooled clusters were focused through a set of electrostatic lenses and

deflectors into the first quadrupole. Each cluster species was individually selected in the

first quadrupole and directed through a second set of electrostatic lenses into the octopole

reaction cell. Carbon monoxide reactant gas was added to the reaction cell and the

pressure, ranging from 0-12 mTorr, was monitored by a MKS baratron. The products

were directed through a third set of lenses to a second quadrupole where they were mass

analyzed and finally detected using a channel electron multiplier. Studies were also

conducted employing inert nitrogen gas in the octopole reaction cell under the same

110

conditions of pressure and energy as carbon monoxide. These studies verify the chemical

reaction products with carbon monoxide, since both CO and N2 are of the same nominal

mass. Products with nitrogen also serve to elucidate whether oxygen is bound atomically

or molecularly to the iron atoms, which can then be compared to the calculated structures

and oxygen bond strengths. Also, collision induced dissociation (CID) experiments were

conducted to study the fragmentation patterns of the iron oxide clusters. In these

experiments, xenon gas was introduced into the octopole reaction cell at single collision

conditions of 0.09 mTorr while the kinetic energy in the collision cell was slowly

increased from 0 to 40 eV lab-frame energy. Slightly higher collision pressures of 0.2

mTorr Xe were also conducted to find sequential fragmentation patterns at multiple

collision conditions.

In order to help interpret the findings, theoretical studies were carried out by the

Khanna research group using two different numerical schemes that were developed

within a density functional formalism (11). The exchange and correlation effects were

incorporated through the generalized gradient approximation (GGA) via a functional

proposed by Perdew, Burke and Ernzerhof (12). The electronic structure was determined

using a linear combination of atomic orbitals molecular orbital approach. The wave

function for the cluster was constructed by a linear combination of Gaussian type orbitals

centered at the atomic positions in the cluster. The actual calculations employed two

different numerical programs. Most calculations were carried using the Naval Research

Laboratory Molecular Orbital Library (NRLMOL) set of codes developed by Pederson

and co-workers (13-15). For these calculations, a 5s, 4p and 3d basis set for the C and O

atoms and 7s, 5p and 4d basis for the Fe atom was employed (15). In each case, the basis

111

set was supplemented by a diffuse Gaussian. For more details, the reader is referred to

the original papers (13-15). Supplementary calculations were also carried out using the

deMon2K software (16) in order to eliminate any uncertainties associated with the choice

of basis set or the numerical procedure. In these studies, a gradient corrected density

functional (12) and the double zeta valence polarized (DZVP) basis sets (17) for C and O

and the Wachters-F basis set (18) was employed for Fe. The GEN-A2 auxiliary function

set for C and O and the GEN-A2* auxiliary function set for Fe were used. For each

cluster structure, the configuration space was sampled by starting from several initial

configurations. Then the geometry was optimized by moving atoms in the direction of

forces until they dropped below a threshold value. Since transition metal atoms are

marked by non-zero spin multiplicities, the calculations included optimizing the spin

multiplicities of each cluster.


Various cluster formation conditions were examined at the exit of the source

region and experiments conducted with no expansion nozzle produced monomer iron

species dominated by Fe+. Figure 7-1(a) shows a typical iron oxide cation cluster

distribution obtained with the 27 mm conical expansion nozzle in which monomer iron

oxide clusters were predominant. Figure 7-1(b) shows the distribution obtained with a

longer, 51 mm conical nozzle. The longer nozzle allowed for more third-body collisions

and better clustering which produced a higher intensity of dimer iron oxide clusters. It is

interesting how the source conditions greatly affected the mass distribution.

112

Figure 7-1: Typical mass distribution for iron oxide cation clusters produced when employing a) a 27 mm conical nozzle and b) a 51 mm conical nozzle at the exit of the source.

0 50 100 150 200

0 50 100 150 200

Fe+

FeO

+

FeO

2+

FeO

3+

FeO

4+

FeO

3+

Fe 2

O+

FeO

5+

FeO

6+

FeO

8+

Fe 2

O2+

Fe 2

O3+

Fe 2

O4+

FeO

8+

Mass (amu)

Inte

nsit

y (a

rb. u

nits

)

Mass (amu)

Inte

nsit

y (a

rb. u

nits

)

(a)

(b)Fe

O4+

0 50 100 150 200

0 50 100 150 200

Fe+

FeO

+

FeO

2+

FeO

3+

FeO

4+

FeO

3+

Fe 2

O+

FeO

5+

FeO

6+

FeO

8+

Fe 2

O2+

Fe 2

O3+

Fe 2

O4+

FeO

8+

Mass (amu)

Inte

nsit

y (a

rb. u

nits

)

Mass (amu)

Inte

nsit

y (a

rb. u

nits

)

(a)

(b)Fe

O4+

113

Therefore both nozzles were employed for these experiments depending on the cation

cluster selected for study. Cationic FexOy+ clusters with a stoichiometry of x=y along

with species that were oxygen rich with a ratio of x/y <1 were produced. The following

studies were conducted on FeO1-10+ and Fe2O1-4,8,10

+ cationic species.

To understand the specific reaction pathways producing each of the products

observed, CID experiments and DFT calculations were undertaken to aid in structure

identification of cationic iron oxide clusters. A discussion of the observed reaction

products with CO follows.

Structures. Figure 7-2 shows the ground state structures of FexOy+

clusters and

geometries of Fe(CO)1-2+ and FexOyCO+. Similar to the anions in Chapter 6, cationic iron

oxide clusters containing a single Fe atom have a maximum coordination number of four

oxygen atoms bound directly to the metal in the tetrahedral form of FeO4+. FeO+ has a

bond length of 1.64 Å. For FeO2+ and FeO3

+ the bond length is reduced to 1.56 Å and

FeO4+ presents three bond lengths of 1.58 Å and one bond length of 1.64 Å. For FeO5

+,

the fifth oxygen attaches to an oxygen atom forming an O2 unit with an elongated Fe-O2

bond at 2.16 Å. The spin multiplicity changes from sextet in FeO+ to doublet in FeO2-5+.

The iron dimer has a bond length of 2.18 Å and a spin multiplicity of octet. In

Fe2Oy+ clusters, a basic ring structure composed of Fe2O2

+, bridging each Fe atom, is

formed. For Fe2Oy+ with y>3, the oxygen atoms attach to iron outside the ring. Maximum

coordination was obtained for the Fe2O6+ cluster with two extra oxygen atoms bound

directly to each Fe atom. The Fe-Fe bond length increases to around 2.50 Å with the

exception of the Fe2O2+ cluster that contains a Fe-Fe bond length of 2.31 Å. The Fe-O

bond length ranges from 1.76 to 1.81 Å within the ring and 1.50 to 1.57 Å outside. The

114

spin multiplicity is octet of Fe2O+ and changes to doublet and quartet states. An anti-

ferromagnetic spin coupling was found for Fe2O2+, Fe2O3

+ and Fe2O4+ cluster species and

the Fe sites in Fe2O5+ and Fe2O6

+ were found to be coupled ferromagnetically.

The structures for carbon monoxide bound to the iron oxide clusters were also

calculated. CO binds more strongly to an iron atom with a bond length of 1.91 Å in

Fe(CO)2+ whereas if CO binds to an oxygen atom the bond length is stretched to 2.04 Å

in FeO2CO+. These bond strengths are shown to be important in the observed reactions

which reflect the stability of the intermediates.

7.3.1 CID Studies

A systematic energy resolved CID study employing inert xenon gas was

undertaken. Table 7-1 lists the fragmentation products in order of increasing kinetic

energy for CID studies undertaken with xenon gas under single collision conditions. The

fragmentation patterns of these clusters help to identify common structural features. The

major product identified was loss of atomic oxygen. Most clusters lost an oxygen atom at

higher collision energies, however Fe2O3+ lost an O atom at near thermal energy. This is

evidence that Fe2O3+ contains a weakly bound O atom in the ground state structure.

Further implications of this weakly bound oxygen atom will be discussed in terms of the

reactivity below.

Loss of molecular oxygen from oxygen rich clusters having a stoichiometry of

FexOy+ (y≥x+2) was another major fragmentation product. If enough kinetic energy is

input into the cluster through raising the DC voltage of the octopole, then vibrational

115

bond rearrangement and exothermic fragmentation occurs, given the large gain in energy

for forming O2. This is likely the case in iron clusters not saturated with oxygen atoms.

Oxygen recombination was observed for vanadium oxide clusters containing only metal-

oxygen atom bonds in previous CID experiments (19). Extremely oxygen-rich clusters,

for example FeO10+ and FeO9

+, lost successive O2 molecules at near thermal energy

uncovering a FeO+ core structure. During mass selection of these clusters it was observed

that an oxygen molecule would dissociate without the addition of a collision gas, which

suggests that these clusters were meta-stable with loosely associated O2 units. Indeed,

Figure 7-2 shows a maximum of 4 oxygen atoms can bind directly to the iron metal

center.

For the dimer cluster series, specifically Fe2O2+ and Fe2O4

+, fragmentation

involved a Fe-Fe bond dissociation to form FeO2+ and FeO4

+, respectively. For Fe2O2-4+

subsequent dissociation with increasing kinetic energy revealed a basic core of FeO2+.

For oxygen-rich clusters, Fe2O8+ and Fe2O10

+, a stable core structure of Fe2O2+ was

revealed after collisional loss of oxygen molecules. There was no Fe-Fe fragmentation in

these clusters and, therefore, it is assumed that the larger oxygen rich clusters dissociate

O2 units and, subsequently have similar reactivity as the underlying core cluster. The

building blocks of small iron oxide cations were identified as FeO+, FeO2+, and Fe2O2

+.

General trends can be established between the calculated DE and the order of

experimental CID products. The graphs in Figure 7-3 show that as the number of oxygen

atoms surrounding iron increases, the dissociation energies for neutral loss of O and O2

decreases. It is interesting to note Schwarz and coworkers showed that neutral O, FeO,

and FeO2 loss was dependent on the iron to oxygen ratio of the parent cluster (6).

116

Figure 7-2: The ground state geometries for a) FexOy+ clusters and b) for Fe(CO)1-2

+ and FexOyCO+ clusters. The bond lengths are given in Angstroms and the superscripts indicate the spin multiplicity. The arrows indicate spin polarization at the Fe atoms for Fe2Oy

+ and Fe2OyCO+. The Mulliken charges are marked below each atom.

(a)

(b)

117

Table 7-1: Results of CID studies showing the fragmentation channels for selected Fe1-2Oy+

clusters

FexOy+


Neutral(s) Lostb

FexOy+


Neutral(s) Lostb

1,1 1,0 0,1 2,1 2,0 0,1 1,2 1,0

1,1 0,2 0,1

2,2 2,1 2,0 1,2

0,1

0,2 1,0c

1,3 1,2 1,1 1,0

0,1 0,2 0,3

2,3 2,2 2,1 2,0 1,2

0,1d

0,2

0,3c 1,1c

1,4 1,2 1,3 1,1 1,0

0,2d 0,1 0,3 0,4

2,4 2,2 2,3 1,4 1,2

0,2 0,1

1,0 1,2

1,5 1,3 1,4 1,2 1,1

0,2d

0,1 0,3 0,4

2,8 2,6 2,4 2,2 2,3

0,2d 0,4 0,6 0,5

1,6 1,4 1,2 1,5 1,1

0,2d 0,4 0,1 0,5

2,10 2,8 2,6 2,4 2,2

0,2d 0,4d 0,6 0,8

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

0,2d 0,4 0,5 0,6

1,8 1,6 1,4 1,2 1,7

0,2d

0,4d 0,6 0,1

1,9 1,7 1,5 1,4 1,3 1,1

0,2d 0,4d 0,5d 0,6 0,8

1,10 1,8 1,6 1,4 1,2 1,9 1,1

0,2d

0,4d

0,6d 0,8 0,1 0,9

aFragmentation channels are shown in order of observation with increasing collision energy. bThe neutral loss is assigned based on the difference between the selected cluster and fragment ion formed. cDescribes channels taking place under multiple collision conditions. dRepresents that dissociation occurs at near thermal energies.

118

In oxygen-rich clusters with a metal to oxygen ratio of 1>x/y>2/3, neutral loss of FeO2

and O2 became more dominate. For Fe2O2+ and Fe2O3

+, it can be seen in Figure 7-3, that

the energy needed to fragment an oxygen atom is less than 1 eV greater than the energy

needed to fragment O2. Consequently, under the low kinetic energies in which these

products were experimentally observed, the overall energy needed to fragment two O

atom bonds is not overcome. These trends are used to describe the CO oxidation and

oxygen replacement by CO reaction pathways for FexOy+ clusters.

7.3.2 Reactivity Studies

The products of reactions of FeO1-5+ and Fe2O1-5

+ cluster species in the presence

of either carbon monoxide or nitrogen are shown in Table 7-2. The dissociation energies

associated with the CO product channels are also listed. For iron clusters with up to three

oxygen atoms, oxygen atom transfer was an observed product channel. Iron species

FeO+ and Fe2O+, with one O atom, were selective for CO oxidation. This was confirmed

in the nitrogen studies which showed no loss of an oxygen atom under the same

experimental pressures as CO reactant gas. Figure 7-4 displays the branching ratios for

FeO1,2+ and Fe2O1,2

+ clusters which all show oxygen atom transfer as the dominant

reaction channel in the presence of CO. Therefore, only clusters with a FexO+ and FexOy

+

(y=x+1) cluster stoichiometry showed atomic oxygen atom transfer to CO producing

neutral CO2 as the major reaction pathway. Considering the energy required to remove

an oxygen atom from iron, and the gain in energy from forming CO2, shows the reaction

to be overall exothermic for these clusters.

119


+ and b) Fe2Oy+ clusters. c) Graph of the fragmentation energy for neutral metal

and metal oxide loss from Fe2Oy+ clusters.

23456789

1 2 3 4 5 6

# of Oxygen atoms

DE

(eV

)

0

1

2

3

4

5

1 2 3 4 5 6

# of Oxygen atoms

DE

(eV

)

0

1

2

3

4

5

6

1 2 3 4 5 6

# of Oxygen atoms

DE

(eV

)

FeOy+ FeOy-1

+ + OFeOy

+ FeOy-2+ + O2

Fe2Oy+ Fe2Oy-1

+ + OFe2Oy

+ Fe2Oy-2+ + O

Fe2Oy+ Fe2Oy

+ + FeFe2Oy

+ Fe2Oy-1+ + FeO

Fe2Oy+ Fe2Oy-2

+ + FeO2Fe2Oy

+ Fe2Oy-3+ + FeO3

a)

b)

c)

23456789

1 2 3 4 5 6

# of Oxygen atoms

DE

(eV

)

0

1

2

3

4

5

1 2 3 4 5 6

# of Oxygen atoms

DE

(eV

)

0

1

2

3

4

5

6

1 2 3 4 5 6

# of Oxygen atoms

DE

(eV

)

FeOy+ FeOy-1

+ + OFeOy

+ FeOy-2+ + O2

Fe2Oy+ Fe2Oy-1

+ + OFe2Oy

+ Fe2Oy-2+ + O

Fe2Oy+ Fe2Oy

+ + FeFe2Oy

+ Fe2Oy-1+ + FeO

Fe2Oy+ Fe2Oy-2

+ + FeO2Fe2Oy

+ Fe2Oy-3+ + FeO3

a)

b)

c)

0

1

2

3

4

5

1 2 3 4 5 6

# of Oxygen atoms

DE

(eV

)

0

1

2

3

4

5

6

1 2 3 4 5 6

# of Oxygen atoms

DE

(eV

)

FeOy+ FeOy-1

+ + OFeOy

+ FeOy-2+ + O2

Fe2Oy+ Fe2Oy-1

+ + OFe2Oy

+ Fe2Oy-2+ + O

Fe2Oy+ Fe2Oy

+ + FeFe2Oy

+ Fe2Oy-1+ + FeO

Fe2Oy+ Fe2Oy-2

+ + FeO2Fe2Oy

+ Fe2Oy-3+ + FeO3

a)

b)

c)

120

Table 7-2: Reaction products for the reaction between selected iron oxide cluster cationsand carbon monoxide.

FexOy+


DE (eV)

Products with N2

1,1 Fe+ 1.78 No products 1,2 FeO+

FeCO+ 2.72 0.31

No products

1,3 FeO+

FeOCO+

Fe+

FeO2+

5.36 1.66 0.65 2.64

FeO+

1,4 FeO2CO+

Fe(CO)2+

FeO+

3.75 1.99 2.49

FeO2+

Fe+

1,5 FeO3+

FeO3CO+ 1.16 3.06

FeO3+

2,1 Fe2+ 0.69 No products

2,2 Fe2O+

Fe2+

Fe2CO+

0.87 1.56 5.54

Fe2O+

2,3 Fe2O2+ 2.12 Fe2O2

+

2,4 Fe2O2CO+ Fe2O2

+ Fe2O3

+ FeO4CO+

6.17 4.25 2.12

Fe2O2+

Fe2+

2,5 Fe2O3+

Fe2O2+

Fe2O+

Fe2O3CO+

4.7 6.83 7.7

7.78

Fe2O3+

121

Other clusters undergoing oxygen atom transfer were FeO3+, Fe2O3

+, and Fe2O4+;

however, these species had competing reaction channels. In the studies involving

nitrogen, oxygen atom loss products were detected under conditions typically considered

to involve multiple collisions.

The Fe2O3+ cluster species was the most active and selective cluster toward the

transfer of an oxygen atom as shown in Figure 7-5. Studies with inert gas used in CID,

have also produced an atomic oxygen loss product at near thermal energies. This

suggests that there is a very low barrier for removing an oxygen atom from Fe2O3+ and

that the energy gained from forming CO2 allows the reaction to proceed. The increased

activity of Fe2O3+ is a consequence of the structure shown in Figure 7-2 and also the

bonding strength of the oxygen atom to iron. The cationic metal center binds the oxygen

atom more weakly than the anionic iron as shown in Chapter 6 (20). Fe2O3 is the most

studied species because it is stable in the condensed phase and thought to be the active

site for many reactions (21). It is not surprising our experiments show Fe2O3+ promotes

CO oxidation as studies of neutral Fe2O3 nanoparticles have shown iron to work as both a

catalyst in the presence of oxygen and a direct oxidant by losing lattice oxygen in the

absence of oxygen (22).

A secondary reaction observed for iron oxide cluster cations was CO bound to the

positive metal center. In the case of FeO2+, a product species corresponding to FeCO+

could arise from several different pathways including O2 replacement by CO or

association of CO after collisional O2 loss. From the branching ratio for FeO2+ in Figure

7-4(b), the product species FeCO+ occurs at low CO pressures supporting the O2

replacement by CO reaction channel. It is energetically feasible to remove oxygen and

122

bind CO since losing an O2 requires only 2.0 eV, providing for the energy gain in

forming an O2 molecule. The subsequent energy gain in binding CO is greater than 4 eV.

Although the product FeO+ could be an intermediate species in a successive reaction, the

loss of an oxygen atom from FeO+ requires more energy than the gain for binding CO.

This is evidence that these two species occur during separate reactions. Indeed, Figure 7-

6 shows the calculated profile for FeO2+ in which FeCO+ and FeO+ follow separate

reaction paths after the initial binding of CO.

Oxygen molecule replacement by CO was a prominent reaction channel observed

for many oxygen-rich clusters, including FeO3+, FeO4

+, FeO5+, Fe2O4

+. Most O2

dissociation energies for monomer iron oxide clusters were less than 1.4 eV. Figure 7-7

shows this reaction pathway for FeO4+, where the product species, FeO2CO+ resulted.

FeO4+ is also unique in that it also showed two CO molecules attached to the metal

center. Armentrout and co-workers have studied bond energies of gas-phase iron

carbonyls and found the strongest bond dissociation energy (BDE) to be for the second

CO ligand (4,5) This supports the high stability of Fe-CO bonds formed in this reaction.

In the case of FeO2+, FeO4

+, and Fe2O2+, carbon monoxide was observed to be bound to

bare iron, Fe1-2(CO)x+. It is not surprising that bare iron clusters were detected as

products from FeO+, FeO3+, Fe2O

+, and Fe2O2+ species, since iron is easily reduced. No

iron carbides were observed, which suggests that all the carbon present was oxidized

(20).

123

Figure 7-4: Branching ratios for a) FeO+ b) FeO2

+ c) Fe2O+ and d) Fe2O2

+ reacted with CO. Note the left-hand axis corresponds only to the selected ion, parent species as it decreases with CO gas. The right-hand axis tracks all product cluster species that increase with the addition of CO gas.

Figure 7-5: Branching ratio for Fe2O3

+ reacted with CO. Note the right-hand scale for product ions in Fe2O3

+ is double the amount in the branching ratios of Figure 7-4.

0.6

0.8

1

0 2.5 5 7.5 10CO (mTorr)

0

0.1

0.2

0.6

0.8

1

0 2.5 5 7.5 10

CO (mTorr)

0

0.1

0.2

0.8

0.9

1

0 2.5 5 7.5 10CO (mTorr)

0

0.1

0.2

FeO+

Fe+

Rel

.Int

. (p

aren

t io

n)

Rel. Int. (product ions)

a)R

el.I

nt. (

pare

nt io

n)


b)

FeO2+

FeO+

FeCO+

0.8

0.9

1

0 2.5 5 7.5 10

CO (mTorr)

0

0.1

0.2

Fe2O+

Fe2+

Rel

.Int

. (pa

rent

ion)


c)

Rel

.Int

. (pa

rent

ion)

Rel. Int. (product ion

s)

d)

Fe2O2+

Fe2O+

Fe2+

Fe2CO+

0.6

0.8

1

0 2.5 5 7.5 10CO (mTorr)

0

0.1

0.2

0.6

0.8

1

0 2.5 5 7.5 10

CO (mTorr)

0

0.1

0.2

0.8

0.9

1

0 2.5 5 7.5 10CO (mTorr)

0

0.1

0.2

FeO+

Fe+

Rel

.Int

. (p

aren

t io

n)


a)R

el.I

nt. (

pare

nt io

n)


b)

FeO2+

FeO+

FeCO+

0.8

0.9

1

0 2.5 5 7.5 10

CO (mTorr)

0

0.1

0.2

Fe2O+

Fe2+

Rel

.Int

. (pa

rent

ion)


c)

Rel

.Int

. (pa

rent

ion)

Rel. Int. (product ion

s)

d)

Fe2O2+

Fe2O+

Fe2+

Fe2CO+

0.6

0.8

1

0 4 8 12CO (mTorr)

0

0.2

0.4

Rel

.Int

. (pa

rent

ion)


Fe2O3+

Fe2O2+

0.6

0.8

1

0 4 8 12CO (mTorr)

0

0.2

0.4

Rel

.Int

. (pa

rent

ion)


Fe2O3+

Fe2O2+

124

ΔE (eV)

2FeO2+ + 1CO 2OCFeO2

+ -2.33 2OCFeO2

+ 6OCFeO2+ 0.82

6OCFeO2+ 4FeCO+ + 3O2 1.20

2OCFeO2

+ [2FeO2CO+] 1.06

[2FeO2CO+] 2FeO2CO+ -1.80 2FeO2CO+ 6FeO2CO+ -0.90 6FeO2CO+ 6FeO+ + 1CO2 1.26

Figure 7-6: Profile of the reaction energy for FeO2+ with CO including the corresponding

change in energy, ΔE, for each reaction step. The superscripts indicate spin multiplicity. Source: Figure courtesy of the group of Professor S. N. Khanna at VCU.

125

Figure 7-7: Branching ratio for FeO4+ reacted with CO showing at higher pressures two

CO molecules become attached to bare iron.

Rel

.Int

. (pa

rent

ion)

0

0.5

1

0 4 8 12

CO (mTorr)

FeO4+

FeO2CO+

Fe(CO)2+

FeO+

Rel

.Int

. (pa

rent

ion)

0

0.5

1

0 4 8 12

CO (mTorr)

FeO4+

FeO2CO+

Fe(CO)2+

FeO+

126

7.4 Conclusion

Structural calculations and dissociation energies of iron oxide cluster cations

found trends in the bond energy that correlate well with experimental CID fragmentation

products. The observed reaction pathways in the presence of CO were fully accounted

for by the thermodynamic bond energies and the gain in energy from binding CO to the

metal cluster. It is shown for certain stoichiometries that CO oxidation is favorable and

energetically feasible. However, for oxygen rich iron oxide clusters, O2 loss and O2

replacement were observed to be the major reaction channels. Therefore, the reactions of

iron oxide clusters were elucidated on the basis of cluster stoichiometry, oxidation state,

and charge state. These studies provide a foundation for further research in transition

metal oxide clusters as catalytic systems.

127

References

1. Guczi, L.; Frey, K.; Beck, A.; Petõ, G.; Daróczi, C. S.; Kruse, N.; Chenakin, S. Appl. Catal. A 2005, 291, 116.

2. PalDey, S.; Gedevanishvili, S.; Zhang, W.; Rasouli, F. Appl. Catal. B 2005, 56,

241.

3. Schröder, D.; Schwarz, H.; Clemmer, D. E.; Chen, Y.; Armentrout, P. B.; Baranov, V. I.; Böhme, D. K. Int. J. Mass Spec. Ion Proc. 1997, 161, 175.

4. Schultz, R. H.; Crellin, K. C.; Armentrout, P. B. J. Am. Chem. Soc. 1991, 113,

8590.

5. Armentrout, P. B.; Koizumi, H.; MacKenna, M. J. Phys. Chem. A 2005, 109, 11365.

6. Schröder, D.; Jackson, P.; Schwarz, H. Eur. J. Inorg. Chem. 2000, 2000, 1171.

7. Yumura, T.; Amenomori, T.; Kagawa, Y.; Yoshizawa, K. J. Phys. Chem. A 2002,

106, 621.

8. Jones, N. O.; Reddy, B. V.; Rasouli, F.; Khanna, S. N. Phys. Rev. B 2005, 72, 165411.



11. Kohn W.; Sham, L. J. Phys. Rev. 1965, 140, A1133.


13. Pederson M. R.; Jackson, K. A. Phys. Rev. B 1990, 41, 7453.

14. Jackson K.; Pederson, M. R. Phys. Rev. B 1990, 42, 3276.

15. Porezag D.; Pederson, M. R. Phys. Rev. A 1999, 60, 2840.

128

16. Köster, A. M.; Calaminici, P.; Casida, M. E.; Flores-Moreno, R.; Geudtner. G.; Goursot, A.; Heine, T.; Ipatov, A.; Janetzko, F.; del Campo, M. J.; Patchkovskii, S.; Reveles, J. U.; Vela, A.; Salahub, D. R. deMon2k V. 2.2.6., The international deMon developers community 2006, http://www.deMon-software.com; Köster, A. M.; Flores-Moreno, R.; Reveles, J. U. J. Chem. Phys. 2004, 121, 681; Reveles, J. U.; Köster, A. M. J. Comp. Chem. 2005, 25, 1109.

17. Godbout, N.; Salahub, D. R.; Andzelm, J.; Wimmer, E. Can. J. Chem. 1992, 70,

560.

18. Wachters, A. J. H. J. Chem. Phys. 1970, 52, 1033.; Wachters, A. J. H. IBM Tech. Rept. 1969, RJ584; F exponents: Bauschlicher, C. W., Jr.; Langhoff, S. R.; Barnes, L. A. J. Chem. Phys. 1989, 91, 2399.


Phys. Chem. A 1998, 102, 1733.

20. Reilly, N. M.; Johnson, G. E.; Castleman, A. W., Jr.; Reveles, J. U.; Khanna, S. N. Chem. Phys. Lett. 2007, 435, 295.

21. Abdel Halim, K. S.; Khedr, M. H.; Nasr, M. I.; El-Mansy, A. M. Factors affecting

CO oxidation over nanosized Fe2O3, Materials Research Bulletin 2006, doi:10.1016/j.materresbull.2006.07.009

22. Li, P.; Miser, D. E.; Rabiei, S.; Yadav R. T.; Hajaligol M. R. Appl. Catal. B 2003,

43, 151.

Chapter 8

Gas Phase Study of Cobalt, Nickel, and Copper Oxide Cluster Ions: Dissociation

Patterns and Trends in Reactivity with CO

8.1 Introduction

Transition metals offer a wide range of industrial applications, especially in

heterogeneous catalysis, therefore there is considerable interest in understanding their

reactive behavior. Investigations of transition metals in the bulk phase have been carried

out to describe their reactive and structural properties. However, on an atom by atom

basis, the properties and characteristics of transition metals change dramatically at the

nanoscale (1, 2). In this size regime, each atom may alter the structure, geometry and

electronic properties of a species. Gas phase metal cluster experiments are especially

amenable to studying the behavior of nanoscale materials through mass selected size-

specific experiments that have the ability to isolate potential catalytic active sites. Cluster

studies are able to probe the effects of size, oxidation and charge state, composition and

stoichiometry on cluster reactivity.

Numerous past studies have focused on bare transition metals and their

interactions with oxygen (3, 4), nitrogen (4-7) and carbon monoxide (4, 6, 8-16). Efforts

to elucidate small cluster geometries and to characterize surface binding sites have been

conducted through uptake and saturation experiments using either CO or N2, through

collision induced dissociation (CID) studies with xenon, and by various theoretical

130

structural calculations (17-19). Experiments involving CO have used the saturation limits

of metal sites to elucidate complicated transition metal structures through electron

counting of closed shell orbitals (4, 20, 21). Wang and coworkers (22, 23) and Ganteför

et al. (24) studied the electronic structure of several 3d transition metal oxides with

photoelectron spectroscopy (PES). Collision induced dissociation studies also revealed

information on bond strengths and trends in metal cluster structures (25-29). However, it

is difficult to fully determine the nature of bonding in transition metals because there is

limited information available on the ground state structures of metal oxides (23) and the

electronic properties change when oxygen is bound to the metal (25). The reaction of

transition metals with both oxygen and CO (30-33) therefore provides new insight into

catalytic applications, with particular emphasis on environmental pollution abatement.

One study involving CO reactivity with ozone and oxygen found the activity of CO

oxidation was higher in the presence of ozone with ozone decomposition also occurring

(30).

Many interesting structural and reactive properties arise in transition metals due to

the large number of d electrons and their localized bonding character (19,34-36).

Fielicke et al. investigated metal-CO bond strength and metal carbonyl properties based

on cluster size (37). They employed IR-multiple photon dissociation (IR-MPD)

spectroscopy to study the binding of CO on free nickel and cobalt clusters. Their focus

on electrostatic charge and size effects in small clusters helped identify different

structural binding sites of CO and trends in CO binding to various metal clusters. Several

studies have also tracked the effect of changes in the oxidation state of transition metals

toward the observed reactivity for CO oxidation (38-41). Jernigan et al. found for

131

metallic copper and copper (I), CO oxidation proceeds by the Langmuir-Hinshelwood

mechanism, whereas copper (II) was proposed to operate through a redox cycle

mechanism (39). Furthermore, the activity for CO oxidation was found to decrease with

increasing oxidation state of Cu. There is no consensus in the literature on the nature of

the active sites and mechanisms governing the oxidation of CO in the presence of

transition metal oxides due to complications arising from environmental conditions such

as temperature (40) and humidity (42). Therefore, further investigation of isolated

transition metal oxide systems in the presence of CO would be useful.

Experiments investigating several first row transition metals were undertaken to

determine the physical and chemical characteristics of cobalt, nickel, and copper oxides.

Studies on iron oxides covered in Chapters 5-7 are compared here to elucidate the

periodic trends in binding energy and reactivity toward CO oxidation (43). A

comprehensive study of their reactivity and structures through collision induced

dissociation (CID) experiments reveals how additional d electrons in the system and

changes in electron affinity affect the reactivity of the metal oxide clusters.

This chapter focuses on transition metal elements known to be common dopants

or supports for industrial catalysts. Previously, it was believed that the support materials

used in heterogeneous catalysis were inert substances which did not influence the

reactivity of the catalyst. Recently, however, studies conducted on catalysts with various

metal oxide supports have begun to shed light on the role that the support plays in the

catalytic process (44-47). Metal oxide supports were shown to have considerable impact

over the particle size and activity of catalysts in low temperature oxidation of CO (48). It

has also been proposed that CO adsorbs to the catalyst, while transition metal oxide

132

supports, such as CoOx, provide the necessary oxygen to oxidize CO (49). An

understanding of the influence of the metal oxide support may aid in designing enhanced

materials to deposit catalysts onto, which will increase the activity and selectivity toward

a chosen reaction pathway. These pure metal oxide studies lay the groundwork for

further exploration into the effects of catalyst-support interactions and dopant effects.

8.2 Experimental

All metal oxide cluster studies were conducted on a guided ion beam mass

spectrometry coupled to a laser vaporization source described in detail previously (50)

and in Chapter 2. Briefly, the second harmonic of a Nd:YAG laser was used to ablate a

continuously rotating and translating metal rod. Oxygen seeded in helium (~10%) was

pulsed over the ablated metal rod at a predetermine time to form the metal oxide cluster

ions. A 27 mm conical expansion nozzle was employed to aid in the clustering of these

species. After exiting the nozzle, the ions traversed a field free region, were collimated

by a 3 mm skimmer and focused by a set of electrostatic lenses and deflectors into the

mass selecting quadrupole. Individual cluster ions were mass selected and focused by a

second set of lenses and directed into the octopole reaction cell. Carbon monoxide

reactant gas ranging from 0-10 mTorr for cations and 0-20 mTorr for anions was added

into the reaction cell and the pressure of the gas was monitored by a MKS Baratron. The

remaining reactant and product ions were mass analyzed in the second quadrupole and

detected by a channel electron multiplier. For collision induced dissociation (CID)

experiments, xenon was added to the reaction cell at three different pressures, 0.08

133

mTorr, 0.12 mTorr, and 0.22 mTorr, as the DC voltage on the octopole rods was slowly

increased from thermal to 40 eV laboratory-frame energy. The first two pressures were

calculated to be single collision conditions while the last pressure was at multiple

collision conditions. Sequential fragmentation was determined from the results of

multiple collision products at elevated energies and pressures.


Figure 8-1 shows the typical mass distributions for anionic metal oxide clusters

of nickel, cobalt and copper. These are stable metal oxide species formed in the laser

vaporization source and mass selected for the present studies. The corresponding cationic

mass distributions for nickel and cobalt oxides are shown in Figure 8-2. An analysis of

the fragmentation patterns and insight into the structures of the metal oxide ions in this

study will be presented first through CID experiments. Reactivity studies with carbon

monoxide reactant gas and a discussion based on stoichiometry, composition, and charge

state follows.

8.3.1 Collision Induced Dissociation Studies

Table 8-1 and Table 8-2 show the fragmentation products from CID studies for

cobalt, nickel, and copper oxide anions and cations, respectively.

134

Figure 8-1: Typical anionic mass distributions for a) Nickel oxides, b) Cobalt oxides, and c) Copper oxides.

0 100 200 300 400

0 100 200 300 400 500

CoO

2-

CoO

3-

Co 2O

3-

Co 2O

5-

Co 3O

4-

Co 4O

6-

Co 5O

6-

Ion

Inte

nsit

y (a

rb. u

nits

)

Mass (a.m.u.)

b)

0 100 200 300 400

Mass (a.m.u.)

Ion

Inte

nsit

y (a

rb. u

nits

)

NiO

2-

NiO

3-

Ni 2O

3-

Ni 3O

3- Ni 4O

4-

Ni 5O

5-

a)

Mass (a.m.u.)

Ion

Inte

nsit

y (a

rb. u

nits

)

CuO

2-

Cu 2O

3-

Cu 3O

3-

Cu 4O

4-

Cu 5O

5-

c)

0 100 200 300 4000 100 200 300 400

0 100 200 300 400 5000 100 200 300 400 500

CoO

2-

CoO

3-

Co 2O

3-

Co 2O

5-

Co 3O

4-

Co 4O

6-

Co 5O

6-

Ion

Inte

nsit

y (a

rb. u

nits

)

Mass (a.m.u.)

b)

0 100 200 300 400

Mass (a.m.u.)

Ion

Inte

nsit

y (a

rb. u

nits

)

NiO

2-

NiO

3-

Ni 2O

3-

Ni 3O

3- Ni 4O

4-

Ni 5O

5-

a)

Mass (a.m.u.)

Ion

Inte

nsit

y (a

rb. u

nits

)

CuO

2-

Cu 2O

3-

Cu 3O

3-

Cu 4O

4-

Cu 5O

5-

c)

135

Figure 8-2: Typical cationic mass distributions for a) Nickel oxides and b) Cobalt oxides.

0 50 100 150 200Mass (a.m.u.)

Ion

Inte

nsit

y (a

rb. u

nits

)

NiO

+

Ni+

NiO

2+ NiO

3+

NiO

6+

NiO

5+

NiO

4+

Ni 2O

2+

NiO

7+

Ni 2O

4+

NiO

8+

Ni 2O

+

Ni 2O

3+

a)

0 50 100 150 200 250 300

CoO

+

CoO

2+

O2+

2O2+

CoO

3+ CoO

4+

CoO

5+ Co 2O

2+

CoO

6+

Co 3

+

Co 2

O4+

CoO

8+

CoO

9+

CoO

10+

Co 2

O5+

Co 2O

6+

Co 3O

3+C

o 2O7+

CoO

11+

Co 3

O4+

Co 2O

8+

Co 3O

5+

Co 2

O9+

Mass (a.m.u.)

Inte

nsit

y (a

rbit

rary

uni

ts)

Co+

Co 2O

10+

b)

0 50 100 150 200Mass (a.m.u.)

Ion

Inte

nsit

y (a

rb. u

nits

)

NiO

+

Ni+

NiO

2+ NiO

3+

NiO

6+

NiO

5+

NiO

4+

Ni 2O

2+

NiO

7+

Ni 2O

4+

NiO

8+

Ni 2O

+

Ni 2O

3+

0 50 100 150 200Mass (a.m.u.)

Ion

Inte

nsit

y (a

rb. u

nits

)

NiO

+

Ni+

NiO

2+ NiO

3+

NiO

6+

NiO

5+

NiO

4+

Ni 2O

2+

NiO

7+

Ni 2O

4+

NiO

8+

Ni 2O

+

Ni 2O

3+

a)

0 50 100 150 200 250 300

CoO

+

CoO

2+

O2+

2O2+

CoO

3+ CoO

4+

CoO

5+ Co 2O

2+

CoO

6+

Co 3

+

Co 2

O4+

CoO

8+

CoO

9+

CoO

10+

Co 2

O5+

Co 2O

6+

Co 3O

3+C

o 2O7+

CoO

11+

Co 3

O4+

Co 2O

8+

Co 3O

5+

Co 2

O9+

Mass (a.m.u.)

Inte

nsit

y (a

rbit

rary

uni

ts)

Co+

Co 2O

10+

b)

136

Anionic oxygen rich metal oxide clusters, MO>3-, M2O>4

-, M3O>5-, exhibited molecular

O2 loss at near thermal energies and under single collision conditions. Previous

calculations performed on iron oxide clusters found a maximum saturation of four

oxygen atoms on a metal center (43, 51). For oxide clusters with oxygen atoms bound

directly to the metal, CID experiments have indicated that oxygen recombination takes

place (52). However, the facile loss of O2 indicates that oxygen is loosely bound as either

a peroxide or superoxide unit for cluster species with coordinative saturation. Metal

atoms, MO and MO2 were also common neutral fragments of larger anionic clusters

demonstrating the greater electron affinity of the more oxygen rich fragments produced

by dissociation of these neutral fragments. Sequential fragmentation occurred with select

copper oxides at elevated energies and under multiple collision conditions. Cationic

metal oxide clusters also lost O2 at near thermal energy for MO>4+ and M2O>2

+ clusters

which is consistent with the smaller activation of the O-O bond near charge withdrawing

metal centers.

Cobalt. The anionic cobalt oxides studied were CoO2-3-, Co2O3-5

-, and Co3O3-6-.

Oxygen atom loss was found to be the major fragment for monomer clusters, CoO2- and

CoO3-, occurring at slightly elevated energies. This is in agreement with the theoretical

structural calculations of Uzunova et al. who predict atomically bound oxygen units (53).

Neutral metal atom loss was observed for dimer oxides resulting in a stable CoO2-

subunit, while for trimer oxides, Co2O3- was the most stable fragment species. This in

agreement with reactions of cobalt anions and oxygen performed by Kapiloff and Ervin,

who found CoO2- and Co2O3

- to be the major products at high O2 concentration,

indicating that these structures were particularly stable (4).

137

Table 8-1: CID fragmentation channels for selected MxOy- (M=Co, Ni, Cu) anions.

Cobalt Nickel Copper Cluster MxOy

-


Neutral (s) Lostb

Products with Xea

Neutral (s) Lostb

Products with Xea

Neutral (s) Lostb

1,2 1,1 0,1 1,1 0,1 1,1 0,1 1,3 1,2 0,1 1,1

1,2 0,2 0,1

1,1 1,2

0,2c 0,1

2,3 1,2 2,2

1,1 0,1

1,2 2,2

1,1 0,1

1,2 2,2 2,1

1,1 0,1 0,2

2,4 1,3 1,2 2,3

1,1 1,2 0,1

1,2 2,3 2,2 1,3

1,2 0,1 0,2 1,1

2,2 2,3 1,3 1,2

0,2c 0,1 1,1d 1,2d

2,5 2,3 1,3 2,4 1,2

0,2 1,2 0,1 1,3

2,3 0,2

3,3 2,3 1,0 2,3 1,0 3,4 3,3

2,3 0,1 1,1

3,2 3,3 2,3

0,2 0,1 1,1

3,5 3,3 3,4 2,4 2,3

0,2 0,1

3,3 3,4 2,3

0,2c 0,1 1,2

3,6 3,4 3,5 2,4 3,3 2,3

0,2c 0,1 1,2 0,3 1,3

4,4 4,2 3,3

0,2 1,1

3,3 1,1d

4,5 4,3 4,4 2,3

0,2 0,1 2,2

4,3 4,4 3,3

0,2c 0,1 1,2

aFragmentation channels are shown in order of observation with increasing collision energy. bThe neutral loss is assigned based on the difference between the selected cluster and fragment ion formed. cRepresents that dissociation occurs at near thermal energies. dDescribes channels taking place under multiple collision conditions.

138

Table 8-2: CID fragmentation channels for selected MxOy+ (M=Co, Ni) cations.

Cobalt Nickel Cluster MxOy

+


Neutral (s) Lostb

Products with Xea

Neutral (s) Lostb

1,1 1,0 0,1 1,0 0,1 1,2 1,0

1,1 0,2 0,1

1,0 1,1

0,2 0,1

1,3 1,1 1,2

0,2 0,1

1,1 1,2

0,2 0,1

1,4 1,2

1,3 0,2c

0,1c 1,2 1,3

0,2 0,1

2,2 2,1

2,0 0,1 0,2

2,0 1,3 1,2

0,2c

1,1 1,0

2,3 2,1 0,2c

2,4 2,2

2,0 0,2c

0,4

2,6 2,4

2,2

2,1

1,4 1,2

0,2c

0,4c

0,5c

1,2 1,4

aFragmentation channels are shown in order of observation with increasing collision energy. bThe neutral loss is assigned based on the difference between the selected cluster and fragment ion formed.cRepresents that dissociation occurs at near thermal energies.

139

For larger cobalt oxide anions of Co2O5-, Co3O5

- and Co3O6-, O2 loss was a major product

fragment. O2 loss from larger CoxOy- clusters is consistent with the PES results of

Pramann et al. who predict peroxo bound O2 units for these species (54). In Co2O5- the

ordering of CID products and N2 products shown in Table 8-3 was reversed. It should be

noted that nitrogen products were observed at relatively high pressures, approximately 15

mTorr N2 and the intensity was roughly in the same relative porportion. When energy

was added into the system during collisions with Xe, loss of an oxygen molecule was

observed first, in contrast with N2 collisional studies where an oxygen atom was the first

product observed. This may be due to the sticky nature of the collisions between Xe and

Co2O5- which may better facilitate vibrational energy transfer resulting in a larger

fragment species (55).

The cationic cobalt oxides examined were CoO1-4+ and Co2O2,4,6

+. Both CoO+

and Co2O2+ lost an oxygen atom first at slightly elevated energies. The remaining

monomer clusters lost O2 molecules first then O atoms, with CoO4+ losing both O2 and O

at near thermal energy. The cationic dimer cobalt oxides were unique because they

attached only an even number of oxygen atoms. From these studies it can be assumed

that only molecular oxygen was associated to Co2O4+ and Co2O6

+ because O2 loss

occurred readily under thermal energy. For Co2O4+ and Co2O6

+ the major fragments

were O2 loss followed by the loss of two O2 molecules. The stable fragments tended to

have molecular oxygen units and included CoO2+, CoO4

+ and Co2O2+.

Nickel. The following nickel oxide anions were studied: NiO2-3-, Ni2O3-4

-, Ni3O3-5-

and Ni4O4-5-. Previous experimental work in our laboratory utilizing a metal ion fast flow

reactor has shown similar nickel oxide anion and cation mass distribution (56).

140

Table 8-3: Reaction products for selected MxOy- (M=Ni, Co, Cu) anions with CO and N2.

Cobalt Nickel Copper Cluster MxOy

-


Products with N2

Products with CO

Products with N2

Products with CO

Products with N2

1,2 1,1 No products

1,1 No products

1,1 1,0

1,3 1,2 1,2 1,2 No products

1,2

1,1

1,0a

1,1 1,2

2,3 2,2

1,2a 2,1 2,2

1,2a No

products 2,2

2,1a

1,2a

No products

2,4 2,3a No products

2,3

2,2 2,3 2,3

2,2 2,2 2,3

2,5 2,4

2,3

1,3

2,4 2,3 1,3

2,4

2,3 2,1a 1,2a

2,4 2,3

3,3 No Reaction

3,2

3,1a 3,1 2,3

3,4 3,3a

2,4a No

products 3,3 No

products 3,3 3,2

3,3 3,5 3,4

3,3

2,3a

2,4a

No products

3,3

3,4 3,3 3,4

3,6 3,5

3,4

3,3

2,4

2,3a

3,4 3,5

4,4 No Reaction

4,3a

3,3a 4,2 3,3

4,5 4,4

4,3 4,3 4,4

4,4 4,4 4,3 3,3

4,6 4,4a 4,7 4,5

4,6a

aDenotes a minor product channel whose relative intensity is less than 1% of the total reaction products and are not displayed in their respective branching ratios.

141

It is interesting that monomer and dimer anionic species in the present studies are oxygen

rich which is consistent with previous studies showing slightly more oxidized anionic

species. The absence of nickel monomer and dimer clusters in the previous study was

attributed to a slower electron attachment rate. Indeed, the electron affinities of NixOy+,-

increase with increasing size. The fragments occurring at lower energies were mostly O

or O2 units; however, several clusters lost neutral nickel and nickel oxygen. For Ni2O3-

and Ni2O4-, NiO2

- was the first fragment detected. Studies with N2 gas shown in Table 8-

3 verify that this product fragment was initiated through additional kinetic energy added

to the system and did not result from thermal collisions. Since these products were not

observed with N2 at high collision pressures, it suggests that the addition of energy in

CID experiments was needed for any metal bond dissociation. In the case of Ni3O5- and

Ni4O5- the order of CID products follows the same order of N2 products observed.

The nickel oxide cations studied included NiO1-4+ and Ni2O2-3

+. For all of the

clusters except NiO+, O2 was the first fragment observed. The dimer clusters showed this

O2 subunit to desorb at near thermal energy, demonstrating that this fragment was only

weakly bound to the nickel cation. Similar results were observed with N2 studies in

Table 8-4 with the exception of NiO+ and NiO2+. Only Ni2O2

+ exhibited loss of a metal

atom, revealing a stable NiO2+ species at the highest energies. This is consistent with the

results of Vardhan and coworkers who demonstrated Ni atom loss from nickel dioxides

formed at low energy in a molecular beam (57).

Copper. Anionic copper oxide clusters which were studied include CuO2-3-,

Cu2O3-5-, Cu3O3-4

-, and Cu4O4-5-. Oxygen rich clusters that lost O2 readily were CuO3

-,

Cu2O4-, and Cu4O5

-, with Cu2O5- losing molecular oxygen at near thermal energy.

142

Table 8-4: Reaction products for MxOy+ (M=Co, Ni) cations with CO and N2.

Cobalt Nickel Cluster MxOy

+


Products with N2

Products with CO

Products with N2

1,1 (1,0) + No products

(1,0) + (1,0) +

1,2 (1,0)CO+

(1,0) +

(1,1)(CO)2+a

No products

(1,0)CO+

(1,1) + (1,1) + (1,0) +

1,3 (1,1)CO+

(1,1) + (1,0) +

(1,1)(CO)2+a

(1,1)CO+

(1,0) +

(1,2) +

(1,0)CO+a

(1,1) +a

(1,1) +

1,4 (1,2)CO+

(1,0)(CO)2+

(1,2) +

(1,0)CO+

(1,0)CO+

(1,2) +

(1,2)CO+

(1,1) +

(1,0) +

2,1 (1,1)CO+ 2,2 (1,0)(CO)2

+

(1,2)CO+ (1,0)(CO)2

+

(1,2)CO+

(1,0)CO+

(1,1)CO+a

(2,0) +

2,3 (2,0)CO+

(2,1)CO+ (1,1)(CO)2

+ (2,2) +

(1,0)(CO)2+

(1,1)CO+ (1,0)CO+

(2,1) +

2,4 (2,0)CO+

(2,2) + (1,0)(CO)2

+

(1,2)CO+

2,6 (2,4) +

(2,2) + (1,0)(CO)2

+

(1,0)CO+a

aDenotes a minor product channel whose relative intensity is less than 1% of the total reaction products and are not displayed in their respective branching ratios.

143

PES studies conducted by Wang et al. have revealed that once the transition metal

reaches its saturated oxidation state of +3 for copper, then molecular O-O units are

formed (58). This supports the facile loss of O2 from oxygen-rich clusters since

dissociative adsorption of oxygen does not occur on saturated transition metals. Loss of

CuO and Cu were the first fragments observed for Cu2O3- and Cu3O3

-, respectively, with

neutral Cu loss for Cu4O4- occurring under multiple collision conditions. Other metal

atom losses at multiple collisions occurred in Cu2O4- which produced a stable CuO2

-

species. For this stable copper dioxide molecule, it has been proposed by Wang et al. that

two isomers exist, OCuO, and Cu(O2), even though they do not observe CuO- at the

wavelengths employed in the PES experiments (23). For Cu3O3- and Cu4O4

- the first

fragmented species was Cu2O3- and Cu3O3

-, respectively. This agrees with CID studies

conducted by Gord et al. on oxygen-deficient and metal and oxygen equivalent copper

oxide anions (25). The reason we may not observe further dissociation products reported

by Gord et al., however, may be due to the short reaction time of our experiments

compared to the trapping ability of the FT-ICR-MS technique. Several of the copper

oxide CID products fragmented differently with Xe than the N2 collisional fragments

shown in Table 8-3. For example, Cu4O5- produced O2 loss at near thermal energy with

Xe, while it was the second product detected in N2 studies. Again, this may be due to

better kinetic to vibrational energy transfer with Xe as characterized by Armentrout et al.

(55).

The mass degeneracy observed in the spectra of copper oxide cations made

identification of individual cluster species difficult. No copper oxide cationic species are

144

reported herein due to the production of oxygen rich metal clusters having mass overlap

with larger metal clusters.

Trends. Similar dissociation patterns were observed for each cluster size;

however, there were some differences in the energy ordering of products within the metal

clusters. For M2O3-, all species lost neutral MO first, followed by O loss. This sequence

is unique because most often oxygen atoms or molecules are observed to fragment first.

It is likely, therefore, that the M-O bond in these clusters is stronger than M-M bonding.

Similar products are detected for M2O4- clusters, but with a different order. Nickel oxides

lose MO2, cobalt oxides lose MO, and copper oxides lose neutral O2 upon collisional

studies at thermal energy. This suggests that the structures and bond strengths are

different between the various metal oxides of this stoichiometry. For cobalt specifically,

loss of an oxygen atom is the third product fragment, compared to the second detected for

both nickel and copper. Therefore we expect this increased M-O bond energy to reflect

in the reactivity studies, discussed below. Other minor differences occur in M3O4- and

M4O4- between nickel and copper oxides, with nickel losing an O atom and NiO more

easily and copper losing O2 and CuO.

The differences in the N2 and CID dissociation patterns occur mostly in the

oxygen rich metal clusters, Co2O5-, Cu2O5

- and Cu4O5-. This difference suggests that

increasing the kinetic energy of the collision will increase the internal energy of the

reactant species resulting in vibrational excitation and geometric changes. If the energy

within the system allows for rearrangement of the structure and two oxygen atoms come

in close contact, then the energy to form an oxygen molecule and allow for O2

dissociation may exceed the energy required to break the metal-oxygen bonds. Other

145

studies have reported on the slight activity observed between transition metals and N2,

however the reaction products were often molecular adsorption of N2 onto the metal

cluster (6). Armentrout and coworkers have observed FexN+ products in reactions

between Fex+ clusters and N2, but only at elevated energies (59). We do not observe

molecular or dissociative adsorption onto any metal clusters and only collisional

fragmentation with N2 occurs in the present experiments. It is not thermodynamically

feasible to break a N-N bond to form NO and we do not believe this chemical reaction is

occurring in these experiments. However, at the high pressures indicative of multiple

collisions, it may be possible to form N2O.

The dissociation patterns of cationic metal oxide exhibited a main product

fragment of O2. The only exception was in the M2O2+ cluster stoichiometry. Here, O2

loss was followed by MO and M loss in nickel clusters, whereas O loss was followed by

O2 loss in cobalt clusters. Although most studies on the structure of transition metal

oxides revealed metal oxygen bonds and not molecular oxygen bonds until coordinative

saturation (60), oxygen recombination was observed. Previously, in our laboratory, we

observed O2 loss in V2O5+ in CID studies (52). Also, N2 and CID dissociation patterns

were observed to differ for only NiO2+, with O atom loss dominant in N2 studies and O2

loss in CID experiments.

8.3.2 Reactivity Studies

Table 8-3 and Table 8-4 present the products of reactions with CO and collisions

with N2 for each of the nickel, cobalt, and copper oxide clusters studied. Figure 8-3

146

shows the branching ratios for the clusters having the most reactive stoichiometries for

oxygen atom transfer to CO for anionic cobalt, nickel, and copper oxides. For both

cobalt and nickel the same stoichiometry observed in iron oxides of (x, y = x+1) is the

most active. However, in copper oxides, a stoichiometry with the same number of metal

atoms and oxygen atoms is the most reactive. For cationic clusters of nickel and cobalt

oxides non-equivalent numbers of metal and oxygen atoms were the most active toward

oxidation. Shown in Figure 8-4 for NiO3+ and CoO2

+, the major reaction channel was O2

replacement by CO. Detailed below are the individual metal oxides and the observed

products for anion and cation reactions with CO.

Cobalt. Oxygen atom transfer was observed for all anionic monomer through

trimer cobalt clusters. CoO3- exhibited a weakly bound oxygen atom as shown in N2

studies at high pressures, while Co2O4- and Co3O4

- had only a slight reaction with less

than 1% relative product intensity. Molecular oxygen loss was another product shown in

reactions with CO especially in larger oxygen rich clusters, Co4O6- and Co4O7

-. It is

interesting that Co3O5- and Co3O6

-, which are also oxygen rich species, underwent

oxygen atom transfer as the major reaction channel instead of molecular oxygen loss.

For cationic monomer cobalt clusters, O2 replacement by CO was the major

product followed by O2 loss. In the dimers, O2 loss was the dominant pathway with

attachment of CO to the metal cluster following. For Co2O2+, however, CoO2 and Co

replacement by a CO molecule was detected. Also multiple replacement reactions were

observed for Co2O4+ and Co2O6

+ with two CO molecules attached to the metal cluster.

Nickel. Nickel oxide anions underwent mostly oxygen atom transfer to CO

forming neutral CO2.

147

Figure 8-3: Branching ratios for a) Co2O3- b) Ni2O3

-and c) Cu3O3- reacted with carbon

monoxide. Each cluster shows an active reaction pathway for oxygen atom transfer.

Rel

ativ

e In

tens

ity R

eact

ant

0.7

0.8

0.9

1

0 5 10 15

CO (mTorr)

0

0.1

0.2

0.3

Ni2O3-

Ni2O2-

Rel

ativ

e In

tens

ity

Rea

ctan

t Relative Intensity Product

0.7

0.8

0.9

1

0 5 10 15CO (mTorr)

0

0.1

0.2

0.3

Cu3O3-

Cu3O2-

Cu3O-

Rel

ativ

e In

tens

ity

Rea

ctan

t Relative Intensity Products

(b)

(c)

0.7

0.8

0.9

1

0 4 8 12 16CO (mTorr)

0

0.1

0.2

0.3

Co2O3-

Co2O2-

Relative Intensity Product

(a)

Rel

ativ

e In

tens

ity R

eact

ant

0.7

0.8

0.9

1

0 5 10 15

CO (mTorr)

0

0.1

0.2

0.3

Ni2O3-

Ni2O2-

Rel

ativ

e In

tens

ity

Rea

ctan


0.7

0.8

0.9

1

0 5 10 15CO (mTorr)

0

0.1

0.2

0.3

Cu3O3-

Cu3O2-

Cu3O-

Rel

ativ

e In

tens

ity

Rea

ctan


(b)

(c)

0.7

0.8

0.9

1

0 4 8 12 16CO (mTorr)

0

0.1

0.2

0.3

Co2O3-

Co2O2-


(a)

0.7

0.8

0.9

1

0 5 10 15

CO (mTorr)

0

0.1

0.2

0.3

Ni2O3-

Ni2O2-

Rel

ativ

e In

tens

ity

Rea

ctan


0.7

0.8

0.9

1

0 5 10 15CO (mTorr)

0

0.1

0.2

0.3

Cu3O3-

Cu3O2-

Cu3O-

Rel

ativ

e In

tens

ity

Rea

ctan


(b)

(c)

0.7

0.8

0.9

1

0 4 8 12 16CO (mTorr)

0

0.1

0.2

0.3

Co2O3-

Co2O2-


(a)

148

Figure 8-4: Branching ratios for a) CoO2+ and b) NiO3

+ reacted with carbon monoxide. The dominant reaction pathway for both clusters is molecular oxygen replacement by CO.

0

0.25

0.5

0.75

1

0 2 4 6 8CO (mTorr)

0

0.25

0.5

0.75

1

0 2.5 5 7.5 10CO (mTorr)

NiO3+

NiOCO+

Ni+

NiO2+R

elat

ive

Inte

nsit

y

CoO2+

CoCO+

Co+

Rel

ativ

e In

tens

ity

(b)

(a)

0

0.25

0.5

0.75

1

0 2 4 6 8CO (mTorr)

0

0.25

0.5

0.75

1

0 2.5 5 7.5 10CO (mTorr)

NiO3+

NiOCO+

Ni+

NiO2+R

elat

ive

Inte

nsit

y

CoO2+

CoCO+

Co+

Rel

ativ

e In

tens

ity

(b)

(a)

149

Ni2O4- had a weakly bound oxygen atom as nitrogen studies also showed oxygen atom

loss with high collision pressures. For larger metal clusters, only Ni3O4- and Ni4O5

-

showed oxygen atom transfer, evidence that clusters with one more oxygen atom than

metal atom have a particularly active stoichiometry. For Ni3O5-, molecular oxygen loss

was observed followed by oxygen atom loss. No reactions were observed with clusters

containing the same number of metal and oxygen atoms, Ni3O3- and Ni4O4

-.

In cationic nickel oxides, oxygen atom transfer was observed for NiO+, NiO2+,

and Ni2O3+; however, it was the major reaction channel only for NiO+. Molecular

oxygen loss was a minor reaction channel observed for select clusters including, NiO3+

and NiO4+. The major product channel detected was O2 replacement by a CO molecule.

Dimer species also underwent Ni, NiO, or NiO2 replacement by CO and in some cases

two CO molecules were bound to the metal cluster. For Ni2O2+, all the products had

metal atom loss and CO attached to either Ni+ or NiO2+. In studies by Parks and

coworkers on CO binding to nickel, it was observed that adsorption of the CO molecule

weakened the Ni-Ni bonding (18). This may explain why metal atom loss was prevalent

when carbon monoxide was attached to the product species.

Copper. In all the anionic copper oxides studied the major product channel was

oxygen atom transfer, with Cu4O4- being only slightly reactive. This reaction pathway

was followed by O2 loss and metal atom loss for Cu2O3- and Cu4O4

-. For Cu2O5- oxygen

atom transfer was observed with CO; however, this was a very weakly bound oxygen

atom since experiments with N2 showed the same results.

150

The mass degeneracy in oxygen rich cationic copper oxide clusters made the

study of these species difficult due to the resolution of our mass spectrometer and the

results are not reported here.

Reactivity Trends. The reactivity of transition metal oxide clusters was studied by

investigating the effects of charge state, cluster size, composition, and stoichiometry on

the observed reaction pathways. The results for nickel, cobalt, and copper oxides show

that the reaction pathways were similar for each charge state. However, unique

characteristics for each metal composition depending on the stoichiometry and metal to

oxygen ratio were observed. Oxygen atom transfer was the major pathway for negatively

charged clusters and molecular oxygen replacement by CO was the major reaction

channel for positively charged metal oxide clusters.

The reactions of positively charged transition metal oxides are different from the

corresponding anionic clusters, since CO is attached to the metal oxide center for most

cationic clusters. Figure 8-5 shows NiO2-,+ clusters reacted with CO. From these

branching ratios, it can be seen that oxygen atom transfer occurs for both charge states,

however cationic NiO2+ has a competing O2 replacement by CO reaction pathway.

Differences in the reactive behavior of these small transition metal oxides may be

explained by the cluster charge state and effects of electron density towards bonding. For

cationic transition metals, the positive metal charge induces electrostatic effects which

create a strong bond between the charge on the metal and CO (36). The general

mechanism for CO adsorption onto a metal surface is described by the Blyholder model,

which involves σ donation and π backdonation (61).

151

Figure 8-5: Branching ratios for a) NiO2- and b) NiO2

+ reacted with carbon monoxide. The dominant reaction pathway for anions is oxygen atom transfer while cations have two competing reactions, oxygen atom transfer and molecular oxygen replacement by CO.

0

0.25

0.5

0.75

1

0 5 10 15

CO (mTorr)

NiO2-

NiO-

0

0.25

0.5

0.75

1

0 2.5 5 7.5 10

CO (mTorr)

Rel

ativ

e In

tens

ity

Rel

ativ

e In

tens

ity

NiO2+

NiCO+

NiO+

(b)

(a)

0

0.25

0.5

0.75

1

0 5 10 15

CO (mTorr)

NiO2-

NiO-

0

0.25

0.5

0.75

1

0 2.5 5 7.5 10

CO (mTorr)

Rel

ativ

e In

tens

ity

Rel

ativ

e In

tens

ity

NiO2+

NiCO+

NiO+

(b)

(a)

152

Essentially the lone pair electron of CO in the occupied 5σ antibonding orbital donates

into the metal σ valence orbital, while the unoccupied 2π antibonding orbital in CO

accepts charge from the d orbitals of the metal. Therefore electrostatic effects are used to

generate stronger metal-CO bonds. In cationic nickel and cobalt oxide clusters the

positive metal center disperses its charge into the metal-carbon (M-C) bond which creates

more covalent bond character (37). In free CO, electron density is polarized toward the

more electronegative O atom. When CO binds to M+, the positive charge at the metal

center draws electron density away from the O atom and into the bond, increasing

covalency. Also, the M+-C-O- attraction increases the bond strength. Investigations of

late d-block metal cations have also suggested stronger M-CO bonds due to the effect of

coulombic attraction between positive metal centers and CO (62).

The anionic counterparts show weaker M-C bonds for several reasons including

differences in the electronic structure compared to cations. For anions, a non-optimal

overlap exists due to the small 3d metal orbitals and the 5σ orbital of CO. The presence

of the extra electron in anionic systems greatly affects its behavior (17). For example,

CuO- has its extra electron occupying the 2pπ orbital resulting in a closed shell electronic

ground state (22). This anionic configuration does not accept the charge from the lone

pair of CO and therefore we do not see CO attached to the anionic cluster.

Another significant effect that charge state has on the reactivity of transition metal

oxides is on the qualitative reaction rate. From Figure 8-5 it can be seen that the

reactions of cations are more active at lower CO pressures (10 mTorr versus 15 mTorr in

anions). This may be a consequence of the electron deficient cation exhibiting higher

activity as it will be able to accept σ electron density from CO forming a bond. In studies

153

by Russell et al., smaller ionic clusters and clusters with open coordination sites

displayed an increased relative reaction rate (63).

The metal to oxygen ratio is observed to affect the reactivity in the present

studies. Stable cobalt oxides contain a different number of oxygen atoms than nickel and

copper oxides, with more oxygen atoms surrounding cobalt. Oxygen rich species with a

stoichiometry of M3O6- and M4O6,7

- were only detected in the cobalt mass distribution

(Figure 8-1b). It should also be noted that no anionic cobalt or nickel oxides with the

same number of metal and oxygen atoms were reactive in the presence of CO. Only

copper oxides, Cu3O3- and Cu4O4

-, were active towards oxygen atom transfer followed by

O2 loss. In contrast, studies by Huang et al. showed that non-stoichiometric copper oxide

species (ie. Cu3O4- and Cu4O5

-) were more active toward CO oxidation because these

oxygen rich species were able to transport surface lattice oxygen more efficiently (40).

Indeed in oxygen deficient environments one can see the usefulness of oxygen rich

species in order to supply oxygen to the reaction. However for copper oxides, in terms of

active sites, there is a special reactivity associated with metal and oxygen equivalent

species in the present studies.

Cations with the same number of metal atoms and oxygen atoms readily undergo

reactions with CO compared to the anions. From Table 8-4, the results of one oxygen

atom on nickel or cobalt show oxygen atom transfer. Ni2O2+ and Co2O2

+ both show MO2

(M=Ni, Co) replacement by CO followed by M (M=Ni, Co) replacement by a CO

molecule. These clusters are the only species to contain two CO molecules as the major

reaction channel and it is interesting that these species readily lose a metal atom. This

suggests that the CO binding energy is greater than the energy required to break a M-M

154

or M-O bond. For small clusters the energy released in binding CO produces competitive

reaction pathways between unimolecular decomposition and collisional stabilization. In

this case, the charge donation from CO binding causes destabilization in the metal (24)

resulting in weaker M-M bonds then M-C bonds (18). Cox et al. have shown that cobalt

and nickel have the strongest M-C bond strengths, followed by iron and then copper (12).

Size dependence is another important aspect in studying the reactivity of small

transition metal oxide clusters. For the monomer and dimer clusters, all the metal anions

exhibit oxygen atom transfer. However, the major reaction pathways begin to differ

starting with the trimer depending on the metal composition. Tetramer oxide clusters,

Co4O6- and Co4O7

-, do not readily promote the oxidation of CO through the oxygen atom

transfer pathway. In fact, only Ni4O5-, Cu4O4

- and Cu4O5- promote the oxidation of CO.

Size may also play a role in the cationic adsorption mechanism of oxygen. The amount

of oxygen adsorbed onto the metal clusters differs with size. For example, the cobalt

monomer oxide contained dissociated oxygen; however, the dimer species absorbed an

even number of oxygen atoms. This provides evidence that oxygen adsorbs molecularly

to cobalt dimers and may explain the dominant oxygen molecule replacement by CO

channel in these clusters.

Periodic Trends. Previous studies of metal oxide clusters reacted with CO allow

us to make other periodic comparisons (43, 64). All the metals discussed in the present

study are heavier 3d transition metals, which show electron delocalization only for the 4s

electrons (65). This means that the lightest elements, like titanium have 3d electrons that

contribute to delocalization and participate in valence orbitals and bonding. Although not

directly involved in the bonding, metal d orbital contributions strengthen bonds. This

155

effect becomes less critical moving from left to right on the periodic table because the

radial extent of the d orbitals decreases relative to the s orbital (35). In fact, other studies

have shown that the metal d contribution to bonding is higher for CO adsorption on Ni

than on Cu surfaces (11). Iron oxides belonging to the first row transition metals exhibit

a behavior depending on the metal to oxygen ratio similar to nickel and cobalt for the

most reactive stoichiometries toward CO oxidation (51). These metals have open d shells

and may allow more coordination of oxygen resulting in the observed active

stoichiometries. However, copper oxides displayed an equivalent number of metal and

oxygen atoms for the most active species in oxidation. The reason copper may differ

from the other metal oxide is that copper possesses a full d orbital shell and bonding

occurs primarily through s interactions. Copper oxides may become more stable with

less oxygen coordinated.

Differences in the electronic structure and bonding energies of small anionic and

cationic metal oxides are indeed evident from our reactivity studies. The significance of

additional electron density in anionic clusters is also exhibited in the present studies.

Closed shell electron configurations produce more stable clusters, and studies show that

the lone pair electrons on CO readily donate charge to clusters needing additional

electrons to form a closed shell (16). Cations were shown to have more CO binding in

order to gain electron density and better cluster stability.

The nickel group metals include some of the most reactive metals in catalysis

besides the commonly employed platinum and palladium nanoparticles. Studies

involving Pt, Pd, and Ni with CO were conducted and reveal fundamental differences in

the electronic structure of the different metals (24). Compared to Pt and Pd, nickel

156

showed the highest uptake of CO. Also Ni-CO bonds tend to be stronger with more

contributions from π backbonding (13,17). However, studies have not been conducted in

our laboratory on other metal oxides in this group with CO and therefore it is hard make

comparisons based on the reactivity of these metal oxides with CO.

The coinage metals have been studied for their reactivity towards CO. Sueyoshi

and coworker found that the low activation energy for the reaction between adsorbed CO

and oxygen toward the oxidation of CO on Cu(110) surfaces makes this surface more

active than standard Pd, Pt, or Rh catalytic materials (31). They extended their low

temperature studies to include Ag and Au surfaces and suggested a trend for low

activation energy on these surfaces that allows for more energetically favorable CO

oxidation than currently used metal surfaces. Our studies on copper, silver, and gold

oxides show mostly anionic species promote the oxidation of CO. It is interesting that no

anionic trimer gold oxides (64) or dimer silver oxides show oxygen atom transfer

products, even though they are in the same group of the periodic table. These reactions

are highly dependent on the structure of the metal oxide as pointed out in previous studies

(64). However, copper seems to be the most active for all sizes studied and we provide

further evidence that charge transfer between the negative metal center and oxygen

enhances the oxidation of CO.

8.4 Comparison to Bulk

Low temperature oxidation of carbon monoxide has been carried out on mixed

transition metals (66) and on transition metal oxide supports (67, 68). Also, X-ray

157

photoelectron spectroscopy has been used to study the specific composition of these

species. Studies involving cobalt oxide catalytic supports showed these materials to be

composed of Co3O4 and CoO species in the near-surface region (68). Isotope studies of

oxygen found that CO2 was formed over preoxidized cobalt oxide surfaces and that

surface oxygen participated in the reaction (41). There have been numerous other studies

that now show that surface oxygen from metal oxide supports is responsible for

promoting the oxidation of CO to CO2. It has also been shown previously that diffusion

of lattice oxygen enhances the activity of CO oxidation. Thus, the ability of transition

metal oxides to adsorb and release oxygen greatly affects the activity. Oxygen rich

anionic clusters with at least one more oxygen atom than metal atom are shown to

promote the oxidation of CO in the present experiments. This supports the fact that

transition metal oxides composed of iron, cobalt, nickel, or copper do promote CO

oxidation. With further studies on mixed metal systems, the full potential of these

reactive species may be realized.

8.5 Conclusion

The dissociation patterns and reactivity in the presence of CO of several first row

transition metal oxides were investigated as a function of size, stoichiometry, and ionic

charge state. Numerous factors affected the reaction selectivity including oxygen

stoichiometry and metal composition. This comprehensive study of the electron density

changes in charge state moving across the periodic table has revealed similarities and

differences in the observed reaction products. Small anionic metal oxides are more

158

selective toward CO oxidation. However cationic metal oxides are more efficient in their

observed reaction pathways. This study has shed light on the active species for CO

oxidation and provided evidence that transition metal oxides, commonly used in catalytic

supports, may contribute to the overall reactivity. The ability of transition metal oxide

clusters to transfer oxygen to CO forming CO2 potentially mimics the mechanism that

occurs on condensed phase catalyst supports. Further research into the nature of mixed

metal cluster systems would be beneficial to understand the interactions of supported

catalysts at the molecular level.

159

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Chapter 9

General Conclusions

The field of heterogeneous catalysis is ever-evolving with new technologies and

better methods for energy-efficient conversion of fossil fuels. One major concern that

arises from the incomplete combustion of fuel is the environmental pollutant, carbon

monoxide. With increasingly strict standards for automotive emissions and the push for

cleaner burning fuels, improved methods for converting fuel are needed. The studies in

this thesis were aimed at providing a fundamental picture of the oxidation of CO and the

mechanisms that govern the conversion of CO to CO2 in the presence of alternative

metals. Current catalysts employed in industry to reduce pollution consist of costly

platinum or palladium metals which require high working temperatures. Metal catalysts

capable of low temperature conversion are ideal in order to avoid unnecessary energy

consumption for non-viable products. With a better understanding of the mechanisms

governing potentially active alternative metals, the ability to tailor design catalysts to

effect specific oxidation reactions could be realized. There were two main alternative

catalytic materials investigated in this thesis, noble metal oxides in Chapters 3 and 4, and

transition metal oxides in Chapters 5-8.

The combined studies examined in this thesis reveal new information on the

mechanisms of different catalytic compositions toward the oxidation of CO. Also, these

studies provide corroborative evidence of recent findings with a proposed mechanism to

explain observed products, such as the existence of intermediate species with dissociated

164

oxygen (1, 2, 3). These initial studies uncover information based mostly on individual

active sites in catalysis, gleaned from studies based on size and stoichiometry. For

instance transition metal oxides, often employed in catalytic supports, were found to

promote the direct oxidation of CO. This is evidence that support materials actively

participate in the catalytic conversion of CO to CO2, an issue which has stirred much

debate in the literature (4).

9.1 Major Findings

The major theme throughout these studies was the investigation of the ionic

charge state and its effect on the electron density and charge transfer to different metal

atoms at the nano-scale. It was shown that especially in the catalytic oxidation of CO in

noble metal systems that negative charge was more effective in promoting oxidation. It

has been suggested in the literature that negatively charged metal centers donate charge

to O2 thereby elongating the O-O bond and further activating the dissociation (5).

It is interesting that calculations by our collaborators at the Humboldt University

and VCU predicted that positively charged metal centers also would promote the

oxidation of CO. The reason was that metal-oxygen bonds were found to be weaker in

cationic systems compared to anionic systems due to charge transfer effects. Indeed, a

kinetic analysis of the rate constants for the direct oxidation of CO with Fe2O3+ and

Fe2O3- showed the cationic rate constant was 8.5x10-13 ± 0.8 cm3s-1 compared to the

anionic rate contant of 2.9x10-13 ± 0.2 cm3s-1. However, cationic gold oxide clusters of

similar stoichiomerty to anionic gold oxide clusters, found to undergo oxygen atom

165

transfer previously, did not undergo CO oxidation. This is an indication of the different

mechanisms occurring on gold centers compared to transition metal oxide centers which

promoted CO oxidation.

For transition metal oxide systems, oxygen atoms are bound directly to the metal

atom, which enables direct oxidation with the use of oxygen atoms coming from the

metal. In contrast, the Langmuir-Hinshelwood type mechanism is the operative

mechanism for most catalysts (6). A reaction proceeding through this mechanism would

adsorb CO and react with adsorbed oxygen, not oxygen that is already bound to the metal

catalyst. Thus, oxygen dissociation is a barrier needed to be overcome for the reaction to

proceed. This step is eliminated in transition metals that are not saturated with oxygen

and do not contain O-O bonds. Therefore, transition metals tend to be reduced as CO is

oxidized and the cycle continues based on a redox cycle mechanism. The nature of

transition metals is to be stable in several oxidation states, therefore making them

particularly suitable for direct oxidation applications and catalytic supports.

The stoichiometry and oxygen to metal ratio was also an important factor in the

observed reaction pathways. It is not surprising that our findings show oxygen rich

transition metal oxide clusters were more active toward CO oxidation since they employ

their own oxygen atoms in the reaction. Interestingly, AuO- and Au2O-, both oxidize CO

whereas no negatively charged transition metal oxides were even formed with the same

number of oxygen and metal atoms. In comparison, positively charged transition metal

oxides were formed and underwent CO oxidation for clusters containing the same

number of metal and oxygen atoms. This demonstrates the formation mechanisms of

metal oxides and the chemistry of oxygen bonding to different charged metal species.

166

The studies in this thesis were based on dissociated as well as molecularly adsorbed

oxygen since oxygen was added directly in the source region where cluster formation

occurs. This allowed the investigation of potential barriers to CO oxidation after O-O

bond dissociation. Indeed, recent studies conducted in our laboratory found that

dissociated oxygen is necessary but not sufficient for the reaction to proceed on anionic

gold oxide clusters (7). The reaction rates for CO oxidation in the presence of Au2O3-

and Au2O4-, which both contained dissociated oxygen, were on the order of 10-13 cm3s-1.

Size was found to be very important in the overall reaction pathways observed. In

the nano-scale regime, there is much controversy over the smallest active size of metal

nanoparticles that are able to oxidize CO (8). Our studies reveal only larger cationic gold

oxides (clusters with at least four gold atoms) undergo CO oxidation. Whereas, previous

investigations show that only select monomer and dimer anionic gold oxide clusters were

effective in promoting the oxidation of CO (7). Also, cationic gold oxides and larger

anionic gold oxides had other reaction products that decreased the overall efficiency and

selectivity of the oxidation reaction channel.

Although CO oxidation is the most viable reaction channel to understand for use

in industrial catalysts, it is important to identify the mechanisms of other reaction

channels observed. Oxygen replacement by CO was a common reaction channel observed

in cationic noble metal and transition metal oxides. The positively charge metal centers

were more efficient in this reaction channel than CO oxidation was in anionic metal

centers due to weaker oxygen bonding. In a kinetic analysis of the rate constants, the

replacement rate constant for Au2O+ was an order of magnitude greater, 10-12 cm3s-1, then

the oxidation rate constant for Au2O-, 10-13 cm3s-1. It is possible that the replacement

167

reaction pathway may have prevented the slower CO oxidation pathway from occurring

or being the major reaction channel. Also, the strength of M-CO bonds formed between

cationic metal clusters showed that CO was bound strongly and desorption of CO2 was

limited. Therefore, potential intermediates in the oxidation of CO were too stable to react

further.

Thus many factors contribute to the overall reactivity of metal oxides and it is

important to study the role of ionic charge, size, stoichiometry, and composition in order

to understand the governing active sites and mechanisms in catalysis. From the studies

presented in the previous chapters of this thesis, there remain several questions to be

investigated as a result of the findings.

9.2 Future Studies

Further research on bimetallic systems would be beneficial to providing a

complete picture of the operative mechanisms of the oxidation of CO. The production of

larger clusters with different stoichiometries of silver and gold atoms would reveal

interesting information on trends and active sites for CO oxidation. Although the dual

rod source described in Chapter 2 produced mixed cluster species of Ag-Au oxides,

increased signal intensity is needed to further this research. Method development on a

new source may be the most promising for larger bimetallic clusters. Different mixed

metal systems would also be of interest as electron transfer differs in different metal

centers shown in condensed phase studies of increased reactivity for only certain metal

oxide supports.

168

It would be interesting to study the transition metal systems with oxygen added

downstream from the source. The 5-component guided ion beam mass spectrometer

could be employed to study this reaction by adding another reaction cell octopole and

third quadrupole to the chamber. Mass selection would occur in the 1st quadrupole, and

CO could then be added to the first octopole. Investigations with molecular oxygen added

into the first octopole could determine how oxygen is adsorbed to the ionic metal centers.

Next CO added to the second octopole would provide conditions close to condensed

phase adsorption of both oxygen and CO. Product species detected from these studies

would be very interesting to compare to those predicted by the Langmuir-Hinshelwood

catalytic mechanism especially for transition metal oxide clusters. The transition metal

oxide clusters show promise in becoming more active alternatives to costly precious

metal catalysts and more insight into their reaction mechanisms would be beneficial.

169

References

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2. Socaciu-Siebert, L. D.; Hagen, J.; Le Roux, J.; Popolan, D.; Vaida, M.; Vajda, Š.;

Bernhardt, T. M.; Wöste, L. Phys. Chem. Chem. Phys. 2005, 7, 2706.

3. Henao, J. D.; Caputo, T.; Yang, J. H.; Kung, M. C; Kung, H. H. J. Phys. Chem. B 2006, 110, 8689.

4. Negishi, Y.; Nakamura, Y.; Nakajima, A.; Kaya, K. J. Chem. Phys. 2001, 115,

3657.; Arrii, A.; Morfin, F.; Renouprez, A. J.; Rousset, J. L. J. Am. Chem. Soc. 2004, 126, 1199.; Comotti, M.; Li, W. C.; Spliethoff, B.; Schüth, F. J. Am. Chem. Soc. 2006, 128, 917.

5. Böhme, D. K.; Schwarz, H. Angew. Chem. Int. Ed. 2005, 44, 2336.; Yoon, B.;

Häkkinen, H.; Landman, U. J. Phys. Chem. A 2003, 107, 4066.; Stolcic, D.; Fischer, M.; Ganteför, G.; Kim, Y. D.; Sun, Q.; Jena, P. J. Am. Chem. Soc. 2003, 125, 2848.; Chen, M. S.; Goodman, D. W. Science 2004, 306, 252.

6. Baxter, R. J.; Hu, P. J. Chem. Phys. 2002, 116, 4379.



8. Lee, S.; Fan, C.; Wu, T.; Anderson, S. L. J. Am. Chem. Soc. 2004, 126, 5682.; Sanchez, A.; Abbet, S.; Heiz, U.; Schneider, W.-D.; Häkkinen, H.; Barnett, R. N.; Landman, U. J. Phys. Chem. A 1999, 103, 9573.; Wallace, W. T.; Whetten, R. L. J. Am. Chem. Soc. 2002, 124, 7499.

Appendix

Kinetic Analysis of the Reaction between (V2O5)1,2+ and Ethylene

A.1 Abstract

A systematic experimental and theoretical investigation of the influence of

reactant energy on the reactivity of (V2O5)n=1,2+ clusters with ethylene provided evidence

of the rate controlling steps in the reaction (1). Herein, we present further experimental

and theoretical evidence for our recently proposed radical cation mechanism for oxygen

atom transfer from (V2O5)n=1,2+ clusters to ethylene. In particular the results of ab initio

molecular dynamics simulations are found to further support the radical cation

mechanism. Experimental reaction cross sections at the zero pressure limit and rate

coefficients show that the energy dependence of the reaction cross section is in accord

with the Langevin formula. Evidence is presented that ion-molecule association is the

rate determining step, whereas subsequent hydrogen transfer and formation of

acetaldehyde proceed without significant barriers. We propose a kinetic model for the

reaction cross section which fully accounts for the experimental findings. The model

offers the prospect of elucidating the details of the general reaction mechanisms through

a study of the energy dependence of the reaction cross sections.

(Reproduced in part with permission from J. Phys. Chem. B 2006, 110, 3015-3022. Copyright 2006 American Chemical Society.)

171

A.2 Introduction

There is growing interest in the use of metal oxide clusters as model systems for

gaining insight into the mechanisms of various reactions of importance in the field of

catalysis (2-4). Gas phase cluster research is proving to be very valuable in determining

the properties and behavior of transition metal oxides, yielding information on

thermochemistry (5), structure (6-13), and reactivity (14-16). Recently, we have

undertaken a comprehensive joint theoretical and experimental effort to reveal the

reactive behavior of vanadium oxides of various sizes and stoichiometries with a variety

of small organic molecules. Our findings show that among vanadium oxides of widely

differing stoichiometries and sizes, only V2O5+ and V4O10

+ effect oxygen transfer

reactions with ethylene (1, 17-18). In particular, these results have prompted us to devote

considerable attention to a study of these species. Prior evidence from experimental

investigations obtained in our laboratory (1) and theoretical findings revealed evidence of

a facile oxygen transfer mechanism in direct analogy to ones believed to occur on

vanadia surfaces at reaction sites of similar stoichiometry (15, 18). Our earlier theoretical

work, in combination with the experimental studies, proposed the involvement of an

oxygen centered radical reaction, proceeding subsequent to formation of a metal oxide

cluster-molecule complex (1).

Specifically, we have obtained evidence of a general radical cation mechanism for

the oxidation of ethylene on cationic vanadium oxide clusters with the composition

(V2O5)n=1,2+, leading to the formation of acetaldehyde as a neutral product according to

the reactions (1):

172

V2O5+ + C2H4 → V2O4

+ + C2H4O

V4O10+ + C2H4 → V4O9

+ + C2H4O

(I)

(II)

The key point is a radical center located at a peripheral oxygen atom on the V2O5+

and V4O10+ clusters. The reaction is considered to proceed through a generic reaction

profile based on energetic data obtained from ab initio calculations as shown in Figure 1-

1. The initial reaction step is believed to proceed via the formation of a complex arising

from an ion-molecule collision of A and B (Figure 1-1). The structural transformation

along the reaction pathway indicates that a single bond is ultimately formed between the

first carbon atom of the ethylene and the oxygen atom (complex C). This leads to the

shift in the radical center, initially located at the oxygen atom, to the second carbon atom

of the ethylene molecule. In the second reaction step, a hydrogen atom is transferred

from the first to the second carbon atom, forming a complex in which the acetaldehyde

molecule is bound to the vanadium atom (complex D). Finally, in the third step, the V-O

bond is broken and an acetaldehyde molecule is released, resulting in an overall oxygen

transfer process between the two reactants, leading to (E+F).

To provide a more complete understanding of the factors influencing the reaction

dynamics of this observed oxygen atom transfer reaction, we have undertaken further

experimental and theoretical studies of the energetics of the reaction. The findings,

which are reported herein, demonstrate consistency with a proposed mechanism of this

reaction that is of considerable interest in the area of fundamental catalysis, as well as in

the industrial production of acetaldehyde (19).

173

Figure 1-1: General scheme of the energetic profile and structural transformation of the mechanism for the reaction between (V2O5)n=1,2

+ and C2H4. Source: Figure courtesy of the group of Professor V. Bonačić-Koutecký at the Humboldt Unisversity in Berlin.

A+B

C

D

E+F

A+B

C

D

E+F

174

A.3 Experimental and Computational Methods

The experiments performed for this study were conducted using a guided ion

beam mass spectrometer coupled to a laser vaporization source, described in detail

previously (20) and briefly explained here. Metal oxide species were generated using a

laser vaporization source, whereby a vanadium rod was ablated using the second

harmonic of a Nd:YAG laser while rotating and translating to continually expose a fresh

surface. At a predetermined time, a small amount of oxygen seeded in helium, ~8%, was

pulsed over the surface of the rod, creating a plasma in which the vanadium oxide

clusters were formed. The clusters were then cooled in a region of supersonic expansion

and guided into the first quadrupole through a skimmer and a set of electrostatic lenses.

The ion of interest was then mass selected and focused through a second set of

electrostatic lenses into a reaction cell incorporated within an octopole guide. The

effective path length of the octopole reaction cell was calculated to be 12.9 cm using the

trapezoidal pressure fall-off approximation (20). The pressure of the ethylene in the

reaction cell was kept constant while the energy of the octopole was increased from 0 to

20 volts (laboratory frame) in single volt increments. The lab-frame energies (Elab) were

converted to center-of-mass frame energy (ECM), ECM=Elab[M/(M+m)], in order to enable

a comparison of calculated and experimental values. Here, M is the mass of neutral

target gas and m is the mass of the selected ion. The described procedure was repeated

for pressures ranging from 0.1 to 0.6 mTorr of ethylene, chosen to effect single as well as

multiple collision reaction conditions. The products generated in the octopole reaction

175

cell were focused into the second quadrupole through a third set of electrostatic lenses,

mass analyzed, and detected using a channel electron multiplier.

The calculations performed by the Bonačić-Koutecký group and presented here

were carried out at the density functional level of theory using the Becke’s hybrid three

parameter non-local exchange functional combined with the Lee-Yang-Parr gradient

corrected correlation functional (B3LYP) (21) and employing the all-electron triple-ζ

valence plus polarization basis set (TZVP) developed by Ahlrichs and co-workers (22).

As shown in our previous work, this method allows for the accurate description of

structural properties of cationic vanadium oxide clusters and their interaction with

ethylene (1). The reaction pathways were studied employing ab initio molecular

dynamics (MD) with forces calculated by using density functional theory (23). For the

solution of Newton’s equations of motion, the Verlet algorithm was used. The MD

simulations were initiated from the activated stable complexes of (V2O5)n=1,2+ with

ethylene by randomly distributing the energy among all degrees of freedom. This

allowed us to verify reaction steps deduced from the calculation of the stationary points

and obtain support for the proposed mechanism.

A.4 Results and Discussion

Ab initio MD simulations at constant energy, corresponding to the stability of the

initial complex with respect to the reactants, were carried out to gain further evidence for

the operative reaction mechanism. This corresponds to the limiting case of a single

176

collision condition in the experiment (1) in which the energy of the initial complex is not

released by collisions with gas molecules.

The snapshots of the MD trajectory for the reaction between V2O5+ and C2H4 are

shown in Figure 1-2. The formation of complex C is a barrierless process. The

simulation beginning with this complex shows that the association of ethylene to V2O5+ is

followed by a very rapid (20 fs) hydrogen transfer from the carbon atom bound to oxygen

toward the terminal carbon atom. After the hydrogen transfer has taken place (complex

D), the V-O bond is broken within the next 40 fs, leading to acetaldehyde and V2O4+ as

the final products (E and F noted in Figure 1-1). Analogous snapshots for the reaction

between V4O10+ and C2H4 are shown in Figure 1-2. Similar to the case of V2O5

+, the

reaction with V4O10+ also involves very rapid hydrogen transfer (~30 fs) followed by the

formation of acetaldehyde which occurs typically within 100 fs.

In addition to ab initio MD simulations, experiments were performed by

systematically raising the energy of the reaction under conditions of constant selected

pressure of ethylene within the reaction cell. Figures 1-3 and 1-4 show the branching

ratios of the reactions of V2O5+ and V4O10

+ with ethylene, respectively. The relative

intensities of the ions are plotted as a function of the energy added to the octopole

reaction cell. Minor molecular oxygen loss products, V2O3+ and V4O8

+, are observed and

are insensitive to the energy addition. The decrease in reactivity with increasing energy

is evident from the data in Figure 1-3, which show that there is comparatively less

conversion of V2O5+ reactant ion to the major product, V2O4

+, at higher collision energies.

A similar situation arises for V4O10+ and V4O9

+ as seen in Figure 1-4.

177

Figure 1-2: Snapshots of the ab initio MD trajectory calculated using DFT methods for the reaction of a) V2O5

+ with C2H4 and b) V4O10+ with C2H4 leading to the formation of

acetaldehyde. Source: Figure courtesy of the group of Professor V. Bonačić-Koutecký at the Humboldt Unisversity in Berlin.

178

Figure 1-3: Branching ratios for the energetic analysis of V2O5+ with ethylene conducted

under the following pressures: a) 0.1 mTorr, b) 0.2 mTorr, c) 0.3 mTorr, d) 0.4 mTorr,e) 0.5 mTorr and f) 0.6 mTorr of ethylene. These are defined by the ratio of each product ion divided by the total ion signal. Notice that the minor contribution of V2O3

+ is essentially independent of the center-of-mass reaction energy.

179


conducted under the following pressures: a) 0.1 mTorr, b) 0.2 mTorr, c) 0.3 mTorr, d) 0.4 mTorr, e) 0.5 mTorr and f) 0.6 mTorr of ethylene.

180

In both cases, as the pressure of the ethylene is increased from 0.1 to 0.6 mTorr,

the energy under which equal amounts of the parent and the oxygen transfer product ions

are attained is increased. This mimics a negative temperature dependence and is not the

behavior expected in the case of collision induced dissociation reactions. In fact,

collision induced dissociation experiments have shown that the observed oxygen

deficient products, V2O4+ and V4O9

+, were not the result of a collisional process because

they were not observed in separate experiments using an inert gas at near thermal

energies (6). Accordingly, the pressure corresponding to the crossing point of the

branching ratios can be calculated as follows:

),)(

exp(0 Tk

pEII

B

CMp

lσ−= (A-1)

where I0 is the sum of the reactant and all product ion intensities and Ip is the final

intensity of the product ion, in this case VxOy-1+. The total reaction cross section is

σ(ECM) and kB is the Boltzmann constant. The effective path length of the collision cell is

ℓ, while T and p are the temperature and the pressure of the reactant gas, respectively.

Cross Section Anaylsis. The experimental cross sections of both V2O5+ and

V4O10+ clusters with ethylene are calculated according to eq A-1. In order to account for

the pressure dependence of the cross section measurements, data from several pressures

are plotted and then extrapolated to zero pressure in the customary way (24). At the low

pressure limit, eq A-1 becomes

181

Tk

pI

I

E

B

p

CM l0)( =σ

(A-2)

yielding the cross section from the slope of the line in Figure 1-5 for the intensity ratio

Ip/I0 as a function of pressure. These plots are then repeated for every energy input into

the reaction cell.

For V4O10+, charge transfer and associated ion reaction products, C2H4

+ and

C3H6+, can be seen in the branching ratios of Figure 1-4. These products are accounted

for in the cross section calculations through the sum of product intensities via the analysis

using eq A-2. The secondary reaction products, which are a consequence of multiple

collisions, become significant only at 0.6 mTorr. The experimental cross sections are

presented and discussed below with the kinetic model.

Reaction Rates. The reaction rate coefficients, which are dependent on the

lifetime of the ion-molecule complex, were also calculated for both oxygen transfer

reactions. The experimental cross section data are converted to an expression for the

phenomenological rate coefficient (24) by eq A-3

)()( 00 CMEvvk σ= (A-3)

where the relative velocity (ν0) of the reactant ion is

2/1

0

2⎟⎟⎠

⎞⎜⎜⎝

⎛=μ

CMEv

(A-4)

182

It should be emphasized that this equation is valid only in the case that the energy

distribution of the reactant ions is very narrow (monoenergetic ions). The energy input

into the reaction cell (ECM) is in the center-of-mass frame, and μ is the reduced mass.

Figure 1-6 shows the phenomenological rate constants calculated from zero

pressure cross section data and therefore does not account for multiple collision

conditions in our higher pressure experiments. Another approach to calculate the rate is

via the usual method given by eq A-5

[ ][ ],BAkdt

dA −= (A-5)

and employed in analyzing data acquired in a flow tube reactor operated at low pressures

and low electric drift fields. Use of eq A-5 requires knowledge of the interaction time in

the reaction cell. This time is determined by means of calculation in the present case.

The slope of Figure 1-7, which is a plot of ln[Ip/I0] versus the concentration of the

reactant ion, is equal to –kt. A series of graphs, similar to Figure 1-7, are plotted for each

energy value.

The time is found by estimating the sum of the velocity due to the supersonic

nozzle expansion (25) and the velocity acceleration from the voltage field. The

supersonic expansion velocity is given by eq A-6

2/1

2

11 ⎥

⎥⎥⎥

⎦

⎤

⎢⎢⎢⎢

⎣

⎡

⎟⎠⎞

⎜⎝⎛ −+

=T

Ts

Mm

RTMv

γγ

(A-6)

where MT is the mach number, m is the mass of the ion, and γ is the specific heat capacity

ratio.

183

Figure 1-5: a) Ratio of product ion intensity, Ip, to total ion intensity, I0, versus the reaction cell gas pressure for reactions of V2O5

+ with ethylene at ECM= 0 eV and b) at ECM= 2.0 eV.

Figure 1-6: a) The rate constant versus energy for V2O5

+ with ethylene calculated from cross section data. b) The rate constant versus energy for V4O10

+ with ethylene calculated from cross section data. The error bars represent the relative uncertainty in the cross section measurements.

y = 1.77x

R 2 = 0.95 0

0.2 0.4 0.6 0.8 1.0 1.2

0 0.2 0.4 0.6 0.8 Pressure (mTorr)

a) y = 0.70x

R 2 = 0.93 0

0.1 0.2 0.3 0.4 0.5 0.6

0 0.2 0.4 0.6 0.8

b)

I p/I

0

Pressure (mTorr)

2.0x10-10

4.0x10-10

6.0x10-10

8.0x10-10

1.0x10-9

1.2x10-9

0.0 1.0 2.0 3.0 EnergyCM (eV)

2.0x10-10

4.0x10-10

6.0x10-10

8.0x10-10

1.0x10-9

1.2x10-9

0.0 0.3 0.6 0.9 1.2 1.5 EnergyCM (eV)

Rat

e C

onst

ant

(cm

3 s-1

) a) b)

184

In this case, the heat capacity was calculated for atoms because the carrier gas was mostly

comprised of helium being pulsed through the nozzle. The velocity due to the voltage

applied to the octopole in the reactant gas cell is deduced from eq A-7

,2

m

qEv LAB=

(A-7)

where the charge of the ion is q, and ELAB is the electric field voltage applied to the

reaction cell in laboratory frame. The electrostatic lens immediately before the reaction

cell is set at 0 volts, creating a voltage field difference when energy is added to the

octopole. The effective path length of the reaction cell divided by the sum of the

velocities provides reaction times ranging from approximately 0.3x10-4 to 1x10-4 seconds

for V2O5+, over the energy range studied, and 0.4x10-4 to 1.4x10-4 seconds for V4O10

+.

Compared to the effect of the electric field, the contribution due to supersonic expansion

is minimal. As the energy in the reaction cell is increased the time decreases because the

ions are moving faster through the cell. These times are within the experimental time

frame used in collision induced dissociation experiments and are due to the short distance

between the reaction collision and the quadrupole entrance (20). Finally, the negative

slope of Figure 1-7 divided by the time corresponding to the energy input yields the rate

coefficient.

Figure 1-8 shows the rate coefficients plotted versus energy. These values were

calculated from the experimental data, which include effects due to the pressure and

concomitant multiple collision conditions. It can be seen from Figure 1-8 that the rate of

V2O5+ oxygen transfer is on the order of high-10-10 cm3s-1 and is slightly higher than the

rates for V4O10+ oxygen transfer which are in the mid-10-10 cm3s-1 range.

185

Figure 1-8: a) The rate constant versus center-of-mass energy for the reaction of V2O5+

with ethylene calculated from the velocity of the ions and reaction time and b) The rate constant versus center-of-mass energy for the reaction of V4O10

+ with ethylene calculated from the velocity of the ions and reaction time. The error bars are the relative uncertainties of measurements considered in the calculation.

Figure 1-7: a) Logarithm of the intensity ratio [V2O5

+/V2O5++ΣIp] versus the

concentration of C2H4 calculated as an ideal gas at ECM=0 eV and b) at ECM= 2 eV.

Rat

e C

onst

ant

(cm

3 s-1

)

5.0x10-11

3.5x10-10

6.5x10-10

9.5x10-10

1.3x10-9

0.0 1.0 2.0 3.0 EnergyCM (eV)

1.0x10-11

3.1x10-10

6.1x10-10

9.1x10-10

0.0 0.3 0.6 0.9 1.2 1.5 EnergyCM (eV)

a) b)

y = - 6E - 14x - 0.11 R 2

- 1.4 - 1.2

- 1 - 0.8 - 0.6 - 0.4 - 0.2

0 0 1E+13 2E+13

[C 2 H 4 ] mc/cc

Ln r /I ]

y = - 2.8E - 14x + 0.02 R 2 = 0.94

- 0.6 - 0.5 - 0.4 - 0.3 - 0.2 - 0.1

0 0 1E+13 2E+13 3E+13

[C 2 H 4 ] mc/cc -1.4 -1.2 -1.0 -0.8 -0.6 -0.4 -0.2

0 1x1013 2x1013 3x1013

Conc. [C2H4] mc/cc -0.6 -0.5 -0.4 -0.3 -0.2 -0.1

0 0 1x1013 2x1013 3x1013

Conc. [C2H4] mc/cc

0

R2=0.92 R2=0.94

a) b)

y=-6x10-14x-0.11 y=-2.8x10-14x+0.02

ln[I

r/I0]

186

The rate constant for V2O5+ at lower energies first increases, and then slowly

decreases at values above ~1 eV center-of-mass energy. This decrease in the rate

constant is what is expected for an exothermic reaction with negative temperature

dependence and no potential barriers. The trend in V4O10+ is similar with the rate

constant increasing and then leveling off at higher energies. It is interesting to note that

there is reasonably good agreement in the rate constants derived by the two methods, and

presented in Figures 1-6 and 1-8, respectively. The data in Figure 1-6 have been obtained

by analyzing data based on single collision conditions to acquire cross sections, and then

treated on the basis of a narrow velocity distribution to acquire rate constants.

Conversely, the data in Figure 1-8 are acquired through calculations that describe our

experimental conditions of higher pressures in a manner typically employed in flow tube

kinetics. Although branching ratios are taken into account, the latter method involves an

average rate constant under multiple collision conditions.

Kinetic Model. A kinetic model for the reaction cross section for the oxygen

transfer rate follows from eq A-1. Combining with eq A-3, and as kt→0 for short times

and low pressure, it follows that

tTk

pk

I

I

B

R )(1 0

0

ν−= (A-8)

where IR is the intensity of the reactant ion and t is time. This is rearranged in terms of

the rate constant and gives eq A-9

tTk

pI

I

k

B

p

00 )( =ν .

(A-9)

187

Employing eq A-3 and the relationship, l=t0ν , along with the crossing points, Ip/I0 =

0.5, shown in Figures 1-3 and 1-4, it follows that

l)(

2ln2/1

CM

B

E

Tkp

σ=

(A-10)

and the crossing point for a given energy is a direct measure of the reaction cross section.

If we assume that the initial step involves formation of an ion-molecule association

between ethylene and the mass selected reactant oxide ion, understanding the reaction

cross section requires consideration of the kinetics of several potentially operative

reaction steps in the overall conversion mechanism.

In order to derive the expression for the reaction cross section as a function of

energy we write the reaction symbolically in accordance with the mechanism depicted in

Figure 1-1 as

FEDCBAkk

k

k

k+→⇔⇔+

53

4

1

2

. (III)

By using the steady state approximation for the intermediates C and D, shown in Figure

1-1, the change in C and D with respect to time are obtained by the following equations

54

3 ][][

kk

CkD

+=

(A-11)

and

54

4332

1 ]][[][

kk

kkkk

BAkC

+−+

= (A-12)

The individual rate constants shown in the equations are also defined in Figure 1-1. It

readily follows that the total reaction rate d[E]/dt, is given by eq A-13,

188

== ][][

5 Dkdt

Ed]][[

))(( 435432

531 BAkkkkkk

kkk

−++

and the overall rate constant by eq 14

(A-13)

435432

531

))(( kkkkkk

kkkk

−++=

(A-14)

The first step of the reaction can be modeled after a long-range ion-molecule association,

and the reaction cross section can be calculated from the Langevin cross section, σL,

2/1

0

2

4 ⎟⎟⎠

⎞⎜⎜⎝

⎛=

CML E

e αε

σ (A-15)

where the collisional energy of the reactant gas is ECM, α is the polarizability of ethylene,

and ε0 is the vacuum permittivity. In principle, the rate constants k and k1 can be

obtained from the energy dependent reaction cross sections by averaging over the

velocity distribution. However, because the energy distribution of the ions under the

considered experimental conditions is narrow, it can be assumed that the velocity

distribution is also narrow. In this limiting case, as an approximation, the rate constants k

and k1 are simply products of the velocity of the ions and the corresponding cross

sections. This yields a total reaction cross section as

435432

53

))(()(

kkkkkk

kkE LCM −++

= σσ (A-16)

The rate constants for the elementary reaction steps can be approximated using the RRK

expression, yielding the following four equations for k2, k3, k4 and k5.

1

122 )( −

Δ+= N

EE

EAk

(A-17)

189

1

1

#1

33 )( −

Δ+−Δ+= N

EE

EEEAk

(A-18)

1

2

#1

44 )( −

Δ+−Δ+= N

EE

EEEAk

(A-19)

1

2

3255 )( −

Δ+Δ−Δ+= N

EE

EEEAk

(A-20)

In the above equations, N is taken to be the number of vibrational degrees of freedom.

The parameters ΔE1, ΔE2, ΔE3, and E# are defined in Figure 1-1 and A2, A3, A4, and A5

represent the respective frequency factors. The frequency factor A2 is approximated by

the frequency of the normal mode which leads to the formation of complexes V2O5+-

C2H4 and V4O10+-C2H4 (Figures 1-2(a) and 1-2(b), respectively, at t=0). This results in

values of ν = 992 cm-1 for V2O5+ and ν = 502.0 cm-1 for V4O10

+. Frequency factors A3

and A4 are calculated using theoretical normal mode frequencies for the complexes C and

D. Complex D is related to the structures shown in Figure 1-2(a) (t=20 fs) for V2O5+-

C2H4 and Figure 1-2(b) (t=50 fs) for V4O10+-C2H4 and the transition state (1) according

to:

#1)4(3i

Ni

iNiA

νν

−ΠΠ= .

(A-21)

A5 is approximated by the frequency of the V-O stretching vibration which leads to the

products and has a value of ν = 484 cm-1 for V2O5+-C2H4 and ν = 397.3 cm-1 for V4O10

+-

C2H4 complexes.

To apply RRKM theory to the reaction mechanism given in Figure 1-1, it is

necessary to identify the transition states of the association, hydrogen transfer, and

190

acetaldehyde formation. No transition states are available because the association step

engages an ion-molecule reaction without a barrier and the formation of acetaldehyde

involves a barrier-less V-O bond dissociation. Therefore, RRK theory is used to

calculate the rate constants for these two reaction steps using approximate frequency

factors of the normal modes which lead to the association and emanation of acetaldehyde.

For the hydrogen transfer step, RRK theory is also used because no low frequency modes

are present for which the harmonic approximation underlying RRK theory fails.

Moreover, because the barrier for the hydrogen transfer is significantly lower than the

total internal energy acquired in the association step, quantitative differences between the

RRK and RRKM theory are not expected to influence the final reaction cross section.

Figure 1-9 compares the experimental cross sections at the zero pressure limit

and the cross sections calculated from the Langevin model corresponding to the center-

of-mass energy. The theoretical cross sections for the reaction of V2O5+ with ethylene in

Figure 1-9(a) employ eqs A-16 to A-20 with the following values for parameters defined

in Figure 1-1: ΔE1 = 3.85 eV, ΔE2 = 5.19 eV, ΔE3 = 2.66 eV, and E# = 1.45 eV obtained

from DFT calculations (1,26). Figure 1-9(b) is a plot for the reaction of V4O10+ and

ethylene calculated in the same way with parameter values as follows: ΔE1 = 2.18 eV,

ΔE2 = 4.17 eV, ΔE3 = 2.16 eV, and E# = 0.13 eV.1 As can be seen from eq A-16, the total

reaction cross section consists of the product of two terms. The first term is the pure

Langevin cross section for the association of ethylene with the metal oxide, and the

second term is dependent on the rate constants for the subsequent steps (A-16 to A-20).

191

The second term involving the rate constants is found through calculations to be unity for

these particular parameter values due to the low reaction barriers. Therefore, based on

the above assumptions the total reaction cross section is predicted to be accounted for

exclusively by the Langevin cross section. It is interesting to note that Hanmura et al.

have found that the reaction of benzene with nickel cluster ions is also governed by

Langevin kinetics (27).

Comparing the cross sections obtained theoretically using eq A-16 with those

from experiment, we find good agreement between experiment and theory for both

systems as shown in Figure 1-9. It is reasonable that the experimental cross sections are

systematically lower, showing that not every collision is an effective collision. Both

V2O5+ and V4O10

+ with ethylene have decreasing reaction cross sections with increasing

energy, characteristic of an exothermic reaction process with a negative temperature

dependence.

This provides consistent evidence that subsequent oxygen transfer and formation

of acetaldehyde proceed through a mechanism where the barriers are significantly lower

than the energy of the initial association complex. On the basis of the theoretical model

presented here, together with the experimental findings, we show that the reaction

kinetics are comparatively insensitive to the detailed energetics of the involved steps.

This applies as long as the barrier for the hydrogen transfer is significantly lower than the

energy of the initial complex. This is in contrast to what has been recently claimed for

the energy profile differences between the 2A'' ground state and the 2A' ground state in

V2O5+ (9). Although the reaction energies for V2O5

+ + C2H4 change for the 2A'' ground

192

state, it is insignificant. The barriers for the reaction are lower in energy than those for

the initial rate-limiting association step, showing that the reaction is governed by

Langevin kinetics. Hence subsequent reaction steps readily proceed.

Subsequent to the above study, we conducted an investigation to ascertain under

which conditions deviations from the Langevin cross sections would be expected to

occur. In Figure 1-10 the reaction cross sections are plotted for four different values of

the hydrogen transfer barrier, 1.45 eV, 3.75 eV, 3.85 eV, and 3.88 eV. Whereas the cross

section for the situation of a barrier of 1.45 eV is indistinguishable from the Langevin

cross sections, deviations start to occur when the barrier reaches a value of 3.85 eV which

corresponds to the energy of the initial complex. For barriers higher than 3.85 eV, the

behavior is qualitatively different, giving rise to a cross section which initially increases

with energy, reaches a maximum, and then decreases. This is illustrated in Figure 1-10

for only a slightly increased barrier with respect to 3.85 eV. These findings will allow for

insight into mechanisms of related reactions through a study of the energetics of the

reactive cross sections.

A.5 Conclusion

A theoretical and experimental analysis of the reaction of V2O5+ and V4O10

+ with

ethylene provides additional information that demonstrates consistency with the

previously proposed energetic profile and mechanism of the oxygen transfer radical

cation reaction.

193

Figure 1-9: a) Energy dependence of the total reaction cross sections for V2O5+ with

ethylene for the theoretical model (solid line) and experimental measurements (diamonds). b) Energy dependence of the total reaction cross sections for V4O10

+ with ethylene. The error bars represent one standard deviation of the experimental cross section calculated from the root mean square value of the least squares analysis.

Figure 1-10: Dependence of the reaction cross section (σ(E)) on the energy for different values of the barrier for hydrogen transfer: 1.45 eV (black line), 3.75 eV (blue line), 3.85 eV (green line), 3.88 eV (red line). Source: Figure courtesy of the group of Professor V. Bonačić-Koutecký at the Humboldt Unisversity in Berlin.

0 10

20

30 40 50 60

0 0.5 1.0 1.5 2.0 2.5 3.0 0

10

20

30 40 50 60

0 0.5 1.0 1.5 2.0 2.5 3.0 0

10

20

30 40 50 60

0 0.5 1.0 1.5 2.0 2.5 3.0 0

10

20

30

40

50

60

0 0.5 1.0 1.5 2.0 2.5 3.0 0

10

20

30

40

50

60

70

0 0.5 1.0 1.5 2.0

a) b)

EnergyCM (eV) EnergyCM (eV)

σ[Å

2 ]

194

Comparison between experimental and theoretical results has given evidence that all

steps following the formation of the initial complex as well as the hydrogen transfer

reaction proceed with barriers which are significantly lower than that of the initial

association complex. This finding supports our earlier proposed reaction mechanism,

furthermore showing that the reactions effectively proceed according to the Langevin

model. This established that the reaction is driven by the initial rate constant and

associated cross section that governs the encounter of the reactants and leads to the

formation of an intermediate collision complex. Moreover, the kinetic model proposed

allows one, in principle, to estimate the barriers from experimental data on the energy

dependence of the reaction cross section, providing important information about cluster

reactivity.

Acknowledgements

NAM, DRJ, and AWC, Jr. gratefully acknowledge the U. S. Department of

Energy, Grant No. DE-FG02-92ER14258, for financial support of the experimental work

reported herein, and RM and VBK acknowledge support from Deutsche

Forschungsgemeinschaft (DFG). We thank Dr. Michele Kimble for helpful discussions

during the course of the study.

195

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VITA

Nelly Moore Reilly

Education The Pennsylvania State University, University Park, Pennsylvania Doctorate of Philosophy in Chemistry, May 2007.

Ursinus College, Collegeville, Pennsylvania

Bachelor of Science in Chemistry, May 2000 Experience 2002-2007 Graduate Research Assistant Department of Chemistry The Pennsylvania State University, University Park, PA 2000-2002 Associate Scientist

Nutritional Research Wyeth Pharmaceuticals, Collegeville, PA

Awards 2006 Geiger Fellowship 2006 Graduate Student Travel Award for APS Conference 2005 Braucher Fellowship 2004 Graduate Student Travel Award to Humboldt University 2002 Roberts Fellowship Professional Affiliations 2000-present American Chemical Society, Analytical Division 2006-present American Physics Society Select Publications

1. Reilly, N. M., Johnson, G. E., Reveles, J. U., Khanna, S. N., and Castleman, A. W., Jr. “Experimental and Theoretical Study of the Structure and Reactivity of Fe1-2O≤6

- Clusters with CO.” J. Phys. Chem. A 2007, Accepted.

2. Reilly, N. M., Johnson, G. E., Reveles, J. U., Khanna, S. N., and Castleman, A. W., Jr. Chem. Phys. Lett. 2007, 435, 295.

3. Moore, N. A., Mitrić, R., Justes, D. R., Bonačić-Koutecký, V., and Castleman, A. W., Jr. J. Phys. Chem. B 2006, 110, 3015-3022. Select Presentation

1. Moore, N. A., Kimble, M. L., Johnson, G. E., and Castleman, A. W., Jr. “Investigation of Gas Phase Gold Oxide Cations towards the Oxidation of Carbon Monoxide” American Physics Society, Baltimore, MD (March 2006).

fundamental study of noble metal and transition metal …

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