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THE ACTIVATION OF CARBON MONOXIDE AND CARBON 
DIOXIDE BY TRANSITION METAL CARBONYL COMPLEXES 



BY 
KEITH D. WEISS 



A DISSERTATION PRESENTED TO THE GRADUATE SCHOOL 

OF THE UNIVERSITY OF FLORIDA IN 

PARTIAL FULFILLMENT OF THE REQUIREMENTS FOR 

THE DEGREE OF DOCTOR OF PHILOSOPHY 

UNIVERSITY OF FLORIDA 

1986 



To my mother and father 



ACKNOWLEDGEMENTS 

I am very indebted to Dr. Drago for the guidance he has 
given me over the past five years. I would like to thank 
him and his wife for making my stay in graduate school both 
an enjoyable and a rewarding experience. I must confess 
upon departing from Florida that his definition of a "silver 
fox" may be correct. 

I would like to thank the members of the Drago group 
for their support over the past few years. I owe a special 
thank you to Jim Miller for his friendship through my stay 
at the University of Florida. I would also like to thank 
Leigh Ann Files, Cindy Goldstein, Cindy Bailey, Maribel Lisk 
and Nancy Miller for giving me over the past several years 
many reasons to smile. 

The motivation to excel at anything usually stems from 
experiences that have happened early in one's life. In this 
respect, I would like to acknowledge my high school 
chemistry teacher, Mr. Russell Hull, for instilling within 
me an interest in chemistry. 

The assistance and expertise of several people have 
been essential to the research presented in this 
dissertation. I would like to acknowledge Tim Koloski for 
his contribution as an undergraduate research student to the 



chapter on the activation of carbon dioxide. Professor Roy 
King and Dr. Tom Gentle are acknowledged for their 
help in obtaining, respectively, the GC/MS and ESCA data 
presented in this dissertation. It also is acknowledged 
that the typing of this dissertation could not have been 
completed without the thoughtful diligence of Sharon Decker. 

Finally, the devotion and understanding of my parents 
have been indispensable during the years I have been in 
school. It primarily has been their support that has 
enabled me to finish this degree. It was with them in mind 
that this dissertation was written. 



TABLE OF CONTENTS 



Page 



ACKNOWLEDGEMENTS m 

LIST OF TABLES ix 

LIST OF FIGURES xi 

ABSTRACT xvi 

CHAPTER 

I INTRODUCTION 1 

II ACTIVATION OF CARBON DIOXIDE 4 

Background 4 

Experimental 2 

Reagents 20 

Instrumentation 21 

Preparation of Potassium Tetracarbonyl- 

cobaltate(l-) 21 

Preparation of Potassium Tetracarbonyl- 

ferrate(2-) 22 

Preparation of Sodium Pentacarbonyl- 
manganate (1-) and Sodium Penta- 

carbonylrhenate (1-) 23 

Preparation of Potassium y-hydridobis- 

(pentacarbonyltungsten(O) ) 23 

Preparation of a Mixture of Rhenium 

Carbonyl Hydrides 25 

Reaction of the Transition Metal 

Carbonyl Salts with Carbon Dioxide. . . 25 



Reaction of Carbon Dioxide and Hydrogen 
with Transition Metal Hydrides in 
Alcohol Solvents 26 

Results and Discussion 27 

The Binding of Carbon Dioxide by 
Transition Metal Carbonyl Anions. ... 27 

Formation of Alkyl Formates at Low 
Pressures and Temperatures 36 

Summary 57 

III ACTIVATION OF CARBON MONOXIDE 60 

Background 60 

Experimental 76 

Reagents 7 6 

Instrumentation 76 

Fixed Bed Flow Reactor 7 8 

Preparation of Dicarbonylchloro- 
(p-toluidine) iridium(I) . . .' 78 

Preparation of a Phosphinated Support. . 80 

Preparation of Supported Mono- and 
Di-phosphine Substituted Tetrairidium 
Carbonyl Clusters 81 

Preparation of Supported Tri-phosphine 
Substituted Tetrairidium Carbonyl 
Clusters 82 

Preparation of Other Supported 
Phosphine Substituted Metal Carbonyl 
Complexes 8 2 

Preparation of Iridium Complexes 
Impregnated Onto a Support 8 3 

Reaction of Catalysts with Carbon 
Monoxide, Hydrogen and HCl(g) 83 

Reactions Involving Carbon-13 
Isotopically Labelled Gases 84 

Results and Discussion 85 



Reproduction of the Previously Reported 
Supported Iridium Carbonyl Catalyst 
System 85 

Catalyst Deactivation is a Valid 
Observation 91 

Investigation of the Methyl Chloride 
Formation Observed in the Control 
Reactions 9 5 

Solvent Decomposition Can Explain 
Other Reported and Observed Results . . 101 

The Reduction of Carbon Monoxide 
Still Occurs 107 

Minor Impurity Routes to Methyl 
Chloride 114 

Structural Determination of the 
Supported Iridium Clusters by 
Infrared Spectroscopy 121 

Examination of the Decomposition of the 
Supported Clusters by Infrared 
Spectroscopy 128 

Verification That a Discrete Molecular 
Complex Exists in the Activated 
Iridium Svstems 132 

Investigation by Infrared Spectroscopy 
of the Discrete Iridium Species Present 
in the Activated Systems at 75°C. . . . 138 

Reevaluation of the Supported Cluster/ 
AlCl^-NaCl System 150 

Proposed Mechanism for the Formation 
of Methyl Chloride 152 

Investigation of a Phosphine Substituted 
Triosmium Carbonyl System 161 

Investigation of Supported Cobalt and 
Iron Carbonyl Systems 165 

Investigation of a Phosphine Substituted 
Triruthenium Carbonyl System 169 

Summary 188 



IV CONCLUSION 192 

APPENDICES 

A Infrared Spectra for Chapter II 194 

B Infrared Spectra for Chapter III 211 

REFERENCES 2 20 

BIOGRAPHICAL SKETCH 234 



LIST OF TABLES 



2-1 Enthalpy Changes in Reactions Involving 
Carbon Dioxide 



Page 



2-2 Modes of Carbon Dioxide Binding for 

Complexes with Geometry Determined by 
X-ray Crystallography , 



2-3 Acid Dissociation Constants for Metal 
Carbonyl Hydrides and the Nucleo- 
philicities of the Corresponding Anions.. 10 

2-4 A Summary of Catalyst Activity for the 
Reduction of Carbon Dioxide to Alkyl 
Formates 15 

2-5 A Summary of Catalyst Activity for the 
Reduction of Carbon Monoxide to Alkyl 
Formates 18 

2-6 Characterization of Prepared Transition 
Metal Carbonyl Anions by Infrared 
Spectroscopy 2 4 

2-7 Color Changes for Reactions Between 
Carbon Dioxide and Transition Metal 
Carbonyl Anions 26 

2-8 A Summary of Infrared Data Obtained for 
Reactions Between Carbon Dioxide and 
the Transition Metal Carbonyl Anions.... 29 

2-9 A Summary of the Quantitative Data 

Obtained for the Reaction Products by 

Gas Chromatography 3 8 

2-10 A Summary of Infrared Data Obtained for 
the Reaction of Re„ (CO) ^„ with CO and 
H2 in Methanol 7 46 

2-11 A Summary of Infrared Data Obtained for 
the Reaction of Re (CO) with Carbon 
Monoxide in Methanol 49 



Page 



3-1 A Comparison of the Advantages and Dis- 
advantages of Using Homogeneous and 
Heterogeneous Catalysts 63 

3-2 A Summary of the Infrared Data Reported 
for the Supported Iridium Carbonyl 
Catalysts 69 

3-3 A Summary of the Reported Activity 
for the Supported Iridium Carbonyl 
Catalyst System 73 

3-4 The Initial Activity of Control 

Supports Prepared in 2-methoxyethanol . . . 95 

3-5 Factors That Enhance the Adsorption of 
Polar Organic Compounds onto the 
Surface of a Hydroxylated Support 97 

3-6 The Activity of Alumina Supported 
Tetrairidium Clusters for Various 
Metal Loadings 102 

3-7 Thermodynamic Data for Reactions 

Involving Synthesis Gas Reduction 107 

3-8 Thermodynamic Data Concerning the 

Decomposition of Toluene 117 

3-9 A Summary of the Infrared Data Obtained 
for the Supported Iridium Carbonyl 
Clusters 123 

3-10 Infrared Data for Complexes of 

Triosmium Carbonyl 163 

3-11 Infrared Data for Supported Phosphine 

Carbonyl Complexes of Iron 167 

3-12 A Summary of Infrared Data Obtained 
for the Supported Ruthenium Cluster 
System 179 



LIST OF FIGURES 



Page 



2-1 Modes of Binding for Carbon Dioxide 

with a Metal Center 6 

2-2 A Proposed Mechanism for the Reduction 
of Carbon Dioxide to Alkyl Formates 
by Group 6B Metal Carbonyl Hydrides 13 

2-3 A Proposed Mechanism for the Reduction 
of Carbon Monoxide to Methyl Formate 
by Tungsten Carbonyl Complexes 19 

2-4 Infrared Spectrum of Precipitate 

Obtained from Reaction of Carbon Dioxide 

and NaMn(CO)^ 31 

2-5 Solution Infrared Spectrum of the 

Carbonyl Region for the Reaction Between 

Carbon Dioxide and NaMn(CO)c in 

Acetonitrile 32 

2-6 Solvent Subtracted Infrared Spectrum 

for the Reaction Between Carbon Dioxide 

and NaMn(CO)c in Tetrahydrofuran 34 

2-7 Mass Intensity Report for Methyl 

Formate 39 

2-8 Mass Intensity Report for Dimethyl 

Ether 40 

2-9 Mass Intensity Report for Dimethoxy- 

me thane 41 

2-10 Mass Intensity Report for Hexane 42 

2-11 Gas Chromatogram of Gas Sample Taken 
During Reaction of Re^CCO),^ with CO- 
+ H- in Methanol 7 43 



Page 



2-12 Infrared Spectrum of Reaction Products 

for CO /H Reaction with Re (CO) 47 

2-13 Infrared Spectrum of Reaction Products 
for Rhenium Hydride and CO Reaction 
with Re (CO) and Methanol Subtracted 
Out . 51 

2-14 Infrared Spectrum Obtained for the 
Reaction of Carbon Monoxide with a 
Mixture of [Re2 (CO) g ( y-OCH^) ^l ~ and 
[H2Re (CO) ^] ~ in Methanol 54 

2-15 A Comparison by Gas Chromatography of 

Reaction Products in Liquid Samples for 
Carbon Monoxide Reactions with Re„(CO)^_, 
KH[W(CO) ]2 and KOCH 7 .. 56 

3-1 A Proposed Mechanism for the Formation 
of Methyl Chloride from Synthesis Gas 
and HCl Over a Supported Iridium 
Carbonyl Catalyst System 74 

3-2 A Diagram of the Fixed Bed Flow Reactor. 79 

3-3 A Gas Chromatogram of the Product Gases 
(Poropak Q Column, Attenuation = 8, 
Column = 130°C) 87 

3-4 Gas Chromatography Separation of 

Methanol and Methyl Chloride 89 

3-5 A Graph of Catalyst Activity Versus 

Reaction Temperature 90 

3-6 A Graph of Catalyst Activity Versus 
Reaction Time. (Residence Time and 
HCl(g) Concentration was Observed to 
Increase with Time) 92 

3-7 A Graph of Catalyst Activity Versus an 
Extended Reaction Time. (Residence 
Time and HCl(g) Concentration Held 
Constant) 93 

3-8 NMR Spectrum of 2-methoxyethanol That 
Condensed at the Top of the Reactor 
System 99 



Page 

3-9 A Graph of Methyl Chloride Activity 
Versus Residence Time for a Reaction 
Involving a Control Support 104 

3-10 A Graph of Initial Methyl Chloride 
Activity Versus Time for a Catalyst 
Prepared in Toluene Ill 

3-11 Mass Intensity Report for Acetylene 112 

3-12 Infrared Spectrum of the C=C Vibrations 
in the Phenyl Groups of the Silane 
Linkage 119 

3-13 Infrared Spectrum of 3- (PPh2) 3lr^ (CO) g 
Obtained From the Reaction of 3-PPh„ 
With Ir^(C0)^2 • ^^^ 

3-14 A Comparison of the Infrared Spectra of 
the Mono- and Di-phosphine Substituted 
Tetrairidium Carbonyl Clusters Prepared 
in 2-methoxyethanol and Toluene 127 

3-15 Infrared Spectra for the Supported 

Clusters After Exposure to the Reactant 
Gases at 75°C 129 

3-16 Infrared Spectra Obtained for the Low 
and High Load Iridium Systems at 
Various Temperatures 131 

3-17 Infrared Spectrum Obtained for Metallic 
Iridium/Alumina After Exposure to CO, 
H and HCl (g) at 200°C 136 

3-18 X-ray Photoelectron Spectrum for the 
Supported Clusters Before and After 
Activation 137 

3-19 A Comparison of the Infrared Spectra 
Obtained for the Supported Clusters 
(Low % Ir) and for IrCKCO)^ on a 
Phosphinated Support 140 

3-20 Infrared Spectrum of IrCl (CO) ^ + 3-PPh2 

at Various Temperatures 141 



3-21 A Comparison of the Infrared Spectriun 
of IrCl (CO) -/Alumina With the Spectra 

Obtained for the Activated Iridium 

Clusters at Various Temperatures 144 

3-22 Infrared Spectrum of IrCl (CO) ^/Alumina 

at Various Temperatures 145 

3-23 Infrared Spectrum of a Mixture of the 
Supported Multinuclear Complex and 
IrCl(CO)^ Exposed to Carbon-13 Carbon 

Monoxide 147 

3-24 Infrared Spectrum of Vaska's Complex 

Supported on Alumina After Exposure to 

CO/H /HCl(g) at Various Temperatures.... 149 

3-25 Infrared Spectrum of the Supported 

Iridium Clusters in an AlCl_-NaCl Melt.. 151 

3-26 Proposed Mechanism for the Initial 
Formation of Methyl Chloride in the 
Supported Iridium Cluster System 153 

3-27 Proposed Mechanism for the Formation of 
Methyl Chloride From Synthesis Gas and 
HCl (g) 156 

3-2 8 Overview of Proposed Mechanism for the 
Formation of Methyl Chloride in the 
Supported Iridium Cluster System 162 

3-29 Comparison of the Infrared Spectrum 

of 3-PPh Os (CO) ^^ Exposed to CO, H^ 

and HCl(g) with that of Os2(CO)^2 '^^^ 

3-30 Infrared Spectrum of B-PPh^Fe (CO) ^ Before 

and After Exposure to CO, H_ and HCl(g) 

at 75°C 7 168 

3-31 Gas Chromatogram of Two Carbon products 
Obtained in the Supported Ruthenium 
Cluster System 170 

3-32 Mass Intensity Report for Acetaldehyde. . 172 

3-33 Mass Intensity Report for Ethyl 

Chloride 173 



3-34 Mass Intensity Report for Ethyl 

Formate 174 

3-35 Mass Intensity Report for Diethyl 

Ether 175 

3-36 Mass Intensity Report for 1,1- 

Dichloroethane 176 

3-37 Mass Intensity Report for Ethyl 

Acetate 177 

3-38 Infrared Spectrum of a Mixture of the 
Supported Ruthenium Clusters Exposed 
to Air 180 

3-39 Infrared Spectrum of the Supported 

Ruthenium Clusters Exposed to CO, Hp 

and HCl(g) at75°C T 182 

3-40 A Comparison of the Infrared Spectrum 
of the Supported Ruthenium Clusters 
with that of Supported [RUCI2 (CO) ^1 2 1^3 

3-41 Infrared Spectrum of Supported 

[RuCl-(CO) ] Before and After Exposure 

to CO, H and HCl (g) at 75°C 184 

3-42 Fragmentation of Ru^ (CO) ^2 °^ ^ Silica 

Gel Support 186 



Abstract of Dissertation Presented to the Graduate School of 

the University of Florida in Partial Fulfillment of the 

Requirements for the Degree of Doctor of Philosophy 

THE ACTIVATION OF CARBON MONOXIDE AND CARBON 
DIOXIDE BY TRANSITION METAL CARBONYL COMPLEXES 

By 

Keith D. Weiss 

August 1986 

Chairman: Russell S. Drago 
Major Department: Chemistry 

The development of new carbon resources to be used 
either as fuels or as chemical feedstocks has been of major 
concern to the industrial community since the oil shortage 
in the early 1970s. In this respect, a lot of activity has 
been observed over the past decade concerning the binding 
and activation of carbon dioxide and carbon monoxide. The 
emphasis in this work has been upon the activation of carbon 
dioxide and carbon monoxide by transition metal complexes. 

The first study involved the binding of carbon dioxide 
to a variety of transition metal carbonyl anions. It was 
found that the nucleophilicity of the metal center greatly 
affected the coordination of carbon dioxide. This work then 
was extended into an investigation of the low pressure 
reduction of carbon dioxide to methyl formate by a rhenium 
carbonyl catalyst. Infrared spectroscopy was used to 
characterize the active species and probe the mechanism of 
the reaction. It was found that this rhenium carbonyl 



system was more effective towards the reduction of carbon 
monoxide than carbon dioxide. 

The second study dealt with the conversion of synthesis 
gas and HCl(g) to methyl chloride under mild temperatures 
and pressures by a supported iridium carbonyl catalyst. 
Several different routes to methyl chloride were identified 
within the system. Infrared spectroscopy was used to 
identify the active species in the reaction and investigate 
their various modes of deactivation. It was shown that 
discrete iridium carbonyl complexes existed under the 
employed reaction conditions. Finally, a mechanism for the 
formation of methyl chloride in this system was proposed. 
This study concluded with an examination of other supported 
metal carbonyl catalysts. It was found that a change in the 
composition of the metal catalyst could alter the activity 
and selectivity of the reaction. In this respect, a 
supported ruthenium carbonyl catalyst was observed to be 
active towards the formation of ethyl chloride and various 
other two-carbon products. 



CHAPTER I 
INTRODUCTION 
With the oil shortage in the 1970s, there was great 
interest in the development of alternative carbon resources 
to be used as fuels and chemical feedstocks. Although the 
oil shortage has receded for the time being, the problem has 
not been resolved. The reoccurrence of a fuel shortage is 
predicted for the economy of the future. One new carbon 
resource would involve the gasification of coal into a 
mixture of carbon monoxide and hydrogen known as synthesis 
gas. Since there is an abundance of coal reserves, the 
utilization of synthesis gas as a chemical feedstock or fuel 
could hold future economic advantages. Another untapped 
carbon reserve is the carbon dioxide released into the 
atmosphere as an industrial waste product. Carbon dioxide 
has the ability to absorb infrared radiation which is 

predicted to give rise to an increase in global temperature 

2 3 
commonly referred to as the "Greenhouse Effect". ' 

Although other gases can initiate a similar temperature 

rise, the most abundant gas in the atmosphere to produce 

4 
this effect is carbon dioxide. Recycling carbon dioxide 

waste by utilizing it as an inexpensive chemical feedstock 

would decrease the amount of the gas released into the 

atmosphere. 



2 

There are several methods that can be used to 
investigate the activation of carbon monoxide and carbon 

dioxide. These methods range from mechanistic studies of 

5 
biological or natural systems, such as hemoproteins , 

carbonic anhydrase and photosynthesis to the investigation 

of the interaction of carbon monoxide and carbon dioxide 

with metal complexes in organometallic reactions. Currently 

over 90% of the commercial chemical processes are catalytic 

g 

in nature. Since many of these commercial processes employ 
either metals (heterogeneous) or metal complexes 
(homogeneous) as catalysts, a logical starting point for the 
development of carbon monoxide and carbon dioxide as a 
chemical feedstock or fuel is a study of their interaction 
with transition metal complexes. 

Reported here are the results of two studies involving 
an evaluation of the feasibility of binding and activating 
carbon dioxide and carbon monoxide by transition metal 
carbonyl complexes. The first study involves the binding of 
carbon dioxide to transition metal carbonyl anions with 
varying degrees of nucleophilic metal centers. The results 
obtained aid in the understanding of the interaction of 
carbon dioxide with transition metal catalysts, as well as 
the effect that residual contaminants, such as water, may 
have on carbon dioxide activation. This study concludes 
with an investigation of the catalytic behavior of the 
corresponding transition metal carbonyl hydrides towards 
carbon dioxide reduction. Primarily, the low temperature 



3 

and low pressure reduction of carbon dioxide and hydrogen in 
alcohol solvents to alkyl formates was investigated. 

The second study involves the low temperature and low 
pressure reduction of carbon monoxide by "heterogenized" 
homogeneous catalysts. The mechanism of a novel system 
employing a supported tetrairidium carbonyl cluster as 
catalyst for the conversion of carbon monoxide, hydrogen and 
hydrogen chloride to methyl chloride was studied by infrared 
spectroscopy. The commercial feasibility of this system was 
evaluated through the optimization of the different system 
parameters, such as temperature and residence time. 



CHAPTER II 
ACTIVATION OF CARBON DIOXIDE 
Background 
The thermodynamic stability of carbon dioxide is the 
primary reason it is an "oxygen sink" or waste product in 
many commercial chemical processes. However, the two 
unsaturated double bonds in carbon dioxide make it 
theoretically possible to convert carbon dioxide into 

organic products. This is supported by the thermodynamic 

9 
feasibility of many reactions involving carbon dioxide as 

shown in Table 2-1. Furthermore, carbon dioxide is used as 



Table 2-1. Enthalpy Changes in Reactions Involving Carbon 
Dioxide 



Reaction AH° 

C02(g) + H2(g) >► CH^OHd) + H20(l) -31.3 

C02(g) + 4H2(g) ^^CE^iq) + 2H20(1) -60.5 

2C02(g) + 6H2(g) ^►CH^OCH^ig) + 3H20(1) -60.9 

C02(g) + H2(g) + CH^OHd) •►HC02CH2(1) + H20(l) -7.7 

C02(g) + H2(g) + CH^OHd) •►CH2C02H(1) + H20(l) -33.0 

C02(g) +CH^(g) ^CH2C02H(1) -3.8 

C02(g) + H2(g) + C2H2(g) •►C2H^C02H(1) -40.6 



AH° data given as kcal mole 



a chemical feedstock in the commercial production of 

1 • 1 • 'J 10,11 12,13 -, ^ ^u -, • -0 14 
salicylic acid, ' urea ' and terepthalic acid. 

Unfortunately, the existence of a kinetic barrier in many 

reactions involving carbon dioxide prevents these reactions 

from spontaneously occurring. These reactions may occur if 

the activation energy associated with this kinetic barrier 

can be lowered through the binding of carbon dioxide to a 

metal catalyst. 

There have been a variety of literature review articles 

dedicated to the binding and activation of carbon 

9 15-22 
dioxide. ' Several possibilities exist for the 

18 2 3 
interaction of carbon dioxide with a metal center. ' One 

mode of binding is a donor-acceptor type complex termed 

"end-on" formed by electron pair donation from the oxygen's 

highest occupied molecular orbital (HOMO) to the 

corresponding metal d orbital as shown in Figure 2-1. A 

"c-coordination" mode of binding involving a dative 

interaction or transfer of two electrons from the metal atom 

to the lowest unoccupied molecular orbital (LUMO) of carbon 

dioxide also may occur. Finally, a third possible mode of 

binding referred to as "side-on" is a combination of the 

interaction of the HOMO of carbon dioxide with a vacant 

metal d orbital and the simultaneous transfer of metal a 

electrons to the vacant LUMO of carbon dioxide. The high 

electron affinity of carbon dioxide suggests that the 

existence of the "side-on" and "c-coordination" modes of 

binding should be more favorable than the "end-on" mode. 



"End-on" "C-coordination" "Side-on" 



iyN-Np=C=(l M-C'< >-M-c' M I 

01 

Figure 2-1. Modes of Binding for Carbon Dioxide with a 
Metal Center 



This is supported by the favorability of "side-on" > 

"c-coordination" >> "end-on" reported in several molecular 

2 4 — 2 8 
orbital investigations. These studies point out that 

the favorability of "c-coordination" may be enhanced by the 

presence of a counter ion which can interact with the 

electron density surrounding the oxygen atoms of the bound 

carbon dioxide. Further evidence is provided by the limited 

number of x-ray structures that have been obtained for bound 

carbon dioxide complexes as shown in Table 2-2 . Many other 

complexes have been reported to bind carbon dioxide as 

9 
suggested by infrared spectroscopy. 

Although the "side-on" mode of coordination is the most 

favorable, the "c-coordination" mode is very important for 

catalysis. Metal-carbon bond formation may lead either to 

the growth of carbon-carbon chains through insertion 

reactions or to the catalytic formation of formate species 

through hydrogenolysis. To date there has been only limited 

success involving the catalytic reduction of carbon dioxide. 

The first example of the catalytic reduction of carbon 

dioxide was the conversion of an amine to a formamide 



c-coordination 


29 


c-coordination 


30 


c-coordination 


31 


c-coordination 


32 


side-on 


33 


side-on 


34 


Combination of modes 


35 



Table 2-2. Modes of Carbon Dioxide Binding for Complexes 
with Geometry Determined by X-ray 
Crystallography 



Bound Carbon Dioxide Complex Mode of Binding References 
(py)Co(salen)K(C02) 
Rh(diars) ^ (C02)C1 
[HOS3 (CO) ^Q (02C)0Sg (CO) ^7] " 
[(OC)5Re(C02)Re(CO)4]2 
Ni(C02) (PCy3)2 

Nb(n-C^H^Me) 2 (CH2SiMe2) (CO2) 
IrCKC 0^) (PMe )2-0.5«CgHg 

a = Complex not formed directly from carbon dioxide (g). 

employing either IrCl (CO) (PPh^) 2 OJ^ CuCl(PPh2)3 as 
catalyst. It was proposed that the mechanism for this 

reaction proceeded through the insertion of carbon dioxide 

9 
into a metal-hydrogen bond. 

The interaction of carbon dioxide with a transition 

metal hydride may proceed through two different reaction 

pathways. Carbon dioxide insertion into a metal-hydrogen 

bond will lead to either the formation of a formate complex 

or to a metallocarboxylic acid derivative as shown in 

Equation 2-1. Although the formation of metallocarboxylic 

acid complexes has been reported for reactions between 

37 
carbon monoxide with a metal hydroxide complex and for 

38-42 ^, , 
hydroxide ion with a metal carbonyl complex, there has 

been no direct evidence for the formation of a metallocar- 



boxylic acid derivative through the insertion of carbon 
dioxide into a metal hydride bond. The isolation of organic 
products in several reactions has suggested the possibility 

» A 

M-O-CH or M :CH 



CO2 + M-H ^ 

N. (2-1) 

II 
M-C-OH 

43 
of a metallocarboxylic acid intermediate. On the other 

hand, the formation of formate intermediates has been 

observed for carbon dioxide reactions with metal hydride 

T 44,45 
complexes. 

It is possible that metal-carbon bond formation can be 

enhanced by altering the nucleophilicity of the metal center 

in a transition metal carbonyl hydride complex. It has been 

shown by infrared and carbon-13 NMR spectroscopy that alkali 

metal salts of transition metal carbonyl hydrides can bind 

carbon dioxide through the "c-coordination" mode as shown in 

4 6 
Equation 2-2. There are several added advantages for the 

interaction of carbon dioxide with a transition metal 

/■•• 

2Li'^ [W(CO)J^" ^2 ►Li^ + (CO).W-C: Li (2-2) 


complex. First, many carbonyl complexes activate hydrogen 
under mild conditions to produce metal carbonyl hydride 
complexes. Since the reduction of carbon dioxide requires 



9 

a source of hydrogen, the ability to bind carbon dioxide and 
activate hydrogen gas by the same transition metal complex 
would be advantageous. There also have been recent reports 

of transition metal carbonyl complexes reducing carbon 

48 49 

dioxide to carbon monoxide and to alkyl formates. 

An evaluation of the feasibility of binding carbon 

dioxide through metal-carbon bond formation can be 

accomplished by an investigation of the interaction of 

carbon dioxide with the alkali metal salts of transition 

metal carbonyl hydride complexes whose K values or metal 

nucleophilicities are known. The K and nucleophilicity 

data for the complexes to be investigated are summarized in 

Table 2-3. The available literature has indicated that 

several of these anions do interact with carbon 

dioxide. ' ' For instance, the formation of iron 

pentacarbonyl and sodium carbonate has been reported to 

48 
occur for the reaction of carbon dioxide with Na-FeiCO)^. 

It also has been reported that carbon dioxide reacts with 

NaMn(CO)c- to form sodium bicarbonate and an unidentified 

50 
manganese complex. More recently, preliminary solution 

infrared data were interpreted in a Russian report to 

suggest that both NaRe(CO)_ and NaMn(CO)^ stabilize the 

formation of a "c-coordination" bound carbon dioxide complex 

as shown in Equation 2-3. The addition of methyl iodide 

to this carbon dioxide bound complex was reported to result 

in the formation of [ (CO) ^M (CO2CH2) ] 2 by methyl cation 



10 
addition to an oxygen of the bound carbon dioxide. After 
evaluation of the feasibility of binding carbon dioxide to 
transition metal anions, a logical extension of this work 



Table 2-3. Acid Dissociation Constants for Metal Carbonyl 
Hydrides and the Nucleophilicities of the 
Corresponding Anions 

Hydride Complex K (H-O) Anion Nucleophilicity 

HCo(CO)^ <2 Co(CO)^" 1 

H Fe(CO) 3.6 x 10"^(K ), Fe(CO).^~ 

-14 
1.0 x 10 (K2) 

HMniCO)^ 8 x 10~^ MnlCO)^" 77 

HRelCO)^ "Very weakly acidic" Re (CO) 5" 25,000 

a = K data obtained from reference 52; 

b = Nucleophilicity data obtained from reference 53. 



would be an investigation into the catalytic behavior of 
these complexes towards carbon dioxide reduction. 



CO 
NaM(CO) ^(M = Mn, Re) ^[ (CO) ^M (C02Na) ] 2 (2-3) 

Recall that the first example of the catalytic 
reduction of carbon dioxide was the conversion of an amine 
to a formamide. It was found that replacement of the 
amine with an alcohol produced a formic ester as the primary 
product. Since then there have been several reports 
indicating the formation of metalloformate derivatives 
through the interaction of carbon dioxide with group 6B 



11 

metal carbonyl anions and hydrides. ' ' ~ The 
reduction of carbon dioxide to carbon monoxide by group 6B 
metal carbonyl anions, Li2[M(C0)^], has been shown to occur 

by the formation of lithium carbonate and the corresponding 

46 
group 6B metal hexacarbonyl complex, M(CO),. Carbon-13 

b 

labeling studies involving the reversible binding of carbon 
dioxide to group 6B metal carbonyl hydrides, as shown in 
Equation 2-4 (A-B) , demonstrate that carbon dioxide is not 

reduced to carbon monoxide at atmospheric pressure by these 

57 
complexes. A mechanism for the intramolecular conversion 

of a metalloformate complex to a metallocarboxylic acid 

complex would most likely proceed through the reduction of 

the bound carbon dioxide to bound carbon monoxide followed 

by the addition of hydroxide ion to a carbonyl ligand. 

Therefore, these carbon-13 labeling studies also suggest 



}iM(-^^CO)f + ^^C0 2< »H^^CO^M(-^^CO)^" (2-4A) 



HMC-^^CO)^" + ^^CO^l^ ^ H^^C02M(^^C0) ^" (2-4B) 



that the metalloformate complex does not intramolecularly 

5 8 
convert to a metallocarboxylic acid species. Although the 

formation of metallocarboxylic acid species has not been 

reported to occur in reactions between carbon dioxide 

and metal hydrides, they have been reported to form as 

intermediates in reactions between group 6B metal 



12 

hexacarbonyls and hydroxide ions en route to the formation 

57 
of a metal hydride anion and carbon dioxide. It is 

possible that the relative stability of M(COOH), MIO^CH) or 

M(OCHO) could be influenced by an alteration in the 

nucleophilicity of the metal center. 

Equation 2-5 illustrates that at elevated pressures in 

alcohol solvents the group 6B metal carbonyl hydrides 

catalytically reduce carbon dioxide to alkyl formates and 

49 
water. The predominate species in solution during 

[Catalyst] 
CO2 + H2 >► HCO2R + H2O (2-5) 

ROH 

(250 psig) (250 psig) 



catalysis were determined through the interpretation of 
infrared data to be M(CO)g and HC02M(C0)^~. The existence 
of hydrogen bonding between the alcohol solvent and the 
metalloformate intermediate also was suggested in the 
interpretation of these infrared data. The HC02M(CO)5 
intermediate was proposed to be the catalytically active 

species since M(CO), was found to be catalytically inactive. 

49 
The proposed mechanism for this reaction shown m Figure 

2-2 suggests that it is actually formic acid which is 

produced catalytically. The alcohol solvent then reacts 

with this formic acid to form the alkyl formate and water. 

This mechanism is supported in the identification of formic 

acid by gas chromatography in reactions where benzene has 

replaced the alcohol as solvent. A decrease in activity was 



13 






Q) 




73 




•H 




X tn 




OJ 




•H re 




Q -H 




M 




C -C 




>i 




ja X 




^ 




(C iH 




u >, 




c 




M-l 




XI 




S-l 




C (C 




u 




•rH 




-P nH 




U (0 




3 4-1 




-C 0) 




<u s 


hi 


tf 


1 


PQ 


\ 


<U vo 


1 (N 


£1 


ac 


■p a 


^ 


3 




M 


o 


u 


u 


M-l O 






s 


e >i 


<N 


m ja 


o 


•H 


u 


c tn 


a 


(0 Q) 




x: -p 


1 


u to 




0) E 




s u 









TJ fc, 




(U 




m rH 




>i 




04^ 




rH 




^ < 




a< 









< -p 




c\ 




I 




r>j 




<D 




M 




3 




t7> 




•rt 




Ph 



14 
observed for the formation of ethyl formate as compared to 
methyl formate in the corresponding alcohol solvent. This 
effect has been proposed to be due to the increased 
coordinating ability of ethanol inhibiting the oxidative 
addition of hydrogen by the metalloformate intermediate. 

The catalytic activity of systems capable of converting 
carbon dioxide into alkyl formates is summarized in Table 
2-4. The activity of group 6B metal carbonyl hydrides 

compares very closely to the activity for carbon dioxide 

59 54 
reduction to alkyl formates using ruthenium or iron 

carbonyl hydrides as catalysts. Similar activity also has 

been established for systems utilizing group VIII metal 

phosphine complexes with either BF- or a tertiary amine as 

cocatalyst.^^'^^ Although HFe^ICO)^^" and HFe (CO) ^" have 

not been reported to bind carbon dioxide at atmospheric 

pressure these species catalyze the formation of methyl 

54 
formate. The catalytic formation of methyl formate in 

alcohol solvents was found to follow a general trend of 

increasing activity with an increase in temperature, 

pressure or reaction time. Infrared spectroscopy was used 

to identify the formation of carbonate and iron penta- 

carbonyl during the reaction. This reaction was proposed to 

be very selective for methyl formate formation since no 

other low molecular weight products could be identified by 

54 
gas chromatography. The predominant species present m 

reactions involving ruthenium carbonyl hydrides as catalysts 

was identified by infrared spectroscopy to be 



15 

Table 2-4. A Summary of Catalyst Activity for the Reduction 
of Carbon Dioxide to Alkyl Formates 



Catalyst 
u-H[W2(C0)^q]" 
HC02W(C0) 5 
W(CO)g 

y-H[Cr2(C0)^Q]" 
HC02Cr(C0) ^~ 
HRu^ (CO) , ,~ 
HCO2RU2 (CO) ^Q~ 

H3RU4(CO)^2" 
Ru^ (CO) ^2 
HFe3(C0)^^" 
HFe(CO)^" 

Pd(Ph2PCH2CH2PPh2) ^ 
y-H[W2(C0)^Q]" 
HC02W(C0) 5 
HCO2RU3 (CO) ^~ 



Reaction Conditions 

Ref. 49: CO2 (250 psi) + H2 (250 psi) at 125°C for 24 hours 

Ref. 59: CO2 (250 psi) + H2 (250 psi) at 125°C for 24 hours 

Ref. 54: CO2 (300 psi) + ^^ (300 psi) at 150°C for 24 hours 

Ref. 60: CO2 (350 psi) + H2 (350 psi) at 140°C for 21 hours 



Turnover = mole of HCO-CH^/mole of catalyst; (CH^)^N used 
as cocatalyst; Addition of CO (100 psi) to reactant gases, 



Turnover Number 


Reference 


14.7 


49 


16.4 


49 


- 


49 


14.5 


49 


14.6 


49 


4.1 


59 


5.7 


59 


7.3 


59 


<0.3 


59 


5.2 


54 


2.0 


54 


23.0^ 


60 


5.1^ 


49 


3.8^ 


49 


4.1^ 


59 



16 

- 59 
H-Ru.(CO)^_ . It was suggested that these reactions also 

could lead to the formation of carbon monoxide through the 

reverse of the water-gas shift reaction as shown in Equation 

2-6. It is possible that the observed formation of the 

alkyl formate may have resulted from the reduction of carbon 

monoxide. However, the addition of carbon monoxide to the 

carbon dioxide-hydrogen gas mixture in these reactions was 



CO + E^o:^lS^^^lXl^l^co, + H, 



(2-6) 



59 
observed to inhibit the formation of alkyl formates. A 

similar retarding effect upon carbon monoxide addition was 

observed in the activity of the group 6B metal carbonyl 

hydrides towards alkyl formate formation. It was concluded 

that this observed decrease in activity towards the 

formation of alkyl formates demonstrated that the reduction 

of carbon dioxide in these reactions did not proceed through 

49 
the formation of carbon monoxide. This was further 

substantiated by gas chromatographic detection of less than 

0.0 5% carbon monoxide in the CO„-H_ reaction gas mixture. 

This amount was reported to be far below the equilibrium 

distribution of carbon monoxide expected for the reverse of 

the water-gas shift reaction. 

The preceding conclusion is in contradiction to recent 

reports which indicate that carbon monoxide will react with 

either tungsten carbonyl or ruthenium carbonyl hydrides 

in methanol to yield methyl formate as shown in Equation 



17 
2-7. As shown in Table 2-5, the activity for a reaction 
utilizing carbon monoxide is vastly increased over the same 
reaction using a CO„-H„ mixture. In fact, the addition of 
carbon dioxide or hydrogen to the carbon monoxide reactant 
gas has been found to inhibit the formation of methyl 



CO + R0H:^li^^^^iZ^^l=^HC02R (2-7) 



formate. Tungsten hexacarbonyl was identified by infrared 

spectroscopy to be the predominant carbonyl species present 

in the reactions involving a tungsten carbonyl hydride and 
6 2 

carbon monoxide. Although tungsten hexacarbonyl is 
inactive as a catalyst precursor, the addition of potassium 
methoxi-de to this reaction produces methyl formate in high 
yields. A mechanism consistent with the preceding 
observations has been proposed for the reduction of carbon 
monoxide to methyl formate in a methanol solvent. This 
mechanism, shown in Figure 2-3, strongly suggests the 
interaction of a methoxide anion with W(CO), to produce the 
active catalytic species, CH^OWlCO)^ . 

The synthesis of methyl formate from carbon dioxide or 
carbon monoxide is important because methyl formate is used 
to synthesize several organic chemicals, such as formic 
acid, acetic acid, formamide, ethylene glycol and 
formaldehyde. ~ Since most of these are important 
commercial commodity chemicals, the formation of methyl 
formate from carbon monoxide or carbon dioxide could be 



Table 2-5. A Summary of Catalyst Activity for the Reduction 
of Carbon Monoxide to Alkyl Formates 



Catalyst 
y-H[W2(C0)^Q] 
HC02W(C0) 5 
CH W{CO) ^ 
W(CO). 



W(CO) /KOCH^ 
KOCH^ 



H3RU3(CO)^^ 
H,Ru, (CO) 



'3 3 



12 



H3RU3(CO)^^ 
W(CO) g/K0CH3 



Turnover Number 


Reference 


269 


62 


185 


62 


305 


62 


- 


62 


333 


62 


50 


62 


106 


59 


88 


59 


40" 


59 


_c 


62 



Reaction Conditions 
Ref. 62: CO (250 psi) at 125°C for 24 hours 
Ref. 59: CO (250 psi) at 125°C for 24 hours 



Turnover Number = mole of HC02H3/mole of catalyst; 
^Addition of E^ (250 psi) to reactant gas; 
'Addition of CO2 (25 psi) to reactant gas. 



19 






o 


1 fN 


u 


r— 1 


■^-' 


in 


s 


— O 




O U 


+ 


U 




^-' — ^^ 


1 in 


s 


^-» 


' — ' 


o 


K 


C) 


3L 






20 

industrially useful. This industrial importance coupled 
with a fundamental interest in the activation of CO^ and CO 
justifies further research into the formation of alkyl 
formates from the reduction of carbon dioxide or carbon 
monoxide by metal carbonyl complexes. The final 
justification for further careful fundamental work in this 
area arises from the contradictory results that have been 
reported in the literature. 

Experimental 
Reagents 

All metal complexes were used as purchased unless 
otherwise stated. The Re^iCO)^-, Mn-CCO)^^-, C02 (CO) „ and 
W(CO), were purchased from Strem Chemical Co. and the 
Fe(CO)^ and KOCH- from Alfa-Thiokol . All solvents were 
dried prior to use by distillation over CaH_ or in the case 
of alcohols over magnesium metal. The carbon monoxide grade 
C. P. 9 9.5% was purchased from Matheson and the hydrogen and 
carbon dioxide were purchased from Strate Welding. The 
carbon dioxide was of "bone dry" grade. Even though the 
presence of water' could not be identified by infrared 
spectroscopy or gas chromatography, the carbon dioxide was 
dried by passing the gas through two 2 1/2" x 2' glass 
columns of 3A molecular sieves prior to use. A trace amount 
of carbon monoxide was observed by gas chromatography to be 
present in the carbon dioxide. 



21 
Instrumentation 

All air sensitive manipulations were performed in a 
Vacuum Atmosphere Co. model HE-43-2 inert atmosphere box or 
in an Aldrich inert atmosphere glovebag. All experiments 
were performed under either a nitrogen, carbon dioxide or 
carbon monoxide atmosphere. GC analyses were performed on a 
model 3700 FID Varian gas chromatograph equipped with a 
Hewlett-Packard 3390A integrator and a 1/8" x 8' stainless 
steel 5% diethylene glycol adipate on chromosorb P (80/100) 
column. GC mass spectrometry was performed by Dr. R. King 
of the Microanalytical Laboratory, University of Florida, 
Gainesville, Florida. Samples were run on an AEI MS30 mass 
spectrometer with a KOITOS DS55 data station. The system 
was equipped with a PYE Unicam 104 gas chromatograph 
containing a 1/4" x 5' glass 10% diethylene glycol succinate 
on chromosorb W-AW (80-100) column. Infrared spectra were 
obtained on either a Nicolet 7199/170SX FT spectrometer or a 
Nicolet 5DXB FTIR spectrometer. All solution samples were 
run using 0.025 mL path length NaCl cells. All solid 
samples were run as mulls using KBr salt plates. All 
catalytic or high pressure experiments were performed using 
a 250 mL Parr pressure bottle equipped with a brass Swagelok 
pressure head. This reactor system could withstand a 
maximum of 120 psig of pressure. 
Preparation of Potassium Tetracarbonylcobaltate (1-) 

The KCo(CO) salt was prepared by a procedure similar 

fi 7 

to that reported by Edgell and Barbetta. Inside an inert 



22 

atmosphere box, a total of 4.80 g of powdered KOH was added 
to 50 mL of tetrahydrofuran. Another solution containing 
2.0 g of Co_ (CO) „ in 20 itiL of tetrahydrofuran was prepared. 
The solutions were mixed together and stirred. After one 
hour the red-black solution had turned yellow in color and a 
pink precipitate had formed. The solution, which contained 
the KCo(CO) ., was filtered away from the precipitate and 
used in the carbon dioxide experiments. The KCo(CO) . salt 
was characterized by infrared spectroscopy as shown in Table 
2-6. 
Preparation of Potassium Tetracarbonylferrate (2-) 

The K„Fe(CO) . salt was prepared by a procedure similar 
to that reported by Krumholz and Stettiner. A solution of 
0.80 g of KOH and 1.25 g of Ba(OH) in 50 mL of distilled 
water was degassed with N„ for one hour. Then 1.0 mL of 
Fe (CO) I- was syringed into the stirred solution. After one 
hour the yellow solution had turned orange in color. After 
an additional two hours the solution was red in color and a 
white precipitate had formed. Inside a glovebag, the red 
solution containing the K„Fe (CO) . was filtered away from the 
precipitate. The red solution was placed onto a vacuum line 
and the solvent evaporated to yield a brown solid. The 
solid was dried and stored under vacuum until used in the 
carbon dioxide experiments. The iron salt was characterized 
by infrared spectroscopy as shown in Table 2-6. 



23 

Preparation of Sodium Pentacarbonylmanganate ( 1-) and Sodium 
Pentacarbonylrhenate (1-) 

The NaM(CO)c. salt (M = Mn , Re) was prepared by a 

69 
procedure similar to that reported by King and Eisch. 

Seven milliliters of mercury was added to a nitrogen purged 

reaction flask. A total of 0.50 g of sodium metal was slowly 

added to the stirred mercury. Inside an inert nitrogen 

glovebag a solution containing either 2.98 g of Mn- (CO) , q or 

1.00 g Re„ (CO) ^ „ in 50 mL of acetonitrile or tetrahydrofuran 

was prepared. The solution was quickly added to the stirred 

amalgam. After 2 1/2 hours the stirring was stopped and the 

excess sodium-mercury amalgam was removed from the reaction 

flask. The resulting army green NaMn(CO)^ and orange 

NaRe(CO) solutions were used in the carbon dioxide 

experiments. The manganese and rhenium salts were 

characterized in solution by infrared spectroscopy as shown 

in Table 2-6. 

Preparation of Potassium y-Hydridobis (pentacarbonyl- 
tungsten (0) ) 

The KH[W(CO)j.]„ was prepared by a procedure similar to 

70 
that reported by Grillone and Kedzia. Inside a glovebag a 

solution containing 5.63 g of W(CO)g and 6.3 g KOH in 7.5 mL 

water, 30 mL methanol and 70 mL tetrahydrofuran was prepared. 

The solution was heated to 45°C for five hours and then 

continued stirring at room temperature for an additional 14 

hours. The reaction solution was filtered and the resulting 

filtrate placed onto a vacuum line. The solvent 

was removed to yield a yellow paste. A total of 82 mL of 



24 



Table 2-6. Characterization of Prepared Transition 
Metal Carbonyl Anions by Infrared 
Spectroscopy 



Compound 



KCo(CO) 



KCo(CO) 



K(HFe(CO) J 
4 

HFe(CO) " 



Fe(CO)^ 
H,Fe(CO) 



Fe{CO), 



Infrared Data (cm~ ) Environment Reference 



1890 (vs) 




THF 


a(I) 


1890(vs) 


, 1857(w) 


THF 


71 


1915(m) , 


1887(vs) 


acetonitrile 


a(II) 


2015(w) , 
1897 (vs) 


1937(sh) , 


water 


72 


1786(vs) 




water 


72 


2121(w) , 


2111 (vw) , 


hexane 


73 



2053 (TO) , 2042(s) , 
2029(vw), 2010(m) 

2020(vs) 



NaMn (CO) 


1910(vs), 1860(vs) 


acetonitrile 


a (Fig. 


2-5) 


KMn(C0)5 


1896(s) , 1862(s) , 
1830(m) 


THF 


71 




«"2<C°>10 


2045(s) , 2009(vs) , 
1978(s) 


THF 


a (III) 




HMn(CO) 


2117, 2043, 2015(vs) , 


cyclohexane 


75 





2008(vs), 1981 (vs) , 
1966 



NaRe(CO) ^ 


1900(m), 1860(m) 


THF 


a (IV) 


KRe(C0)5 


1911 (s), 1864(s) , 
1835(sh) 


THF 


71 


Re^fCO)^^, 


2008(ms), 1972(s) 


THF 


a (IV) 


HRe(CO) 


2131, 2123, 2053, 
2042, 2015(vs) , 
2005(vs), 1982(vs) 


cyclohexane 


75 



a - this work; (I-IV) location of spectrum in appendix A; 

vs - very strong; s - strong; ms - medium strong; m «= medium; 

w " wealc; vw « very weak; sh = shoulder. 



25 

water was added and the mixture stirred for 2 4 hours at room 
temperature followed by 23 additional hours at 0°C in an ice 
bath. The mixture was filtered to obtain white crystals 
which were dried and stored under vacuum until used in the 
carbon dioxide experiments. 
Preparation of a Mixture of Rhenium Carbonyl Hydrides 

The mixture of rhenium carbonyl hydrides was prepared 

by a procedure similar to that reported for the formation of 

7 fi 
H2Re-(C0) . A solution containing 2.0 g of Re_ (CO) and 

1.6 g of NaBH. in 50 mL of tetrahydrofuran was refluxed for 

4 hours. The solution went through a sequence of color 

changes from yellow to orange and finally to red after 4 

hours at reflux temperature. The solution was decanted away 

from the NaBH and the solvent evaporated. The resulting 

solid compound was dried under vacuum for several days. A 

solution containing 80 mL of cyclohexane and 10 mL of H^PO. 

(deaerated and dried by adding several drops of Na-Hg 

amalgam) was added to the solid compound in the reaction 

flask. After 6 hours at reflux temperature, the solution 

was extracted several times with hot cyclohexane. Cooling 

the cyclohexane solution did not precipitate the desired 

product as reported. Thus the solvent was evaporated to 

yield a brown solid which was characterized by infrared 

spectroscopy to be a mixture of H Re. (CO) , H Re (CO) and 

Re2(C0)^Q. 

Reaction of the Transition Metal Carbonyl Salts with Carbon 
Dioxide 

Acetonitrile or tetrahydrofuran solutions of the 

transition metal carbonyl salts were reacted with carbon 



26 

dioxide at atmospheric pressure by bubbling the gas through 

the solution. Higher pressure experiments were performed by 

ft ft 
using a Parr pressure bottle system containing the 

solution of the transition metal carbonyl salt and carbon 

dioxide. The resulting color changes for these reactions 

are summarized in Table 2-7. The reaction products were 

examined by infrared spectroscopy. 



Table 2-7. Color Changes for Reactions Between Carbon 

Dioxide and Transition Metal Carbonyl Anions 



Anion 



Initial Solution Color 



CO, 



-^•Final Solution Color 



KCo(CO) ^ 
K2Fe(C0)^ 
NaMn(CO) ^ 



NaRe(CO) 



Yellow 
Pink-red 
Army green 
Orange 



Yellow 
Orange 
Orange-red 
Yellow-green 



Reaction of Carbon Dioxide and Hydrogen with Transition 
Metal Hydrides in Alcohol Solvents 

A Parr pressure bottle reactor system containing a 

-4 
solution of 1.5 x 10 moles of the catalyst in 20 mL of 

methanol was charged with either carbon dioxide, hydrogen, 

carbon monoxide or some mixture of the three gases in equal 

parts while maintaining the total pressure at 20 psig. The 

solution was stirred and allowed to react within a 

temperature range of 125-150°C. The gaseous reaction 



27 

products were characterized by gas chromatography, while the 

liquid reaction mixture was monitored by both gas 

chromatography and infrared spectroscopy. 

Results and Discussion 

The Binding of Carbon Dioxide by Transition Metal Carbonyl 
Anions 

Following the increasing trend in the nucleophilicity 

(Table 2-3) of the transition metal carbonyl anions, the 

alkali metal salt of each anion was prepared and reacted 

with "bone dry" carbon dioxide. The reaction products were 

characterized by infrared spectroscopy as summarized in 

Table 2-8. The reaction of carbon dioxide with K[Co(CO)^] 

in tetrahydrofuran at atmospheric and elevated pressures 

(<70 psig) substantiated the literature reports of the 

50 
occurrence of no reaction. 

Characterization of the iron salt in acetonitrile by 

infrared spectroscopy (Table 2-6) determined the complex to 

be K[HFe(CO) .] instead of K Fe(CO)^. It is a common 

procedure to form hydridometal complexes from the 

protonation of metal complex anions with water as reported 

for several phosphine substituted metal carbonyl 

7 7 7 8 
anions. Although unsubstituted metal carbonyl anions 

usually require acidification to form a hydrido complex, it 

is reasonable to assume that the trace quantities of water 

observed by infrared spectroscopy to be present may be a 

strong enough acid in acetonitrile to partially protonate 

K2Fe(C0). to form K[HFe(CO)^]. The presence of this trace 



28 

quantity of water is a result of the synthesis of the iron 
salt in a water solution, as well as the difficulty in 
drying acetonitrile . Upon reaction of this acetonitrile 
solution of K[HFe(CO) .] with "bone-dry" carbon dioxide for 
several hours at atmospheric pressure, several infrared 
absorptions (Table 2-8) indicative of a reduced carbon 
dioxide species were observed. These new infrared 

absorptions correspond very well with the formation of a 

79 
small quantity of potassium bicarbonate. During the 

protonation of K_Fe(CO) . by trace quantities of water, the 

formation of potassium hydroxide is inevitable as shown in 

Equation 2-8. Carbon dioxide can be neutralized by this 

potassium hydroxide to form potassium bicarbonate. Carbon 



CH^CN 
K2Fe(C0)^ + H2O >-K[HFe(CO) ^] + KOH (2-8) 



dioxide also can be hydrated in the presence of water to 
form carbonic acid which can dissociate into bicarbonate and 
carbonate as shown in Equation 2-9. It previously has been 
suggested by the interpretation of solution infrared data 

that the reaction of Na^Fe (CO) . with carbon dioxide yields 

48 
iron pentacarbonyl and sodium carbonate. The preliminary 



CO2 + ^2'^< — ^'H^cor^:^ h"^ + Hrn "^— » 7w'^ + co^ (2-9) 



nature of these reported results has precluded any direct 



29 



Table 2-8. A Summary of Infrared Data Obtained 
for Reactions Between Carbon Dioxide 
and the Transition Metal Carbonyl Anions 



Compound + COj Infrared Data (cm"^) 
KCo(CO) 1890(vs) 



Assignment 
KCo(CO). (V) 



Environment 



K(HFe(CO) 



1915(m) , 1887(vs) 



K(HFe(CO).] (VI) 



3600 (mw), 3190 (mw), 1629 (m), KHCO, (VI) 
1345(w) , 699(m) ^ 



CH^CN 
^3' 



3^ 
CH,CN 



NaMn (CO) 



2044(m), 2015(6), 1985(s) 
1910(vs) , 1860(vs) 



Mn2(C0)jg (Fig. 2-5) CH^CN 

NaMn (CO) g (Fig. 2-5) CH CN 

1656(s), 1623(vs), 1047(m), NaHCO, (Fig. 2-4) nujol 
1033(m) , 996(s) , 833(s) , ^ 



703(s) 



NaMn (CO 
Mn- (CO) 



NaRe (CO) ^+ 



Re, (CO) 



10 



2045(m) , 2010(s) , 1975(s) 
1895(vs), 1855(vs) 

3460(w) , 2030(s) , 1935(w) , 

1667 (w) 



Mn2(C0)^Q (Fig. 2-6) THF 

NaMn (CO) J (Fig. 2-6) THF 

"not identified', THF 
(Fig. 2-6) 



1985(s) 
1880 (s) 



Re^iCO)^^ (VII) THF 
NaRe (CO), (VII) THF 



(V-VII) •: location of spectrum in appendix A; vs = very strong; s «= strong; 
m = medium; mw = medium veetk; w = wea)c. 



30 

comparisons of the reaction conditions and infrared data 
with the corresponding conditions and data observed for the 
reaction of K[HFe(CO) .] with carbon dioxide. The 
determination through infrared spectroscopy (Table 2-8) that 
K[HFe(CO) .] was the only iron species present after 
completion of the reaction supports the formation of 
potassium bicarbonate from the reaction of carbon dioxide 
with a reaction contaminant, such as water or potassium 
hydroxide . 

Characterization of an acetonitrile solution of 

53 
NaMn(CO)c (neucleophilicity = 77) by infrared spectroscopy 

(Table 2-6) showed the absence of Mn2(C0),j., as well as the 

absence of any residual water. This solution was reacted 

with "dry" carbon dioxide at atmospheric pressure to yield a 

solution color change coinciding with the precipitation of a 

solid. This solid was identified by infrared spectroscopy 

to consist of primarily sodium bicarbonate with possibly the 

presence of a trace amount of sodium carbonate as shown in 

Figure 2-4. The reaction between carbon dioxide and 

NaMn(CO)c in tetrahydrofuran previously has been reported to 

yield sodium bicarbonate and an unidentified manganese 

carbonyl complex. Interpretation of the infrared spectrum 

of the reaction solution, which is shown in Figure 2-5, 

suggests that another manganese carbonyl complex is present 

along with NaMn(CO)j.. The infrared absorptions (Table 2-8) 

assigned to this manganese carbonyl complex correspond to 

those of Mn„(CO),„. An increase in the quantity of 



31 



Sanple spectrun 



a 


= 


1656 


cm 


b 


= 


1623 


cm 


c 


= 


1047 


cm 


d 


= 


1033 


cm 


e 


= 


996 


cm 


f 


= 


833 


cm 


g 


= 


703 


cm 




-1 
-1 
-1 
-1 
-1 
-1 
-1 



wavena±»ers (cm ) 



NalEO 




-1 



Figure 2-4, Infrared Spectrum of Precipitate 
Obtained from Reaction of Carbon 
Dioxide and NaMn(CO)i. 



32 



CO, 



48 Hours 
CO, 



UJ 

o 

z 
< 

CC 
O 
<0 
00 

< 



JL 




V_^ 



2000 



2000 
WAVENUMBERS 



2000 



d = 1910 on 
e = 1860 cm 
f = 2044 an; 
g = 2015 an' 
h = 1985 an" 



-1 
'-1 

'-1 
'-1 
'-1 



Figure 2-5. Solution Infrared Spectrum of the Carbonyl 
Region for the Reaction between Carbon 
Dioxide and NaMn(CO)_ in Acetonitrile 



33 
Mn_ (CO) ^ . present in the reaction solution was observed in 
these infrared data to coincide with an increase in reaction 
time. The reaction of NaMn(CO)c. and carbon dioxide in 
tetrahydrofuran instead of acetonitrile produced similar 
results. In this case, the presence of a third unidentified 
manganese complex along with Mn„ (CO) , „ and NaMn(CO)j- was 
observed by infrared spectroscopy as shown in Figure 2-6. 
The infrared absorptions at 2030 and 1935 cm assigned to 
this unidentified complex were observed to dissappear upon 
the replacement of the carbon dioxide atmosphere with 
nitrogen. The weak 'V^u ^^^ l^r-m absorptions that were 

OH ' C,0^ 

observed could be due to either a small quantity of 
solubilized bicarbonate or to a bound formato, bicarbonate, 
or metallocarboxylic acid complex of manganese. The 
identification of a metallocarboxylic acid derivative, 
[ (OC) ^M(C02Na) ] 2 (M = Mn, Re) from solution infrared data 
has been reported for the reaction of carbon dioxide with 
NaM(CO)i- (M = Mn, Re). Attempts to isolate a bound carbon 
dioxide complex have been unsuccessful. 

The interaction of carbon dioxide with NaMn (CO) could 
proceed through a variety of different pathways. First, it 
is possible that carbon dioxide directly interacts with 
NaMn (CO) c- forming a metallocarboxylic acid derivative as 
previously suggested. The observed reaction also could 
proceed through the disproportionation of carbon dioxide 
into carbon monoxide and carbonate as shown in Equation 
2-10. It is possible that the insertion of carbon dioxide 



34 




'g'g'g'g'i'g'g'i'g 

oinooLninininr^ 
■voooo>crioooo«3 

II II II II II II II II II 

(0X1 UTD QJiM Cnj::-H 



3 



4 HDNvaaosav % 







c 






(U 






0) 






5 






■p 






OJ 






XI 


- 




c 













•H 






■p 






o 


o 




(d 


o 




(U c 


CJ 




ex; (0 






■P 

> 

9 -P 

U 0) 


8 




1^^ 


o 




(U c 


€>4 




fi-H 




CO 


CO 




nr 






UJ 


-a — . 




m 


,_ <1> o 




2 


1 (TJ "-' 




D 


5 ^ c 




LU 

> 






< 




o 


^ 


TT -d 


o 




OJ c 


00 




+J (0 


CM 




o 
Id 0) 

U 'G 
■P -H 

U3 ^ 
P o 

CO -H 
Q 

c c 

<U 
>J3 


o 




rH ^ 


o 




<a 


(O 




en. cj 


CO 




3 
CJ1 

•H 



35 
into the sodium-manganese bond could form an intermediate 
similar to that isolated for IrCl (C„0 . ) (PMe^ ) --0 . 5»C,H, . '^^ 

^4 J J DO 

The disproportionation of this intermediate into carbonate 



2CO2 + 2e~ »^C0 + CO^^" (2-10) 



and carbonyl complexes could explain the observed infrared 
data. A third alternate way in which a carbonate or 
bicarbonate species could be formed is by the interaction of 
carbon dioxide with water as shown in Equation 2-9. 
However, the absence of any observable 0-H absorptions in 
the infrared data obtained for the starting solution (Table 
2-6) suggests that water is initially not present in the 
reaction. It is possible that the reverse of the water-gas 
shift reaction as shown in Equation 2-6 could produce the 
water necessary to initiate the formation of bicarbonate and 
carbonate species. Hydrogen was found to be present in the 
reaction as a low level impurity arising from the "dry" 
carbon dioxide feed gas. It is impossible to ascertain from 
the available data which of these mechanisms is 
predominantly responsible for the observed results. 



Reaction between carbon dioxide and a mixture of 

53 . 
Re2(C0)^Q and NaRe(CO)c (nucleophxlicity = 25,000) m 

tetrahydrofuran could not be detected by infrared 

spectroscopy. Infrared spectroscopy has been inconclusive 

in ascertaining the existence of any rhenium complex besides 

Re (CO) or NaRe(CO) . Attempts to stabilize or isolate 



36 

any reduced carbon dioxide species has been unsuccessful. 

Even though the results concerning the reaction between 

carbon dioxide and NaRe (CO) j. have been inconclusive, the 

formation of sodium carbonate and an unidentified metal 

carbonyl species has been reported for the reaction of 

carbon dioxide with the more nucleophilic complex, 

Na[CpNi(CO)] (nucleophilicity = 7,500,000)^"^. 

Formation of Alkyl Formates at Low Pressures and 
Temperatures 

The activation of hydrogen by a transition metal 

carbonyl complex is necessary to effectively utilize the 

corresponding transition metal carbonyl hydride as a 

catalyst for the reduction of carbon dioxide. The 

activation of hydrogen by Mn„ (CO) ^ - has been reported to 

occur only under extreme conditions of pressure and 

8 
temperature . On the other hand, Re^ (CO) ^_ has been 

reported to activate hydrogen at atmosphereic pressure under 

mild temperature conditions to form a mixture of H Re (CO) _ 

8 1 
and H.Re.(CO),-. Low temperatures and pressures have been 

reported to be effective for the activation of hydrogen by 

Ru,(CO),- and neutral group 6B metal complexes, such as 

W[P (OCH- ) ^] j.H„ . However, the current method for the 

formation of group 6B metal carbonyl hydrides is the 

reduction of the hexacarbonyl metal complex by either 

82 57 

NaBH. or two equivalents of KOH in aprotic solvents. 

Dodecacarbonyl dirhenium was tested as a catalyst for the 

conversion of carbon dioxide and hydrogen to methyl formate 



37 
in methanol at mild temperatures (125-150°C) and pressures 
(20 psig) . It should be noted that the major difference 
between the reactions conducted in this work and those 
previously reported in the literature is the 
utilization of substantially lower pressures in the present 
work. The reaction products were monitored by gas 
chromatography and infrared spectroscopy. 

The quantitative results of the experiments conducted 
are summarized in Table 2-9. The formation of methyl 
formate was observed by gas chromatography and identified by 
GC/MS as shown in Figure 2-7. Since only a trace amount of 
methyl formate was discovered in the reaction of Re_(CO),„ 
with carbon dioxide and hydrogen (run 1), no quantitative 
data were obtained. A control reaction (run 2) involving 
only methanol, carbon dioxide and hydrogen showed no 
activity for the formation of methyl formate. Analysis by 
GC/MS of the reaction mixture (run 1) identified the 
existence of several additional low molecular weight 
products, such as dimethyl ether, dimethoxymethane and 
hexane. The observed mass spectra of these compounds are 
shown in Figures 2-8, 2-9 and 2-10. Gas chromatography has 
been used to obtain quantitative data for these compounds, 
as well as discover the presence of methane and an 
unidentified substituent at 1.21 minutes as shown in Figure 
2-11. Trace amounts of dimethyl ether and dimethoxymethane 
were observed by gas chromatography to be the only products 
formed in the control reaction (run 2) . The hexane observed 



38 



Table 2-9. A Summary of the Quantitative Data 
Obtained for the Reaction Products 
by Gas Chromatography 



Seactjon Reaction 

Bun Time in 
Number (Hrs.) Methanol* 

1 24 "^2(00130 

* CO, ■» H, 



Jles X 10~ moles x 10~ moles x lo'" 



41 
2.8 



weak trace 



SRe^CCOIjo 


. CO, . H, 


"^z'^^o'io 


+ CO 


Re, (CO. J „ 


+ CO + H, 


Re, (CO) JO 


+ CO + KOCH 


Re2(C0)jj, 


+ CO + H, 



weak trace 



KOCH, ■► CO 



X y 

♦ CO 



weak trace 
1.5 



+ CO 

48 (H,Re(CO)^]" 15 

+ CO + 
140 |Re2(CO)g(i,-OCH3)^)" 19 



• • In all reactions the catalyst concentration was 7.5 x 10 Boles/liter. 



39 



IOOt 31 

observed spectrum 



0) 

2 50 



^ 1 



I I I I I I I I I 1 I I I I I I M I I I I I I I I I I I I I 

10 ' 30 

m/e 



60 



'' M I I I I I I I I I I I I I I I I I I I I I I I I 



50 



70 



reference spectrvm^"^ 



10 30 ro 70 90 

To/e 



Figure 2-7. Mass Intensity Report for Methyl Formate 



40 



100n 



I 



0) 



45 



observed spectnan 



29 



'l I I I I 1 I I I 1 I I I I ' I I I I 



10 



20 



30 

itv/e 



I I I I I I I I I I I I I I 
40 



50 



reference spectrum' 



83 



J 



10 30 50 70 



150 



Figure 2-8. Mass Intensity Report for Dimethyl Ether 



41 



100 



I 



50 



45 



observed spectrum 



29 



|TTTT 

10 



flili |liii i iill|| i i ii i i M | n il 
30 



75 



H"| i iii ii i i i|iiii i ii i i[ in i 



50 

la/e 



70 



90 



100 



reference spectrum 



83 



10 30 50 70 90 

n/e 



Figure 2-9. Mass Intensity Report for Dimethoxymeth 



ane 



42 



100 



57 



50- 



observed spectrum 



41 



86 



'|i I ml li 1 1 1 1 II II il I] 1 1 li II 1 1 1| I ii 1 1 1 11 l|i 1 1 II II II lui 1 1 1 1 1 1 |i 1 1 ii 1 1 1 1| 
20 30 40 50 60 70 80 90 

nv/e 



100 



ij 



reference spectrum' 

. i 



83 



20 40 60 80 100 

va/e 



Figure 2-10. Mass Intensity Report for Hexane 



43 




M-l 









a 









•H 




+J 




O 




IT) 




(U 




tf 




tP 




c 




■H 




^ 




3 




P 




C 




0) 


iH 


>i 





03 


c 


H 


(0 




x: 


Q) 


+j 


cH 


a) 


frS 


g 




(0 


c 


CO 


-H 


tn 


CN 


ns 


K 


C!) 






+ 


IH 




n 


CN 




O 


e u 


rtl 




u 


A 


en 4-1 





•H 


-p 


s 


rtl 




g 


o 


o 


rH 


u 


,-^ 


JZ. 


o 


u 


u 




^— ^ 


tn 


CN 


(0 


fl; 


o 


a 




44 
in the reaction mixture (run 1) is an impurity that arises 

from the use of hexane solvent in the commercial 

84 
recrystallization of Re^(CO),„. Since bulk grade hexane 

is used, the unidentified peak at 1.21 minutes can be 

assigned as another hydrocarbon impurity. This is supported 

by the observed proportional increase in this peak at 1.21 

minutes along with the hexane peak as the quantity of 

Re_(CO)^- used in the reaction is increased. Furthermore, 

both these peaks remain constant throughout the reaction 

period. The dimethyl ether and the dimethoxymethane that is 

observed can be considered as reaction products since they 

increase in concentration as the reaction time progresses 

(run 1). It also is observed that the Re^ (CO) - ^ 

concentration has no effect upon the quantity of dimethyl 

ether or dimethoxymethane produced during the reaction (run 

3) . 

It is possible to speculate that methyl formate in this 

reaction could be produced through a mechanism similar to 

49 
that previously described in Figure 2-2. Recently, the 

formation and characterization by x-ray crystallography of a 

rhenium carbonyl metallocarboxylic acid complex, 

8 5 
Re (CO) ^ .COOH, was reported. This metallocarboxylic acid 

complex of rhenium was formed as a minor product in the 

photolysis of Re (CO) ^ _ in the presence of nitric oxide. It 

is possible that an intermediate such as this could 

eliminate formic acid within a catalytic cycle. The 

interaction of this formic acid with methanol would yield 



45 
the observed methyl formate. The infrared data obtained for 
these carbon dioxide reactions are summarized in Table 2-10. 
The presence of Re„ (CO) and other rhenium carbonyl 
complexes is suggested by the interpretation of this 
infrared data (run 1). The presence of Re„ (CO) ^ . in 
solution was confirmed by the infrared characterization of a 
white solid that precipitated out of solution upon the 
addition of water. Subtraction of the Re_ (CO) ^ . component 
in the infrared spectrum (run 1) as shown in Figure 2-12 
allows for accurate determination of the absorptions which 
can be assigned to the appearance of new rhenium carbonyl 
complexes. The major components in this subtracted infrared 
spectrum have strong absorptions at 2009 and 1892 cm . It 
was noticed that during the reaction a pink film was 
observed to form along the glass reactor walls. Partially 
dissolving this film in carbon tetrachloride gave an 
infrared spectrum similar to the solution data (run 1) 

except for the absence of the absorption at 1892 cm 

8 1 
Since Re2(C0),Q has been reported to activate hydrogen 

within the reaction conditions employed, the infrared 

absorptions observed for this pink film can be assigned to a 

mixture of Re_ (CO) ^ _ and a rhenium carbonyl hydride. 

6 2 
Recently it was reported that the reduction of carbon 

monoxide by tungsten carbonyl hydrides in methanol 

catalytically produced methyl formate. The moles of methyl 

formate formed per mole of catalyst used in these 

experiments were shown to be approximately two orders of 



46 



c o 

O 10 
J3 U 
U 4J 
US V 
O 4J 







^^ 


U) 


,-^ 




^^ 


^^ 




U) 







^^ 




to 




tfl 







3 


^^ 






6 


> 


e 




in 


E 




> 




3 


_E 




> 




> 




_E 


> 


E 












































10 




ir> 


t— 1 


00 




CT^ 


a\ 




^r 




a\ 


^j< 




(N 




in 




m 


00 


f^ 


•p 




^ 


<-H 


^^J 




o 


(N 




r-H 




(N 


r^ 




,-H 




o 




00 




00 


ID 
Q 




o 


o 


- 




o 


(TV 




O 




O 


Ol 




o 




o 




o 

(N 


o 
rg 


ON 


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* 


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—^ 


« 




•. 


^^ 




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^ 


CI 




— » 


— 


— ^ 


— ^ 


-^ 


'^ 


(0 


.— k 


.—^ 


^m. 


10 


^^ 


„^ 


..^ 


(n 


u> 


^_^ 


U) 


^^ 


u 


— * 


e 


e 


E 


u 


E 


5 


> 


<n 


n 


e 


^ 


_e 


a 


a 


^ 


> 


B 


> 


m 


10 










































L^l 




CN 


rH 


f-~t 


•— 1 


«-* 


<T> 


fN 


o 


VS 


cH 


■<»■ 


^ 


(N 


,_( 


IT) 


c^ 


n 


o 


00 


*4-< 


E 


r^ 


m 


r~- 


CTl 


fn 


<J\ 


Oi 


r~ 


r^ 


r- 




m 


i~- 


r^ 




00 


eyv 




o 


c 


o 


o 


o 


o\ 


00 


o 


17\ 


00 


o 


<T\ 


o 


o 


<j> 


o 


ON 


O 




o 


o 


o 


H 


~ 


fN 


(N 


•-' 


'-' 


IN 


•-' 


"-< 


<N 


•-I 




<N 




(V 






^ 


rM 




tN 



4J rt 


o 


R 




•U T! H 


o — 




3 






<a o 


X 't-i 


l-i 




a m 4J 


tJ u 


U 


•y 


• 4) 


•H 10 


jj — 




u 


4J 


c 3 


J3 <N 


•a 4J 


(V 


■H 


ID 


3 01 


C 3 


a 


J= 


U 0<<H 


w « 


■0 


01 


a 


a 3 



I II 



47 




cc 
ai 

OQ 

3 



3 



4 NOlSSIWSNVai % 



•H 

c 
o 

•H — ^ 



o a 
It) o 

(U 
(N4-J 

K O 
\ fO 
<N ^( 

o -u 



M-l C 

(1) 

W 0) 
-P J2 

u 

-d > 

O (T3 

^^ jC 
o 

C <H 

o ^ 

■H O 

-u u 
u ^ 

(1) 0) 







2 


M-l T) 




CO 


LU 


c 




o 


> 


(0 




o 


< 


£ 




C4 


5 


=1 K 
U O 
-P (^ 
U K 
0) U 








c 

T) rH 

0) ^ 




O) 




U O 




<o 




(0 u 




^" 




M '-' 




o* 




14-1 CN 

c a; 
H a 


u 


5 


CN 


<o 




1 

u 








Di 



48 
magnitude greater than those reported for the corresponding 
reactions run under carbon dioxide and hydrogen. The small 
observed activity in the Re2(C0),Q reactions run using 
carbon dioxide and hydrogen can be explained by the 
reduction of either the metal bound carbonyl ligands or of 
carbon monoxide produced by the reverse of the water-gas 
shift reaction as shown in Equation 2-6. This latter case 
is possible since both Re^iCO) ^^ and H^Re^fCO),- have been 

reported to catalyze the water-gas shift reaction under 

8 7 
basic conditions. Although the formation of methyl 

formate was still too small to quantitate, it was 

demonstrated by gas chromatography that more methyl formate 

was formed under a carbon monoxide atmosphere (run 4) . The 

yield of .methyl formate was found to increase with the use 

of a mixture of carbon monoxide and hydrogen over an 

extended reaction time (run 5) . The catalytic ability of 

the Re2(C0)..Q system at low pressures will not be discussed 

_7 
since the 2.6 x 10 moles of methyl formate formed per mole 

of Re- (CO) J- used is far below one catalyst turnover 

assuming Re2(C0)^„ to be the active catalyst. 

Infrared spectroscopy was used in an attempt to 

determine the active species in the Re„ (CO) , ,, reaction 

systems. All infrared data obtained for reactions using 

carbon monoxide are summarized in Table 2-11. The infrared 

spectrum of the reaction run under carbon monoxide (run 4) 

shows the presence of only one absorption at 1890 cm that 

cannot be assigned to Re^ (CO) .. . This suggests that the 



49 



Table 2-11. A Summary of Infrared Data Obtained for 
the Reaction of Re2(C0)iQ with Carbon 
Monoxide in Methanol 



Run 
4 



Reaction 

or 
Complex 

Re_(CO) 



CO 



'2 '""'10 

Re, (CO),. + CO + KOCH. 
2 10 J 



Re, (CO) 
3 



-2 
KOCH 



KOCH, + CO 



H Re (CO) 
X y ; 



Subtract 
out 
Re, (CO) 



Infrared Data, 



2071(8) , 2013(V3) 
1971(3), 1891(10) 

2071(m) , 2013(s) , 

1972(in) , 1888(s) , 

1732 (w) , 1717(w) , 
1605(vs) 



2071(w) , 2013(8) , 
2000(s) , 1971 (w) , 
1888(V8),1606(vs) 

1735(ni) , 1718(m) , 
1605(vs) 

2072(m) , 2031(s) , 
2010(vs) ,1971(m) , 
1927(vs) ,1891(vs) 

2031(s) , 2008(3) , 
1928(vs) ,1892(vs) 



-2' '10 
[(CO)3Re(M-OCH3)3Re(CO)3]' 



Environment 
methanol 

methanol 



methanol 



Reference 
a (XII) 

a (XIII) 



a (XIV) 

a (XV) 

a (XVI) 

a (Fig. 2-13) 



[H2Re(C0)^) 

[Re(CO),OCH,) 
-3^- "'3' 



1990 (s), 1875 (vs) dichloromethane 88 
dichloromethane 88 



2020(vw) ,1995(w) , 
1930(V3) ,1895(s) 



3 "3 '4 
[Re,(u-H),(vi-OCH,) (CO) 



2036, 1935 



THF 



89 



2096(w) , 2020 (m) 
2000(vs) ,1985(sh) , 
1957(vs) ,1935(vs) , 
1888(3) 



dichloromethane 90 



a = This wor)c; (XII-XVI) = Location of spectrum in appendix A; 

vs = very strong; s = strong; sh » shoulder; m = medium; w = weak; 

vw = very wea)c. 



50 
observed absorptions at 2031 and 1927 cm in reactions run 
in the presence of hydrogen can be assigned to the formation 
of a rhenium carbonyl hydride complex such as H Re_(CO)^„. 
Since methyl formate formation was enhanced in the presence 
of hydrogen over an extended reaction time (run 5) , it is 
proposed that the active species in the reaction is a 
rhenium carbonyl hydride. This proposal is further 
supported by the enhanced formation of methyl formate in the 
reaction of carbon monoxide with a mixture of rhenium 
carbonyl hydrides (run 9) . The resulting infrared spectrum 
for this reaction, which is shown in Figure 2-13, resembles 
those obtained for reactions between Re2 (CO) ^^ with hydrogen 
and either carbon dioxide (run 1) or carbon monoxide (run 
5) . 

Recall that in the previously proposed mechanism 
shown in Figure 2-4 for the carbonylation of methanol to 
methyl formate by tungsten carbonyl hydrides, the active 
catalytic species was suggested to be a methoxy tungsten 
carbonyl complex. Recently the bridging methoxy compound 

[ (CO)2Re(y-OCH2)^Re(CO)^] [N(C2H^) ^] was reported to be 

8 8 
formed by the addition of methanolic KOH to Re^{CO) ^q. 

Along with the hexacarbonyl tri-y-methoxydirhenate (1-) the 

reaction was found to form [H2Re(C0)^] as a coproduct. The 

complete conversion of this rhenium hydride to the isolated 

bridging methoxy compound [ (CO) gRe2 ( y-OCH^) ^ ] ~ was observed 

to occur at elevated temperatures. The formation of other 

alkoxide rhenium carbonyl complexes, such as 



51 




^NOISSIWSNVai % 



52 
[Re(CO) ^OCE^] ^,^^ [Re^iM-E) ^iu-OCE^) (00)^^]",^° 

[ (CO) gRe2(u-OC2H^) ^(U-OCH^) ]~ ^^ and [Re^H^ (V3-OC2H5) - 

- 92 
(CO)q] also have been reported. The preparation of 

these alkoxy rhenium carbonyl hydride complexes was 

92 
reported to proceed through a reaction of a rhenium 

carbonyl hydride with the corresponding alcohol. The 

similarities between the infrared spectra for these 

complexes and the infrared data obtained for the Re„(CO),„ 

reactions (Table 2-11) suggests that the 1890 cm" 

absorption can be assigned to the formation of a methoxy 

rhenium carbonyl complex such as [ (CO) ^Re„ (y-OCH-) _ ] ~. This 

is further supported in that no new infrared adsorptions are 

observed for the addition of several equivalents of KOCH^ to 

the Re„(CO)^- reactions- run under either carbon monoxide 

(run 6) or carbon monoxide and hydrogen (run 7), Since gas 

chromatography could not identify the formation of any 

methyl formate in these reactions (runs 6, 7), the methoxy 

rhenium carbonyl complex is most likely catalytically 

inactive. The infrared spectrum of the reaction performed 

under carbon monoxide and hydrogen (run 7) was observed not 

to exhibit any absorptions that could be assigned to a 

rhenium carbonyl hydride. This suggests that deactivation 

of the active rhenium carbonyl hydride results in the 

formation of an inactive methoxy rhenium carbonyl complex. 

Attempts to isolate this methoxy rhenium carbonyl complex 

have been unsuccessful. 



53 
A mixture of [ (CO) Re (y-OCH ) Re (CO) ] ~ and H Re(CO) ." 

in methanol was prepared by the previously reported 

8 8 
procedure and identified by infrared spectroscopy as shown 

in Figure 2-14. The formation of methyl formate was 

observed to occur for the reaction (run 11) of carbon 

monoxide with this mixture of rhenium complexes. The 

activity observed for the formation of methyl formate was 

found to decrease with an increase in reaction time. The 

resulting solution was observed by infrared spectroscopy to 

contain only the [Re- (CO) , (y-OCH^) ^] ~ complex as shown in 

Figure 2-14. It is proposed that a mixture of rhenium 

carbonyl hydrides, such as H^Re^ (CO) , _ ^"d H_Re(CO) ." are 

the active species responsible for the formation of methyl 

formate from carbon monoxide and methanol. These hydride 

species can be converted under the employed reaction 

conditions to the resulting inactive rhenium alkoxy carbonyl 

complex [Re2 (CO) g (p-OCH^) t] ~. The only infrared absorption 

in Figure 2-14 that cannot be assigned to either a rhenium 

carbonyl hydride or alkoxy complex is the medium strength 

band at 1605 cm . It has been observed by infrared 

spectroscopy that a similar absorption results from the 

addition of KOCH- to methanol (run 8) . 

At elevated pressures the carbonylation of methanol to 

methyl formate has been reported to occur using sodium 

93 94 
methoxide as catalyst. ' A comparison of the low 

pressure formation of methyl formate by this reaction (run 



54 



CO 




a = 2038 C3n 



-1 



b = 2009 C3T1 



-1 



c = 1992 cm 



-1 



d = 1971 an 



-1 



e = 1926 cm 



'-1 



f = 1880 an 



-1 



g = 1605 an 



-1 



2301 



2019 



wavenumbers (an 



1788 
-1. 



1557 



Figure 2-14. Infrared Spectrum Obtained for the Reaction 
of Carbon Monoxide With a Mixture of 
[Re2{C0) 6(y-OCH3) 3]- and [H2Re(CO)4]- in 
Methanol 



55 
8) with the Re2 (CO) ^ (run 4) and KH[W(C0)^]2 (run 10) 
systems was done. The results as shown in Figure 2-15 
indicate that Re^ (CO) , „ > KH [W (CO) -l- > KOCH, in activity 
for the carbonylation of methanol at low pressures to form 
methyl formate. This sequence of activity at low pressures 
parallels a recent report which indicates that KH[W(CO)c.]p 
is more active than KOCH, for the carbonylation of methanol 
to methyl formate at elevated pressures. An investigation 
of the activity of the Re (CO) ^ system at elevated 
pressures was not done because of the lack of a high 
pressure reactor. 

The utilization of low pressures in the Re (CO) 
system has allowed the identification of dimethyl ether and 
dimethoxymethane which may be key intermediates in the 
formation of methyl formate. A discussion concerning the 
mechanism for the formation of methyl formate in the 
Re (CO) /CH OH/CO system would be speculative and premature 

at this time. However, it should be noted that formaldehyde 

95 
has been reported to form dimethoxymethane in methanol and 

96 
methyl formate in the presence of a nickel catalyst 

Although a rhenium carbonyl bound formaldehyde complex has 

not been observed, both its precursor, a formyl complex, 

— 97 9 8 
such as [Re2 (CO) g (CHO) ] ' and its product, an 

99 
alkoxymethyl complex, such as Re (CO) ^CH20CH2 have been 

100 
reported. Another way in which dimethoxymethane, 

dimethyl ether^*^*^ and methyl formate ' ' have been 

reported to be formed is through the direct oxidation of 



56 




t^ 




Tune 



Timej^ 




BB^iCO)^^ + CO K[HW2(C0)^q] + CO 



RDCH3 + CO 



a = 0.77 minutes - diinsthyl ether 
b = 1.95 minutes - methyl formate 
c = 1.23 minutes - unidentified 



Figure 2-15. A Comparison by Gas Chromatography of 

Reaction Products in Liquid Samples for 
Carbon Monoxide Reactions with Re2(C0)-,^ 
KH[W(C0)5]2 and KOCH^ 



57 
methanol over a variety of different catalyst substrates. 
Carbon-13 carbon monoxide was reacted with Re_ (CO) ^ „ in 
methanol to ascertain if the methyl formate was being formed 
by the carbonylation of methanol or by methanol oxidation. 
GC/MS results so far have been inconclusive in obtaining the 
extent of carbon-13 incorporation in dimethyl ether, 
dimethoxymethane or methyl formate because of the relatively 
small amounts of products observed. It is proposed that the 
methyl formate observed in the Re (CO) systems is from the 
carbonylation of methanol as reported for the analogous 
KH[W(C0)c]2^^ ^^'^ KOCH^^"^' systems. This proposal is 
based upon the differences observed in reactions (runs 1-3) 
performed under carbon dioxide and those reactions (runs 4, 
5) performed under carbon monoxide. If the formation of 
methyl formate was governed by the oxidation of the methanol 
solvent, there should have been no differences in the 
observed results. 

Summary 
The main goal of this investigation was to evaluate the 
feasibility of binding and activating carbon dioxide by 
transition metal carbonyl complexes. The first study dealt 
with the interaction of carbon dioxide with a variety of 
transition metal carbonyl anions that differed in the 
nucleophilicity of the metal center. Although no reaction 
with carbon dioxide was observed to occur for KCo(CO)^, the 
formation of potassium bicarbonate was observed in the 
K[HFe(CO) .] system. This result was clouded by the reaction 



58 
of carbon dioxide with residual contaminants of water and 
potassium hydroxide. Interaction of carbon dioxide with 
NaMn(CO)t- was found to form sodium bicarbonate, Mn^ (CO) , „ 
and an unidentified manganese carbonyl complex. The 
occurrence of a reaction between carbon dioxide and 
NaRe(CO)j. was not observed. These results suggest that it 
is possible to bind carbon dioxide to a transition metal 
carbonyl anion. Although no conclusion concerning the 
preferred coordination mode of binding can be made, it does 
seem that the interaction of carbon dioxide with transition 
metal carbonyl anions is affected by the nucleophilicity of 
the metal center. 

The second study dealt with an investigation of the 
activity of Re (CO) with respect to the reduction of 
carbon dioxide in methanol to form methyl formate at low 
pressures. Reactions performed under carbon dioxide and 
hydrogen produced trace quantities of dimethyl ether, 
dimethoxymethane and methyl formate as identified by GC and 
GC/MS. Although these products could be formed directly 
from carbon dioxide, it is more likely that the reduction of 
carbon monoxide produced from the reverse of the water-gas 
shift reaction has occurred. This contention is supported 
by the increased activity for the formation of all products 
in reactions run under carbon monoxide. 

Infrared spectroscopy was used to investigate the 
active species in these reactions. It has been proposed 
through an interpretation of this data that the active 



59 
species is a mixture of rhenium carbonyl hydrides, such as 
E^Re^(CO) ^^ and [H2Re (CO) ^] ~. These hydride complexes 
slowly decompose during the course of the reaction to form 
the inactive methoxy rhenium carbonyl complex, 
[Re- (CO) g (y-OCH^) -] . Since an accurate measurement of the 
amount of the active species present during the reaction 
could not be obtained, a discussion of the Re_(CO)-„ 
system's catalytic ability was not undertaken. The 
formation of dimethoxymethane and dimethyl ether was 
observed during the course of the reaction. A comparison at 
low pressures of the Re- (CO) ^q system with systems, such as 
K[HW2(C0)^q] and KOCH^ , that are known to carbonylate 
methanol and form methyl formate was done. The results 
indicate that the activity of Re2(C0),Q > K[HW (CO)^-] > 
KOCH- in the carbonylation of methanol to form methyl 
formate. The possibility of increasing the activity of the 
Re-(CO)^. system at high pressures similar to that 
previously shown for KH[W(C0)^]2 ^^^ KOCH, may allow for a 
more thorough investigation of the reaction mechanism. 



CHAPTER III 
ACTIVATION OF CARBON MONOXIDE 
Background 
The conversion of synthesis gas (CO + H ) into organic 
substrates has been an active field of research since the 

initial work of Sabatier and Sendrens in the early 

104 
1900 's. Many articles have reviewed various processes 

that activate carbon monoxide, such as the water-gas shift 

and Fischer-Tropsch reactions. ~ . The latter has 

evolved for the hydrogenation of carbon monoxide and is 

112 
described in Equation 3-1. The Fischer-Tropsch synthesis 

usually employs a heterogeneous catalyst consisting of 



Catalyst 
(Fe,Co,Ru) 



CO + H^ *- Hydrocarbons + Oxygenates (3-1) 



either Fe, Co or Ru metal. This reaction takes place over a 
wide range of temperatures and pressures. Moderate 
temperatures and high pressures seem to favor the formation 

of oxygenated products while milder pressure conditions 

113 
increase the ratio of hydrocarbon products. In both 

cases the ratio of products obtained follows a simple 

polymerization model (Schulz-Flory distribution) as 



60 



61 
described by Equations 3-2 and 3-3. In Schulz-Flory 



W^ = n (l-a)^a^ ^ (3-2) 



Q = J^p + r^ = 1 (3-3) 

Zj. l-ot 



kinetics the weight fraction, W , of carbon number, n, is 

related to the probability of chain growth, a, which is 

defined in terms of the average degree of polymerization, Q. 

The value of Q, determined from the rate of polymerization, 

r , and the rate of chain termination, r^ , is influenced by 
p t "^ 

the characteristics of the metal catalyst and the reaction 
conditions, such as temperature and pressure. The inherent 
lack of selectivity as demonstrated by Schulz-Flory kinetics 
is the major disadvantage to a Fischer-Tropsch type 
conversion of synthesis gas. 

In order to deviate from this Schulz-Flory product 
distribution, thereby increasing the selectivity of the 

Fischer-Tropsch synthesis, research efforts have 

117 
concentrated on new loading techniques and the use of 

118 119 
shape selective supports. ' There is considerable 

evidence that the probability of polymerization in the 

Fischer-Tropsch synthesis is influenced by the size of the 

^ -1 -1 -^ 116,120,121 ^ . ^ ^ . 

metal crystallites. ' A variety of reasons for 

this particle size effect ranging from differences in the 

electronic band structure of small particles as compared to 



62 

that of the bulk metal to a stronger support interaction and 



higher degree of unsaturation with small particles have been 

122 
suggested. The final outcome has been the development of 

new methods for the preparation of small metal particles, 

123 
such as the solvated metal atom dispersed catalyst method 

124 
and the thermal decomposition of metal carbonyl clusters 

onto inorganic oxide supports. The utilization of these 

techniques has led in several instances to catalysts that 

exhibit higher activities and selectivities for synthesis 

gas conversion to C - C hydrocarbons as compared to 

125 
conventionally prepared catalysts. These dispersed metal 

catalysts deposited on high surface area supports are 

considered a new class of catalysts that lie between the 

boundries of traditional heterogenous and homogeneous 

catalysts . 

As shown in Table 3-1 there are several advantages and 

disadvantages associated with using either homogeneous or 

1 2 fi 
heterogeneous catalysts. The major disadvantage of 

heterogeneous catalysts besides their non-selective nature 

has been the lack of physical techniques to adequately 

characterize these systems. Recent advances in surface 

techniques, such as ESCA, SEM, XPS, Auger, etc. are 

beginning to aid in understanding and characterizing these 

catalyst systems. Homogeneous catalysts, on the other hand, 

are usually well characterized and reproducible. The major 

industrial concern for these catalysts is the additional 



63 

Table 3-1. A Comparison of the Advantages and Disadvantages 
of Using Homogeneous and Heterogeneous 
Catalysts 

Advantages Disadvantage's 

Homogeneous Catalysts 

1. Relatively resistant to 1. Necessary process 
catalyst poisoning step for catalyst 

separation 

2 . High activity 

2 . Temperature 

3. No mass-transfer problems sensitive 

4. High selectivity 

5. Characterization, reproducible 

Heterogeneous Catalysts 

1. Catalyst easily separated 1. Sensitive to 

from substrate catalyst poisoning 

2. Insensitive to high 2. Mass transfer 
temperatures problems 

3 . Characterization 



process step that is necessary to separate the reactant/ 
product/catalyst mixture. 

Another method to obtain a more selective process 
focuses around the utilization of homogeneous catalysts for 
the hydrogenation of carbon monoxide. A vast body of 

literature has developed for the conversion of synthesis gas 

,,.-,■ ^ ^ -1 ^ 127-131 
to oxygenated products by solubilized catalysts. 

These systems usually operate under extreme pressure 

conditions (>1000 atm.). Recently there have been several 

reports of anionic ruthenium carbonyl complexes being 



64 
effective in the homogeneous reduction of carbon monoxide to 

ethylene glycol at moderate temperatures and 

132-133 
pressures. Moderate temperatures and pressures have 

also been reported for the conversion of synthesis gas to 

methanol using a neutral metal complex, Ru(CO) , as 

134 
catalyst. The addition of carboxylic acids to this 

reaction promoted the formation of glycol esters. Most of 

the homogeneous systems reported have formed oxygenated 

products from the hydrogenation of carbon monoxide. 

Although several non-catalytic systems have been 

reported that reduce a carbonyl ligand to hydrocarbon 

products, ~ there have been very few reports concerning 

the homogeneous catalytic reduction of synthesis gas to 

hydrocarbons. The first report, in 1976, employed selected 

metal carbonyl cluster catalysts, such as Os- (CO) and 

13 8 
Ir,(CO),^. Substitution of several of the carbonyl 
4 12 

groups in Ir . (CO) , „ by triphenylphosphine was found to 
increase the hydrocarbon production rate. This rate was 
further enhanced by dissolving the Ir^(C0)-j^2 i" ^" 

• r. ^- ■, A 139,140 

AlCl^-NaCl melt solvent as shown m Equation 3-4. 

This reduction of carbon monoxide was done under very mild 

conditions, 125-210°C and one atmosphere of pressure. 



Ir (CO) 1 atm. 

CO + H, — >■ C-.-C. alkanes (3-4) 

"^ AlCl^-NaCl 125°-210°C 

The introduction of metallic aluminum to this system was 

21 
found to increase the yield of hydrocarbon products. It 



65 
was proposed that the addition of this aluminum metal 
enhanced the formation of HAlX which bifunctionally induced 
the reduction of carbon monoxide. 

A kinetic investigation of the Ir . (CO) melt salt 
system at 175°C discovered the formation of a sustained low 
level concentration of methyl chloride which was proposed to 

be an intermediate in the formation of the hydrocarbon 

141 
products. This kinetic study proposed that the active 

catalytic species was a chlorocarbonyl iridium complex, such 

as IrCl(CO)^. A similar system using Os^ (CO) ^ _ in a BBr- 

melt to convert synthesis gas to hydrocarbon products has 

142 
been reported. In this case, the formation of a low 

level concentration of methyl bromide was detected. It was 

suggested that Os_(CO)^Br. was the active catalytic species 

in the reaction. The most significant contribution of these 

reports was to demonstrate the importance of a Lewis acid 

adduct with a metal carbonyl ligand. This bifunctional 



M-C50«'«'A1C1- 



activation leads to a weakening of the carbon-oxygen bond 
which allows the carbonyl group to be reduced under very 
mild conditions. 

The functionalization of reactive supports with 
discrete molecular catalyst systems also could utilize this 
concept of bimetallic synergism by exploiting the support as 
cocatalyst. These supported catalysts may stabilize and 



66 
increase the concentration of the catalytic active species 

which would allow the reduction of carbon monoxide to occur 

143 
under milder conditions. The immobilization of 

homogeneous transition metal catalysts on various polymer 

supports is currently a very active field in catalysis 

research. The techniques for the covalent or ionic 

attachment of discrete metal complexes to various types of 

144 145 
supports are well documented. ' These supported 

catalysts can be considered as "hybrid" catalysts which 

offer the advantages, such as high activity and selectivity, 

of homogeneous catalysts as well as the ease of 

product/catalyst separation associated with heterogeneous 

catalysts. Several reports have indicated an increase in 

activity and selectivity using a polymer bound catalyst as 

146 
compared to its molecular analog. For instance, the 

polymer bound analog of Vaska's complex, @ - (PPh^) ^'^rCl (CO) , 

catalyzes the hydrogenation of 1 , 5-cyclooctadiene at a 

faster rate than observed for the homogeneous reaction as 

147 
shown in Equation 3-5. 



IrCl(CO) (PPh^)^ X \ ^ / \ (3-5) 






H^ 170°C 



Prior work done concerning the catalytic behavior of 
supported complexes in Dr. Drago's research group led to the 
discovery of a unique system for the selective catalytic 
conversion of synthesis gas and HCl to methyl chloride by a 



67 

supported tetrairidium cluster under very mild temperature 

148 
(25°-100°C) and pressure (1 atm.) conditions as shown in 

Equation 3-6. The support of choice was an inorganic oxide 



3-OSiC^H .PPh„Ir . (CO) , , 

2H„ + CO + HCl ^—1 = ii* H-0 + CH-,C1 (3-6) 

^ 25 - 100°C, 1 atm. 



(alumina or silica gel) because of its high thermal 
stability and the availability of Lewis acid sites to 
promote a bifunctional interaction with carbon monoxide. 
The support was functionalized through a condensation 
reaction involving the hydroxyl groups of the support and 



the ethoxy substituents of the phosphinated silane linkage, 

149 
(C-HcO) ^SiC-H.PPh- as shown in Equation 3-7. The letter 

y represents the number of support 3-0-Si bonds between 



3-OH + (C2H50)3Si(CH2)2PPh2 benzene/p-dioxane > 3(-0-) ySi (CH.,) ^PPh^ + BC^H^OH 

(3-7) 



the support surface and the silane linkage. Evidence 
indicates that this number is dependent upon the 
concentration of the organosilane used. The remaining 
ethoxy groups have been reported to be hydrolyzed by the 
solvent to yield ethanol and Si-OH functionalities. The 
tetrairidium carbonyl cluster was immobilized on the support 



68 



through covalent attachment to the phosphine of the silane 
linkage as previously reported. ' The major difficulty 
in this synthesis, as shown in Equation 3-8, was to maintain 



(OH) 



3-y 



3-tOf SilCH )2PPh2 (OH), 

' CO, Zn, H,0 I -""y 



■^ 3i(OfySi(CH2)2PPh2l,Ir4(CO)^ ^^_g^ 



IrlCO^CKH^N-^-CH, 



CH OCjH.OH (z = 1, 2; X = 11, 10) 



adequate stirring during the reaction. A poorly active 
catalyst was reported in cases where complete mixing was not 
obtained. An infrared investigation of the supported 
tetrairidium cluster, as summarized in Table 3-2, was 
reported to result in the identification of a mixture of 
mono-phosphine and di-phosphine substituted clusters, 
3-OSi(CH2)2PPh2lr4(CO)^^ and 3- [OSi (CH^) 2PPh2] 2lJ^4 (CO) ^^ , 
respectively. 

The catalyst was initially tested in a 3:1 AlCl^-NaCl 
melt salt under similar conditions as reported by 
Meutterties et al. and Collman et al. for the 
Ir . (CO) , ^/AlCl^-NaCl system. A typical catalyst run 
consisted of using 0.7 g of the supported iridium catalyst, 
8 g of A1C1_ and 1.8 g of NaCl in a glass fixed bed reactor 
system. The catalyst was exposed to a 3:1 mixture of H2:C0 
at 145°C. The major products were identified by gas 
chromatography to be methane, ethane, and chloromethane 
which are similar to those previously reported 



69 



(U 




s: 




■p 




u 









<H 




T) 




0) 




■u 




u 


CO 


-P 1 


a m \ 





>1 


« 


rH 




(0 


(0 


+J 


4J 


m 


(0 


u 


Q 






iH 


-0 


>, 


Q) 


c 


U 





(0 JQ L 


>-i 


u 


m 


(0 ' 


c 


u •, 


H 






E 


Q) 


3 


x: 


H <* 


+J ^3 1' 




>iT3 



o ,^ 



70 
for the homogeneous system. ~ When the supported 
tetrairidium catalyst was filtered from the molten salt and 
the AlCl^-NaCl retested, there was no decrease in activity 
observed. It was proposed that the supported tetrairidium 
catalyst leached off the support to give the homogeneous 
Ir . (CO) ^ _/AlCl--NaCl system. However, it was observed that 
prior to melting of the AlCl^-NaCl, methyl chloride was 
produced. The production of methyl chloride at 25°C was 
seen to decrease with time. This decrease in production was 
assumed to be caused by the depletion of the AlCl^-NaCl. 
The addition of anhydrous HCl(g) to the reactant gas stream 
rejuvenated the activity of the system for chloromethane 
formation. The cycling between the addition of HCl (g) and 
the absence of HCl (g) was done several times with no 
detrimental effects to the catalyst. It was reported that 
exposure of the activated catalyst to oxygen caused 
permanent deactivation of the system. 

The presence of HCl(g) in the reactant stream should 
initiate an interaction with the remaining hydroxyl groups 
of the support making either Al-Cl or Si-Cl bonds and water. 
If this is the case, then the presence of AlCl^-NaCl may not 
be necessary for the reaction to occur. Both silica gel and 
alumina supported clusters were tested in the presence of 
HCl(g) and the absence of AlCl^-NaCl. In both cases after a 
15-20 minute incubation period methyl chloride was observed 
at 25°C with the same activity and selectivity as seen 
previously. This induction period was suggested to be due 



71 

to the interaction of HCl with the support hydroxyl groups. 

These Al-Cl or Si-Cl groups are believed to behave in a 

similar manner as that of AlCl^. Gas chromatography and 

GC/MS identified trace quantities of methane, ethylene, 

methyl chloride, ethyl chloride, acetaldehyde and methyl 

formate as reaction products. No difference was noticed 

H H 

CI CI 

I I + 2HC1 *► I I + 2H2O 

■ y \ \ ^ ^« « ^ "J" 



between the aluimina and silica gel bound systems in regards 
to activity or selectivity. 

The activities of the silica gel and alumina catalyst 
systems were shown to be dependent upon temperature. The 
activity of both systems increased with increasing tempera- 
ture. A slight deactivation of the catalyst was observed to 
occur at 100°C over a period of time. At temperatures below 
100°C the catalyst was observed to be stable for several 
days. The activity also was found to be affected by the 
concentration of HCl(g) in the reactant gas stream. The 
concentration of HCl(g) had to be kept at a minimum to 
insure catalyst stability. The comparison of metal loadings 
(% wt.) at various temperatures demonstrated that the 
catalyst activity increased with higher concentrations of 
iridium in the catalyst. It was noted that other factors 
besides the metal loading, such as support interactions, 
deactivation process and phosphine concentration, may 
influence this observed increase in activity. 



72 

Although a complete material balance wasn't obtained, a 
calculation using the amount of chloromethane produced 
relative to the other products at 100°C showed the reaction 
to be at least 99% selective for chloromethane. This wasn't 
considering a polar product that condensed along with water 
at the top of the reactor tube. Even though the existence 
of this compound was discovered using gas chromatography, 
the identity of the complex was not reported. 

All control reactions run with or without AlCl--NaCl 
showed either little or no activity for chloromethane 
production as summarized in Table 3-3. A conclusion drawn 
from these control experiments was that the tetrairidium 
cluster had to be supported through a phosphine linkage for 
catalytic methyl chloride production to occur. It was shown 
that Vaskas' complex bound to a support, 3MPPh ) IrCl (CO) , 
was slightly active for methyl chloride in the temperature 
range 25-100°C. This activity was far below that observed 
for the supported tetrairidium cluster. 

It was found that other halide sources, such as CI , 
HBr(g) and HCl(aq) could be substituted for the HCl(g) with 
no decrease in initial activity or selectivity. In the case 
of HCl(aq) the activity was observed to decrease with time. 
It also was found that chloromethane production could be 
changed to methyl bromide by substitution of HBr(g) for 
HCl(g) in an active system. However, the reformation of 
methyl chloride by the reverse substitution of HCl(g) for 
HBr(g) was observed not to occur. 



73 

Table 3-3. A Summary of the Reported Activity for the 
Supported Iridium Carbonyl Catalyst System 



- 


CH C H 
CH3CI ^ ^ 


- 


N.A. 


- 


N.A. 


Trace CH CI 


- 


N.A. 


- 


N.A. 


- 


N.A. 


- 


N.A. 


- 


Trace CH^Cl 


- 



Catalyst Activity 

25°C 100°C 145''C 
Ir^ICO)^^ + AlCl^-NaCl N.A. 

SG + AlCl -NaCl N.A. 

SG^>'PPH2 + AlCl^-NaCl N.A. 

AlCl^ 

Ir^(CO) iiPPh3 + AICI3 
Ir^ (CO) ^2 + Al + HCl 
Ir^ (CO) ^^PPh, + Al + HCl 
Ir (COCKPPh^) 2 + Al + HCl 
Al /^PPh2lr (CO)PPh 



N.A. = no activity; SG = silica gel; Al = alumina, 



A speculative mechanism for this reaction was proposed 
as shown in Figure 3-1. Presumably oxidative addition of 
hydrogen by the supported catalyst would generate a 
dihydride species. Interaction of the Al-Cl or Si-Cl groups 
with a bound carbonyl ligand of the iridium dihydride 
species could induce rapid hydride migration to form a 
formyl complex. This is similar to the alkyl migration onto 
a bound carbonyl group in a (CH^) Mn (CO) ^/AlBr^ system as 
shown in Equation 3-9. ' The reaction then could 



74 



5- = 



0/ 




/ 



-2-1 

Ol 









m 




e >i 




iH 




(-1 (C 




IW -P 




(0 




QJ U 




TJ 




•H rH 




V^ >i 




c 




rH 




x; ^ — 




U ^-1 00 




fl 'S' 




.H U <-! 




>1 




x: e 0) 




■P 3 u 




(U -rH C 


< 


S T) OJ 


1 
1 




1 

r-o-i 


H "W 


1 " 


0) 


2-0 


C 'V u 




Q) 


a 


•H 4-> 4-1 


^~L~. 


-P V^ 


tf 


(0 


r 


e a c 




s^ a 




3 -rH 




&4 w en 




en 




(1) (C -H 


n 


x: e 




+j (-1 M 




Q) 0) 




^ > cu 




o 




«4-i x: 




rH +J 




e u -H 




m ffi s 




•H 




C tJ Tl 




(0 C (1) 




x: (0 -p 




u c 




(U en -H 




S tJ n 




o a 




•a (u 




oj m OS 




en -H "-^ 




en 




a 0) e 




x: 0) 




)-l 4J 4J 




di C en 




>i >1 




<C M en 



75 
proceed through the formation of a Lewis acid stabilized 
formaldehyde complex or a hydroxymethyl type intermediate. 
Finally, the formation of support 3-0CH_ groups or support 
3-(CH^0H) groups in the presence of HCl(g) would produce the 
observed methyl chloride. 



CH 
(CO).Mn + AlBr^ 



,CH 
-*► (CO) .Mn-C 

Br ^0 

^Al^ 
Br'^ ^Br 



(3-9) 



The possibility of a methanol intermediate parallels 
the fact that the current industrial process for methyl 

chloride production involves the chlorination of methanol by 

155 
HCl as shown in Equation 3-10. Methyl chloride is a 

major commodity chemical with consumptions in the range of 



CH^OH + HCl 



280°C 



Alumina 



^ CH^Cl + H2O 



(3-10) 



hundreds of thousands of metric tons annually. The major 

uses of chloromethane include the production of methyl 

155 
chlorosilanes, tetramethyl lead and butyl rubbers. The 

formation of methyl chloride from synthesis gas and HCl 

under extremely mild conditions may be industrially useful. 

For this reason, a further in-depth investigation into the 

mechanism, as well as the optimization of this unique system 

is warranted. 



76 

Experimental 
Reagents 

All metal complexes were used as purchased unless 
otherwise stated. The IrCl-* 3H and all the metal carbonyl 
complexes were purchased from Strem Chemical Company. All 
solvents, except 2-methoxyethanol and 2-ethoxyethanol, were 
dried prior to use by distillation over CaH . All solvents 
were degassed with N2 prior to use. The alumina, acid 
Brockman Activity I (80-200 mesh) and the mossy zinc metal 
were purchased from Fisher Scientific Company. The alumina 
was dried at 140°C prior to use. This alumina was determin- 

9 ICC 

ed to have a specific area of 180 m /g. The silica gel, 
Davison grade #62, purchased from W. R. Grace, was dried 
under vacuum at 300°C prior to use. It had a specific area 
of 340 m /g, a pore diameter of 14 mm and a pore volume of 
1.1 cm"^/g. The zeolite, LZY-82, was purchased from Alfa- 
Thiokol. All silanes were purchased from Petrarch Chemical 
Company and used without purification. The carbon monoxide 
C. P. grade 99.5% and the hydrogen chloride technical grade 
99.0% or semiconductor grade 99.995% were purchased from 
Matheson Gas Products. The hydrogen was obtained from 
Strate Welding. All carbon-13 isotopically labelled gases 
were purchased from Merk, Sharp and Dohme Isotopes. 
Instrumentation 

All air sensitive manipulations were performed in a 
Vacuum Atmosphere Co. model HE-43-2 inert atmosphere box or 
in an Aldrich inert atmosphere glovebag. All syntheses were 



77 
performed under either a nitrogen or carbon monoxide 
atmosphere. GC analyses were performed on either a model 
3700 FID Varian gas chromatograph equipped with a Hewlett- 
Packard 3390A integrator and a 1/8 inch x 8 foot stainless 
steel 5% diethylene glycol adipate on chromosorb P (80/100) 
column or on a model 940 FID Varian gas chromatograph 
equipped with a 1/8 inch x 8 foot stainless steel poropak Q 
(100/120) column. GC mass spectrometry was performed by Dr. 
R. King of the Microanalytical Laboratory, University of 
Florida, Gainesville, Florida. Samples were run on an AEI 
MS 30 mass spectrometer with a KOITOS DS55 data station. 
The system was equipped with a PYE Unicam 104 gas chromato- 
graph containing a 1/4 inch x 5 foot poropak Q column. 
Nuclear Magnetic Resonance spectra were obtained on a Varian 
EM360L NMR spectrometer. Infrared spectra were obtained as 
mulls on a Nicolet 5DXB FTIR spectrometer using KBr salt 
plates. All elemental analyses for carbon, phosphorous and 
iridium were performed by Galbraith Laboratories, Knoxville, 
Tennessee. All ESCA data was obtained through the courtesy 
of Dr. Tom Gentle, Dow Corning Corporation, Midland, 
Michigan. The samples were run in a Perkin-Elmer Model 551 
stainless steel ultra-high vacuum chamber equipped with a 
dual magnesium anode x-ray source and a double pass 
cylindrical mirror electron analyzer. Data acquisition was 
controlled by a Digital PDP computer. All high pressure 
experiments were performed using a 50 mL Parr pressure 



78 

bottle equipped with a brass or stainless steel Swagelok 

ft fi 
pressure head. 

Fixed Bed Flow Reactor 

A glass flow system as shown in Figure 3-2 was 

assembled. This was modified from the previously described 

148 
system to allow the entire gas mixture to flow through 

the catalyst. The individual gas flow rates were controlled 

by three teflon needle valves (A, B and C) . The CO and Yi^ 

were bubbled through mineral oil while the HCl bubbler 

contained sulfuric acid. The gases were allowed to flow 

over the catalyst which was supported on a glass frit and 

held in place with glass wool. The overall flow rate of the 

gas mixture was monitored by a bubble flowmeter. The 

temperature was regulated by a model 123-8 Lindberg thermo- 

stated tube furnace that surrounded the reactor tube. Gas 

samples for GC analyses could be obtained through two sample 

ports, one prior to and one after the catalyst. Gas samples 

were collected using a pressure-lok 2 mL syringe purchased 

from Precision Sampling Corporation. Gases could be trapped 

out through the addition of a glass spiral trap to the glass 

reactor system. The spiral trap allowed for maximum contact 

of the gas flow with the dry ice/acetone slush. 

Preparation of Dicarbonylchloro (p-toluidine) iridium (I) 

The IrCl (CO) - (p-toluidine) was prepared by a procedure 

157 
similar to that reported by Klabunde. Inside an inert 

atmosphere glovebag a pressure bottle system containing 1.0 g 

of IrCl 'SH 0, 0.30 g of lithium chloride and 50 mL of 



79 




80 
degassed 2-methoxyethanol was assembled. The pressure 
bottle was charged with 45 psig of carbon monoxide and 
allowed to react for several hours at 130°C. When the 
initial black color had changed to yellow, the pressure 
bottle system was cooled to room temperature. The pressure 
bottle was dismantled under a nitrogen atmosphere and 0.3 5 g 
of p-toluidine added. After several minutes of stirring, 
the yellow solution was poured into a beaker containing 250 
mL of distilled water. A purple precipitate was formed 
immediately upon the mixing of the two solutions. The 
precipitate was collected by vacuum filtration and dried 
under vacuum for 2 4 hours. The purple solid was dissolved 
in a minimum amount of benzene. Then a small amount of 
anhydrous sodium sulfate was added to the stirred brown 
solution. After several hours the solution was filtered. 
The solvent was evaporated from the filtrate to give a 
purple compound which was characterized by infrared spectro- 
scopy to be IrCl(C0)2(p-toluidine) . A typical yield was 
approximately 85% based on the initial IrCl^'SH^O complex. 
Preparation of a Phosphinated Support 

The phosphinated supports were prepared by a procedure 

144,145 
similar to that previously reported in the literature. 

Under a nitrogen atmosphere a total of 5.0 g of a dried 

support, such as alumina, silica gel, or a zeolite was added 

to a stirred solution of 150 mL of toluene. The mixture was 

heated to reflux temperature prior to addition of 0.45 mL of 



81 

2- (diphenylphosphino) ethyltriethoxysilane, 

(C^Hj-O) ^SiC„H .PPh~ , by a syringe method. This reaction was 
allowed to continue for 12 hours prior to collecting the 
functionalized resin by vacuum filtration. The 
functionalized support was dried under vacuum at room 

temperature for 2 4 hours before use. This reaction gave a 

_3 
phosphinated support containing 1.25 x 10 moles of 

accessible phosphine substituents . Supports with different 

phosphine concentrations were prepared in an analogous 

manner. In experiments where the rest of the surface was 

silanated with dichlorodiphenylsilane, an appropriate amount 

of the silane was added by syringe 6 hours after the 

2- (diphenylphosphino) ethyltriethoxysilane had been added. 

Preparation of Supported Mono- and Di-phosphine Substituted 
Tetrairidium Carbonyl Clusters 

The phosphine substituted tetrairidium carbonyl cluster 

was supported by a procedure similar to that reported by 

Struder et al.''"^"'" and Castrillo et al.. ' This 

procedure was adopted from one reported by Stuntz and 

159 
Shapley for the formation of Ir^ (CO) ^^PPh^ . Inside a 

glovebag a total of 130 mL of 2-methoxyethanol and 5 mL of 

water was added to a pressure bottle containing 5.0 g of a 

phosphinated support (1.25 x 10~ moles of phosphine) and 

0.057 g of dicarbonylchloro(p-toluidine) iridium(I) . The 

amount of IrCl (CO) _ (p-toluidine) used changed according to 

the concentration of phosphine on the support that was used. 

The 15.0 g of mossy zinc metal was placed into a teflon 



82 

basket suspended in the solution above the cylindrical 1/2 

inch long stirbar. The pressure bottle was charged with 45 

psig of carbon monoxide and heated to 95°C. The reaction 

was allowed to proceed for 12 hours. The slightly yellow 

resin was collected by vacuum filtration, washed with 

approximately 7 5 mL of toluene and dried under vacuum for 2 4 

hours. The catalyst prepared in other solvents, such as 

2-ethoxyethanol or toluene was done in an analogous manner. 

The characterization of each catalyst by infrared 

spectroscopy is discussed in the results section. 

Preparation of Supported Tri-phosphine Substituted 
Tetrairidium Carbonyl Clusters 

The tri-phosphine substituted tetrairidium carbonyl 
cluster was supported by a procedure similar to that report- 
ed by Karel and Norton. ''" ° A total of 1.3 g of Ir^(C0)-^2 
was added to a stirred toluene solution containing 5.0 g of 
a phosphinated support (1.25 x 10 moles of phosphine) . 
The reaction was allowed to proceed at reflux temperature 
for 2 4 hours. The yellow resin was collected by vacuum 
filtration and dried under vacuum for 2 4 hours. The 
characterization of the catalyst by infrared spectroscopy is 
discussed in the results section. 

Preparation of Other Supported Phosphine Substituted Metal 
Carbonyl Complexes 

All other metal carbonyl complexes, such as Ru2(CO)^2' 

0S3(C0)^2' ^^6^^°^6' Mn2(C0)^Q, Re2{C0)^Q, CO2(C0)g, 

Fe(CO) , IrCKCO)^ and IrCl (CO) (PPh^) 2 were supported in an 

analogous manner to the preparation of the supported tri- 



83 

phosphine substituted tetrairidium cluster. Appropriate 

amounts of the metal carbonyl complexes were added to 

stirred toluene solutions containing 5.0 g of a phosphinated 

support. The reaction was allowed to proceed at reflux 

temperature for 2 4 hours. The resins were collected by 

vacuum filtration and dried under vacuum for 24 hours. The 

characterization of each catalyst by infrared spectroscopy 

is discussed in the results section. 

Preparation of Iridium Complexes Impregnated Onto a Support 

The iridium complexes were impregnated onto a support 

through the incipient wetness impregnation of the support 

with a solution containing the metal complex. An 

appropriate amount of an iridium complex, such as Ir.(C0),2/ 

Ir(CO)^Cl or IrCl-«3H was added to a stirred cyclohexane 

solution containing 5.0 g of the support. The amount of 

metal complex added was dependent upon the concentration of 

iridium desired on the support. The mixture was allowed to 

stir at room temperature for several hours. The supported 

complexes were collected by vacuum filtration and dried 

under vacuum for 24 hours. The characterization of each 

catalyst by infrared spectroscopy is discussed in the 

results section. 

Reaction of Catalysts with Carbon Monoxide, Hydrogen and 
HCl(g) 

Prior to running the catalyst experiments, the blank 

reactor tube was tested for any residual activity towards 

methyl chloride formation. Then a total of 1.0 g of a 



84 

catalyst was placed into the glass fritted reactor tube. 
The catalyst was held in place with glass wool. The reactor 
tube was placed into the fixed bed flow reactor system 
previously described in Figure 3-2. A typical reaction was 
run at 75°C with the individual H2:C0:HC1 gas flows at a 
ratio of 3:1:0.5 combining to give an overall flow rate of 
1 iTiL/40 seconds. The reactant and product gases were 
monitored by gas chromatography. Investigation of the 
active catalyst by infrared spectroscopy is disucssed in the 
results section. 
Reactions Involving Carbon-13 Isotopically Labelled Gases 

Carbon-13 incorporation into alkyl chlorides was 
investigated by the reaction of carbon-13 labelled CO or CH^ 
gas with the catalyst in a stagnant reactor. The stagnant 
reactor was either a pressure bottle fit with a stainless 
steel Swagelok head or a closed reactor tube. In both 
cases, the reactor containing the catalyst was evacuated at 
room temperature. Then the gases were placed into the 
system in the appropriate ratios with the overall pressure 
never exceeding one atmosphere. The reaction was allowed to 
continue at 75°C for 24 hours. The extent of carbon-13 
incorporation into the products was determined by GC mass 
spectrometry . 



85 

Results and Discussion 

Reproduction of the Previously Reported Supported Iridium 
Carbonyl Catalyst System 

The supported tetrairidium carbonyl cluster was 

148 
prepared as previously described by Miller. First, the 

surface of the inorganic oxide (alumina, silica gel or 

zeolite) was modified by phosphine functionalization through 

a simple condensation reaction between 2-(diphenyl- 

phosphino) ethyltriethoxysilane and the surface hydroxyl 

groups as shown in Equation 3-11. Then, as shown in 

Equation 3-12, the supported phosphine tetrairidium carbonyl 

cluster was assembled through the reduction of IrCltCO)-- 

(p-toluidine) by zinc metal in the presence of carbon 



toluene (OH) ^ 

3-OH + (CjH^Oj^SilCHjJjPPhj 1 *" 3" (O) -Si (CH2) jPPhj * ^C^H^OH 

(3-11) 



+ CO, Zn, H^O 3+(0)y-Si(CH2)2PPh2)x^'^4'C°'z 



IrCKCOjlHjN-^^CH^ 



2-methoxyethanol, A x « 1, 2; z = 11, 10 

(3-12) 



monoxide and a 2-methoxyethanol solvent. The catalysts were 
characterized by infrared spectroscopy to be a mixture of 
supported mono- and di-phosphine substituted tetrairidium 



86 

carbonyl clusters. A detailed discussion of the infrared 
data is presented later. These supported clusters were 
tested for catalytic activity in the presence of a 1:3:0.5 
ratio of CO:H-:HCl(g) at 75°C in a modified fixed bed flow 

reactor system as described in Figure 3-2. 

148 
Similar products to those previously reported for 

this reaction were identified by GC and GC/MS. A typical 
gas chromatogram using a poropak Q column at 130°C is shown 
in Figure 3-3. Although not shown in Figure 3-3, gas 
chromatography also was used to detect the presence of 
residual solvents, such as toluene and 2-methoxyethanol with 
retention times at 16 and 24 minutes, respectively. In 
addition, the presence of H2O, CO- and a trace amount of 
acetylene was confirmed by GC/MS. A sample of the gas mix- 
ture taken prior to the catalyst was found through the use 
of gas chromatography to contain methane, ethylene and ethyl 
chloride as the only observable impurities. The presence of 
these impurities was confirmed from gas specification data 
obtained from Matheson Gas Company for CO(g) and HCl (g) 
(technical grade) . ■'"^■'' The utilization of HCl(g) "(semi- 
conductor grade) eliminated the presence of ethylene and 
ethyl chloride in the pre-gas mixture. However, the 
presence of both ethylene or ethane and ethyl chloride was 
observed by gas chromatography to remain as reaction 
products in the post-gas mixture. It was found that if the 
amount of HCl (g) present in the system was reduced after the 



87 









<u 












TJ 




Q) 








•H 




TJ 








>-l 


0) 


•H 











•o 


M 








•-{ 


>1 


C5 








JS 


x: 


f-i 




(1) 




o 


0) 


s: 


Q) 


c 






-a 


o 


C 


<u 




iH 


iH 




rt 


rH 




>i 


10 


r-i 


ji 


>1 




x: 


4J 


>1 


4J 


x: 


<-i 


-p 


0) 


X 


<1) 


-p 


D 


(U 


u 


■P 


E 


(U 


K 


e 


rO 


0) 


0) 


en 


cn 


(0 


W 


01 


fl) 


a) 


(y 


0) 


0) 


(U 


4J 


-p 


-P 


+J 


-P 


-P 


3 


3 


3 


3 


3 


3 


C 


C 


C 


c 


C 


C 


•H 


•tH 


-H 


•H 


•H 


-H 


E 


e 


e 


E 


E 


fc 


o 


in 


o 


o 


o 


O 


in 


00 


fS 


CN 


o 


H 


o 


o 


fH 


CM 


fO 


in 


II 


II 


II 


II 


II 


II 


ta 


XI 


o 


Xi 


0) 


M-l 





u 
o 

(d 
a 



>^ 
o 



(0 

<u 

CO 
(0 

o 

3 U 

r^O 
O O 

P^ H 

(U II 

4-1 C 




O 
E U 

^1 - 

CT>00 

o 
■p II 

(0 

x; 
u 



:^J^^S^- 



system's initial activation, the identification of methanol 
by gas chromatography as shown in Figure 3-4 was possible. 
The initial activity for methyl chloride formation at 
various temperatures by the catalyst (0.31% Ir) supported on 

alumina was found to parallel the activity observed in the 

148 
previously reported alumina bound catalyst system, 

3xvPPh Ir (CO) (0.75% Ir) . Actually, the 0.31% iridiiim 

catalyst that was tested showed greater initial activity 

than the reported 0.75% iridium catalyst as shown in Figure 

3-5. This is contrary to the previously reported results 

which indicated that an increase in activity was associated 

with an increase in the metal loading of the catalyst. The 

consequence of this discrepancy will be discussed at a later 

time. It is important at this" time to observe the decrease 

in catalyst activity associated with the second day of 

testing as shown in Figure 3-5. The observation of this 

decrease in activity was made possible through the use of a 

modified flow bed reactor system as described in Figure 3-2. 

Since this modified flow bed reactor minimized the amount of 

148 
waste gas with respect to the previous flow bed reactor , 

the reaction system was allowed to operate 24 hours a day. 

The observed decrease in activity during this time period is 

most likely associated with catalyst decomposition. This is 

supported by a first order deactivation of the catalyst 

system observed by Dow Chemical Company in an independent 

^ ■, ^162 
investigation of the supported catalyst. 



89 




a = methane 

b = methyl chloride 

c = methanol 



1.0 3.0 

Time (minutes)^ 



HZl 



Al B-OSiiCH^) ^Ph^Ir^iCO) ^ + 00 + H^ j^^ »► CH3CI 



75 C 



ICl depleted 



CH CI + Ca^CB 



Figure 3-4. Gas Chromatography separation of 
Methanol and Methyl Chloride 






90 



o= Initial Activity (0.31% Ir) 
n= Initial Activity (0.75% Ir)^ 
• = 2nd Day Activity (0.31% Ir) 




Temperature ("C) 



Lteita obtained fran reference 148 



Figure 3-5. A Graph of Catalyst Activity Versus 
Reaction Temperature 



91 
Catalyst Deactivation is a Valid Observation 

The catalyst deactivation observed between the initial 
activity (day one) and the recorded activity on the second 
day was monitored over an extended reaction time. It was 
observed that the activity of the catalyst seemed to 
stabilize during the third day as shown in Figure 3-6. 
However, this apparent stabilization was found to inherently 
coincide with an increase in both the HCl(g) concentration 
in the reactant gas and the residence time of the gas over 
the catalyst sample. It was found that if both of these 
parameters were stringently maintained at their original 
values, the catalyst's activity did not stabilize until 
after approximately ten days as shown in Figure 3-7. 

The activity associated with several control reactions 
over an extended reaction time also is shown in Figure 3-6. 
The alumina control supports were prepared in the absence of 
any iridium precursor through a procedure analogous to the 
preparation of the catalyst sample. The first control 
reaction showed a minimal initial activity and a rapid 
decrease in activity with time. However, the second control 
reaction showed greater initial activity than the alumina 
supported iridium catalyst. The activity for this second 
control experiment seemed to level off during the third day 
in a similar fashion to that of the catalyst sample. The 
only noticeable difference in the preparation of both the 
alumina control supports was in the degree of solution 
stirring. A similar effect concerning the catalyst 



92 



60 



3 50 
to 



40 





en 
o 


30 


^ 


o 






U 


20 




n 






g 






tn 






<l) 





g 1(T 




catalyst (0.57% Ir) 
cxjntrol #1 (poor stirring) 
control #2 (excellent stirring) 



[? stabilization ?] 
a _ o D o- 



72 120 
Tine (hours) 



168 



Figure 3-6. A Graph of Catalyst Activity Versus 

Reaction Time (Residence Time and HCl(g) 
Concentration was Observed to Increase 
with Time) 



93 



I 



m 
o 






0= catalyst (0.47% Ir) 




II - 9 X 10 noles CH^Cl/sec./grn catalyst] 



15 25 35 
Reaction Time (days) 



45 



Figure 3-7. A Graph of Catalyst Activity Versus an 
Extended Reaction Time (Residence Time 
and HCl(g) Concentration Held Constant) 



94 

148 
preparation was reported in which poor stirring during 

preparation resulted in a decrease in catalyst activity. 

This stirring problem has been adequately disposed of by 

placing the desired quantity of mossy zinc in a fiberglass 

mesh basket suspended in the solution above the 1/2 inch 

cylindrical stirbar. A large number of control supports 

using phosphinated alumina or silica gel have been prepared 

under excellent stirring conditions and tested for catalytic 

activity towards methyl chloride formation. In each case 

similar activity to that observed for the supported iridium 

catalyst was recorded. A summary of the activity of the 

control supports that were run is presented in Table 3-4. 

When a control support was run over an extended reaction 

time, its activity towards the formation of methyl chloride 

was found to decrease to a negligible amount during the 

first ten days. However, the activity of the supported 

catalyst was observed to stabilize after ten days and 

-12 
persist for a period of several weeks at 1-9 x 10 moles 

CH^Cl sec"''' g~ as shown in Figure 3-7. It should be noted 

that this level of activity is very close to the detection 

limit of the gas chromatograph. A buildup of methyl 

chloride over a period of twelve hours could be observed in 

a stagnant reactor that was formed by closing the valves on 

the flow reactor system. Although the occurrence of a 

decrease in the catalyst's activity has been verified, the 

cause of this decrease, as well as the cause of the activity 

observed in the control reactions still remains unanswered. 



Control 
Supports 


Al- 


■PPh2 


Al- 


■PPh2 


Al- 


-PPh2 


Al/ 


no zinc 


SG 




SG- 


-PPH2 


Al- 


-PPh2 



Day (1) 




2.80 X 


10-1° 


5.10 X 


10-' 


1.57 X 


10-^ 


5.18 X 


10-5 


9.07 X 


10-" 



95 

Table 3-4. The Initial Activity of Control Supports 
Prepared in 2-methoxyethanol 

Activity (moles CH^Cl sec g ) 

Day (2) Day (*) 



2.25 x 10-1°(6) 
1.14 X 10-^ 1.80 x 10-1°(5) 
1.01 X 10"^ 1.53 X 10-1°(4) 



2.45 X 10-1° ^_^g ^ 10-1° g_45 ^ 10-11(4) 
3.99 X 10-1° ^^^3 ^ 10-1° j^^ ^_ ^^Qj 



Al = alumina; SG = silica gel; (*) = day activity was last 
monitored; N. A. = no activity. 



Investigation of the Methyl Chloride Formation Observed in 
the Control Reactions 

Since zinc metal was used as a reducing agent in the 

synthesis of the supported iridium cluster, a small amount 

of zinc is inherently present in the support. To ascertain 

if this zinc metal could be responsible for the observed 

formation of methyl chloride, control supports were prepared 

in the absence of zinc, as well as the iridium precursor, 

through a procedure analogous to the preparation of the 

catalyst samples. Since these control reactions showed 

similar activity as compared to the supported catalyst, it 



96 
is suggested that the presence of zinc metal has no apparent 
effect upon the activity of the system. 

Another possibility that could explain the activity 
observed for the control reactions is the presence of a 
residual amount of the 2-methoxyethanol solvent that was 
used in the synthesis of the supported catalyst and control 
supports. It is possible that an interaction between 
physically adsorbed or chemisorbed 2-methoxyethanol with 
HCl(g) could be responsible for the observed formation of 
methyl chloride. The cleavage of ethers and alcohols by 
hydrogen halides to form the corresponding alkyl halides is 
enhanced in the presence of a Lewis acid. This is believed 
to proceed through a mechanism as shown in Equation 3-13, 
where the Lewis acid (i.e., Lucas reagent) coordination to 
the hydroxyl oxygen allows attack of the halide source upon 



ZnCl„ 

CI^^hZ-O >- RCH„C1 + HOZnCl + CI (3-13) 

I ^ \ 

R H 



the electrophilic carbon through a weakening of the 
carbon-oxygen bond. Furthermore, the current industrial 
process for the production of methyl chloride has previously 
been shown in Equation 3-10 to involve the chlorination of 
methanol over an inorganic oxide support at elevated 
temperatures. 



97 
Several factors as shown in Table 3-5 can enhance the 

physical adsorption of a polar organic compound onto the 

164 
surface of a hydroxylated support . Since 2-methoxy- 

ethanol meets several of the criteria shown in Table 3-5, 



Table 3-5. Factors That Enhance the Adsorption of Polar 
Organic Compounds onto the Surface of a 
Hydroxylated Support 



(1) The presence of strong electron-donor groups (oxygen, 
nitrogen, etc.) 

(2) Increasing size and number of hydrocarbon groups in 
molecule. 

(3) Soluble salts (NaCl, etc.) in aqueous phase. 

(4) Multiple polar groups leading to multiple surface 
interactions . 

(5) Low temperature. 

(6) A high concentration of compound to be adsorbed. 



the adsorption of this -solvent onto the surface of the 
support should occur quite readily. The procedure of drying 
the catalyst at room temperature under vacuum is not 
sufficient to remove this adsorbed 2-methoxyethanol . 

It is well documented that the hydroxyl groups of a 
support, such as alumina or silica gel, will react in the 
presence of alcohols and ethers at temperatures exceeding 

190°C to form the corresponding support bound alkoxide 

^ , , 165,166 ^ 
groups and water as shown in Equation 3-14. An 

analogous reaction could be predicted for the chemisorption 



98 

of 2-methoxyethanol onto the surface of the phosphinated 
support. 

Recall that the presence of residual 2-methoxyethanol 
solvent in the catalyst reactions was observed by gas 
chromatography. It previously was reported that the 

condensation of a liquid occurred near the top of the 

148 
reactor tube. This liquid was composed of water and two 

unknown compounds with GC retention times of 4.16 minutes 

and 6.07 minutes on a DEGA column at 100°C. The larger of 




3-OH + ( \ 3-OH (3-14) 



3-OR + ROH 



the two peaks at 4.16 minutes has been identified by gas 
chromatography to be 2-methoxyethanol. Further 
identification of 2-methoxyethanol in this condensed liquid 
was obtained by proton NMR spectroscopy as shown in Figure 
3-8. Gas chromatography results obtained by Dow Chemical 
Company during an in-depth investigation of the catalyst 

system also indicated the presence of residual 

16 2 
2-methoxyethanol and toluene solvents. Since these 

solvent peaks were not observed by gas chromatography on 

the second day of catalyst activity, Dow Chemical Company 



99 



Sample spectrum 



CH^-O-CH CH -OH 
a bed 



e = CDCl^ solvent 



f = impurity found in the 
CDC] 

g = HCl 



CDCl^ used. 




H 



TMS 



LU 



10 9 



6 5 4 3 2 10 
PPM (6) 



2-methoxyethanol reference 
a 



cb 



TMS 



> — A Ji^ 



5 4 3 
PPM (6) 



Figure 3-8. NMR Spectrum of 2-methoxyethanol that 

Condensed at the Top of the Reactor System 



100 
concluded that the solvent did not play a major role in the 
formation of methyl chloride. However, Dow Chemical Company 
was using higher pressures and shorter residence times than 
could be obtained in a glass flow reactor system. The 
length of time in which 2-methoxyethanol is observed is 
dependent upon a variety of factors, such as residence time 
and temperature. The presence of 2-methoxyethanol in the 
catalyst and control reactions employing the glass flow bed 
reactor has been observed by gas chromatography for periods 
of one to ten days. Anywhere up to a 100-fold decrease in 
the amount of 2-methoxyethanol present between days one and 
two of the catalyst's activity has been observed by gas 
chromatography. This decrease in activity corresponds very 
closely to the drop in activity for methyl chloride 
formation that also was observed to occur between days one 
and two. The cracking of physically adsorbed and 
chemisorbed 2-methoxyethanol present in the supported 
catalyst and control supports is a feasible explanation for 
the initial activity observed for the formation of methyl 
chloride . 

The inability to identify by gas chromatography or 
GC/MS other products, such as ethylene glycol or 
1,2-dichloroethane that may result from the decomposition of 
2-methoxyethanol does not deter this proposal. The 
inability to detect these compounds may be related to the 
decreased sensitivity of the gas chromatograph for these 
products, long retention times on the GC column, the 



101 

stabilization of unidentified surface species or to the 

further decomposition of these compounds over the inorganic 

oxide supports to the observed products of methane, 

ethylene, methyl chloride, ethyl chloride, acetylene, 

methanol, carbon dioxide and water. 

Solvent Decomposition Can Explain Other Reported and 
Observed Results 

In order to investigate the validity of the decrease 

observed in the initial activity (Figure 3-5) upon an 

increase in metal loading, catalysts containing different 

iridium concentrations were prepared and tested. The 

results as shown in Table 3-6 indicate that an increase in 

the percentage of metal on the support does decrease the 

initial activity of the system. However, this result also 

can be explained by the cracking of physically adsorbed and 

chemisorbed 2-methoxyethanol. An increase in the metal 

loading of the phosphine bound catalyst must coincide with 

an increase in the phosphination of the support. Since the 

phosphination of the support proceeds through a condensation 

reaction of (C^H 0) Si (CH^) 2PPh2 with the hydroxyl groups of 

the support, an increase in the extent of surface 

phosphination would result in a decrease in the number of 

hydroxyl groups remaining on the support. The smaller 

number of hydroxyl groups present in the catalyst supports 

containing higher metal loadings would limit the amount of 

2-methoxyethanol that could be physically adsorbed or 

chemisorbed on the support. The overall effect would be a 



102 

Table 3-6. The Activity of Alumina Supported Tetrairidium 
Clusters for Various Metal Loadings 





uUm 


Activity (moles 


CH^Cl sec ■*• g ^) 






Irid: 








Loading 


Initial 




Decrease Factor 


(Wt. 


%) 


(Day 1) 


Day 2 


Day 


1 to Day 2 


0.14 




2.14 X 10"^ 


5.37 X 10-^° 




40 


0.31 




8.90 X 10"^ 


1.11 X 10-^° 




80 


0.57 




2.49 X 10"^ 


2.46 X 10"-'-° 




10 


2.34 




1.04 X lO"^ 


3.10 X 10"-'-° 




3 



decrease in the activity for methyl chloride formation 
associated with the cracking of the smaller quantity of 
2-methoxyethanol adsorbed on the support. The decrease in 
activity observed for the catalysts between days one and two 
as shown in Table 3-6 corresponds very closely to the 
decrease in adsorbed 2-methoxyethanol observed over the same 
time period by gas chromatography. Any small discrepancies 
in the activity calculations can be explained by the 
inherent errors that arise from the difficulty in 
maintaining the appropriate gas concentrations, a constant 
residence time and through interpretation of the gas 
chromatography calibration curves for methyl chloride. 

Similarly, the differences in the initial activity of 
the control reactions (Figure 3-6) that arose from the 
difficulty in controlling the rate of stirring during the 
preparation of the control support can be explained through 
the adsorption of 2-methoxyethanol by the support. An 



103 
inadequately stirred reaction would decrease the number of 
hydroxyl groups that would be exposed to the 
2-methoxyethanol solvent. Since these inadequately stirred 
reactions would result in supports containing less 
physically adsorbed and chemisorbed 2-methoxyethanol, they 
should show a smaller initial activity for the formation of 
methyl chloride. Since the cracking of 2-methoxyethanol 
governs the amount of methyl chloride initially formed, the 
amount of methyl chloride observed in these control 
reactions should be dependent upon the amount of HCl (g) 
present in the system and the length of the residence or 
contact time the gas mixture has over the control sample. 
As expected, an increase in the HCl (g) concentration or an 
increase in the residence time as shown in Figure 3-9 was 
found to increase the amount of methyl chloride observed in 
the system. It should be noted that the detection limit of 
the gas chromatograph for methyl chloride requires a 
residence time of at least 25 seconds/gram of catalyst to 
form an easily detectable amount of methyl chloride. 
Finally, the separation of methanol observed in 
reactions which contained a small amount of HCl (g) as shown 
in Figure 3-4 could be derived from the cracking of 
2-methoxyethanol to methyl chloride. Recall that the 
formation of water was verified by GC/MS in the catalyst 
reactions. The relationship between methyl chloride, water, 
methanol and HCl is described in Equation 3-15. A decrease 
in the concentration of hydrogen chloride would shift the 



104 




> u 

o 
-p a. 

•H 3 

> Ui 



-H U 
U 



U C 

•H 
rH > 
>irH 

Si o 
-p > 

0) c 

S H 

m c 
o o 

•H 
X -P 

a o 

U 0) 

o « 
<C to 



uot:;osCut -[ui 2/_0T x id^hO saioui 



105 
equilibrium of the reaction towards the formation of 
methanol. The presence of this equilibrium is further 
supported in that the addition of water vapor to the feed 



CH^OH + HC1< ^CH^Cl + H^O (3-15) 



gas was observed to decrease the activity of the system 
towards the formation of methyl chloride. 

The cracking of residual 2-methoxyethanol can explain 

several aspects of the catalyst system that was previously 

148 
reported. For instance, all control reactions were 

previously reported to be inactive towards methyl chloride 

formation. This can be explained by the realization that 

all control reactions previously run were never exposed to 

the 2-raethoxyethanol solvent during preparation. It also 

148 
was reported that the active catalyst was oxygen 

sensitive. It is obvious that the cracking of the 

chemisorbed or physically adsorbed solvent by HCl(g) should 

not be affected by oxygen exposure. All catalysts that were 

run were found not to be air sensitive. However, the 

solvent can still explain this phenomena when one examines 

1 fi 7 
the procedure for the reported experiment. The 

experiment was run at room temperature in the presence of 

CO, H„ and HCl{g) over several days for a period of 

approximately eight hours per day. The system was closed 

and allowed to sit for ten days. A tremendous amount of 

methyl chloride was observed by gas chromatography to be 



106 
present at the end of this ten day period. The catalyst was 
exposed to air and a large decrease in activity was observed 
upon restarting the system in the presence of CO, H and 
HCl(g). It is suggested that the presence of 
2-methoxyethanol was never depleted in the initial running 
of the catalyst. When the system was restarted after ten 
days and exposed to oxygen, the decrease in the observed 
activity was caused by the depletion of adsorbed 
2-methoxyethanol that resulted during the ten day exposure 
to HCl(g) in the stagnant reactor. It seems that the 
cracking of adsorbed 2-methoxyethanol is a feasible 
alternative explanation for the reported oxygen sensitivity 
of the catalyst. 

Finally, the short incubation period that was 

14 8 
reported to occur prior to methyl chloride formation may 

be related to the cracking of 2-methoxyethanol. This 

incubation period was found to be related to the time 

required to build up an adequate concentration of methyl 

chloride in the system which could be observed by gas 

chromatography. This is supported by the observation that 

the length of time associated with this incubation period 

could be altered by changing either the HCl(g) concentration 

in the system or the contact time between the gases and the 

catalyst. Thus, the cracking of physically adsorbed and 

chemisorbed 2-methoxyethanol can adequately explain the 

observed initial activity, as well as the reported oxygen 



107 
sensitivity and incubation period of the catalyst and 
control reactions. 
The Reduction of Carbon Monoxide Still Occurs 

Recall that the stabilization of the activity at 1-9 x 

-12 -1-1 

10 moles CH CI sec g as shown in Figure 3-7 occurred 

after approximately the tenth day of running the catalyst 

system. It is possible that this stabilization in activity 

results from the conversion of synthesis gas and HCl (g) over 

the supported iridium catalyst to methyl chloride. This is 

supported by the thermodynamic feasibility of the conversion 

of synthesis gas and HCl(g) to methyl chloride as shown in 

Table 3-7. 



Table 3-7. Thermodynamic Data for Reactions Involving 
Synthesis Gas Reduction 

° -1 

Reaction AH 293 (kcal mole ) 



CO + 2H + HCl •►CH^Cl + H^O 

CO + 2H •►CH^OH 

CO + 3H2 *'CH^ + H^O 

2C0 + 2H2 *"CH^ + CO2 



-28 


9 


-21 


7 


-49 


3 


-59 


1 



In order to investigate if the iridium bound catalyst 
was reducing carbon monoxide, experiments were conducted 
with the goal of eliminating the activity associated with 
the 2-methoxyethanol solvent. First, elimination of the 
solvent was attempted by washing the supported catalyst and 



108 
the control supports with 2-ethoxyethanol , ethanol or 
methylene chloride. In each case, a 10 to 100-fold decrease 
in the initial activity was observed. This diminished 
activity is associated with the partial removal of the 
adsorbed 2-methoxyethanol solvent. In the case of the 
methylene chloride wash, the separation of adsorbed metal 
carbonyl complexes from the supported iridium carbonyl 
clusters was observed by infrared spectroscopy. The 
complete removal of the adsorbed 2-methoxyethanol was 
attempted by washing the catalyst with water and by heating 
the catalyst under vacuum. It was shown by infrared 
spectroscopy that washing the catalyst with water hydrolyzed 
the bound cluster off of the support through attack upon the 
silane linkage. Subjecting the catalyst to elevated 
temperatures under vacuum was found by infrared spectroscopy 
to decompose the tetrairidium carbonyl cluster. 

An experiment designed to block the reaction between 
2-methoxyethanol and the surface hydroxyl groups was 
performed. The phosphinated support prior to contact with 
2-methoxyethanol was reacted with dichlorodiphenylsilane in 
an attempt to silanate the rest of the support's surface. 
This silanation process proceeds through the reaction of the 
support hydroxyl groups with the silicon-chloride linkages 
of the silane as shown in Equation 3-16. The decrease in 
methyl chloride activity observed for this silanated support 
can account only for the partial elimination of the adsorbed 



109 
2-niethoxyethanol. It is possible that complete silanation 
of the support's surface was not obtained because of the 
steric effect associated with the phenyl groups of the 



-OH 



-OH 



+ CI SiPh2 



-0 Ph 

\/ 

Si 

\ 
Ph 



-/ 



+ 2HC1 



(3-16) 



silane. Data concerning the duration of the activity 
associated with this silanated support was not obtained. 

In order to substantiate the cluster's activity, 
attempts were made to synthesize the supported cluster in 
the absence of 2-methoxyethanol . The infrared data obtained 
for these supported iridium complexes will be discussed at a 
later time. Testing the supported iridium complex (0.99% 

Ir) made in 2-ethoxyethanol in either a flow reactor or a 

-12 
stagnant reactor system showed an activity of 4.34 x 10 

moles CH3CI sec" g~ . It was found that using supports 

containing different iridium loadings, such as 1.92% and 

0.99%, showed little affect upon the system's activity. Gas 

chromatography also was used to identify large amounts of 

methane and either ethane or ethylene and trace quantities 

of acetaldehyde and ethyl chloride as reaction products. 

The separation of ethane from ethylene by gas chromatography 

was not obtained. It was noticed that the concentration of 

acetaldehyde was found to increase with higher reaction 

temperatures . 



110 
Testing the supported iridium complex (1.18% Ir) 
prepared in toluene showed similar activity to that observed 
for the 2-ethoxyethanol prepared catalyst. An initial surge 
in activity was observed during the first day of activation 

as shown in Figure 3-10. The activity was found to level 

-12 
off after the first day at approximately 7.80 x 10 moles 

— 1 -1 
CH^Cl sec g . A similar surge in activity was observed 

for the supported cluster exposed to a reactant gas stream 

consisting of only H„ and HCl(g). This suggests that this 

surge in activity can be associated with the conversion of 

coordinated carbonyl groups during the decomposition of the 

supported cluster. This is further substantiated by the 

identification of methyl chloride, ethyl chloride, water, 

carbon dioxide and a trace amount of acetylene in 

experiments using isotopically labelled carbon-13 carbon 

monoxide. The mass intensity report for acetylene is shown 

in Figure 3-11. A total of 1.4% carbon-13 incorporation 

above natural abundance into methyl chloride, as well as 

between 32-45% incorporation of carbon-13 into carbon 

dioxide was observed by GC/MS. The incorporation of any 

carbon-13 into ethyl chloride could not be detected by 

GC/MS. The small amount of incorporation of carbon-13 into 

methyl chloride is not surprising since the original 

supported iridium clusters contain entirely carbon-12 

carbonyl ligands. The 55-68% carbon-12 carbon dioxide that 

was observed may arise from either the decomposition of the 

cluster or a water-gas shift mechanism. This water-gas 



Ill 



>1 

4-1 




o= catalyst (1.18% Ir) 



10 20 30 
Time (hours) 



40 



Figure 3-10. A Graph of Initial Methyl Chloride 
Activity Versus Time for a Catalyst 
Prepared in Toluene 



112 



26 



loor 

80- 

60' 
40- 

20' 



Observed spectrum 



I ' " ' ■ 1 ' ' ' ' I ' ' ' ' I " ' ■ ' I ' ' ' ' I ' ' ' ' J ' 
10 15 20 25 30 35 40 



Reference spectnm' 



83 



10 30 50 90 



Figure 3-11. Mass Intensity Report for Acetylene 



113 
shift mechanism would involve the substitution of carbon-13 
carbon monoxide for carbon-12 carbonyl groups in the cluster 
followed by the subsequent activation of the resulting 
carbon-12 carbon monoxide over the hydroxylated inorganic 
oxide support. The absence of carbon-13 incorporation into 
ethyl chloride suggests that the ethyl chloride is formed 
from HCl(g) attack upon surface ethoxide groups. The 
presence of surface ethoxide groups arises from the reaction 
of the surface hydroxyl groups with the ethanol that is 
inherently produced during the phosphination of the support 
as shown in Equation 3-11. 

Another method of catalyst preparation involved the 
substitution of a carbonyl ligand in Ir (CO) with a 
phosphine group of the functionalized support. It has been 
reported that this reaction performed in toluene produces 
entirely the triphosphine substituted tetrairidium carbonyl 
cluster.''" This supported cluster (1.48% Ir) was observed 
by gas chromatography to produce methyl chloride, 
acetaldehyde and ethyl chloride. The cluster's activity for 
the formation of methyl chloride was found to be 6.3 5 x 
lO"''"^ moles CH CI sec" g~ at 75°C. This activity 
increased to 2.11 x 10~ moles CH^Cl sec" g~ upon raising 
the temperature to 200 °C. The activity observed for this 
supported cluster and the other supported clusters prepared 
in the absence of 2-methoxyethanol supports the formation of 
the methyl chloride through the reduction of carbon 
monoxide. It is proposed that this activity of 1-9 x 



114 

_-\ -^ -1-1 

10 moles CH-,C1 sec g associated with the supported 

clusters is due to the conversion of bound carbonyl groups 

and the reduction of carbon monoxide. Since this activity 

is considerably less than previously thought, the total 

amount of methyl chloride produced in a 40-day period is 

below one catalyst ( 3 -P Ir (CO) ) turnover. Therefore, a 

discussion concerning the catalytic ability of the system 

will not be undertaken. 

Minor Impurity Routes to Methyl Chloride 

Several control reactions employing inorganic oxide 

supports either prior to or after phosphination in solvents 

other than 2-methoxyethanol were run. A large amount of 

methane, ethylene or ethane, ethanol, ethyl chloride and 

acetaldehyde was observed by gas chromatography for a 

control support which had been prepared in 2-ethoxyethanol. 

A trace amount of methyl chloride also was observed by gas 

chromatography over a short period of time at 125 °C. In 

general, it seems that whenever ethanol or ethyl chloride is 

present, small quantities of acetaldehyde, ethylene or 

ethane and methyl chloride are observed. It also was 

noticed that the presence of HCl (g) is required for the 

formation of these products. The formation of these 

products was observed to be more dramatic for inorganic 

123 
oxides of increasing acidity, ^^'^2 ^ ^■'•2°3 ^ Zeolite. 

The quantity of these products was observed to depend upon 
the amount of ethanol or ethyl chloride initially present. 
A similar observation has been reported for control 



115 

reactions involving HCl(g) and AlCl at elevated 
temperatures. This suggests that the ethylene, 
acetaldehyde, as well as the trace quantity of methyl 
chloride arises from the decomposition of either ethanol or 
ethyl chloride. The ethanol and ethyl chloride most likely 
arises from the surface ethoxide groups formed during 
support phosphination, as well as from adsorbed 
2-ethoxyethanol solvent. 

A reaction scheme consistent with these observed 
results can be proposed from the available literature. 
Several reports indicate that ethanol can be converted at 
elevated temperatures over inorganic oxides to 

ethylene, "'"^^~''"^"'" ethane and acetaldehyde or over 

172 
supported metal catalysts to ethylene and 

acetaldehyde. ■'■'^^'"'' This reaction scheme is complicated in 

that these decomposition products of ethanol can 

interconvert under the reaction conditions employed. For 



instance, ethylene over aluminum trichloride in the presence 

174 
of HCl(g) has been reported to form ethyl chloride. It 

also has been found that acetaldehyde is a by product in the 

174 
direct hydration of ethylene to ethanol. Finally, 

hydrocarbon cracking reactions were suggested to be 

140 
occurring in the homogeneous Ir^ (CO) ^^/^^'^^S'^^'^^ system. 

Similar reasoning gives a possible explanation for the 
observed formation of methyl chloride. It has been reported 
that ethylene decomposes to acetylene over various metal 
surfaces. ■'■'^^"''"^'^ Recall that a trace quantity of acetylene 



116 
was identified by GC/MS in the reactions involving the 
supported cluster. Acetylene chemisorbed onto a support has 

been reported to decompose to carbon atoms between 

177 
25-125°C. A similar carbon bond rupture has been 

postulated for the chemisorption of acetylene onto a 

178 
hydrogenated iridium surface. The formation of other one 

carbon species, such as formaldehyde and carbon monoxide has 

been reported for the oxidation of acetylene over palladium 

179 
black. It is possible that either formaldehyde or other 

reduced carbon species on the inorganic oxide support could 

form methyl chloride in the presence of hydrogen and HCl (g) . 

This is supported by the observation that formaldehyde and 

HCl(g), over a zeolite, produces a mixture of methyl 

180 
chloride and methylene chloride. 

Several control reactions using a phosphinated support 

prepared in toluene showed the formation of a trace quantity 

of methyl and ethyl chloride. The lifetime and the amount 

of activity of these control resins was observed to be 

reduced as compared to the control resins prepared in 

2-ethoxyethanol. The formation of the ethyl chloride is 

assumed to arise from the ethoxide groups on the support 

which were formed during the phosphination procedure. The 

formation of the trace quantity of methyl chloride is 

proposed to arise from the decomposition of toluene adsorbed 

onto the support. Recall that the presence of residual 

toluene in the supported cluster reactions was confirmed by 

GC/MS. It previously has been reported that toluene is 



117 

converted in 50% yield to benzene and methane at 400°C over 

181 
a supported iridium metal catalyst. A similar conversion 

of toluene to benzene and methane has been reported for a 

18 2 
supported platinum metal catalyst. In this case, an 

increase in activity was observed by pretreating the support 

with HCl(g). Furthermore, it is thermodynamically possible 

that toluene could decompose totally to methane or to carbon 

181 
atoms as shown in Table 3-8. It is possible that an 

interaction of these reduced carbon species with hydrogen 

and HCl(g) could produce the trace quantity of methyl 

chloride observed over a short period of time in these 

control reactions. 

The proposed explanation for the low level formation of 

the products observed in the control reactions prepared in 



Table 3-8. Thermodynamic Data Concerning the Decomposition 
of Toluene 

Reaction ^^298 ^^"^ Mole" ) 



\J-^^3 + H^ ^<^ + CH^ -42.0 

<Q>-CH2 + lOH^ ► 7CH^ • -574.3 

<^CH3 >-7C + 4H2 -49.8 



2-ethoxyethanol or toluene is supported by the presence of 
several impurities inherent to the reactant gases. It has 

been reported that a 600 ppm impurity of oxygen is present 

18 3 
in the carbon monoxide gas that was used. A low level 



118 



impurity of methane in this CO feed gas also was identified 
by gas chromatography. The technical grade HCl (g) has been 
reported to contain a 0.4 weight percent hydrocarbon 

impurity which consists of a mixture of ethylene, 

183 
1,1-dichloroethane and ethyl chloride. On the other 

hand, the semiconductor grade HCl(g) was reported to contain 

183 
only a 10 ppm methane impurity. The formation of 

ethylene, acetaldehyde, ethyl chloride and methyl chloride 

still was observed for the reaction of the phosphinated 

support with carbon monoxide, hydrogen and the semiconductor 

grade HCl(g). In this case, ethylene was found not to be 

present in the reactant gas mixture. Schwartz and Kitajima 

have reported the conversion of methane and HCl (g) to methyl 

184 
chloride over supported rhodium complexes. The supported 

13 
iridium clusters were tested in the presence of CH^, HCl, 

H and CO. The incorporation of carbon-13 into methyl 

chloride could not be substantiated by the results obtained 

by GC/MS. 

One last possible carbon source for methyl chloride 

formation that must be considered is the ethyl chain and 

phenyl groups of the phosphinated silane linkage. The 

(C^H^O) ^Si(CH-) _PPh„ support linkage has carbon-carbon 
2 5 3 2. 2. £■ 

stretches in the phenyl groups which exhibit infrared 
absorptions at 1587 (s) and 1572 (w) cm"-*- as shown in Figure 
3-12. These infrared absorptions have been observed in the 
infrared spectra obtained for supports that have been 
phosphinated with the silane linkage and subsequently 



119 



A= (C2H50)3Si(CH2)2PPh2 

B = 3>^Ph2lr^(CO) j^j^ + CO/K2/HCI at 100*0 



(film) 




1680 



1618 
Wavenuinbers 



1556 
(cm-l) 



1494 



Figure 3-12. Infrared Spectrim of the C=C 

Vibrations in the Phenyl Groups 
of the Silane Linkage 



120 
functionalized with the iridium cluster. However, these 
infrared absorptions observed for the supported silane 
linkage are much broader and not as intense as those 
observed for the homogeneous complex, (C„Hj.O) -Si (CH- ) „PPh„ . 
This effect is associated with the occurrence of a broad 
absorption in this region due to the support, as well as 
water vapor. Upon activation of the supported cluster with 
CO, H„ , and HCl (g) , the infrared absorptions assigned to the 
phenyl groups of the phosphine linkage were observed to 
increase in intensity. The infrared spectrum in this region 
of an activated supported iridium carbonyl cluster is shown 
in Figure 3-12. The presence of these absorptions at 1587 
and 1572 cm~ in the activated sample of the supported 
iridium carbonyl cluster suggests that the phosphine linkage 
is not attacked by the HCl(g) feed gas. Therefore, it is 
doubtful that decomposition of the phosphine linkage is 
responsible for the formation of any of the observed 
reaction products. 

The preceding observations indicate that there exists a 
variety of inherent minor impurity routes through which 
methyl chloride may be formed. These minor routes can 
account for some of the initial activity observed in the 
supported cluster reactions for the formation of methyl 
chloride and the other observed products. It is proposed 



that after approximately the first several days the activity 

-12 -1-1 

of 1-9 X 10 moles CH^Cl sec g observed for the 

supported cluster prepared in the absence of 



121 

2-inethoxyethanol results predominantly from the conversion 

of bound carbonyl groups and the reduction of carbon 

monoxide. 

Structural Determination of the Supported Iridium Clusters 
by Infrared Spectroscopy 

The supported iridium carbonyl cluster prepared through 

the reduction of IrCl (CO) _ (p-toluidine) in 2-methoxyethanol 

was identified by infrared spectroscopy to be a mixture of 

the mono- and di-phosphine substituted tetrairidium carbonyl 

clusters. All infrared data obtained for the synthesis of 

these iridium clusters in various solvents is summarized in 

Table 3-9. The infrared spectra for the supported iridium 

clusters prepared through the reduction of IrCl(C0)2- 

(p-toluidine) in 2-ethoxyethanol or toluene did not 

identically match the spectra obtained for the supported 

clusters prepared in 2-methoxyethanol. A mixture of mono- 

di- and tri-phosphine substituted tetrairidium carbonyl 

clusters were identified by infrared spectroscopy for the 

supported clusters prepared in 2-ethoxyethanol. The 

tri-phosphine substituted iridium carbonyl cluster was the 

predominant species identified in the reactions involving 

toluene as the solvent. The remaining infrared absorption 

at 1960 cm" in both cases can be assigned to a supported 

mononuclear or dinuclear iridium species, such as 

3 -PPh IrCl(C0)2 or 3 - [PPh2lr (CO) ^1 2 • ^he formation of a 

mononuclear species most likely could arise through the 

substitution of the para-toluidine ligand in IrCl (CO) 2" 



122 
(p-toluidine) with the supported phosphine , 3~0Si(CH ) PPh . 
It also has been reported that the reduction of IrCl (CO) - 

(p-toluidine) in the presence of excess triphenylphosphine 

158 
leads to the formation of [Ir (CO) ^PPh^l 2 • ^^ analogous 

reaction could be suggested to form a supported dinuclear 

species, such as 3- [PPh2lr (CO) ^1 2 • 

The formation Of this supported dinuclear species, may 

be a result of trapping an intermediate complex that 

normally would lead to the formation of a tetranuclear 

iridium cluster. It has been proposed that the conversion 

of IrCl (CO) 2 (p-toluidine) into phosphine substituted 

tetrairidium carbonyl clusters proceeds through the 

159 
formation of iridium dinuclear complexes. This proposal 

is supported by the reported conversion of [Ir (CO) ^PPh^] ^ 

into Ir. (CO) ^(PPh^) 3,"'^^^ as well as the formation of 

Ir.(C0)^2 ^^°^ Ir2 (CO) g."*"^ Similarly, the formation of a 

mononuclear species, such as 3-PPh2lr (CO) 2CI could be 

related to the trapping of a reaction intermediate. This is 

suggested by the reported formation of Ir^(C0)^2 f^oni the 

— 187 
anion [Ir (00)2012! • 

The predominant species identified by infrared 
spectroscopy for the supported cluster prepared in toluene 
through phosphine substitution of the carbonyl ligands in 
Ir (CO) „ was the tri-phosphine substituted tetrairidium 
carbonyl cluster. The observed infrared spectrum for this 
supported cluster is shown in Figure 3-13. Attempts to 
isolate either mono- or di-phosphine substituted 



123 



•d 

4J 

u 
o 

04 

3 
w 

q; 

4J 

V^ 
O 
M-l 

0) 
C 

-H 
(0 

4J 

XI 

O 

rO 
4-1 
(0 
Q 

0) 
U 
rO 



-.__0 (-I'M— o -< ~ 



o o 

u 

V — 



m rncnrhrhrn mrnm rn 



xrli 



> > > 3 



e J in J 



o — 

U 0) 

— c 



— u — 



o — — 



3 ►-< 



O -o 

c 

O CP 



T3 >, 
<U C 

u o 



124 




+j 

s 
o 
u 

-a 

0) 

c 

•H 
rO 
■P 
jQ 
O 



n 


o 


— . 


CJ 


CN 




x: 


'S 


0^ 


i-i 


04 


H 


^— ' 




1 


£1 


Ml 


-P 




•H 


<u 


5 








u m 



luoTssTuisueaj, % 



I 

m 

n 

3 

-H 
lit 



125 

tetrairidium carbonyl clusters by this procedure have been 
unsuccessful. The isolation of predominantly the tri- 
phosphine substituted iridium carbonyl cluster is a result 
of the rate acceleration for CO dissociation observed for 

progressive phosphine substitution of carbonyl ligands in 

1 90—1 92 
Ir.(CO),-. Interpretation of the infrared data shown 

in Table 3-9 suggests that the formation of the mono- 

phosphine tetrairidium carbonyl cluster has been obtained 

only when the cluster was prepared in 2-methoxyethanol . It 

has been reported that the mono-phosphine substituted 

tetrairidium cluster shows greater activity than the di- or 

tri-phosphine substituted clusters towards olefin 

193 
hydrogenation. Analogously, it could be argued that the 

mono-phosphine substituted tetrairidium carbonyl cluster is 

more active than the di- or tri-phosphine substituted 

clusters for the conversion of synthesis gas and HCl(g) into 

methyl chloride. 

To ascertain the validity of this argument, further 

attempts to prepare the supported mono-phosphine iridium 

carbonyl cluster in a solvent other than 2-methoxyethanol 

were done. It was found that a mixture of mono- and 

di-phosphine tetrairidium carbonyl clusters could be 

obtained in toluene using a phosphinated support in which 

the remaining hydroxyl groups of the surface had been 

silanated with (CgH^) 2SiCl2 as shown in Equation 3-18. 

It seems that the size of the surface phenyl groups 



126 



sterically hinders the phosphine substitution of more than 
two carbonyl groups in Ir.(CO)^_. The comparison of the 
infrared spectrum for the supported clusters prepared in 



OSiC.H.PPh, 




(3-18) 



this manner with the spectrum of the supported clusters 
prepared in 2-methoxYethanol is shown in Figure 3-14. 
Although a high percentage of the di-phosphine substituted 
tetrairidium carbonyl cluster is present in this sample, the 
existence of the mono-phosphine substituted tetrairidium 
carbonyl cluster is suggested by the presence of the 
bridging carbonyl infrared adsorption at 1845 cm . It is 
predicted that utilizing a smaller ratio of support 
phosphination: support silanation that the mono-phosphine 
tetrairidium carbonyl cluster may be formed exclusively. 
This was not investigated since the resulting iridium 
carbonyl cluster concentration would be below the available 
infrared detection limits. Testing this mixture of 
supported mono- and di-phosphine iridium carbonyl clusters 
for synthesis gas conversion resulted in an activity of 6.21 
X lO"'''^ moles CH^Cl sec"''' g~ at 75°C. This activity is 



127 



A = Supported clusters prepared in 
2-inethoxyethanol 

B = Supported clusters prepared in 
toluene 

• = 3-PPh2lr^(CO) j^j^ absorptions 

O = 3-(PPh2)2lr^(CO) j^Q absorptions 




2301 



2084 



1917 



1750 



Wavenumbers (cm ) 



Figure 3-14. A Comparison of the Infrared Spectra 

of the Mono- and Di-phoshine Substituted 
Tetrairidium Carbonyl Clusters Prepared 
in 2-inethoxYethanol and toluene 



128 

very similar to the activity previously observed for the 

supported clusters, 3-PPh2lr^ (CO) ^ „ and B-PPh^Ir^ (CO) g , 

prepared in either 2-ethoxyethanol or toluene. Therefore, 

it is proposed that the mono-, di- and tri-phosphine 

substituted tetrairidium carbonyl clusters under the same 

reaction conditions either show similar activity or act as 

precursors to iridium species that show similar activity for 

the formation of methyl chloride. All data discussed in the 

following sections concern supported complexes or control 

supports that have been prepared in the absence of 

2-methoxyethanol . 

Examination of the Decomposition of the Supported Clusters 
by Infrared Spectroscopy 

Infrared spectroscopy was used to investigate the 

possibility that the supported iridium clusters may have 

been converted to other iridium species upon exposure to 

synthesis gas and HCl(g) at 75°C. All of the activated 

supported clusters, either the mono-, di- or tri-phosphine 

substituted iridium carbonyl clusters, showed a predominant 

absorption in the infrared spectrum at either 2069 or 2040 

cm" after exposure to the reactant gas mixture. The 

positioning of the predominant absorption in the infrared 

spectrum at 2069 or 2040 cm was observed to be dependent 

upon the weight percentage of iridium present and not upon 

the nature of the initial cluster. The typical infrared 

spectra that were obtained are shown in Figure 3-15. If the 

iridium weight percentage was below approximately 1.00% (low 



129 



(A) 
High Load System - (2.34% Ir) 




2301 2129 1884 

Wavenumbers (cm ) 



1701 



(B) 
Low Load System - (0.37% Ir) 



97 T 



96 



95 



94 



93 



92 




a = 2040 cm 



c = 1719 cm 



2301 2129 1884 

Wavenumbers (cm~ ) 



1701 



Figure 3-15. Infrared Spectra for the Supported Clusters 
After Exposure to the Reactant Gases at 
75 "C A) High Load Svstem B) Low Load 
System 



130 
loading) , the predominant absorption observed in the 
infrared spectrum was at 2040 cm . If the iridium weight 
percentage was above approximately 2.00% (high loading), the 
predominant absorption observed in the infrared spectrum was 
at 2069 cm . It also was found that an infrared spectrum 
containing both the absorptions at 2069 and 2040 cm could 
be obtained at 75°C for supported clusters containing 
between 1-2% iridium by weight. Since similar infrared 
spectra were obtained for the various supported clusters 
after exposure to the reactant gases, it is suggested that 
the clusters are converted into the same active species or 
precursor to the active species. Therefore, all of the 
supported clusters are expected to exhibit similar activity 
towards carbon monoxide reduction. 

Infrared absorptions at 1734 and 1719 cm as shown in 
Figure 3-15 were observed in the spectra obtained for the 
activated clusters at 75°C. These absorptions are more 
noticeable in the infrared spectra of the supported systems 
containing a lower weight percentage of iridium. Since 
these infrared absorptions suggest the presence of bridging 
carbonyl ligands, it is possible that multinuclear iridium 
carbonyl species are still present in the activated systems 
at 75°C. 

A change in the infrared spectra as shown in Figure 
3-16 was observed upon raising the temperature in the high 
and low iridium loaded systems. A shifting of the infrared 



131 





1= 


1 


1 


1 




0) 


e 


b 


fc 




■U 


o 


o 


U 




en 








^^ 


>. 


o 


rr 


a> 


m 


en 


■^ 


m 


rH 


N..^ 




o 


r- 


r~ 




Ti 


(N 


fH 


M 




m 













II 


II 


II 




k4 












(0 


^ 


u 




? 










O 










^A 










UOTSSTUISUBJI, % 



^^ 


1 


1 


1 


1 




M 


e 


e 


B 


e 




M 


o 


u 


o 


o 




dP 


VO 


o 


^ 


a^ 




^ 


<N 


en 


n 


H 




ro 


O 


<Ti 


r- 


t~~ 




. 


CM 


rH 


r-H 


H 




04 












^ 


II 


II 


II 


II 




e 


M-l 


Di 


x: 


•r-T 




0) 












-p 












en 


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iH 


r-{ 


H 


iH 


— >i 


1 


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— ' 


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in 


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n 


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\o 


IT) 


.J 


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r-l 


r^ 


O 


O 




CN 


CM 


<N 


rs 


CN 


J= 












Cn 


II 


II 


II 


II 


II 


•H 












K 


10 


XI 


u 


TI 


0) 



■p 

(0 

0} 

6 
Q) 
4J 

tn g 

en -P 
^^ tn 

H e >i 

•H m 
tn s-i O 

oj '"''"' 

3 (0 

q Oh:; 
S xi m 

•H 

"e 

C -P 
IT3 tn 


a T! 

to 

0) 

x: J 

x: 

O -H 

IH IE 

T) -- 

c — 

■H 
(0 

+J 
ja 
o 




UOTSSTUISUPJCJ, % 



O QJ 

OJ Ou 

a S 

en OJ 

EH 

TI 

0) tn 



132 

-1 -1 

absorption at 2150 cm to 2137 cm and a decrease in the 

intensity of the bridging carbonyl absorptions at 1734 and 

1719 cm" were observed for the high load iridium system 

upon raising the temperature from 75°C to 125°C. At 

temperatures exceeding 200 °C, the infrared spectrum for this 

system exhibited only one absorption at 20 55 cm . The 

infrared spectrum for the low load iridium system upon 

raising the temperature above 200 °C showed the disappearance 

of the bridging carbonyl absorptions and a broadening in the 

primary absorption at 2040 cm" . Since the infrared spectra 

for the high and low load iridium systems change upon 

raising the temperature, it is suggested that further 

decomposition of the iridium species occurs at these 

elevated temperatures. 

Verification That A Discrete Molecule Complex Exists in the 
Activated Iridium Systems 

It is evident from the change that has occurred in the 

infrared spectra upon raising the temperature that the 

supported clusters are being converted to either another 

iridium carbonyl complex or to a carbonylated iridium metal 

surface. Vannice has reported that halogenated hydrocarbons 

can be produced from synthesis gas and HCl(g) over inorganic 

194 
oxide supported iridium metal as shown in Equation 3-19. 

It should be noted that this reaction is claimed to 

occur between the temperature range of 200-1000°C, It also 

has been reported that methyl chloride is temporarily 

observed during the conversion of synthesis gas to 



133 

hydrocarbons at 250°C over iridium metal supported on an 

195 
inorganic oxide support. The chloride content was found 

to originate from the initial impregnation of the support 

with IrCl-' 3H_0 followed by its thermal deposition. It has 



group VIII mtjtal/ inorganic oxicje 

CO * Hj * m ■ >► CH3X > CH^Xj + CHX3 * CX^ - CjH^X H. CjHg 

200-1000°C, 0.1-500 atm 



(3-19) 



been reported that the phosphine substituted tetrairidium 
carbonyl clusters supported on inorganic oxides decompose to 
iridium metal above 180°C under an atmosphere of carbon 
monoxide and hydrogen. ■'■^■'" Similarly, Ir^ (CO) ^2 physisorbed 

on an inorganic oxide support has been shown to begin to 

196 
decompose by CO dissociation between 50-350°C. 

Furthermore, the predominant infrared absorption assigned 

for carbon monoxide absorption onto an iridium metal surface 

has been reported to range from 2010 cm at low metal 

-1 197-199 

coverage to 2093 cm at metal saturation. . Since 

the location of this broad infrared absorption has been 

reported to be influenced by a variety of parameters, such 

19 7—199 
as the extent of metal loading , the metal particle 

199 200 

size , the strength of the support interaction , the 

■,200-202 , ^, 
oxidation state of the indium metal and the 



134 

201 
temperature of the system , it is possible that the 

changes observed to occur in the infrared spectra of the 

supported clusters upon exposure to the reactant gases 

(Figure 3-15) and an increase in reaction temperature 

(Figure 3-16) could be explained by the formation of 

metallic iridium. 

Several control reactions using metallic iridium 

adsorbed onto an inorganic oxide support were run in order 

to ascertain if metallic iridium could account for the 

formation of the methyl chloride observed at 75°C in the 

supported cluster systems. Either Ir^ (00)^^2 °^ IrCl2»3H20 

was physically adsorbed onto alumina to give supports 

containing 1.69% Ir and 2.30% Ir, respectively. Each sample 

was pretreated by calcination at 250°C under hydrogen for at 

least five hours. The resulting gray supports showed the 

formation of only a trace amount of methyl chloride at 75°C 

under CO, H_ and HCl(g). The formation of this methyl 

chloride at 75°C was observable only upon using a stagnant 

reactor. An increase in activity into the range of 8.9 x 

10~ - 2.6 X lO" moles CH_C1 sec~ g~ was observed for 

these metallic iridium systems upon raising the reaction 

temperature to 200°C. Deactivation of these metallic 

iridium systems was observed to occur upon exposure to air. 

This is not surprising since it has been reported that 

metallic iridium upon exposure to air undergoes oxidation 

197 2 03 
and agglomeration to large crystallites of Ir02. ' The 

infrared spectra observed for these activated metallic 



135 
iridium systems contained one broad, weak absorption as 
shown in Figure 3-17. The location of this absorption was 
found to depend upon the extent of the metal loading. It is 
suggested from the interpretation of these results that the 
formation of metallic iridium in the supported iridium 
cluster systems cannot account for the observed activity of 
1-9 X lO""""^ moles CH^Cl sec"""" g""*" at 75°C. 

In order to substantiate the existence of a discrete 
molecular complex in the activated supported cluster systems 
at 75°C samples of these systems taken prior to and after 
exposure to CO, H and HCl(g) were examined by x-ray 
photoelectron spectroscopy (ESCA) . No change in the binding 
energy of the iridium core electrons as shown in Figure 3-18 
was observed for analysis of the supported cluster (2.34% 
Ir) before and after activation. This suggests that the 
iridium species present after activation is a discrete 
molecular complex. This is further supported by a similar 
conclusion reported by Dow Chemical Company in an 
independent analysis of the supported iridium cluster by 
X-ray photoelectron spectroscopy. The increased width of 
the signal in the iridium photoelectron spectrum as compared 
to molecular standards suggests that the iridium may be in 
several oxidation states. 

The activity of the supported phosphine substituted 
tetrairidium carbonyl clusters towards the formation of 
methyl chloride was monitored over the temperature range of 
75-300°C in order to determine if the temperature of the 



136 






p 




•rH 




Ti U 




-H 




S-l O 




H O 




CN 




u 




•H 4-1 




.H (0 




iH 




(C ^ 


,_^ 


-P Cji 


H 


0) — 




r -H 


R 


u 


O 


s-l m 









iw fO 


tn 


c 


u 


^3 (0 


0) 


<U 





C fv 




■H ffi 


3 


(0 


c 


4J - 


0) 


X! O 


> 


O U 


(0 




:? 


§ 




3 4-1 




U 




+J QJ 




U H 




OJ 3 




W 




CO 




a 




T3 X 




(1) W 




U 




(0 Vj 




^ OJ 




IW 4-1 




C 4-1 




H < 



lUOTSSTUISUejJ, % 



137 



A = initial supported cluster 
(2.34% Ir) 

B = Supported cluster (2.34% Ir) 
exposed to CO/H^/RCl at 100 *C 




-76 -72 -68 -64 

Binding Energy, eV 



-60 



-56 



Figure 3-li 



X-ray Photoelectron Spectrxnn for the Supported 
Clusters Before and After Activation 



138 

system would have any effect on the formation of metallic 

iridium. The activity at 75°C was observed to average 3.27 
,-11 .-_ ^., r.. ___-! _-l 



X 



10 moles CHoCl sec g . A decrease in activity to 



-12 -1 -1 

9.20 X 10 moles CH-.C1 sec g was observed between the 

temperature range of 125-200°C. The activity increased to 

4,23 X 10" moles CH^Cl sec" g~ upon raising the 

temperature above 200°C. It is proposed that between 

75-125°C under CO, H and HCl (g) the supported clusters 

decompose to form other discrete iridium complexes. The 

decrease in activity observed between 125-200 °C can be 

explained by the gradual decomposition of these discrete 

iridium complexes to the less active metallic iridium. At 

temperatures exceeding 200 °C most of the discrete iridium 

complexes are decomposed to metallic iridium. The activity 

of the system between 200-300°C was observed to increase 

because the metallic iridium becomes more active towards the 

formation of methyl chloride in this temperature range. 

Investigation by Infrared Spectroscopy of the Discrete 
Iridium Species Present in the Activated Systems at •75°C 

Collman et al. proposed that IrCl (CO) ^ was either the 
active catalyst or catalyst precursor for carbon monoxide 
conversion in the homogeneous Ir ^{CO) ^^^^^^^2'^^^'^ melt 
system. ■'"^■'" It also has been reported that [IrCKCO)^]^ will 

react with a phosphine substituent to form dinuclear iridium 

20 4 
complexes as shown in Equation 3-10. A supported indium 



[IrCl(C0)3]^ + nKPR2 >-n/2 [ (OC) 3lrPR2] 2 + "KCl (3-10) 



139 
carbonyl complex was prepared by reacting IrCl (CO) , with a 
phosphinated support in refluxing toluene. The infrared 
spectrum for this supported iridium complex is very similar 
to that obtained at 75°C for the activated supported iridium 
clusters containing a low iridium content as shown in Figure 
3-19. The presence of the bridging carbonyl absorptions at 
1734 and 1719 cm" indicates that the reaction between 
IrCl(CO)- and the phosphinated support yields a multinuclear 
iridium species. Unfortunately, the supported iridium 
carbonyl complex responsible for the infrared absorptions at 
2045 (vs), 1990 (shoulder) , 1734 (m) and 1719 (w) cm" as shown 
in Figure 3-19 has not been identified. There is only 
limited data available in the literature concerning the 
characterization of multinuclear phosphine substituted 

iridium clusters which contain bridging carbonyls and 

V--1 • 1 ■ ^ 151,158,189,205,206 ^v,^^r,iw 
possibly chlorine ligands. The only 

iridium complex that was found in the literature to give 

rise to an infrared spectrum which is similar to the 

observed infrared spectrum (Figure 3-19) is the multinuclear 

"? fi 

species, [Ir2 (CO) ^Cl ( (Ph) 2PCH2P (Ph) 2) 2HBPh^] . The 
supported multinuclear iridium species prepared from the 
reaction of IrCl (CO) with the phosphinated support was 



expo 



sed to CO, H„ and HCl (g) at 75°C. The activity of this 



supported complex at 75°C for the formation of methyl 

-12 -1 

chloride was observed to be 9.10 x 10 moles CH^Cl sec 

g" . A change in the infrared spectrum of this supported 

iridium complex, as shown in Figure 3-20, was observed to 



140 



A= Initial IrCl (CO) -./3-PPh- containing 
1.66% Ir -3 / 

B= Supported cluster (Low Load - 0.37% Ir) 
exposed to CO/E^/HCl at 75 'C 




2301 



2068 



1884 
-1, 



1701 



Wavenumbers (cm ) 



Figure 3-19. A Comparison of the Infrared Spectra Obtained 
for the Supported Clusters (Low % Ir) and for 
IrCl (CO) ^ on a Phosphinated Support 



141 



a 


= 


2136 


cm 


b 


= 


2080 


cm 


c 


= 


2068 


cm 


d 


= 


2045 


cm 


e 


= 


1990 


cm 


f 


= 


1734 


cm 


g 


= 


1719 


cm 



-1 
-1 
-1 
-1 
-1 
-1 
-1 




2300 



2067 



1884 



1701 



Wavenumbers (cm ) 



Figure 3-20. Infrared Spectrum of IrCl (CO) + 3-pph, 



at Various Temperatures 



142 
occur after exposure to CO, H- and HCl(g) at various 
temperatures. At 75°C the growth of an infrared absorption 
at 2068 cm' was observed to coincide with the disappearance 
of the shoulder at 1990 cm" . The reappearance of this 



POP 

I II I 

o^', CO 

[CI - Ir Ir ] [BPh^] 



\<^ 



POP 



CO 



2047 (m), 1998 (vs), 1759 (s) cm"''- 



shoulder along with a shoulder at 2080 cm and a new 
absorption at 2136 cm" was observed at 125°C. The primary 
infrared absorption at 200°C was observed to be a broad band 
at 2045 cm"'''. In general, the intensity of the bridging 
carbonyl absorptions at 1734 and 1719 cm" was observed to 
decrease as the reaction temperature was increased. The 
infrared spectrum taken after the supported complex had been 
activated at 125°C was observed to resemble the spectra 
obtained for the supported clusters containing a high 
iridium content (Figure 3-16) and for IrCl (CO) ^ (Figure 
3-21) . 

A sample of IrCl (CO) -. physically adsorbed onto alumina 
gave rise to an infrared spectrum that is similar to that 
obtained for the unsupported IrCKCO)^- This spectrum also 



143 
is very similar to that observed at 75°-125°C for the 
supported clusters containing a high iridium content as 
shown in Figure 3-21. A change in the infrared spectrum as 
shown in Figure 3-22 was observed upon exposing the 
IrCl (CO) _ adsorbed on alumina to CO, H and HCl(g) at 
various temperatures. Primarily, the growth of a new 
absorption at 2109 cm" and the splitting of the 2142 cm 
absorption into two absorptions at 2154 and 2139 cm was 
observed at 75°C. This change which has occurred in the 
infrared spectrum of IrCl(CO)^ on alumina at 75°C is 
suggested to be caused by the formation of different isomers 
of [IrCKCO)^] . The infrared spectrum at 200°C for the 
activated irCl (CO) -/alumina exhibited only a broad 

absorption at 2073 cm . The formation of only a trace 

-14 -1 -1 

amount of methyl chloride (<10 moles CH^Cl sec g 

accumulated in a stagnant reactor for twelve hours) was 

observed for this system at 75°C. The formation of this 

trace amount of methyl chloride most likely occurs through 

the previously discussed minor impurity routes. The 

activity of the system was observed to increase to 6.30 x 

lO""*"^ moles CH^Cl sec""^ g" at 200°C. 

The infrared and activity data obtained for IrCKCO)^ 

adsorbed on alumina and supported on phosphinated alumina 

parallels the results obtained for the supported clusters at 

various temperatures. It is proposed that the active 

species present in the supported cluster systems is a 

multinuclear iridium species characterized by infrared 



144 



A= Supported cluster (high load - 2.34% Ir) 
at 125 'C 

B= Supported cluster (high load - 2-34% Ir) 
at 75 *C 

C= IrCl(C0)3 

D= IrCKCO)^ / Alumina (1.40% Ir) 



c 
o 

-H 
CO 

a 

e 

CO 

c 

10 

u 




2301 



2066 



2021 cm 



1884 
-1, 



1701 



Wavenumbers (cm ) 



Figure 3-21. A Comparison of the Infrared Spectrum of 

IrCl(CO) 3/Alumina With the Spectra Obtained 
for the Activated Iridium Clusters at 
Various Temperatures 



145 



200 



c 
o 

•H 
(0 

CO 

•H 

G 

la 
u 




Initi 



2206 



2066 
Wavenuinbers 



1884 
(cm'^) 



1701 



Figure 3-22. Infrared Spectrin of I rCl (CO) -^/Alumina at 
Various Temperatures 



146 
absorptions at 2045, 1734 and 1719 cm" as shown in the 
infrared spectrum (Figure 3-19) of the activated supported 
cluster containing a low iridium content. Similar infrared 
absorptions (Figure 3-19) and activity was observed for the 
IrCl(CO)T supported on phosphinated alumina. The presence 
of this multinuclear species also is indicated in the 
infrared spectrum (Figure 3-15) of the supported clusters 
containing a high iridium content by the presence of weak 
absorptions at 1734 and 1719 cm" . Thus, the presence of 
this multinuclear species can explain the similarity 
observed in the activity of the supported cluster systems 
containing either a high or low iridium content. This 
proposal is further supported by the observed inactivity of 
the IrCl (CO) -/alumina system along with the absence of any 
bridging carbonyl absorptions in the corresponding infrared 

spectrum (Figure 3-22) . Final support for this proposal 

13 12 
arises from the exchange of CO for CO in this 

multinuclear iridium complex as shown in the infrared data 

portrayed in Figure 3-23. 

It is proposed that this active multinuclear complex is 

gradually decomposed to IrCl (CO) ^ under the reaction 

conditions. The rate of this decomposition is increased as 

the temperature of the system is raised. This proposal is 

based on the temperature induced growth of the infrared 

absorptions in the infrared spectrum (Figure 3-20) of the 

active multinuclear complex which are similar to those of 

IrCKCO)^ (Figure 3-21). It also is suggested that the 



147 



A = Sample exposed to ■'■^C0/H2/HC1 at 75 "c 
B = Sample exposed to "'■^CO/HVHCl at 75 "c 



c 
o 

•H 
(0 
CO 

•H 

e 

CO 

c 
tn 
u 




2300 



2067 1884 

Wavenumbers (cm~-^) 



1701 



Figure 3-23. Infrared Spectr-um of a Mixture of the 

Supported Multinuclear Complex and IrCl (CO) 
Exposed to Carbon-13 Carbon Monoxide 



148 
IrCl(CO)^ is inactive towards the formation of methyl 
chloride. This is supported by the inactivity of the 

IrCl (CO) ^/alumina system. Further support arises from the 

13 
inability to substitute CO for the carbon-12 carbonyl 

absorption at 2067 cm" assigned to IrCl (CO) ^ in the 

infrared spectrum shown in Figure 3-23. As the temperature 

of the supported cluster systems is increased, the decrease 

in observed activity is caused by the formation of IrCKCO)^ 

and iridium metal. The formation of iridium metal is 

completed at temperatures above 200°C. The increase in the 

system's activity observed above 200°C previously has been 

194 
described by Vannice. 

In order to ascertain if similar results could be 

obtained from other supported iridium carbonyl complexes, 

Vaska's complex, IrCl (CO) (PPh^) 2 , was reacted in refluxing 

toluene with a phosphinated support. The supported complex 

was characterized by infrared spectroscopy to contain a 

mixture of products as shown in Figure 3-2 4. The carbonyl 

infrared absorption at 1965 cm" can be assigned to the 

carbonyl ligand in 3 - (PPh2) 2^^ (CO) CI. The remaining 

infrared absorptions at 2041, 2014, 1734 and 1719 cm" 

resemble those previously assigned to a multinuclear iridium 

complex. An activity of 1.38 x lO" moles CH^Cl sec g 

was observed at 75°C for this supported complex. The 

infrared spectrum of this supported complex (Figure 3-2 4) 

exposed to CO, H and HCl(g) at 75°C showed the 

disappearance of the 1965 and 2014 cm" absorptions and the 



149 



c 
o 

•H 
CQ 
(0 

-H 

E 
tn 
c 

^1 
E-i 




2238 



2066 1884 

Wavenioinbers (cm" ) 



1700 



Fig-ure 3-24. Infrared Spectriun of Vaska^s Complex 

Supported on Aliimina After Exposure to 
C0/H2/HCl(g) at Various Temperatures 



150 
growth of a shoulder at 2065 cm" . This shoulder at 2065 
cm" is suggested to be due to the formation of IrCl(C0)2. 
Above 200°C, the predominant infrared absorption was 
observed to be a broad absorption centered at 2041 cm 
Thus it has been shown that Vaska's complex, IrCl (CO) , and 
Ir.(CO)^„ supported on a phosphinated inorganic oxide 
support behave in a similar fashion under identical reaction 
conditions. 

Reevaluation of the Supported Cluster/AlCl , -NaCl System 
Since all early work involving the AlCl^-NaCl melt 
solvent was performed in the presence of residual 

2-methoxyethanol, it was decided worthwhile to reevaluate 

148 
the previously reported results. The supported iridium 

clusters made in the absence of 2-methoxyethanol were tested 

in AlCl.-NaCl under similar conditions to those previously 

-1 4P 
reported. At 25°C, the formation of a trace amount of 



methyl chloride was observed in a stagnant reactor. The 
activity was found to increase at 50°C to 3.97 x 10 moles 
CH3CI sec"""" g""*" and at 75°C to 4.76 x 10~ moles CH^Cl 
sec""*" g" . Allowing the reaction to proceed at 75°C for 
several hours caused partial melting of the AlCl^-NaCl salt. 
Control reactions of either AlCl^ or AlCl^-NaCl run in this 
temperature range were observed not to be active for methyl 
chloride formation. An infrared spectrum as shown in Figure 
3-25 of the activated supported iridium cluster in the 
AlCl--NaCl melt was observed to be similar to the reported 

^ n 140 ^^ 

spectrum (Table 3-2) obtained by Meutterties et al. It 



151 




E 


E 


e 


e 


E 


u 


u 


u 


u 


u 


o 


vo 


vo 


o 


CM 


<T> 


in 


OJ 


iH 


a\ 


ft 


r-i 


H 


H 


o 


CM 


<N 


<N 


CM 


cs 


II 


II 


II 


II 


II 



U -O 0) 




uoTSSTUisueajj % 



(U 

u 

3 
•H 

&4 



152 
should be noted that this infrared spectrum is completely 
different from those obtained for the supported clusters in 
the absence of the melt solvent. This indicates that the 
homogeneous system in the AlCl -NaCl melt solvent is 
different from the heterogeneous system employing HCl(g). 
Finally, it was confirmed that leaching of the supported 
iridium complex into the melt solvent did occur under the 
reaction conditions. Therefore, the use of a supported 
iridium cluster in an AlCl--NaCl melt solvent has no 
inherent advantages over using a discrete homogeneous 
complex. 
Proposed Mechanism for the Formation of Methyl Chloride 

A mechanism for the formation of methyl chloride in the 
supported cluster systems can be divided into two separate 
sections. The first section, which pertains to the gradual 
degradation of the supported cluster and to the decompo- 
sition of impurities that are present in the system, is 

13 
shown in Figure 3-2 6. Since CO studies have shown the 

incorporation of only 1.4% carbon-13 above natural abundance 

into methyl chloride, this cluster degradation and impurity 

decomposition pathway initially accounts for approximately 

98% of the observed methyl chloride. However, it is 

predicted that as the reaction time progresses a greater 

percentage of the methyl chloride observed will arise from a 

second pathway involving the reduction of carbon monoxide. 

The thermal decomposition of Ir^(C0)^2 °" ^" inorganic 

oxide support under an argon or carbon monoxide atmosphere 



153 




0) 




TJ 




■H 




U 









iH 




JS 




o 




iH 




>1 




4= 




+J 




(1) 




s 




14-1 









c 









■rH 




-p 




m 






B 


^ 


<u 





+J 


fo 


U] 




>1 


f-i 


en 


(0 




•H 


(-1 


-P 


0) 


•H 


+J 


c 


tn 


H 


3 




iH 


(1) 


U 


x; 




j-i 


^ 


u 


-H 


-a 


U-l 


-H 




Sh 


e 


H 


tn 




-H 


T) 


C 


a) 


(0 


-p 


£ 


M 


o 


o 


QJ 


a, 


? 


P- 




3 


TJ 


t/3 


(U 




tn 


0) 


.c 


a-p 







u 


C 


cu 


H 



I 

n 

0) 
U 

Cn 

■H 



m 



154 

has been reported to produce carbon dioxide, hydrogen and 

2 7 2 8 
various hydrocarbon products. ' The evidence suggests 

that the production of hydrogen arises from the interaction 

of the surface hydroxyl groups of the support with the 

tetrairidium carbonyl cluster. Therefore, the fragmentation 

of the adsorbed Ir . (CO) ^ „ is initiated by the displacement 

of carbonyl ligands with surface hydroxyl groups as shown in 

20 8 
Equation 3-21. The formation of carbon dioxide and 



Ir^(C0)^2 "^ ^SOH •►Ir^(C0)^2_x^^°"^x "^ ^^° (3-21) 



hydrogen results from a water-gas shift type reaction 
involving carbon monoxide and either adsorbed water or the 
surface hydroxyl groups. Recall that the formation of 
carbon dioxide from the reaction involving the supported 
phosphine substituted tetrairidium carbonyl clusters was 
confirmed by GC/MS. The decomposition of the supported 
phosphine substituted iridium carbonyl clusters is most 
likely initiated through a similar interaction with the 
remaining hydroxyl groups of the support. It is suggested 
that the stabilization of the supported clusters may result 
from the silanation of the remaining hydroxyl groups with 
CI SiPPh as shown in Equation 3-18 prior to exposure to CO, 
H2 and HCl(g) . 

If the concentration of the supported iridium cluster 
was low, the predominant fragmentation product was observed 
to be a phosphine bound multinuclear iridium complex 



155 
containing bridging carbonyl ligands. If the concentration 
of the supported iridium cluster was high, the fragmentation 
products were observed to be a mixture of the multinuclear 
species and a complex similar to IrCl(CO)^. It is suggested 
that this multinuclear species may be similar to the 
previously reported complex, [Ir (CO) CI ( (Ph) PCH P- 



9 n fi 
(Ph) -,) ^] [BPh . ] . The conversion of this multinuclear 



2 ^--" 4- 
1 206 „^ 

complex into IrCl(CO), was observed by infrared spectroscopy 
to occur as the reaction temperature was increased. 
However, raising the temperature above 200°C resulted in the 
formation of metallic iridium, (Ir) , as evidenced by 
infrared spectroscopy and the resulting gray color of the 
supports. The formation of methyl chloride at temperatures 

above 200°C can be accounted for by a Fischer-Tropsch 

194 
mechanism as previously described by Vannice. Similarly, 

a Fischer-Tropsch mechanism may be vital for the formation 

of methyl chloride from the carbon substituents formed in 

the decomposition of inherent minor impurities in the 

system. 

It is proposed that the formation of methyl chloride 

also may result through a second pathway involving the 

reduction of carbon monoxide. A mechanism for the formation 

of methyl chloride from synthesis gas and HCl(g) at low 

temperatures (25-200°C) as shown in Figure 3-27 only can be 

speculated upon at this time. It is proposed that the 

multinuclear iridium species can activate either hydrogen or 

HCl(g) to form a hydride species. It has been well 



156 




=^ /<■ 



•u 









10 


a H 


n 


iJ 


h 


(D 




c 


() 


■H 


(D 


<u 


() 


c 


1-1 


^ 


di 


<8 


P 


a 


£ 


o> 


It) 




O 


H 


U T3 


0) u, 


b. 


c 


T. 


^— 




<i 







157 
established that iridium carbonyl complexes can activate 

molecular hydrogen or hydrogen chloride to form hydride 

209-211 
species. For instance, the activation of hydrogen 

211 
chloride by a dinuclear iridiiom complex has been reported 

to result in the formation of a mononuclear complex as shown 

in Equation 3-22. The formation of another dinuclear 

iridium complex can result through the oxidative addition of 



Ir^ (CO) g (PPh^) 2 + 2HCl(g) 

(3-22) 

hydrogen or HCl as shown in Equation 3-23. The same 

reaction employing HBr in place of HCl leads to the 

211 
stabilization of different iridium complexes. Primarily, 

the presence of carbon monoxide induces the disproportion- 

ation of the dinuclear iridium species into mononuclear 

iridium complexes as shown in Equation 3-2 4. The formation 

of different iridium complexes from similar reactions 

employing HBr(g) or HCl (g) suggests the possibility that 

methyl halide formation from HBr(g) or HCl(g) may proceed 



2Ir(C0) 2(PPh3)Cl + 2HCl(g) ^ I Ir (CO) ( PPh^ ) CI H ] + CO 



(3-23) 



158 
through the stabilization of different reaction 
intermediates. The formation of different discrete 
molecules could explain the inability of the iridium system 

(Ir(CO) (PPhjjHBrjl J + 2HBr(g| + CO ^ I r (CO) (PPh ^ ) ^lIBr^ + H|Ir(CO) Br 1 + H 

(3-24) 



to reversibly utilize HBr(g) and HCl{g) to form the 

148 
corresponding methyl halide as previously noted. 

The formation of an iridium hydride complex is most 

likely followed by hydride migration onto a carbonyl ligand. 

The formation of this formyl intermediate (I) is mediated 

through a bifunctional activation of the carbonyl ligand 

with an acidic site in the support. Hydride transfer to a 

carbonyl ligand previously has been reported to be 

212 
facilitated in the presence of a Lewis acid. It 

148 
previously was proposed that the HCl(g) reacted with the 

support hydroxyl groups to form support-Cl moieties as shown 

in Equation 3-25. However, there have been several reports 

3-OH + HCl(g) >-3-Cl + H^O (3-25) 

that indicate that this reaction does not occur in the 
temperature range of 25-500°C. ' The hydride transfer 
to a carbonyl ligand most likely occurs in a similar fashion 
to the alkyl migration observed in the interaction of Al^O^ 



159 

154 
with Mn(CH^) (CO) c as shown in Equation 3-26. In the 

presence of carbon monoxide the cyclic formyl 



CH, 

/ ' 
Mn(CH2)(CO)5 + 0-Al-O-Al >(OC)^Mn— Q (3-26) 





/ 

-0 _A1— 0— Al 



intermediate (I) can be rearranged to an acyclic formyl 
intermediate (II) by the addition of the carbon monoxide to 

the iridium center similar to that shown in Equation 3-2 7 

154 
for the Mn(CH-) (CO) ^/alumina system. 

The subsequent addition of hydrogen to the acyclic 

formyl intermediate (II) could result in the formation of a 

bound formaldehyde type intermediate (III) . Coleman et al. 

postulated the existence of a formaldehyde intermediate in 

the formation of the hydrocarbon products in the 

141 
Ir . (CO) , -/AlCl^-NaCl system. The formation of a bound 

formaldehyde intermediate in the homogeneous reduction of 

carbon monoxide to oxygenated products also has been 

proposed. A bound formaldehyde complex can be 



CH- 

/ ' 
(OC) ,Mn— C + CO ►(OC) ^Mn-C-CH^ (3-27) 

•\ :l 

•0 

/ I 

-0 — Al— -0-Al— 



160 
transformed into either a metal alkoxide complex, M-OCH_ , or 
a metal hydroxymethyl complex, M-CH„OH. Since the bound 
formaldehyde complex (III) is stabilized by the interaction 
of a Lewis acid with the oxygen atom of the aldehyde, the 
most likely intermediate is the formation of the hydroxy- 
methyl complex (IV) . The formation of a hydroxymethyl 
intermediate should produce at least trace 2-carbon products 
due to migratory insertion of a carbonyl ligand into the 
Ir-C bond. This could account for some of the ethylene, 
acetaldehyde and ethyl chloride that was observed. 

There are several ways in which methyl chloride can be 
formed from this hydroxymethyl intermediate (IV) . First, 
the activation of hydrogen by the hydroxymethyl intermediate 
(IV) could rejuvenate the initial iridium hydride complex 
while forming methanol as a by-product. The methanol could 
react with HCl(g) over an inorganic oxide to form methyl 
chloride and water. It has been observed that in low 
concentrations methanol can be converted in 100% yields to 
methyl chloride under the reaction conditions employed. 
Secondly, the activation of HCl could lead to the formation 
of a chloromethyl iridium complex (V) . Further activation 
of hydrogen could produce the observed methyl chloride and 
rejuvenate the initial iridium hydride complex. The 
formation of a chloromethyl osmium complex from the 

activation of HCl by a bound formaldehyde complex as shown 

215 
in Equation 3-2 8 has been reported. For continuity, the 



161 



PPh 
I j|,CH.OH 

:03- ■^ Hci 



PPh, 



CI '^ oc"! ^ci 



PPh 

OC^I jL-CH,Cl 



PPh, 



(3-28) 

overall mechanism including both subdivisions is shown in 
Figure 3-28. Similar types of mechanisms including formyl, 

formaldehyde and hydroxymethyl intermediates has been 

^ ^ ^^ ^ ^- ^ 1 u 1 216,217 

proposed for the formation of alcohols. 

Investigation of a Phosphine Substituted Triosmium Carbonyl 
System 

Since methyl bromide has been reported to be produced 

in a homogeneous Os, (CO) , ^/^Br, system, the formation of 

methyl chloride was attempted using a supported phosphine 

substituted osmium carbonyl cluster. The supported cluster 

was identified by infrared spectroscopy to be the supported 

mono-phosphine substituted triosmium carbonyl cluster as 

shown in Table 3-10. Upon exposure to the reactant gas 

mixture, yery little activity towards the formation of 

methyl chloride was observed at 75°C. A comparison of the 

infrared spectrum of the supported osmium cluster that had 

been exposed to the reactant gases at 75 °C with the infrared 

spectrum of Os^iCO),- ^s shown in Figure 3-29 indicates that 

Os^(CO),-, is being formed in the reaction. The formation of 

Os,(CO),_ most likely occurs through the substitution of the 

phosphine ligand by the carbon monoxide in the reactant gas. 



162 




0) 




-o 




•H 




U 









iH 




X 




u 




l-l 




>1 




J3 




■P 




(U 




g 




IW 









C 









-rH 




+J 




IC 









fe 






Es 


aj 


QJ 


X 


+J 


-p 


m 




>i 


M 


en 







M-l 


M 




(U 


e 


■p 


(n 


m 


•H 


3 


c 


iH 


(C 


CJ 


s: 




o 


fc; 


(U 


P 


g 


■H 




fO 


Ti 


•H 


0) 


M 


tn 


M 







aTS 





0) 


u 


+J 


eu 


u 







m 


il, 


n 


a, 




3 


:5 


01 


0) 




•H 


QJ 


> x: 


M 


+J 


(1) 




> 


C 


o 


-H 



I 

(U 
■H 



163 



c 
o 

X! 

u 

u 
e 





o 
i< 

r-l 

a 
e 
o 
u 

M 

o 

M-l 

n) 
+> 

R] 

Q 

m 
c 



000<T><T\ 00<T\ 



or^iAfMCD onoco 
i-ioooa\ ^oo<Ti 



ooo<J^ OOON ooc7^o^ 

(NfNOJf— t fNlrsIr-l fNJOJ.— IrH 






m 



164 



3-PPh20s2(CO) j^^ exposed to CO/H2/HCI at 75 "c 




2200 



2100 



2000 
-1, 



1900 



Wavenumbers (cm ) 



Figure 3-29. Comparison of the Infrared Spectriim of 

3-PPh20s3(CO)-|^^ Exposed to CO, E^ and HCl(g] 



With that of Os,(CO),_ 



165 
It has been reported that Os2(CO)^2 ^"*^ H20s2(CO)^q 
physically adsorbed onto an inorganic oxide support 

decompose at 200 °C through the formation of chemisorbed 

220 
Os(CO)- and Os (CO) , fragments. At lower temperatures, it 

is possible that an interaction between Os- (CO) ^„ and the 

hydroxyl groups on the support's surface could lead to the 

formation of a neutral hydrido cluster, such as HOs (CO) -^ — 

220 
(0-E ) . Similar trinuclear osmium hydride complexes have 

been formed through the coupling of mononuclear species as 

221 
shown in Equation 3-29. Furthermore, it has been 



Os(CO)^ Os(CO)^ 
H20s(C0) ^ *► H20S2(C0) g *► H2OS3 (CO) ^2 (3-29) 



reported that osmium hydride complexes, such as H Os (CO) 

and H_Os-(CO),, react under carbon monoxide to yield 

221 
Os^(CO),2- Therefore, it is not surprising that the 

formation of Os_ (CO) ^ „ '^^^ observed by infrared spectroscopy 

in the reaction of H^, CO and HCl (g) with the supported 

phosphine substituted triosmium carbonyl cluster at 75°C. 

Investigation of Supported Cobalt and Iron Carbonyl Systems 

A cost analysis of the various transition metals 

illustrates that the lower group VIII metals, such as 

iridium and osmium are the most expensive metals. 

Therefore, from an economic standpoint it would be 



166 

beneficial to utilize either cobalt or iron as catalysts for 

the commercial production of methyl chloride. For this 

reason, Co„ (CO) „ was supported on a phosphinated support by 
z o 

222 
a procedure similar to that reported by Evans et al. 

Upon exposure to the reactant gases of E^' ^^ ^^'^ HCl (g) at 
25°C an immediate color change of white to blue was observed 
to occur. There was only a trace amount of methyl and ethyl 
chloride observed to be initially formed between 75-100°C. 
It is noted that this blue color is indicative of the 
formation of cobalt chloride. The formation of cobalt 
chloride is further supported in that the infrared spectrum 
of the supported cobalt carbonyl complex after exposure to 
CO, H and HCl(g) did not show the presence of any carbonyl 
absorptions. 

A supported phosphine substituted iron carbonyl complex 
was prepared by the thermal displacement of a carbonyl 

ligand in Fe(CO)i. with a supported phosphine donor as 

223 
reported by Wrighton et al. This phosphine substitution 

reaction has been reported to be facilitated in the presence 

of a CoCl2»2H20 catalyst. The supported complex was 

characterized by infrared spectroscopy to be a mono- 

phosphine substituted iron tetracarbonyl complex as shown in 

Table 3-11. Upon exposure of the supported complex to the 

reactant gases at 75°C the formation of methyl and ethyl 

chloride was observed to occur at rates of 4.56 x 10 

moles CH^Cl sec""*" g""*" and 5.42 x 10~ moles C2H^C1 

sec""*" g"""", respectively. This activity for the formation of 



167 
alkyl chlorides was observed to decrease rapidly with 
increasing reaction time. This deactivation process was 



Table 3-11. Infrared Data for Supported Phosphine Carbonyl 
Complexes of Iron 



Complex 


Infrared Data (cm~ ) 


Environment 


Reference 


A13 -PPh.Fe(CO) ^ 
2 4 


2045(s) , 1967(m) , 
1938(s) 


nujol 


b (Fig. 3-29) 


?e{CO) ^ 


2025, 2000 


isooctane 


225 


Fe(CO) .PPh, 
4 3 


2054, 1978, 1942 


isooctane 


225 


(?) -PPh^FelCO)^ 


2045, 1968, 1932, 
1876 


KBr 


225 



b = This work; Al = Alumina; P = Polystyrene; s = strong, 
m = medium. 



observed to be accompanied by a color change from tan to 
white in the support. The infrared spectrum of this white 
support did not exhibit any carbonyl absorptions as shown in 
Figure 3-30. The infrared absorption observed at 1595 cm 
corresponds to the C=C stretches in the phenyl groups of the 
phosphine linkage on the support. The absence of any 
carbonyl absorptions in the infrared spectrum suggests that 
under a carbon monoxide atmosphere the phosphine ligand in 
3-PPh_Fe(C0) . is exchanged for a carbonyl group. The 
volatility of iron pentacarbonyl results in the metal 
complex being swept out of the system with the reactant and 
product gas stream. This is further supported in the 
discoloration that occurs in the gas lines and the mineral 
oil in the gas exhaust bubbler. Attempts to eliminate the 



168 



Sample exposed' to 
CO/H2/HCI a;t TS^C 




Original 
Sample 



2300 



2000 1800 
Wavenumbers (cm ) 



1600 



Figure 3-30. Infrared Spectrum of Al3-PPh2Fe (CO) 4 Before 
and After Exposure to CO, Ht and HCl (g) at 
75 *C 



169 

volatility problem of Fe(CO) by initially supporting 

Fe^ (CO) - „ have been unsuccessful. It has been reported that 

nucleophilic attack on the iron cluster causes the 

degradation of the cluster into mononuclear carbonyl 

2 2 fi 
complexes. The reaction of a phosphine donor with 

Fe (CO) would proceed through metal-metal bond cleavage 

and result in the formation of Al 3 -PPh2Fe (CO) ^ units. 

Therefore, no advantage exists for supporting an iron 

cluster on a phosphinated support. 

Investigation of a Phosphine Substituted Triruthenium 
Carbonyl System 

A variety of other metal carbonyl complexes, such as 

Re2(CO)^2' ^"2^^°^ 10' ^'^e^'^^^ie ^"^ ^^3^^°' 12 ^^^^ supported 
on a phosphinated support through the substitution of a 
carbonyl ligand for a supported phosphine donor in refluxing 
benzene or toluene. In these trials the activity for methyl 
chloride production ranged from a trace amount for the 
rhodium and ruthenium complexes to an activity comparable to 
that observed for the supported iron complex for the 
manganese and rhenium complexes. It also was observed that 
a large amount of ethyl chloride and other two carbon 
products were formed in the reaction involving the supported 
ruthenium carbonyl complex. Gas chromatography and GC/MS 
identified the formation of methyl chloride, acetaldehyde, 
ethyl chloride, ethyl formate, diethyl ether, 
1 ,1-dichloroethane, ethyl acetate and residual benzene 
solvent as shown in Figure 3-31. The mass spectra obtained 



170 



a = HCl 
b = CH^Cl 
c = CH^CHO 
d = C2H CI 
e = HCO2C2H 
f = (C2H^)0 
g = Cl2CHCH^ 
h = CH3CO2C2H5 




5.36 11.19 17.03 22.46 
Time (minutes) 



28.29 



34.12 



Figure 3-31. Gas Chromatogram of Two-Carbon Products 
Obtained in the Supported Ruthenium 
Cluster System 



171 
for these two carbon products are shown in Figures (3-32) - 
(3-37) . The activity for the formation of ethyl chloride at 

50°C using the supported ruthenium complex (1.01% Ru) was 

-9 -1 -1 

observed to be 1.16 x 10 moles C„H-C1 sec g . If the 

HCl(g) content in the reactor feed gas was kept at a 

minimum, the formation of ethanol and methanol could be 

observed by gas chromatography. The supported ruthenium 

cluster is unique in that it is the only system that has 

shown a preference for the formation of two carbon products. 

This system also showed a much greater activity towards the 

formation of ethyl chloride than the supported iron, osmium, 

manganese, rhodium, rhenium or iridium systems. This 

contention is supported by a recently reported Ir-Ru/Si02 

catalyst system which showed an increase in CO conversion to 

C„-Cj. hydrocarbons with an increase in the ruthenium content 

227 
of the catalyst. Unfortunately, carbon-13 labelling 

13 ... 

studies using CO have been inconclusive in ascertaining 

the extent of carbon-13 incorporation into ethyl chloride. 

However, these studies did show that 3 9% of the carbon 

12 
dioxide observed was C0„. This suggests that the 

supported ruthenium clusters are decomposing under the 

reaction conditions as did the supported iridium clusters. 

The orange supported ruthenium complex was 

characterized by infrared spectroscopy to be a mixture of 

the mono-phosphine substituted triruthenium carbonyl cluster 

and Ru (y-H) (CO) ^^ (y-O-E) • It was found that by altering 

the concentration of Ru, (CO) ^2 ^""^ ^he phosphine on the 



172 



100 
80-1 

60 



40 

o 

* 20 1 



29 



Observed 
spectrum 



27 



44 



42 



■ I ■ M ■ ■ I ■ 



2*0 25 3b 35 40 45 50 

M/E 



Reference spectrum 



83 



6u lOu li»u 
M/E 



Figure 3-32. Mass Intensity Report for Acetaldehyde 



173 



1001 28 



Observed spectrum 
64 



M 


80- 


(0 




CU 




a, 




di 


60- 


m 




(t3 




m 






40- 



20- 



O-Itttttt 



49 



20 40 



■ J......II l[liiiitiii|ll|IHui|ii i ii iiini i n i i|i i jiiniiii i j ii i,|iin irir i ii |[ || 



60 80 

M/E 



100 120 



1 q 



iJl 



Reference spectrum 



83 



2 u E i'o 1 U 

M/E 



Figure 3-33. Mass Intensity Report for Ethyl Chloride 



174 



Observed spectrum 



28 



lOOn 



M 801 

(0 

<u 
(h 

(1) 
m 

(0 

" 40- 

O 

<*= 20- 



60- 



38 



45 



^ I I I'lS r*! I I I I I I I I I I I I' l I I I I I 



l> I I I I I I I 



20 40 



60 80 
M/E 



100 120 



1 oq 



Reference spetrum 



vU 



60 ' I'o ' TV 

M/E 



Figure 3-34. Mass Intensity Report for Ethyl Formate 



175 



loon 



(0 


80" 


(U 




a, 




(U 


60- 


en 




(0 


" 


n 


40- 


iw 







- 



20- 



31 



Observed spectrum 



59 



45 



74 



O-' j I 1 I ' l |* I I | I I 'll^ I I'l M J l I I 'l 

20 4'0 6*0 

M/E 



I"' """I 
50 100 



10 0. 



Reference spectruin^"^ 



JM 



20 60 100 ,„o 

M/E 



Figure 3-35. Mass Intensity Report for Diethyl Ether 



176 



Observed spectrum 



100 

80 

604 

40 
20 
0, 



63 



27 



■ ■ ' ■ U, 



20 40 



I I I I I I m I 'r i' i 111 I i I , ,1, ,, , l ^' (| , , , , I ( , I ■ 



80 
M/E 



100 120 140 



-Jl 



Reference spectrum 



6 10 14 

M/E 



Figure 3-36. Mass Intensity Report for 1 ,1-Dichloroethane 



177 



Observed spectrum 



100 

80 
60 

40 
20 ^ 



43 



29 



-lll l 'inl i 



ttOU 



6|1 



20 



6 



70 



iilii[<<iiiiiii|iinii 
80 

M/E 



'll|MIiJllll|MlllJ|mij|l(|i|[ 

l6o ' 12( 



1 CO, 



Reference spectrum 



83 



-I JJ ' ■ 



20 eu 100 i<<u 



M/E 



Figure 3-37, Mass Intensity Report for Ethyl Acetate 



178 
support that the tri-phosphine substituted triruthenium 
carbonyl cluster could be prepared. It has been reported 
that the formation of entirely the mono-phosphine 

substituted triruthenium carbonyl cluster can be facilitated 

22 8 
in the presence of a NaPh^CO catalyst. A summary of all 

infrared data is provided in Table 3-12. It was observed 

that a color change from orange to yellow occurred upon 

exposure of the supported ruthenium carbonyl cluster to air. 

A similar observation has been reported for the air 

oxidation of Ru (CO) on alumina to mononuclear Ru(II) 

229 
species. This color change was observed to be 

accompanied by a change in the infrared spectrum as shown in 

Figure 3-3 8 which can be explained by the formation of a 

mixture of mononuclear ruthenium species such as 

3-PPh Ru(CO)^ and Ru(CO)^. 

A color change from orange to white was observed upon 

exposure of the supported ruthenium cluster to the reaction 

gas mixture of H^ , CO and HCl(g). An explanation for this 

color change can be derived from the fragmentation of the 

original cluster into mono and dinuclear species. It has 

been reported that an increase in temperature from 60 °C to 

150 °C under a carbon monoxide atmosphere can initiate the 

fragmentation of the triphosphine substituted triruthenium 

carbonyl cluster into phosphine substituted mononuclear 

230 
carbonyl species. It also has been reported that 

Ru^(CO)^^ is oxidized in the presence of halogens to form 

,^^v ^ 231 This 
the corresponding mononuclear species, RuiCOj^X^. 



179 



Table 3-12. A Summary of Infrared Data Obtained 
for the Supported Ruthenium Cluster 
System 



Complex 

or Reaction 

Si 3-{pphj) ^RUj(COIg 
IRUj(COI^(PPh3) j) 
IRUj (CO) j2 + Si 3 -PPhjl II) 
Ru^ (CO)^PPhj 

Ru (u-H) (CO) (i,- ESi) 



Infrared Data (cm" ) Environment Reference 



2066(vw), 2024(w), nu]ol 
1967(VB,br) 



2044(vw), 1978(6h), benzene 
1967(s, br) 



232 



2095 (vw) , 2060 (s) 



2097 (m) , 2046(6) , 
2030(sh), 2023(sh) 

2014 (s) , 1996(sh) , 

1986(m), 1972(sh) 
1960 (sh) 

2107(w) , 2076(s) , 

2066(B) , 2026(s) , 
1991 (m) 



nu jol 
cyclohexane 232 



nu jol 



b (Figs. 3-38, 
3-39) 



(I) + air 


2060(vs), 2030(w) 
1998(VB) 


nujol 


b (Fig. 


3-38) 


(I) + Hj, CO, HCl 


2134(s) , 2063(vs) 
2030(s), 1980(111) 


nuiol 


b (Figs 


3-39, 
3-40) 


Si 3-PPh2Ru(CO) ^ 


2059, 1995, 1952 


toluene 


234 




Si 3-PPhjRu (CO) ^PPh^ 


1900 


toluene 


234 




Ru(CO) ^ 


2035, 1999 


deca lin 


235 




""3'^°' 10 '"^3* 2 


2078, 2024, 1999, 


decalin 


235 




SiS-PPhjRuICO)^ 


2045(m) , 1990(s) , 
1940(sh) 


wafer 


230 




Ru(CO) ^Clj 


2185(w) , 2135(5) , 
2115(m), 2077(s) 


Et.O 


236,231 




Ru(CO) ,Cl2 


2063(s), 1989(s) 


CHClj 


236 




PPh^PuICO) jCl^ 


2133(m) , 2075(s) , 
2033(m) 


^6"l2 


236 




(PPh2Ru(C0) -.Cljlj 


2076(s) , 2016(s) 


CHCl 


236 




RuCljICO^IPPh^)^ 


2062(vs) ,2001 (vs) 


nu:ol 


b (VI) 




iKuCljICO)^)^ 

1 |RuCl2(C0)^Ij + 

Al3-PPh^| (ID 


2145(s) , 2089(m) , 
2067 (vs) ,2025(m) 
2132(s) , 2054(vs) , 
2024(w), ]991(vs) 


nuiol 
nujol 


b (VII) , 
b (Figs. 


231 

3-40, 
3-41) 


(II) t air 


2087(w) , 2059(vs) , 
2024(w), 1993(VE) 


nu:ol 


b (VIII) 




(II) •► llj, CO, HCl 


2133 (s) ,2057(V8) , 
2018(w) , 1'>94(V8) 


nujol 


b (Fig. 


3-41) 



b - ThiB work: (V-VIII) - location of spectrum in appendix B; 
vs - very strong; s - strong) ms - medium strong, m - medium; 
w - wealj; vw - very weak; br - broad; sh - shoulder. 



180 



Sample exposed to air 




2301 



2128 



2006 



1883 



1700 



Wavenumbers (cm" ) 



Figure 3-3 8. Infrared Spectrxmi of a Mixture of the 

Supported Ruthenium Clusters Exposed to 
Air 



181 
mononuclear species, Ru(CO) .X , can react further to produce 
a variety of white mono and dinuclear ruthenium carbonyl 

n -3 /r 

complexes. The infrared spectrum that was obtained for 
the white supported complex as shown in Figure 3-39 can be 
assigned to the mononuclear complex, 3-PPh Ru (CO) -Cl_ . The 
carbonyl absorption observed in the infrared spectrum at 
1980 cm" suggests the presence of a small amount of another 
ruthenium complex, such as RuiCO-Cl^. It is also possible 
that a supported dinuclear carbonyl species is present. 
This contention is supported in that the reaction between 
[RuCl^ (CO) ^] „ and the phosphinated support yields a complex 
with an infrared spectrum which is similar to that observed 
for the supported triruthenium cluster after being exposed 
to the reactant gas mixture as shown in Figure 3-40. The 
infrared spectrum of this supported dinuclear complex as 
shown in Figure 3-41 was observed to remain unchanged upon 
exposure to the reactant gas mixture. This suggests that a 
chloro-ruthenium compound is the resulting carbonyl complex 
present under CO, H- and HCl(g) at low temperatures. 

Interpretation of the previous infrared spectra is 
complicated by the possibility that the stabilization of 
various mononuclear fragments which result from the 
decomposition of the triruthenium cluster may occur under 
the employed reaction conditions. It has been reported that 
Ru (CO) decomposes on a silica gel support to form off 
white colored mononuclear species as shown in Figure 
3-42, ^"^"^ The similarities between the reported carbonyl 



182 



Sample exposed to 
CO/H2/HCI at 75®C 



c 
o 

•H 
(0 

n 

•H 

g 
01 
c 

10 




Origi 
Sarapl 



2301 



2128 



2006 



1883 



1700 



Wavenun±)ers (cm ) 



Figure 3-39. Infrared Spectrirm of the Supported Ruthenium 
Clusters Exposed to CO, H2 and HCl(g) at 75 "C 



183 



A = Ru3(CO)j^2 + 3-PPh2 exposed to 
CO/H2/HCI at 75 *C 

B=(RuCl2(CO)3)2 + 3-PPh2 



•H 

E 

(0 

c 
la 
u 

Eh 




2301 



2128 



2006 



Wavenumbers (cm ) 



1883 
-1, 



1761 



Figure 3-40. A Comparison of the Infrared Spectrum of the 
Supported Ruthenium Clusters With that of 
Supported IRUCI2 (CO) -^] 2 



184 



Sample exposed to 
CO/H2/HCI at 75 "C 




2301 



2128 



2006 



1383 

-1, 



1761 



Wavenumbers (cm ) 



Figure 3-41. Infrared Spectrum of Supported [RuCl- (CO) -.] _ 
Before and After Exposure to CO, H~ 
and HCl(g) at 75 "c 



185 
absorptions for these ruthenium carbonyl fragments and those 
observed for the triruthenium cluster exposed to H , CO and 
HCl(g) suggests the possibility of the presence of a mixture 
of ruthenium fragments and ruthenium chloro-carbonyl 
complexes . 

A mechanism for the formation of 2-carbon products can 
only be speculated upon at this time. It is possible that 
some of these products arise from the presence of surface 
ethoxide groups that were formed during the phosphination of 
the support. These surface ethoxide groups were exchanged 
for 3-0C^H_ groups by stirring the supported ruthenium 
complex in propanol in an attempt to determine the extent in 
which the B-OC^H^ groups affect the activity for ethyl 
chloride formation. It was demonstrated by infrared 
spectroscopy that the supported ruthenium cluster had not 
been affected by this procedure. The activity for ethyl 
chloride formation at 75°C under CO, H^ and HCl{g) for this 

supported complex was found to be similar to the previously 

-9 -1 

reported activity of 1.16 x 10 moles ethyl chloride sec 

g~ . This suggests that the ethoxide groups play a minor 

role in the formation of the ethyl chloride. It should be 

noted that a considerable amount of propyl chloride was 

observed to form in this reaction. It is suggested that 

this propyl chloride arises from attack of HCl(g) upon the 

absorbed propanol present in the support. 

A more likely possibility for the formation of ethyl 

chloride arises from the addition of HCl(g) across the 



186 
double bond in ethylene over the supported cluster. Recall 
that ethylene and ethyl chloride are trace impurities found 



RU2(C0)^2 "•" SiO^ 



Ru^ly-H) (CO)^Q(p-OSiE) 
2107(w) , 2076(s) , 2066(s) , 
130°C X 2026 (s) and 1991 (m) cm" 
CO 



.Air 



Ru"'--'- (CO) 2 + Ru-"- ■'■■'■ (CO) + Ru-'-^(CO) 



2067, 



2075 cm 



-1 



2144 cm 



-1 



2 00 5 cm 



-1 



[RU^^(C0)2]2 
2055, 1990 cm 



-1 



s = strong; m = medium; w = weak, 



Figure 3-42. Fragmentation of RU2(CO)^2 °'^ ^ Silica Gel 
Support 



in the technical grade HCl(g). It is possible that the 
supported ruthenium complex is more efficient at converting 
ethylene to ethyl chloride than the other supported metal 
systems that were investigated. If electronic grade HCl(g), 
which did not contain any ethylene impurity, was used in 
place of the technical grade HCl(g), the formation of 
ethylene was observed to occur as identified by gas 



187 
chromatography. The formation of ethylene is most likely 
occurring by the conversion of bound carbonyl groups during 
the decomposition of the supported ruthenium cluster. It 
has been reported that Ru- (CO) reacts with AlH in THF to 

produce predominantly methane and ethylene in a ratio of 

237 • ■ 

1:1.7. This reaction was shown to be stoichiometric in 

that only approximately 10% of the bound carbonyl groups in 

237 
Ru-.(C0),2 were found to be converted. It was suggested 

that the formation of ethylene proceeded through a carbenoid 

metal intermediate as shown in Equation 3-30. 

One last possibility for the formation of these 

2-carbon products is the reduction of carbon monoxide 

through a mechanism similar to that discussed in Figure 3-27 

for the formation of methyl chloride in the iridium system. 

In this mechanism the formation of 2-carbon products could 

result from the insertion of carbon monoxide into a 

hydroxymethyl intermediate. Attempts to incorporate 



CH^ (3-30) 

/ '•••. 

(CO) M=CH- *-(C0) M ^M(CO)^ ►2M(C0) + C H. 

'CH^ 

13 
carbon-13 into ethyl chloride from CO have been incon- 
clusive. It is proposed that the formation of the two- 
carbon products in the supported ruthenium cluster system 
results from the conversion of carbonyl ligands in the 



188 
cluster during decomposition or fragmentation on the support 
and from the reduction of carbon monoxide. 

Summary 

148 
The previously reported results for the formation of 

methyl chloride from CO, H- and HCl(g) in the supported 

iridium cluster system were shown to be clouded by the 

presence of adsorbed 2-methoxyethanol on the support. This 

adsorbed 2-methoxyethanol was shown to be cracked by HCl(g) 

to methyl chloride at 75°C. Attempts to remove this 

absorbed 2-methoxyethanol from the system by varying the 

washing and drying procedures during the preparation of the 

catalyst were unsuccessful. Since this adsorbed 

2-methoxyethanol was difficult to remove, attention was 

turned to the synthesis of the supported iridium clusters in 

other solvents, such as 2-ethoxyethanol or toluene. 

The clusters synthesized in either 2-ethoxyethanol or 

-12 
toluene were observed to exhibit an activity of 1-9 x 10 

moles CH^Cl sec~ g~ at 75°C. The decomposition of 

inherent impurities in the system were observed to lead to 

the formation of trace quantities of methyl chloride in the 

initial stages of the reaction. It was proposed that the 

major contributor to the initial formation of methyl 

-12 -1-1 

chloride (1-9 x 10 moles CH^Cl sec g ) in the 

supported iridium cluster system is the conversion of 

iridium bound carbonyl groups during the support induced 

fragmentation of the cluster. It was suggested that as the 

reaction time progressed the percentage of methyl chloride 



189 
produced from the reduction of synthesis gas would increase. 

Infrared spectroscopy was used to characterize the 
supported iridium clusters and the fragmentation products 
that resulted upon exposure to CO, H and HCl(g) at various 
temperatures. The synthesis of predominantly the mono- 
phosphine cluster was observed only to result in a 
2-methoxyethanol solvent. It was suggested that this 
cluster could also be formed if Ir.(C0),2 "^s reacted with a 
phosphinated support which had been previously silanated 
with Cl_SiPh_. In this case, multiple phosphine 
substitution in Ir.iCO),- would be sterically hindered by 
the presence of BiSiPh- groups. All of the supported 
iridium carbonyl clusters were observed to exhibit a 
predominant infrared absorption centered at either 2069 or 
2040 cm"""" after exposure to CO, E^ and HCl(g) at 75°C. A 
high iridium concentration of the support was found to lead 
to the observation of the 2069 cm infrared absorption. 
The 2040 cm" absorption was found to correlate with the 
presence of a low iridium concentration. The infrared 
absorptions observed for the supported clusters (Low % Ir^ 
at 2048, 1734 and 1719 cm" were assigned to the formation 
of a phosphine bound multinuclear species containing 
bridging carbonyl and chloride ligands. This multinuclear 
complex was suggested to be similar to 

[Ir (CO) ^Cl(Ph2PCH2PPh2) 2nBPh^] . The infrared absorptions 
observed for the supported clusters (High % Ir) were 
assigned to a mixture of the multinuclear complex and 



190 
IrCl(CO)-.. The conversion of the multinuclear complex to 
IrCl(CO)^ was observed as the reaction temperature was 
increased. Similar results were observed for reactions 
employing either IrCl(CO)^ or Vaska's complex on a 
phosphinated support. It is suggested that the multinuclear 
species is active towards the formation of methyl chloride 
from synthesis gas and HCl (g) . It also was shown that 
IrCl(CO)^ is inactive as a CO reduction catalyst under the 
employed reaction conditions. 

The gradual formation of metallic iridium was observed 
by infrared spectroscopy at temperatures exceeding 200 °C. 
The conversion of synthesis gas and HCl(g) to alkyl chloride 
in the presence of iridium metal catalysts at temperatures 

of greater than 200°C has previously been described by 

194 
Vannice. It was shown that iridium metal is inactive 

towards the formation of methyl chloride below 200°C. Thus 

it was suggested that metallic iridium does not play a major 

role below 200°C in the supported iridium carbonyl cluster 

system. 

Finally, a mechanism was proposed for the formation of 

methyl chloride in the supported iridium carbonyl cluster 

system. The initial activity was dominated by methyl 

chloride formed through the degradation of the supported 

iridium cluster and by the decomposition of inherent system 

impurities. As the reaction time progressed the system's 

activity was suggested to be dependent upon the conversion 

of CO, H and HCl(g). The mechanism for the reduction of 



191 
carbon monoxide was based upon the formation of formyl, 
formaldehyde and hydroxymethyl intermediates. 

This investigation concluded with the examination of 
various other phosphine supported metal carbonyl systems. 
It was observed by infrared spectroscopy that the supported 
osmium, cobalt and iron carbonyl systems formed Os-(CO),„, 
C0CI2 and Fe(CO)^, respectively, upon exposure to CO, Ey and 

HCl(g) at 75°C. A supported ruthenium carbonyl cluster was 

-9 
found to produce ethyl chloride (1.16 x 10 moles C2HCCI 

sec" g~ ) along with a variety of other 2-carbon products. 

The supported triruthenium cluster was observed by infrared 

spectroscopy to form ruthenium chlorocarbonyl complexes upon 

exposure to H , CO and HCl(g) at 75°C. It was proposed that 

the two-carbon products were formed from the bound carbonyl 

ligands during cluster fragmentation and from the reduction 

of carbon monoxide. A mechanism similar to that for the 

supported iridium carbonyl system was suggested. In this 

case, insertion of CO into the Ru-C bond in the 

hydroxymethyl intermediate could lead to the two-carbon 

products. In general, it is suggested that a change in 

either the metal nucleophilicity or the supports acidity may 

alter the activity and selectivity of the system. 



CHAPTER IV 

CONCLUSION 
The primary objective of the preceding two studies was 
to investigate the feasibility of binding and activating 
both carbon dioxide and carbon monoxide. The first study 
demonstrated that the binding of carbon dioxide was 
influenced by the nucleophilicity of the metal center in the 
transition metal carbonyl anion. An extension of this work 
into the low pressure reduction of carbon dioxide by 
Re (CO)^- in methanol to methyl formate concluded that the 
carbon dioxide most likely is initially converted to carbon 
monoxide through a reverse water-gas shift reaction. It is 
this carbon monoxide that is reduced under the employed 
reaction conditions to methyl formate. 

The feasibility of reducing carbon monoxide was 
investigated further in an acidic system employing a 
supported iridium catalyst. It was concluded that the 
conversion of synthesis gas and HCl (g) to alkyl halides 
proceeds under low temperatures and pressures through the 
stabilization of discrete iridium carbonyl complexes. The 
formation of ethyl chloride and other two-carbon products in 
a supported ruthenium system demonstrated that a change in 
catalyst composition could affect the selectivity and 
activity of the reaction. 



192 



193 

Although the preceding two studies are not processes 
which are ready for commercialization, they do further the 
understanding of the interaction between carbon dioxide and 
carbon monoxide with transition metal carbonyl complexes. 
The feasibility of reducing either carbon monoxide or carbon 
dioxide in the presence of transition metal carbonyl 
complexes has been demonstrated. This work suggests that it 
may be desirable to design catalysts for the sole purpose of 
reducing carbon dioxide to carbon monoxide. The further 
reduction of this carbon monoxide could be accomplished 
through a bifunctional interaction with an inorganic oxide 
support and a transition metal complex. In general, it is 
only through more work in this area that the overall goal of 
developing carbon dioxide and carbon monoxide as alternate 
carbon resources for chemical feedstocks and fuels may be 
realized. 



APPENDIX A 

The following infrared spectra are meant as 
supplementary material for Chapter II. A list of the 
spectra included in this section is provided below. 



Page 



I. KCo(CO) 195 

II. K[HFe(CO)^] 196 

III. Mn2(C0)^Q 197 

IV. NaRe(CO)^, Re2(C0)^Q 198 

V. KCo(CO)^ + CO2 199 

VI. K[HFe(CO)]^ + CO2 200 

VII. NaReiCO)^, Re2(C0)-^Q 201 

VIII. Re2(C0)^Q + CO2 + H2 + CH3OH 202 

IX. White Precipitate from Reaction (VIII) .... 203 

X. Pink Film from Reaction (VIII) 204 

XI. Re2(C0)^Q + CH^OH 205 

XII. Re2(C0)^Q + CO 206 

XIII. Re2(C0)^Q + CO + KOCH^ .' . 207 

XIV. Re2(C0)^Q + CO + H2 + KOCH^ 208 

XV. KOCH^ + CO 209 

XVI. H Re (CO) + CO 210 

X y z 



194 



^v-/\^l 



c 
o 

•H 

n 

09 

E 
m 

c 

(0 




2500 



2000 



1800 
-1, 



1600 



Wavenumbers (cm ) 



196 



c 
o 

-H 

m 

CO 
■H 

E 

CO 

c 
u 




1887 



4000 



3000 



2000 
Wavenumbers (cm" ) 



1800 



197 



e 

in 
C 

(0 

u 




(III) 



2009 



2500 



2000 1800 
Wavenumbers (cm" ) 



1600 



198 




2500 



2000 1800 
Wavenumbers (cm~l) 



1600 



199 




2SOO 



2000 



i«oo 



1600 



Wavenumbers (cm ) 



200 




o 
o 

CO 



UOTSSTUISUBJJ, % 



8 



201 



c 
o 

-H 
0] 
03 

-H 

e 

01 

c 
<d 
U 




1985 



1880 



2500 



2000 



1800 



-1, 



Wavenumbers (cm ) 



202 




UOTSSTTUSUCaj, % 



203 




UOTSSTUISUPaj, % 



204 




UOTSSTUISUPJJ, 



205 




UOTSSTUISUejJ. % 



206 




UOTSSTUISUBaj, % 



207 




UOTSSTUISUCaX % 



208 




UOTSSTUISUPJCJ. 



209 




■ UOTSSTUISUBJI, % 



210 




UOTSSTUISUHJi % 



APPENDIX B 

The following infrared spectra are meant as 

supplementary material for Chapter III. A list of the 
spectra included in this section is provided below. 

Page 

I. 3-(PPh2) Ir^(CO) (2-ethoxyethanol) .... 212 

II. 3-(PPh2)jjIr4(CO)y (toluene) 213 

III. IrCl(CO) (PPh2)2 214 

IV. 3-PPh20S2(CO)^^ 215 

V. 3-(PPh2)3RU3(CO)g 216 

VI. RuCl2(CO)2(PPh3)2 217 

VII. [RuCl2(CO)3]2 218 

VIII. [RuCl2(CO)3]2 + 3-PPh2, Exposed to Air . . 219 



211 



212 



I I I I I I I 

g g g g E H g 
o o o o o o o 

iH r>- >H O VD O rH 
IT) rO 04 O CT\ VO (N 
O O O O OS CT> 00 
CN CM CN <N r-l iH iH 



II II II II II II It 




UOTSSTUISUeJJ, % 



213 




UOTSSTUISUBJI, % 



214 




UOTSSXmSUBJi % 



215 




Hr^^^.^^-^^-^r-^r^•^f^ 
I I I I I I I I I I 

gggeeeeeee 
ooooooouoo 

or>r-voinrofMa>oovo 

rHOOOOOOCT><T><Ti 
C>JCNl<N<NCN(NC>JHrHfH 



II II II II II 
(Oja 0*0 Q)H-l Di,C-H-n 



UOTSSTUISUejJ. % 



216 




lUOTSSTUISUPaj, % 



217 




UOTSSTUISUBJl % 



218 




UOTSSTUISUBJi % 



219 




1 UOTSSTUISUPJX % 



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BIOGRAPHICAL SKETCH 
The author was born in Harrisburg, Pennsylvania, and 
grew up in Jamestown, New York. He graduated from Jamestown 
High School with honors in 1977. In 1981, he graduated 
magna cum laude from Clemson University. Upon graduation he 
was awarded the Mark Bernard Hardin Award for excellence in 
chemistry. His graduate studies were started at the 
University of Illinois immediately after receiving his B. S. 
in chemistry. After one year at the University of Illinois 
he transferred to the University of Florida with Dr. Drago 
and other members of his research group. Upon completion of 
his graduate studies he will begin his career in the 
employment of Dow Corning Corporation, Midland, Michigan. 
The author is listed as a co-inventor on U.S. Patent 
4,538,011, entitled "Method for the Preparation of Halogen 
Substituted Methanes and Ethanes". 



234 



I certify that I have read this study and that in my 
opinion it conforms to acceptable standards of scholarly 
presentation and is fully adequate, in scope and quality, as 
a dissertation for the degree of Doctor of Philosophy. 



I'<-A.A ^-^i^'-. ^- '-^y^^y ■ 



Russell S. Drago 

Graduate Research Professor 

of Chemistry 



I certify that I have read this study and that in my 
opinion it conforms to acceptable standards of scholarly 
presentation and is fully adequate, in scope and quality, as 
a dissertation for the degree of Doctor of Philosophy. 

David E. Richardson 
Assistant Professor of 
Chemistry 

I certify that I have read this study and that in my 
opinion it conforms to acceptable standards of scholarly 
presentation and is fully adequate, in scope and quality, as 
a dissertation for the degree of Doctor of Philosophy. 




Robe rt(C. 'Stoufer 
Associate Professor of 
Chemistry 

I certify that I have read this study and that in my 
opinion it conforms to acceptable standards of scholarly 
presentation and is fully adequate, in scope and quality, as 
a dissertation for the degree of Doctor of Philosophy. 




^i 'b. "^artlett 
Professor of Chemistry 



I certify that I have read this study and that in my 
opinion it conforms to acceptable standards of scholarly 
presentation and is fully adequate, in scope and quality, as 
a dissertation for the degree of Doctor of Philosophy. 



David N. Silverman 
Professor of Pharmacology 
and Therapeutics and of 
Biochemistry and Molecular 
Biology 



This thesis was submitted to the Graduate Faculty of 
the Department of Chemistry in the College of Liberal Arts 
and Sciences and to the Graduate School and was accepted as 
partial fulfillment of the requirements for the degree of 
Doctor of Philosophy. 



August, 1986 



Dean, Graduate School 



UNIVERSITY OF FLORIDA 



3 1262 08553 5465