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Trinuclear ruthenium carboxylate complexes as oxidation catalysts

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Title:
Trinuclear ruthenium carboxylate complexes as oxidation catalysts
Creator:
Davis, Leslie Shannon, 1963-
Publication Date:
Language:
English
Physical Description:
viii, 139 leaves : ill. ; 28 cm.

Subjects

Subjects / Keywords:
Alcohols ( jstor )
Alkenes ( jstor )
Carboxylates ( jstor )
Catalysis ( jstor )
Catalysts ( jstor )
Ligands ( jstor )
Oxidation ( jstor )
Ruthenium ( jstor )
Solvents ( jstor )
Trimers ( jstor )
Alcohol -- Oxidation ( lcsh )
Catalysts ( lcsh )
Chemistry thesis Ph. D
Dissertations, Academic -- Chemistry -- UF
Oxidation ( lcsh )
Ruthenium carboxylate ( lcsh )
Genre:
bibliography ( marcgt )
non-fiction ( marcgt )

Notes

Thesis:
Thesis (Ph. D.)--University of Florida, 1988.
Bibliography:
Includes bibliographical references.
General Note:
Typescript.
General Note:
Vita.
Statement of Responsibility:
by Leslie Shannon Davis.

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Source Institution:
University of Florida
Holding Location:
University of Florida
Rights Management:
Copyright [name of dissertation author]. Permission granted to the University of Florida to digitize, archive and distribute this item for non-profit research and educational purposes. Any reuse of this item in excess of fair use or other copyright exemptions requires permission of the copyright holder.
Resource Identifier:
001127180 ( ALEPH )
20082537 ( OCLC )
AFM4350 ( NOTIS )
AA00004969_00001 ( sobekcm )

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Full Text


TRINUCLEAR RUTHENIUM CARBOXYLATE
COMPLEXES AS OXIDATION CATALYSTS
By
Leslie Shannon Davis
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
1988
fgf'OF F LIBRARIES


ACKNOWLEDGEMENTS
An undertaking of this size is rarely accomplished by a single
individual working entirely alone, and this is especially true of
this study. My advisor, Dr. Russell Drago, has been an inspiration,
mentor, and guide throughout this journey. I am indebted to him for
his "idears" and all his encouragement and advice during my sojourn
at Florida. Mrs. Ruth Drago, his kind, gracious wife, opened her
home and welcomed me as family, a gesture I certainly appreciated
and that eased my stay during the past four years. I would also
like to acknowledge Dr. Dave Richardson, Dr. Carl Stoufer, and the
remainder of my committee for their help and support.
The Drago Group as a whole has been an outstanding source of
hope, help and fun during our years together. For all the
camaraderie and aid, I thank each of them. I am especially
grateful, first of all, to my labmates, Alan Goldstein, Tom Cundari,
and Rich Riley, who endured all with happy faces, and were
consistent sources of good humor in the lab. I want to thank Ngai
Wong and Larry Chamusco for their computer and mechanical expertise,
without which much of this work would not have been possible. For
their assistance in various and sundry ways, I thank Jerry Grunwald,
Mark Barnes, and Cindy and Ed Getty. I am also grateful to former


group members Dr. Cindy Bailey and Dr. Iwona Bresinska for the
benefit of their wisdom. To Dr. Carl Bilgrien, the initiator of
this study, I owe a deep debt of gratitude. Very special thanks are
due Mrs. Maribel Lisk for her help, advice, and smiles.
Without the help of many others within the department many
"idears" could not be realized. I am grateful to Dr. Roy King for
his deep understanding of NMR and his willingness to share this
knowledge. The machine shop personnel, Chester, Vernon, and Daley,
were able to make anything I could describe, a talent I am most
grateful for. The creative talents of Rudy and Dick in the glass
shop in deciphering my sketches and still creating what I needed are
greatly appreciated. I also thank Chuck Christ and Paul Sharpe for
their assistance.
The experience of graduate school is not realized entirely in
the laboratory. I am grateful to Fran and Allan Goodman for
illustrating this lesson, and for many, many hours of plain old fun.
I also wish to thank Dr. Linda Lentz and Sasi Kalathoor for their
unswerving encouragement and support.
For first instilling in me an interest in chemistry, I thank
Mrs. Jackie Gay. My love of "things that turn pretty colors" is
entirely due to Dr. Alex Zozulin. I owe Alex an additional debt of
first showing me the joys of research.
My greatest debt is owed my family, without whose love and
support and encouragement I would not have accomplished this feat.
To them--Marcia, Larry and Debbie, and Drew--I dedicate this work.


TABLE OF CONTENTS
B.aqg
ACKNOWLEDGEMENTS ii
KEY TO ABBREVIATIONS vi
ABSTRACT vii
CHAPTERS
I. GENERAL INTRODUCTION 1
Catalytic Oxidations 1
Trinuclear Carboxylate Complexes 6
II. ALCOHOL OXIDATIONS BY TRINUCLEAR
RUTHENIUM CARBOXYLATE COMPLEXES 14
Introduction 14
Previous Work 18
Scope of Catalysis 19
Results and Discussion 24
Experimental 51
III. SYNTHESIS AND CHARACTERIZATION OF A
NOVEL TRINUCLEAR CARBOXYLATE COMPOUND 57
Background 57
Characterization 59
Experimental 78
i v


IV.OLEFIN OXIDATIONS BY A NOVEL TRINUCLEAR
RUTHENIUM CARBOXYLATE COMPLEX 80
Introduction 80
Scope of Catalysis 85
Results and Discussion 86
Experimental 96
V. ALKANE OXIDATIONS BY A NOVEL TRINUCLEAR
RUTHENIUM CARBOXYLATE COMPLEX 98
Introduction 98
Scope of Catalysis 105
Results and Discussion 105
Experimental 125
VI. CONCLUSIONS 126
REFERENCES 129
BIOGRAPHICAL SKETCH 139
v


KEY TO ABBREVIATIONS
Et20 = diethyl ether
OAc = CH3C02'
pfb = CF3CF2CF2CO2"
prop = CH3CH2C02'
PPh3 = triphenylphosphine
py = pyridine
tfa = trifluoroacetate
vi


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
TRINUCLEAR RUTHENIUM CARBOXYLATE
COMPLEXES AS OXIDATION CATALYSTS
By
Leslie Shannon Davis
December, 1988
Chairman: Russell S. Drago
Major Department: Chemistry
The family of trinuclear metal carboxylate complexes has been
known to chemists for over 100 years. General studies in terms of
their classical inorganic chemistry, such as ligand exchange or
electron transfer reactions and reactivity, are well documented in
the literature. However, little application of this knowledge has
been attempted.
The series of trinuclear ruthenium carboxylates is very
intriguing'in light of the extensive electrochemistry demonstrated
in these complexes. The ready accessibility to a variety of
oxidation states, combined with the facile exchange of ancillary
ligands, should make these complexes ideal subjects for catalytic
vi i
studies.


Previous work has shown a series of trinuclear ruthenium
carboxylates [Ru30(02CR)gL3]n to be active catalysts for the
oxidation of alcohols to carbonyl-containing products utilizing
dioxygen as the primary oxidant. Continuation of this study has
revealed a mechanism that utilizes the expected synergism between
the metals in this cluster to explain the unique features of this
oxidation.
A previously unknown member of this family, a complex
containing completely fluorinated ligands, has been synthesized and
characterized. Based on the accumulated evidence, this complex has
been formulated as [Ru30(02CCF2CF2CF3)g(Et20)3](O2CCF2CF2CF3). This
complex has been screened as a catalyst for a variety of organic
transformations and has excelled in initiating the free radical
autoxidation of several olefins again using dioxygen as the primary
oxidant. The oxidation studies were extended to alkane oxidations
as well, and were shown to occur by a slightly different mechanism
than that assumed to operate in the industrial, cobalt-catalyzed
oxidation of alkanes like cyclohexane.
VI 1 1


CHAPTER I
GENERAL INTRODUCTION
Catalytic Oxidations
The oxidation of organic substrates as a field of interest to
chemists has its origins in the beginnings of the history of
chemistry as a science. Lavoisier, the father of modern chemistry,
demolished the phlogiston theory when he explained the results of
Priestley and Sheele's air experiments.1 Air, he claimed, consisted
of two parts, one of which will support combustion (Priestley's
"fire gas") and one of which will not, and not "phlogiston." ^ He
named the "fire gas" oxygen (for acid former) and formulated the
theory of combustion in the late 1700s. In this origin the modern
field of oxidation chemistry has its roots.
Detailed studies of oxidation processes began in the 1800s.
The degradation of natural rubber was linked to oxygen absorption,
and a great deal of research was aimed at discovering anti-oxidants
for the rubber industry.3^ The modern theories of autoxidation
processes (as the free radical oxidation of hydrocarbons by O2 is
known), were developed in the early 1900s. The effects of metal
ions on this process were studied during this period by Haber and
Weiss, who formulated the classical mechanism for metal-catalyzed
autoxidation in use today.^>4 (Figure 1.1)
1


2
In* +
R. +
R02* +
R* +
2 R02
In2
>
21 n
RH
>
R.
02
>
ro2*
RH
>
ro2h +
ro2
>
ro2r
> RO4R > nonradical products
Figure 1.1 Basic autoxidation pathways.^>4


3
Autoxidation as a means for producing oxygenated compounds from
hydrocarbons is a highly desired process, although several serious
flaws exist in present processes. Controlling the selectivity of an
autoxidation process, a key element in terms of its usefulness, is
extremely difficult due to the radical nature of the chemistry. A
high activation energy, related to the spin-forbidden reaction
between dioxygen (a triplet state) and organic molecules (a singlet
state) is a barrier as well. Control of the process once initiated
is another disadvantage--the reaction is often hard to stop short of
CO2 and H2O.
Catalytic oxidations theoretically solve most of these problems
in that the addition of a catalyst should lower the energy barrier,
thus making the reaction easier to start. Product selectivity is
drastically affected by the presence of a catalyst as well. For
these reasons, the "Age of Petroleum" and the "Age of Catalysis" are
inescapably linked.^ Without catalysts to facilitate the conversion
of crude oil to useful products, a petroleum-based economy would not
be possible. Vice versa, without the widespread need for and use of
chemicals and products derived from oil, the study of catalysis
would be relegated to purely academic investigations. Sheldon and
Kochi estimate that today over 90% of the organic chemicals in use
are derived from petroleum, and the majority of petroleum and
petrochemical processes involve the use of catalysts.^ In terms of
the importance of catalytic oxidations, industrial organic
chemicals, including oxygenates from oxidative processes, made up
16.8% of the value of the total chemical industry in 1983.^ Seven


4
of the top fifty chemicals (by volume) were produced directly from
oxidation processes, and several others were produced from
oxidatively generated intermediates.
Obviously, catalytic oxidations are industrially valuable.
Serious study and application of homogeneous, liquid-phase oxidation
began in the 1950s. Before this time, the majority of industrial
processes used heterogeneous or supported catalysts. However,
homogeneous catalyst systems offer several advantages, especially to
the academician, over their heterogeneous counterparts. Generally
milder reaction conditions (i.e., lower temperature and pressure)
are used in homogeneous processes. Temperatures, mixing rates, and
catalyst concentration are more effectively controlled, and most
importantly, the reaction can feasibly be studied using standard
spectroscopic methods. New, improved surface science techniques
have made the study of heterogeneous catalysts easier, but the
relative perspicuity inherent in homogeneous systems still outweighs
these advances. The major disadvantages of homogeneous systems
industrially are the difficulty in separating products from the
reaction mixture and catalyst recovery. This last deterrent becomes
a major problem when dealing with catalysts containing noble metals
like rhodium or iridium due to their expense.
The advent of the Mid-Century Process (Equation 1-1) and the
discovery of the Wacker process (Equation 1-2) heralded a widespread
interest in homogeneous catalysis as well as organometallic
chemistry as fields of study. Emphasis was placed on elucidation


5
p-CH3-(C5H4)-CH3 > C00H-(C6H4)-C00H Equation 1-1
cat = Co(0Ac)2 in HOAc, Br' promoter
200 C
15 30 atm air
CH2=CH2 > CH3CHO Equation 1-2
cat = PdCI2/CuCl2
100 C
10 atm air
of reaction mechanisms and the discovery of new compounds that would
catalyze transformations of organic compounds. Understanding the
chemistrv of these processes eventually would lead to improvements
and enhancements of the catalysis. This understanding led to new
growth in both fields, and formed the basis for new expansions of
the chemistry and technology involved in catalytic oxidations.
The disciplines of homogeneous catalysis and organometallic
chemistry are closely related. So much so, in fact, that a major
justification for the study of organometal1ic complexes has been
their potential use as catalysts. Even though the overwhelming
majority of work in this area has dealt with mono-metallic systems,
the field of multi-metal 1ic catalysts is beginning to emerge as an
area rich in potential for catalytic research. Systems containing
more than one metal have several advantages over their mono-metal
counterparts. Enhanced stability as well as synergistic
interactions between the metals would give multi-metal 1ic systems a
range and versatility unknown in systems containing a single metal.
In theory, the judicious choice of the combination of metals should


6
lead to a "tunable" catalyst system one where selectivity or
conversion is directly related to the metals involved.
The series of trinuclear metal carboxylate compounds is an
ideal choice for carrying out such studies. Their versatility,
combined with the wealth of knowledge available on the coordination
chemistry of these complexes, make them excellent choices as
subjects for the study of homogeneous catalysis.
Trinuclear Carboxylate Complexes
The family of trinuclear carboxylate complexes has been known
in the chemical literature for over 100 years. Only in the more
recent past have these complexes been extensively studied and
characterized. These studies are extremely interesting in light of
the versatility and uniqueness of multi-metallic systems in general.
The synergistic effects of the presence of two or more metals in
close proximity has been widely studied recently'7; several varied
applications of such systems are obvious in biochemistry and enzyme
studies (tryptophan 2,3-dioxygenase, for example, consists of both a-
Cu(II) and an iron porphyrin in the active site) as well as
industrial processes involving transition metals on inorganic
supports (SMSI interactions between Ti and Ru and other platinum
metals in Fisher-Tropsch synthesis), and other commercial
applications (oxidations by Co(11) involving Mn(II) as a cocatalyst,
and the widely studied Ziegler-Natta polymerization system which
involves the combination of Zr or Ti and A1 as the catalytic
species).


7
All of the trinuclear metal carboxylate complexes or "basic
carboxylates"^ discovered to date have virtually the same basic
structure (Figure 1.2). (Although not trimers by the strictest
definition, these trinuclear ruthenium carboxylates will be referred
to as "trimers" for the sake of brevity.) The major differences in
these systems occur in metal-metal distances and the planarity of
the M3-O core. The equilateral triangle formed by the metals (as
the apices of the triangular M3-O core) is bridged above and below
the plane of the triangle by bidentate carboxylate ligands, and each
metal is connected via a central, three-coordinate oxygen atom.
Unlike their dimeric cousins,^ these complexes contain no formal
metal-metal bond. The remainder of the pseudo-octahedral
configuration around each metal atom is completed by the ancillary
ligand L. Obviously a great deal of versatility is inherent in
these complexes not only can the metals used be widely varied, but
the carboxylate bridges and L also increase the permutations
possible. To date, almost all of the first-row transition metals^'
have been isolated as "basic trinuclear carboxylates" (V,15,16
Cr,17-20 Mn,21-24 p6j25-29 an(j qo30-31). others, like Ir,32 Ru,33-37
and Pd, Pt, and Rh38-41 have also been prepared. Titanium^ and
zirconium^ w-¡n also form a trinuclear complex slightly distorted
from the traditional basic carboxylate structure involving a central
hydroxy bridge between the metal centers. The carboxylate ligand
can vary from acetate to butyrate for all of these compounds;
partially chlorinated carboxylates as well as fluorinated ones have
also been used. The ligand L is most often a classical coordination


8
Figure 1.2 Generalized structure of basic trinuclear carboxylates
having the formula [M30(02CR)gL3]n.44


9
ligand such as pyridine, PPh3, or even H2O or diethyl ether. The
last forms of variation take place in terms of the metals involved.
As these complexes are most commonly isolated, the metals are found
in the +3 oxidation state, causing the cluster as a whole to have a
+1 charge. The other most commonly found form of these basic
trinuclear carboxylates is one with one metal in the +2 oxidation
state, rendering the complex neutral. In these systems, the
assignment of oxidation states is truly a formalism. The iron and
ruthenium complexes in particular can be classified as Robin and Day
Class III compounds, indicating complete delocalization of the metal
electrons. This classification is especially important in the
trinuclear ruthenium carboxylates, the subject of this work.
This versatility has made this family of complexes choice
candidates for a wide range of studies. The mixed-valence, neutral
species (primarily the iron complexes) have been extensively studied
in terms of intramolecular electron transfer reactions.45-48 j^g
manganese clusters, as well as similar dimeric systems, have been
studied in hopes of elucidating the role of Mn in photosynthesis as
well as for catalysis.49-51 j^e more classical inorganic chemistry
of these complexes has also been studied, including ligand-exchange
reactions, for example.
Due to the vast information available on these trimers, it is
not unreasonable to expect some studies in terms of their
usefulness. The synergism expected to occur between the metal
centers should manifest distinct differences from their monomeric
analogs. The variety of oxidation states available, combined with


10
the ready exchange of ligands, makes these complexes ideal choices
for catalysts, especially of homogeneous processes. Finally, the
carboxyl ate ligands have been shown to be relatively inert to
oxidation processes, as evidenced by the widespread use of metal
acetates (specifically Co(II) and Mn(II)) and acetic acid in
industrial oxidations.52 For these reasons, the family of
trinuclear metal carboxylate complexes would be expected to be good
catalysts for a variety of homogeneous processes.
Only in the last few years have widespread attempts been made
to utilize these complexes as catalysts. The cobalt acetate trimer
(Co30(0Ac)g(H20)3+ and others) has been proposed to be one of the
active catalytic species in the oxidation of p-xylene to
terephthalic acid (the Mid-Century/Amoco Process).52 It has also
been shown to oxidize toluene and other hydrocarbons under
relatively mild conditions.13,30,31,53 others have been used as
catalysts as well. The rhodium acetate complex uses t-butyl
hydroperoxide to selectively oxidize cyclohexene,5^ while the iron
acetate trimers have been proposed to catalyze a variety of organic
transformations.55-5^ By far the most widely studied (in catalytic
terms) of these trinuclear species is the ruthenium complex.
More literature is available on the ruthenium system in both
catalytic and chemical terms than most of the other trimers. These
complexes were first isolated and characterized22 in 1972; Spencer
and Wilkinson found these trimers to be unique in the family of
basic trinuclear acetates for several reasons. Both mixed-valence
and cationic trimers were readily isolable. These complexes


11
underwent a one-electron non-reversible wave electrochemically, and
readily underwent ligand exchange as well. The most unusual feature
of the ruthenium trimers was the reversible removal of the central
oxygen atom, a reaction unknown for the other trinuclear metal
systems. Later studies by Meyer et al., expanded^^'^ the original
electrochemical studies, revealing for the Ru30(0Ac)g(pyz)3 (where
pyz = pyrazine) complexes a series of five one-electron reversible
waves. Four of these waves were attributed to the metal center,
corresponding to formal oxidation state changes from Ru(II1, 111,11)
to Ru(IV,IV,III). Linking these complexes into multinuclear
oligomers revealed systems that would undergo up to ten one- or two-
electron waves, justifiying the nickname "electron sponge" for these
ruthenium complexes. A generalized MO scheme,shown in Figure
1.3, shows several orbitals of the 7r system of the RU3O core in a
relatively small energy range. For the cationic, Ru(111,111,111)
complexes, all levels up to Ej" are filled; the A2' level is only
partially occupied. The orbitals containing the metal electrons are
virtually indistinguishable, the justification for the Class III
label of delocalization. This depiction also helps explain Meyer's
electrochemistry as well as other spectroscopic properties of these
complexes.
In all likelihood, the electrochemistry revealed for these
complexes prompted the widespread study of the ruthenium complexes
as catalysts. Olefin hydrogenations were first studied*^ by
Wilkinson; further studies were carried out both homogeneously and
heterogeneously supported on a carboxyl ate resin by Rempel and


"O "O "O
12
o
> X
Ru
Figure 1.3 Qualitative molecular orbital description for [Ru30(02CR)6l-3]n
Comdexes, involving only the n system of the RU3O core. After Wilson et


13
others.64-66 Ziolkowski, et al,. have studiedl3>67-69 Nineties
of cumene hydroperoxide decomposition and the exchange of DMF for
H2O using NMR techniques; they also have reported13 some catalytic
work in alkane oxidations. The ruthenium trimers have also been
involved in the Prins reaction,^ oxidative dehydrogenation of
saturated carbinols,^1^ and dimerization of acrylonitrile.^3 In
terms of oxidation catalysis, the ruthenium trimers, in the presence
of hydrogen peroxide, will oxidize substituted phenols to the
corresponding hydroquinone.^ ¡n a mixed solvent system containing
water, carbon tetrachloride, and acetonitrile, the Ru acetate trimer
with periodate will oxidatively cleave alkenes,^3 similar to the
traditional chemistry observed for RUO4. They will also catalyze
the isomerization of allylic alcohols.^
A study of the use of the ruthenium trimers as oxidation
catalysts for a variety of organic transformations seemed
potentially interesting, based on their previous use as catalysts
and the large amount of electrochemical potential to be tapped in
these complexes. The use of molecular oxygen as the primary oxidant
has been an ongoing area of research, and the ruthenium trimers have
not previously been shown to be active as catalysts in such a
system. This work involved the continuation of the study of the
oxidation of alcohols by the ruthenium carboxylate trimers,^ as
well as an extension of these catalytic studies to the oxidations of
alkenes and alkanes by a new member of the trinuclear ruthenium
carboxylate family, [Ru30(pfb)g(Et20)3](pfb), which has been
synthesized and characterized.


CHAPTER II
ALCOHOL OXIDATIONS BY TRINUCLEAR
RUTHENIUM CARBOXYLATE COMPLEXES
Introduction
The oxidation of alcohols is a procedure long known and used in
organic chemistry for the production of aldehydes, ketones and
carboxylic acids. Mild reagents, such as Cr03/pyridine or Mn02,
react with alcohols to give primarily the carbonyl product (aldehyde
or ketone). Stronger oxidants, like RUO4, continue to oxidize
primary alcohols through an aldehyde intermediate to carboxylic
acids. Other high-valent ruthenium-oxo ions such as RuO^' or RUO4'
will oxidize primary alcohols to carboxylic acids, secondary
alcohols to ketones, and will oxidize unsaturated alcohols without
attacking the double bond.^ Autoxidation of alcohols tends to
produce ketone or acid along with hydrogen peroxide. Shell
commercialized a process for the production of hydrogen peroxide by
the autoxidation of 2-propanol (Eqn 2-1).^
(CH3)2C(H)0H > (CH3)2C=0 + H202 (98%) Equation 2-1
The reaction of alcohols with noble metals such as Pd or Pt to
give carbonyl products and a metal hydride species is well-
14


15
documented.a mechanism involving a ^-hydride elimination to
give metal hydrides is generally assumed. This mechanism is also
invoked for the Pd(II)-catalyzed oxidation of secondary alcohols to
ketones with oxygen at 25 C.8
A great deal of literature has been published on ruthenium-
catalyzed oxidations of alcohols.^ Besides the general uses of
RUO4, low-valent Ru(II) complexes have been widely studied as
oxidation catalysts with both O2 and milder oxidants such as
iodosobenzene.8 The most widely studied compound of this type,
RuCl2(PPh3)3, has been used to oxidatively dehydrogenate alcohols
with oxygen.80 with iodosobenzene, RuCl2(PPh3)3 will selectively
oxidize primary alcohols to aldehydes.8 Sharpless et al. have
found N-oxides combined with RuCl2(PPh3)3 and other ruthenium
compounds will also oxidize alcohols to their respective carbonyl
products.81 In benzene solvent, this complex preferentially
oxidizes long-chain primary alcohols over the corresponding
secondary alcohol.82 Using O2 as the primary oxidant, RuCl2(PPh3)3
oxidizes allyl alcohols to a,/?-unsaturated carbonyl complexes in a
variety of relatively poorly coordinating solvents.83 in all of
these oxidations, several general trends arise. All of these
oxidations are shut down in the presence of strong donor solvents
like acetonitrile, indicating coordination of the substrate is
necessary for oxidation to occur. Replacement of a coordinated
nitrile by an alcohol is not highly likely. The mechanism
consistently invoked for these reactions involves the coordination
of alkoxides to a Ru(IV) species with subsequent ^-hydride


16
elimination to give carbonyl product and a Ru(II) hydride. The
hydridic species can be oxidized back to Ru(IV) by the available
oxidant, creating a catalytic cycle.
A wide variety of other ruthenium compounds have also been used
to catalyze the oxidation of alcohols. A ruthenium hydride,
RuH2(PPh3)4, catalyzes the condensation of alcohols to esters and
lactones at elevated temperatures.^ Monomeric ruthenium complexes
containing fluorinated carboxylate ligands have also been shown to
dehydrogenate primary and secondary alcohols via a /3-hydride
elimination pathway.85,86 slightly more active compounds containing
diphosphine ligands have also been prepared and demonstrated to be
catalytic.'7 The mixed-valence ruthenium carboxylate dimer
[Ru2(0CR)4C1] has been shown to dehydrogenate methanol to
formaldehyde under relatively mild conditions. Ruthenium
complexes as simple as commercially available ruthenium trichloride
have also been shown to be active for both the oxidation of
secondary alcohols and amines with oxygen. Ruthenium( 111)
solutions will also oxidize allyl alcohol to acrolein.
Ruthenium complexes containing large, bulky ligands have also
been used for alcohol oxidations. Riley demonstrated a DMSO adduct
of Ru(II), RuX2(DMSO)3L, would catalyze the aerobic oxidation of
thioethers to sulfoxides. This reaction required a reducing
solvent, alcohol, to reduce the Ru(IV) species back to the active
Ru(II) complex, generating a carbonyl product.*
Bidentate imines have also been used with Ru(II) to oxidize
coordinated alcohols in conjunction with 03.^ In these systems, a


17
Ru(IV) to Ru(II) cycle is again proposed as the pathway of the
oxidation, and a disproportionation step enabling an escape from an
inactive Ru(III) species to active Ru(II) and Ru(IV) complexes is
also invoked. An unusual account of a Ru(III) complex containing
l,3-bis(2-pyridylimino)isoindoline (BPI) ligands is also involved in
the oxidation of alcohols.^ The use of Ru(III) is unusual in that
Ru(III) complexes, generally low-spin t2g^, tend to be
substitutionally inert.94,95 since the availability of open
coordination spaces is a requirement for a feasible, selective
homogeneous catalyst, Ru(III) complexes would not be expected to be
vary active catalytically. Gagne's system was active for alcohol
oxidations, producing around 60 turnovers (moles of product per
moles of catalyst used) in 24 hours when a strong, noncoordinating
base is present. Secondary alcohols formed ketones which were inert
to further oxidation. Primary alcohols were oxidized initially to
aldehydes (the primary product) which could react further giving
acetals and other products. Again a disproportionation of Ru(III)
to Ru(II) and Ru(IV) is proposed, with a Ru(IV)-coordinated alkoxide
species as the active intermediate. Hydridic ruthenium(II) may be
an intermediate in this reaction as well, arising from the /5-hydride
elimination of the Ru(IV)-alkoxide species.
T. J. Meyer has also contributed to this area with his well-
studied ruthenium polypyridyl complexes. Extensive kinetic and
mechanistic studies on alcohol oxidations by these high valent
ruthenium-oxo complexes have been carried out.96-98


18
In light of the extensive, ongoing research into oxidations by
ruthenium complexes in general, and the high potential for catalysis
demonstrated by the trinuclear ruthenium carboxylate complexes,
these particular complexes were chosen to screen as catalysts for
the oxidation of alcohols by molecular oxygen.
Previous Work
Bilgrien discovered that Ru30(prop)6(H20)3+ would catalyze the
selective oxidation of primary and secondary alcohols to the
corresponding carbonyl product using O2 as the primary oxidant.44 A
wide number of alcohols were active in this system, and in all cases
the only product formed was the aldehyde or ketone, with no traces
of carboxylic acid observed. Several different trimeric ruthenium
carboxylate complexes were found to be effective catalysts as well.
These oxidations exhibit a slight rate dependence upon acidity, as
demonstrated by the inhibition of the reaction upon the addition of
acids. On the other hand, bases had a curious effect on the
reaction. Sodium ethoxide enhanced the catalysis, 2,6-lutidine
inhibited the reaction, and NaOH caused precipitation of the
catalyst.
In mechanistic terms, Bilgrien found that for every mole of
carbonyl product produced, a mole of water was formed as well,
implying the four-electron reduction of oxygen to water.44 Hydrogen
peroxide, a likely intermediate in this process, was never detected
in the reaction mixture. These complexes would also oxidize
alcohols with H2O2 in place of O2 as the primary oxidant. A rough
calculation using the pressure drop of the pressure gauge for the O2


19
consumption showed that for each mole of O2 consumed, two moles of
product are produced. The rate of the reaction in terms of oxygen
pressure was found to be .25. The catalyst did not seem to
decompose during the reaction, as indicated by both IR and *H NMR
results. Bilgrien also found that the mixed-valence trimer was
readily oxidized by O2 in alcohol solution to the Ru(111,111,111)
complex, but the reduction of this species by alcohol did not occur.
Bilgrien's mechanism for the ruthenium trimer-catalyzed
oxidation of alcohols is shown in Figure 2.1.^ This scheme invokes
the Ru(III,III,II)-alcohol species as the active intermediate, which
undergoes intramolecular disproportionation to form a Ru(IV,II,II)
ruthenium species. The decomposition of this intermediate could
occur via a number of pathways, the most likely of which involves a
two-step reduction of the alcohol by the trimer. This reduction
would generate the Ru(II,II,II) species, without the central ¡j.3-
oxygen first observed by Spencer and Wilkinson,33 which would
readily be oxidized back to the Ru(III,III,II) intermediate in the
presence of O2.
Scope of Catalysis
Apparatus. All pressurized oxidations were carried out in
slightly modified Parr hydrogenation setups (Figure 2.2). This
apparatus has previously been described in detail by Zuzich and
Bilgrien.44,99 por these oxidations, stainless steel pressure
heads, constructed from Swagelok fittings and equipped with standard
sample valves, gas gauges, were also equipped with pressure relief
valves as a safety precaution. This apparatus was directly


20
Ruin
I
0
RI \ul
/\
2
RCH2OH
RuII
I
'\
Ru111 Ruin(RCH2OH)
Ru
II
Ru
II
Ru11(RCH2OH)
RuII
H I
Ru1I RuiV(0CH2R)
RCHO + H20
rch2oh
Figure 2.1 Bilgrien's proposed mechanism for the oxidative
dehydrogenation of alcohols by [Ru30(02CR)g(L)3]n. Other ligands
have been omitted for clarity.^


21
Gas
Outlet
1/4" Silicone
Sceptum
1/4" Tu be to 1/8" NPT
Adapter
On/Off Ball Valve
1/8"NPT
Pressure
Gage
1/4"El bow.
1/4"Cross
d?Wing Nut
Pressure
Bottle
Figure 2.2 Schematic diagram of a standard pressure head without
safety release valves.


22
connected to an oxygen tank by copper tubing. The direction and
path of exit gases were controlled by a length of tygon tubing
attached to the exit valve which ran to the back of the hood. A
glass, 250ml, Parr hydrogenation bottle (the reactor vessel) was
attached to this apparatus by a #6 silicone gum rubber stopper and a
metal cage. The bottle is surrounded inside of this cage by an
aluminum sheath, designed to theoretically reduce the amount of
glass shards that would be produced in an explosion.
Sampling techniques have been previously described in detail by
Bilgrien.44 Briefly, a 1-mL gastight stringe, equipped with a Leur-
lok syringe valve and a 12-inch needle, is inserted through the
septum at the top of the pressure head with the valve closed. The
needle is guided through the ball valve into the reaction mixture.
The syringe valve is opened, a small aliquot withdrawn ( -.2 mL),
the valve closed, and the needle withdrawn. With practice, this
procedure can be accomplished quickly, safely, and with no
observable pressure loss. The aliquots are analyzed using GC,
GC/MS, and GC/IR.
Oxidation procedure. Typical oxidations involved 50 mL of
alcohol (as both solvent and substrate), 1 mL ketone standard, and
10~5 moles of catalyst. Reactions were carried out in a 65 C
silicone oil bath monitored by an Omega 6100 temperature controller
and thermocouple under initial pressures of 40 psig of 02- Stirring
rates of the solutions were controlled by magnetic stirrers beneath
the oil bath; the oil bath was circulated by an overhead stirrer.


23
A slightly modified version of Bilgrien's technique44 was used
for the alcohol oxidations. The pressure bottle was charged with
all components of the reaction except catalyst (i.e., substrate,
standard, and a stirbar), covered with Parafilm, and placed in the
oil bath to equilibrate for 20 30 minutes. The catalyst was added
to the warm solution, the apparatus assembled and pressurized,
placed in the oil bath, and a sample withdrawn. This sample was
denoted as time zero and the start of the reaction. The reaction
was stirred as rapidly as possible to ensure the saturation of the
solution by 02-
Safety precautions. CAUTION! Combinations of warm organic
liquids and dioxygen are potential 1v explosive. Great care should
be taken, especially during setup and dismantling of oxidation
reactions, to avoid sparking the reaction mixture and causing a
violent explosion. General safety precautions to follow include (1)
let the reaction mixture cool to room temperature before
depressurizing; (2) be sure outlet gases are directed away from any
source of sparks; and (3) become aware of the explosion limits of
solvent, substrates, and the oxidant (whether air or O2) before
beginning an oxidation.
Calculations. Amounts of products formed were determined by GC
in all cases. Calibration curves relating moles of product to
relative peak areas were constructed for all products formed.
Standard procedure involved making up a series of solutions
containing a varying, known amount of product, and a constant amount
of standard in the solvent (alcohol) used. Repeated (at least five)


24
injections of each of these solutions gave a statistically valid
value for the area percent of the product peak. Knowing the number
of moles of product and standard in each solution gave a mole ratio
of product to standard, which can be plotted against the ratio of
the area percents of the product and standard. The area percents
are obtained electronically from the integration of the peak areas
of the GC chromatogram by an integrator. From the graph of mole
ratio to area percent ratio, the number of moles of product can be
obtained, if the amount of standard added is known.
Results and Discussion
Although Bilgrien's proposed mechanism adequately described the
experimental data,44 further work remained to be done to
substantiate this proposal. Areas to be addressed included the
differences in terms of activity between the mixed-valence and the
cationic trimers, and the fate of the catalyst during the reaction.
Other substrates should be tested, and more kinetic data should be
accumulated as mechanistic support. For these reasons, this
research project was continued.
Bilgrien found the activity of the ruthenium trimer catalysts
varied greatly depending upon the amount of purification of the
complex.44 Liquid chromatography on a four-foot Sephadex LH-20
size-exclusion gel gave the best results, with dramatic effects on
the catalysis, as shown in Figure 2.3. A general activity curve for
the ruthenium propionate trimer, the standard catalyst for most of
the remaining reactions, is shown in Figure 2.4.


TURNOVERS
25
Figure 2.3 Chromatographed vs. unchromatographed
[Ru30(prop)g(H20)3](prop) in isopropanol oxidations.


1600
1400
1200
1000
800
600
400
200
0
igure 2
sopropa
26
iiii|iiii|iiii|iiir
50 100 150 20
TIME (HRS.)
Activity curve for [Ru30(prop)g(H20)3]+ catalyzed
oxidations.


27
Other alcohol substrates were screened to further test the
versatility of these trimers as catalysts (Table 2-1). Benzyl
alcohol, as expected, produced only benzaldehyde, and allyl alcohol
was exclusively oxidized to acrolein. Both of these substrates were
oxidized significantly slower than the isopropanol oxidation used as
the common standard for comparison in Bilgrien's work. While
isopropanol oxidations resulted in 147 turnovers in 12 hours, these
substrates only produced 40. Bilgrien also found the rate of
reaction slowed as the substrate varied from primary to secondary
alcohols. These substrates follow this general trend, as the rate
of reaction for both benzyl and allyl alcohol is slower than that
for either primary or secondary alcohols. This reduction in rate
for benzyl alcohol is most probably due to steric bulk and
subsequent hindrance in binding the substrate to the metal center.
The oxidation of allyl alcohol is slower due to a different mode of
substrate binding similar to that proposed by Taqui Khan.90 In the
RuCl3-catalyzed oxidation of allyl alcohol by O2, the substrate is
bound in two sites around the octahedral ruthenium center once by
the double bond and once at the OH moiety. A /5-hydride transfer
creates a Ru(III) hydride-alcohol(+) species which is quickly
oxidized by O2 to give acrolein and the regenerated catalyst. No
hydride species is postulated for the trimers, but the relative
slowness of the reaction could be attributed to the inability of the
alcohol to bind at the olefinic site (vide infra), and a loss of
stability in the reduced ruthenium-alcohol intermediate.


28
Table 2-1
Alcohol Substrates3
substrate
T (C)
product
to/12 hrs^
to/24 hrs
ethanolc
25
No reaction
65
acetaldehyde
198
313
isopropanolc
65
acetone
147
254
100
acetone
685
1015
n-propanolc
65
propanal
430
645
n-butanolc
65
butanal
d
d
cyclohexanolc
65
cyclohexanone
d
d
t-butanolc
65
No reaction


benzyl alcohol
65
benzaldehyde
40
e
allyl alcohol
65 '
acrolein
42
e
50%
65
acetone
75
e
isopropanolf
phenolS
65
No reaction


a) reaction conditions are as outlined under "Scope of Catalysis"
b) to = turnovers defined as moles of product/moles of catalyst used
c) from Bilgrien^
d) not quantified
e) reaction run for only 12 hours
f) auxiliary solvent used was acetonitrile as 50% by volume
g) solvent used was acetonitrile


29
Another congener of the ruthenium carboxylate trimer family,
[Ru30(prop)g(py)3](PFe), was synthesized. A bar graph comparing all
of the different trimers used in shown in Figure 2.5. The pyridine
adduct is completely unreactive in the oxidation of isopropanol,
indicating that coordination of the substrate in place of the
ancillary ligand L is necessary for catalysis to occur. When
graphed in terms of turnovers, the differences in the cationic and
the mixed-valence trimers becomes even more striking than Bilgrien
reported. The mixed-valence compounds are greater than three times
more active than their cationic counterparts. These differences are
made more enigmatic by the known reaction chemistry of these
complexes. The Ru(111,II1,11) trimers are readily oxidized to the
Ru(111,111,111) complexes by C>2.^ However, the non-lability of
Ru(III) centers towards substitution is wel 1 -documented5^>9E>; ^he
Ru(111,111,111) trimer would be expected to exchange H2O (or L) for
alcohol ligands very slowly. Assuming coordination of substrate is
necessary for oxidation to occur, the Ru(111,111,111) system should
oxidize alcohols more slowly than the more labile Ru(III,III,II)
counterparts. The exchange of ligands in the Ru(III, III, II) trimer
would be faster, so that even if the oxidation of the trimer from
the Ru(III,III,II) to the Ru(111,111,111) did occur, a molecule of
alcohol would already be present in the coordination sphere of the
catalyst. This would explain some of the differences in the
catalytic activity of these complexes.
A comparison of the ruthenium trimers with other complexes
reported in the literature to oxidize alcohols would be informative


TURNOVERS/24. HOURS
30
CATALYSTS
Figure 2.5 Comparison of [Ru30(02CR)g(L)3]n catalysts for
isopropanol oxidations. (A) [Ru30(0Ac)6(H20)3]+ (B)
[Ru30(prop)5(H20)3]+ (C) [Ru30(prop)6(py)3](PFg) (D)
[Ru30(prop)g(H20)3] (E) [Ru30(prop)g(PPh3)3].


31
in terms of gauging the activity of this system. A graphical
comparison is shown in Figure 2.6. As mentioned in the introduction
to this chapter, both RUCI3 and RuCl2(PPh3)3 have been shown to
oxidize alcohols using molecular oxygen as the primary oxidant.
Since the trimers also operate using oxygen, these systems should be
enlightening for comparing relative reactivities of the catalysts.
The trimers are approximately 10 times more active than the other
ruthenium complexes attempted, on the basis of turnovers in 12
hours. Even taking into account that the trimers contain 3 moles of
ruthenium per mole of catalyst, while the others only have one, the
trimers are still over three times more active.
The mechanism of these oxidations could safely be assumed to
not involve autoxidation pathways, due primarily to the selectivity
observed in the reaction. If free radicals were involved in these
oxidations, the further oxidation of aldehydes to carboxylic acids
would be expected. However, acid products are not observed under
our conditions, leading to the assumption the trimers are selective
oxidants. To further justify this claim, typical reactions designed
to prove or disprove free radical chain mechanisms were carried out
(Figure 2.7). The addition of benzoquinone, a free radical trap, to
a typical oxidation has no effect on the reaction. A free radical
initiator, AIBN (azobis(iso-butyronitrile)), was added to the
reaction in place of the catalyst and achieved approximately 10
turnovers in 1 hour and ceased to function. These experiments
emphasize the non-radical nature of these oxidations.


TURNOVERS
32
Figure 2.6 Comparison of various ruthenium catalysts in isopropanol
oxidations.


Figure 2.7 Free radical experiments in isopropanol oxidations.


34
Bilgrien noted that for every mole of product formed, one mole
of water was also produced.44 If the assumption the substrate must
coordinate in order for oxidation to occur is valid, the effects of
adding or removing water in the reaction should prove useful in
determining a mechanism (Figure 2.8). The addition of 5A activated
molecular sieves to the reaction greatly accelerated the rate, while
a reaction run in a 50/50 mixture of isopropanol and water showed a
drop in activity after about five hours. This curvature, indicative
of catalyst deactivation, is not observed in the activity curve
until after 180 hours of reaction time. Seemingly, the presence of
water slowly inactivates the catalyst.
The catalyst does not seem to decompose during catalysis,
according to NMR and IR.44 UV-Visible spectroscopy has been very
informative in determining the active species in solution. Since no
induction period is observed for these oxidations, either the
trinuclear carboxylate complexes is the active catalytic species, or
it is a precursor that converts rapidly to the active species in
solution. The lack of an induction period also indicates that the
two different versions of the trimer (Ru(III,111,111) and
Ru( III, II1, 11)) perhaps perform the oxidation by slightly different
pathways. Bilgrien noted that while the Ru(III,III,II) could be
oxidized to the Ru(111,111,111) in solution, alcohol was not a
strong enough reducing agent to perform the reverse reaction.44
However, a distinct color change is observed when an alcoholic
solution of the catalyst is heated to 43 C under an inert
atmosphere. The changes were monitored via UV as shown in


TURNOVERS
Figure 2.8 Effects of H2O on isopropanol oxidations.


36
Figure 2.9. These changes correspond to the conversion of the
Ru(111,111,111) to the Ru(111,111,11) complex as reported by
Wilkinson.33 Jo effect this change, the solution had to be heated
for 18 hours. However, oxidations were performed at 65 C, so this
conversion may well occur under typical oxidation conditions. If
this conversion were accompanied by production of ketone, the amount
produced (assuming either a stoichiometric conversion either per
mole of catalyst or per mole of ruthenium) was too small to be
detected by GC. This change is reversible; the addition of O2, 30%
H2O2, or air to the warm alcohol solution immediately oxidizes the
Ru( III, 111, II) back to the Ru(111,111,111) with the corresponding
color change. The color change corresponding to this conversion is
not observed under our catalytic conditions; if present, the
Ru( 111,111, II) complex would be a transient species at best. These
UV-visible studies indicate a Ru(111,III,II) intermediate created
from a Ru(III,III,III) precursor would be a very slow but possible
process. They give little or no information about the pathway used
by a Ru(III,III,II) precursor, however.
The role of H2O2 in these oxidations was also pursued further.
Figure 2.10 shows the effects of adding H2O2 to typical alcohol
oxidations. Hydrogen peroxide is a potent oxidant by itself, as
demonstrated by the upper curves. However, the ruthenium trimer
catalyst will use peroxide in the absence of O2 to oxidize alcohols
to the same carbonyl products. If hydrogen peroxide is an
intermediate in the reduction of O2 as postulated by Bilgrien,^
these graphs indicate the peroxide would be consumed as a co-oxidant


Irt (T >
37
X (nm)
Figure 2.9 UV-Vis studies of [Ru30(prop)g(H20)3](prop) (A)catalyst
in ethanol under N2, 25 C (B) catalyst in ethanol under N2, 43 C
(C) solution (B) exposed to 02-


TURNOVERS
38
Figure 2.10 Role of H2O2 in isopropanol oxidations.


39
in the oxidation reactions. The amount of peroxide formed would, in
all probability, be small and would be consumed as rapidly as it
formed. A low steady-state concentration of peroxide would be one
explanation for the failure to identify peroxide in the reaction
mixture as wel1.
Determining kinetics in this system was based on the method of
initial rates from initial concentrations.The rate expression
was assumed to take the form of Equation 2-2.
dx/dt = k0bs [cat]3 (P02)^ [substrate]0 Equation 2-2
In the alcohol oxidations, the substrate alcohol is present in much
higher volume and the conversion of alcohol to product is relatively
small. Therefore, the substrate concentration was assumed to be
relatively constant, giving the rate equation 2-3.
dx/dt = k'obs [cat]3 (P02)^ Equation 2-3
To obtain the order of the reaction with respect to each remaining
component, one variable was held constant while the other varied.
The rate of the reaction (dx/dt) was assumed to be the slope of the
straight line obtained from a plot of mole of product formed vs.
elapsed time. The appropriate mathematical manipulations gives a
ratio of the rate laws which will yield a value for the reaction
order.(Equations. 2-4,2-5,2-6).


40
(dx/dt)i = k'obsi [cat]ai (PO2)bi Equation 2-4
(dx/dt)1 k'obsi [cat]ai (P02)bi
- = --- Equation 2-5
(dx/dt)2 k'obs2 [cat]a2 (P02)b2
Holding one variable constant (for example, (P02)) gives Eqn 2-6.
log (dx/dt)1 log (dx/dt)2
a = Equation 2-6
log [cat]i log [cat]2
Several reactions were run where each of the variables was changed
in turn; this data is given in Table 2-2 and graphically in Figures
2.11 to 2.14.
Varying the concentration of the ruthenium catalyst (numbers 1,
4, and 5 in Table 2-2) lead to essentially first-order kinetics
(Figures 2-11 and 2-12). Varying the oxygen pressure was slightly
more demanding in that the total pressure had to be kept at 40 psig
for comparison purposes (numbers 1, 2, and 3 in Table 2-2). The
remainder of the pressure was made up of argon. The reaction order
was found to be approximately .2 in 02, very close to the value of
.25 reported by Bilgrien (Figures 2.13 and 2.14).44 For all
practical purposes, however, the reaction could be considered zero-
order in oxygen, considering the amounts of cumulative error in the
analysis, calibration curves, calculations, and the differences in
the values obtained mathematically and graphically.
A proposed mechanism for these oxidations is shown in Figure
2.15. This scheme differs significantly from that given by Bilgrien
in several areas. The mechanism, beginning with the more labile


41
Table 2-2
Kinetic Data for Alcohol Oxidations
by Ru30(prop)6(H20)3+
rate
law = dx/dt
= k'obs
[Ru]a (P02)
b
Exd.
[Ru]J
rx io'4i
Dsiqb
P02
n/vc
dx/dt^
x 10-3
ae
bf
1
8.89
44.0
.108
8.60
1.10
.185
2
1.01
16.3
.040
7.16

.259
3
9.24
27.5
.0675
7.62

.119
4
2.80
45.5
.112
2.42
1.19
5
218.2
45.0
.110
22.3
1.33
a) Concentration calculated in moles/liter using 50 mL as the total
volume.
b) Initial pressure of reaction in psig
c) N/v calculated from the ideal gas law (PV = nRT) assuming a
volume of 270 mL and 65 C.
d) dx/dt has units of molarity/hour; calculated as explained in text
e) a = 1.20 .12
f) b = .187 .07


Figure 2.11 Kinetics: varying catalyst concentration.


43
1.00 -I
0.90 -
0.80 -
0.70 -
0.60 -
X
T>
0.50 -
o>
o
I 0.40 -
0.30 -
0.20 -
0.10 -
0.00
m = 1.04
T 1 1 1 1 1 1 1 1 1 r
2.5 3.0 3.5
- log Ru
t r
4.0
Figure 2.12 Kinetics: order in catalyst


44
Figure 2.13 Kinetics: varying initial O2 pressure.


45
Figure 2.14 Kinetics: order in O2.


46
L-RuIII
\
Ruin
/
0
Ruin
L-Ru111
\
RuIH-L
/
RuIH-L
/
0
Ru'I
I
L
Figure 2.15 Proposed mechanism for the [Ru30(02CR)6(L)3]n -
catalyzed alcohol oxidations. Carboxylate ligands have been omitted
for clarity.


47
Ru(III,III,II) species, involves first replacement of the ligand L
by a substrate molecule with concomitant loss of a proton. This
Ru(III,III,II)-alcohol species is postulated to be the active
intermediate in this cycle. Oxidation of this species by O2 (or
later in the cycle, H2O2) gives a species that can be formulated as
a Ru(III,III,III)-alkoxy radical or a Ru(II1,111,IV)-alkoxide
species, depending on the placement of the extra electron. In Robin
and Day Class III systems, this placement is more or less semantics.
Reductive elimination from this species gives carbonyl product and a
coordinatively unsaturated Ru(111,III,II) species. Solvation of
this species by another mole of alcohol regenerates the active
Ru (111,III,II) species.
The Ru(111,111,111) complex is slightly different in that to
reach the active species it must undergo a one-electron reduction
and replace L by a mole of alcohol. Ruthenium(111) species are, in
general, substitutionally inert,94,95 so replacement step would
be expected to be very slow. The UV-vis studies have demonstrated
the reduction process to be slow as well. These two reasons help
explain the differences between the two congeners.
The kinetics observed experimentally can be verified
mathematically using the mechanism proposed in Figure 2.15. Each
step in the mechanism can be written out and a rate expression
derived for each step (Equations 2-7 through 2-17) using standard
procedures and assuming the steady state approximation is valid for
Equations 2-14, 15, and 16.


48
Ru^3,2 0H2 + ROH v Ru332 OR + H2O Equation 2-7
A k_i B H
k2
Ru3,3,2 OR
B H k_2
Ru332 OR + H+
C
Equation 2-8
Ru332 OR + O2
C
k3
4 Ru333 OR + 022'
D
Equation 2-9
ru3,3,3 or -
D
Ru332 + ROH
E
k4
^5
-) Ru332 + R2C=0 Equation 2-10
E P
-> Ru332 OR
B H
Equation 2-11
dP/dt = k4(D)
Equation 2-12
dA/dt = -ki(A) + k_1(B)
Equation 2-13
dB/dt = -k2(B) + k_2(C) k_2(B) + ki(A)
k-2(C) + ki(A) = k2(B) + k_1(B)
k2(B) = k_2(C) + ki(A) k_1(B)
0 Equation 2-14
Equation 2-14a
Equation 2-14b
dC/dt = -k3(C)(02) + k2(B) k_2(C)
k3(C)(02) = k2(B) k_2(C)
Equation 2-15
Equation 2-15a
dD/dt = k3(C)(02) k4(D) = 0
k3(C)(02) = k4(D)
Equation 2-16
Equation 2-16a
Rearrangement and subsequent substitution of Equations 2-14,
15, and 16 into the expression for dP/dt as shown below give the
rate expression for dP/dt in Equation 2-20. This expression can be
reduced to pseudo first-order in catalyst if k.¡(B) is assumed to be


49
small. Under our conditions, a large excess of alcohol, the reverse
reaction in Equation 2-7 should only occur to a small extent by Le
Chatelier's Principle, so the assumption seems to be valid.
dP/dt = k4(D)
= k3(C)(O2) (from 2-16a)
= k2(B) k_2(C) (from 2-15a)
= [k_2(C) + ki(A) k_i(B)] k.2(C)
(from 2-14b)
Equation 2-12
Equation 2-17
Equation 2-18
Equation 2-19
= ki(A) k_1(B)
Equation 2-20
Equation 2-21 is the rate expression for the oxygen dependence
obtained from the proposed mechanism. Through appropriate
substitution from Equation 2-16a, this expression takes the form of
Equation 2-22. This equation can be reduced to pseudo zero-order in
oxygen pressure by assuming the concentration of D is constant
throughout the reaction by the steady state approximation.
-d02/dt = k3(C)(02) Equation 2-21
= k4(D) (from 2-16a) Equation 2-22
= k4' Equation 2-23
This mechanism also accounts for the product/02 and
product/water ratios previously observed by Bilgrien. An entire
reaction, consisting of two complete cycles, will produce two moles
of product while reducing one mole of 02 to two moles of water.
Hydrogen peroxide is most probably an intermediate in this
reduction, although never positively identified because it is
consumed as rapidly as it is formed.


50
A major driving force in this reaction is the large excess of
alcohol available. Ordinarily, the replacement of ligands such as
H2O or PPh3 by the poorly coordinating alcohol would be highly
unlikely. However, with the large excess of alcohol available, the
substitution occurs to a small extent. The low conversion rates
observed in this oxidation (about 2%) are also explained by the
small amount of substitution occurring in these systems.
Interestingly enough, exchange of deuterated methanol for water in a
mixed-metal (Ru2Rh) acetate trimer has been observed in *H NMR.^^
Few detailed NMR studies of these complexes have been
reported, 13,66,68 so this exchange may be more extensive than
previously expected. The complete failure of the pyridine adduct to
catalyze the oxidation, even after 24 hours, lends support to the
idea of slow substitution by the alcohol substrate.
None of these theories, however, explain the surprising
activity of the PPh3 adduct. Of all the trimers screened as
catalysts, the mixed-valence Ru30(02CCH2CH3)6(PPh3)3 complex
demonstrated the highest activity. Triphenyl phosphine is expected
to be a reasonably strong donor ligand toward Ru(II) (more than
H2O), so the substitution by alcohol should be significantly slower
than for the aquo adducts. However, triphenylphosphine is very
easily oxidized to the oxide, a very poor ligand. If all three
phosphine ligands are removed and subsequently oxidized to
triphenylphosphine oxidej the trimer would be essentially naked, and
alcohol coordination would occur rapidly. The presence of
triphenylphosphine oxide was never observed in the reaction mixture;


51
however, if this hypothesis is true, the quantities of the oxide
would be minute (10'^ to 10"^ moles) and difficult to detect.
The question of nuclearity of the catalyst has yet to be
addressed. The phosphine oxide hypothesis leads to the question of
the number of ruthenium atoms active in the oxidation. In the case
of the triphenylphosphine adduct, theoretically all three atoms
could be involved in the oxidation. A mechanism similar to the
proposal outlined in this chapter could be operating for each metal
center, by virtue of the extensive delocalization over the RU3-O
core. The synergism and interactions between the metals could
support such reactions, as evidenced by Meyer's oligomers.61 The
spectral data show the catalyst is essentially the same before and
after catalysis, and literature evidence is also available to
support the assumption that the complex remains intact. Considering
the volume of literature available on these complexes with no
reports of decomposition during reaction, it is reasonable to assume
that even under these stringent conditions the cluster retains its
nuclearity. The only physical evidence available is the differences
observed in the catalytic activity of the trimer compared to
monomeric ruthenium systems. The large difference indicates the
chemistry is somehow affected by three metals in close proximity, as
was expected from the outset.
The series of trinuclear ruthenium complexes has not failed in
its promise of producing highly intriguing chemistry. These
complexes have been shown to catalyze the selective oxidation of
alcohol to aldehydes and ketones by dioxygen. These oxidations are


52
presumed to occur via a standard Ru(II)-Ru(IV) cycle, but the cycle
involves the reductive elimination of a Ru(111)-alkoxy radical.
Based on these reactions, these complexes have upheld the potential
promised by their electrochemistry. The unique role of three metal
centers, intimately involved in a chemical transformation, has been
demonstrated, and these complexes manifest unusual catalytic
properties compared to monomeric species. Another enigma is their
catalytic activity, considering that Ru(III) centers are
traditionally inactive species in oxidations. The interactions
between the metals in the trimers can also be supposed to overcome
this trend, and in all probability, actually enhance the catalytic
activity of this system. However, the versatility of these trimers
has not been extensively tested.
Bilgrien found initially these complexes would not oxidize
olefins in alcohol solvent.^ However, changing the solvent to
acetonitrile, widely used in oxidation studies for its inertness,
drastically changed the chemistry. Under 40 psig of O2, cyclohexene
was oxidized to numerous products in the 12 hours. The volume of
products formed generally is indicative of free radical chemistry,
which is antipodally related to selectivity (vide infra). The
oxidation of a substrate inert to free radical process, norbornene,
was a complete failure. The lack of success in this area led to
branching out into other trimers containing different ligands.
Experimental
Reagents and equipment. All reagents used were reagent grade
or better and were, for the most part, readily available from


53
Aldrich Chemical Company. All alcohols were passed through a column
of neutral alumina, purity checked by GC, and stored over activated
molecular sieves. If necessary, the substrates were further
purified by standard techniques. Prior to use, the alcohols were
again passed through an alumina column.
GC analysis was performed on a Varian 3300 instrument utilizing
packed, 8-ft, stainless steel columns and both FID and TCD
detectors. Analyses and calibration curves were obtained using 15%
DEGS (diethylene glycol succinate) on Chromosorb W (80/100 mesh). A
Varian 4290 integrator automatically calculated peak areas and
retention times. GCMS was performed on a service basis by Dr. R. W.
King at the University of Florida. All IR spectra were recorded
either as Nujol mulls or KBr pellets on a Nicolet 5DXB spectrometer
and were background corrected. A Perkin-Elmer model 330 UV-visible
spectrometer equipped with a circulating thermal bath was used to
collect UV-vis spectra; all spectra were background corrected.
Elemental analyses were performed on a service basis by the
microanalytical laboratory at the University of Florida.
Synthesis. Trisaquohexakis(propionato) -/i3-oxotriruthenium-
(111,111,111) propionate, [Ru30(02CCH2CH3)5(H20)3](O2CCH2CH3), was
prepared as modified by Bilgrien.^ a mixture of 50 mL propionic
acid, 50 mL ethanol, and 1.2 g NaOH were warmed under N3 until the
NaOH dissolved. Two grams of "RuC13x(H20)3" were added and the
solution refluxed under nitrogen for four hours until deep green-
black. The solution was cooled to -78 C for 3-4 hours and filtered


54
to remove impurities including excess sodium propionate and NaCl.
The filtrate was evaporated on a rotary evaporator and vacuum dried
12 hours at 50 C to give the crude catalyst.
For chromatography on the Sephadex column as described by
Bilgrien, 1 g of crude trimer was dissolved in 100 mL of methanol
and chromatographed in approximately 25 mL fractions. The middle,
blue green fraction was collected, discarding the first and third
"bands," stripped of solvent, and rechromatographed in smaller (5 -
10 mL) fractions This treatment yielded a product whose spectra
matched the reported data. Again, the presence of trace nitrogen in
the elemental analyses of this complex is an enigma. Interestingly,
commercial RuCl3x(H20) from Aldrich also analyzes for trace
nitrogen, while "pure" RuCl3x(H20) from Johnson Matthey does not.
Using RuCl3 from Johnson Matthey eliminates the trace nitrogen in
the analyses as shown in the table below. It should be noted that
commercial RUCI3 is an ill-defined, heterogeneous mixture of mono-
and polymeric ruthenium complexes, including oxochloro,
hydroxochloro, and occasional nitrosyl complexes. The average
oxidation state is closer to Ru(IV) than Ru(III), and the main
constituent of RuCl3xH20 is considered to be a Ru(0H)Cl3
species.102 Although this does not definitively isolate the source
of the nitrogen in the analyses, this data leads to the conclusion
the nitrogen is most probably inherent in the starting material and
is carried through the reaction.


55
Table 2-3
Elemental
Analyses
%C
%H
%N
theoretical for
[Ru30(prop)6(H20)3](prop)
28.49
4.68
0.00
Chromatographed
(Aldrich)
27.89
4.29
0.50
crude trimer
(Aldrich)
29.63
4.65
0.57
crude trimer
(Johnson Matthey)
28.33
4.34
0.00
theoretical for
RuCl3x(H20)
0.00
2.29
0.00
Aldrich
RuC13*x(H20)
0.75
2.16
0.79
Johnson Matthey
RuCl3x(H20)
0.24
1.62
0.00
Tri s(pyr i di ne)hexakis( prop i onato) -/3-oxotr i ruthenium-
(111,111,111) hexafluorophosphate, [Ru30(02CCH2CH3)5(05^)31 (PFg),
was prepared using a modification of Wilkinson's procedure.33 Crude
[Ru30(prop)6(H20)3](prop), (.79 g) was dissolved in 5 mL methanol,
2.5 mL of pyridine was added, and the solution was stirred for 1
hour. A solution of 1 g NaPFg in 1 mL methanol was added to the
mixture and the resulting solution stored at -40 C for 48 hours.
Dark blue crystals were filtered from the cold solution, washed
three times with diethyl ether, and dried under vacuum at room


temperature for 12 hours. The IR and UV-visible of this complex
matched the reported values. Calculated for
[Ru30(02CCH2CH3)6(C5H5N)3](PF6): %C = 34.77, %H = 3.95, %N = 3.69;
Found %C = 33.86, %H = 3.86, %N = 3.51.


CHAPTER III
SYNTHESIS AND CHARACTERIZATION OF A NOVEL
TRINUCLEAR CARBOXYLATE COMPOUND
Background
Although trinuclear metal carboxylate complexes have been
widely studied (as mentioned in Chapter I), little variation in the
nature of the bridging carboxylate ligands has been attempted. The
literature reports only two examples where trinuclear carboxyl ates
have been synthesized using ligands other than alkyl carboxyl ates -
a ruthenium trimer having dichloroacetate ligands^ and more recent
reports detailing the synthesis of Fe, Cr, and V trifluoro-
acetates. ^ The iack 0f such reports, especially for the
ruthenium complexes, most probably stems from Wilkinson's failure to
prepare the trifluoroacetate derivative of the [Ru30(02CR)g(L)3]n
system.33
As reported in the last chapter, a variety of complexes having
the basic structure Ru30(02CR)6L3n are catalysts for the selective
oxidation of alcohols employing molecular oxygen as the oxidant.
However, Ru30(prop)6(H20)3+ did not catalyze the reaction of O2 with
norbornene, even after 48 hours of reaction time. This failure
prompted an investigation into routes to a selective olefin
epoxidation catalyst of this general type. Initial attempts
centered around creating a catalyst containing fluorinated
57


58
carboxylate ligands, thus increasing the acidity of the metal
centers and making the metals more likely to bind an olefin.
Binding the substrate directly to the metal would also provide a
method of selectively oxidizing the substrate to the desired
product.
Previous work in our laboratory^ had shown Rh2(0Ac)4 to have
a much lower acidity than Rh2(tfa)4, primarily due to the
differences in the electronic nature of the carboxylate ligand.
Doyle and othersl07-109 extended these observations to the area of
olefin binding. These workers showed that Rh2(tfa)4 will bind
olefins while Rh2(0Ac)4 will not, and studied the stability
constants for these reactions. Also, the fluorinated ligands should
be harder to oxidize, making the Ru(III) center a better oxidant. In
light of these discoveries, the exchange of fluorinated for non-
fluorinated carboxylate ligands in the Ru30(prop)g(H20)3+ system
would be an interesting extension of the previous studies.
Heptafluorobutyric acid (pfb acid) was chosen as the exchange
medium due to Doyle's reports^ that the perfluorobutyrate rhodium
dimer bound olefins three times better than the trifluoroacetate
complexes, as well as the fact Wilkinson was unsuccessful^3 in his
attempts to perform this exchange with trifluoroacetic acid. The
method used was an adaptation of a previously reported synthetic
route for the conversion of Rh2(0Ac)4 to the trifluoroacetate
analogue.106 a typical synthesis involved refluxing crude
Ru30(prop)5(H20)3+ in a 10:1 mixture of pfb acid and pfb anhydride,
stripping away the solvent, and dissolving the residue in diethyl


59
ether. The solution was then filtered and evaporated, leaving dark
black crystals which were dried i_n vacuo for 12 hours at 50 C.
Characterization
Several spectroscopic methods were used to identify the nature
of this complex. FTIR showed a decided difference between this
complex and the starting propionate trimer (Figures 3.1 and 3.2).
The water absorbance at 3400 cm'* is absent, and the uqq stretch has
shifted from 1567 to 1704 cm'*. Other significant differences occur
as well in the CH3 and CF3 regions. As a reference, the uqq stretch
for neat perfluorobutryic acid occurs at 1774 cm'l.
Proton NMR, shown in Figures 3.3 and 3.4, indicate the absence
of the distinctive peaks representative of the starting material.
The resonances observed are undoubtedly due to a slight impurity,
either in the complex or the solvent since they cannot be attributed
to either coordinated ether or residual starting material. Fluorine
NMR, on the other hand, gives the expected splitting pattern for a
trinuclear carboxylate containing both bound and ionic carboxylates
(Figure 3.5). The resonances are broadened slightly at the base,
indicative of the paramagnetism of the RU3O core. Again, as a
comparison, the free acid gives rise to three resonances at 82.5,
121, and 128 ppm. The relative insolubility of the complex,
combined with parameters inherent to the program used to transform
the data and the presence of fluorine-containing polymers in the
probe, make precise integration virtually impossible. The best
integrated ratios obtained were 16:3, 11:2, and 10:2, not


Figure 3.1 FTIR of [RU3O(pfb)g(Et2)3](pfb) in KBr.


2 2 03 7 4 3 13 <_ 7 000
or.
.D
-O
-n
O')
J")
¡ ; 1 ;
2083 0 1909 2 i 735 5 1561 8 1388 0 1214 3
WAVENUMBERS (CM-l)
en
1040 5
I
866 80
693 06
5 19 3


Figure 3.2 FTIR overlay of [RU3O(prop)g(HpO)3](prop) (
[Ru30(pfb)g(Et20)3](pfb) (--) as Nujol mulls.
) and


^TRANSMITTANCE
30.528 38 695 46.861 55.027 63.194 71.360
WAVENUMBERS (CM-l)
cr>
co


igure 3.3 lH NMR of crude [Ru30(prop)6(H20)3](prop) in CD30D. An
marks residual solvent peaks.


65
Figure 3.4 *H NMR of [Ru30(pfb)6(Et20)3](pfb) in CD3OD. An marks
residual solvent peaks.


Figure 3.5 19F NMR of [Ru30(pfb)6(Et20)3](pfb) in CD3OD.


en


68
significantly different from the 6:1 ratio that would be expected
for a complex having the formula [Ru30(pfb)g(Et2)3](pfb).
The UV-visible spectrum of the perfluorobutyrate complex shows
a similar shift with the absorbances at 610 and 670 nm (of the
original complex) moving to 575 and 760 nm. A new absorbance
appears at 950 nm as well (Figure 3.6). As further evidence for the
existence of the trinuclear species, a titration with pyridine shows
distinct changes in the spectrum upon the addition of three
equivalents of base (Figures 3.7 and 3.8). The shoulder of the
charge-transfer band at 375 shifts to 415 nm with a subsequent
decrease in intensity (e = 2500). The peak at 575 becomes more
distinct as well. These peaks and shoulders, along with the
epsilon values, are given in Table 3.1, as are other spectral data
of interest from IR and NMR spectra.
Molecular weight determinations using the Signer method^0 were
quite unsuccessful. Even after 3 weeks of equilibration, a constant
volume for the complex solution was not obtained, indicating the
complex was probably not stable in solution over extended periods of
time. Hovever, FAB mass spectroscopy, a useful technique for
obtaining molecular weights of materials having a high molecular
weight, gave a parent ion peak at 1675 mass units, corresponding to
a protonated Ru30(pfb)g(Et20)3 species. Other significant peaks in
the mass spectrum correspond to the successive loss of coordinated
ether and pfb ligands. (Figures 3.9 and 3.10). Elemental analysis
data further supports the proposed structure of
[Ru30(pfb)g(Et20)3](pfb). Analysis by Galbraith Laboratories gave


) cr >
69
2-
X (ran)
Figure 3.6 UV-Vis overlay of (A) [Ru30(prop)6(H20)3](prop) and (B)
[Ru30(pfb)6(Et20)3](pfb) in methanol.


Figure 3.7 UV-Vis titration of [RU3O(pfb)g(Et2)3](pfb) in
acetonitrile with pyridine. (A) [RU3O(pfb)5(Et20)3](pfb) (B) one
equivalent of pyridine (C) two equivalents of pyridine (D) three
equivalents of pyridine.


i/> cr >
X (nm)


<-0 10
72
Figure 3.8 Expansion of the UV-vis titrations with pyridine. (A)
[Ru30(pfb)g(Et20)3](pfb) (B) one equivalent of pyridine (C) two
equivalents of pyridine (C) three equivalents of pyridine.


73
Table 3.1
Spectral Data for
[Ru30(pfb)6(Et20)3](pfb)
NMR (referenced to internal CFCI3 at 0 ppm)
80.8
116.7
126.7
81.2 (triplet)
117.4 (quartet)
127.2 (singlet)
FUR (Nujol mull)
1704(s) 974(m)
1342(m) 936(m)
1224(s) 821(m)
1120(s)
UV-Vis
nm
e
375
(sh)
3383
575
(sh)
1574
760
(sh)
1312
950
1444


Figure 3.9 FAB positive ion mass spectrum of
[Ru30(pfb)6(Et20)3](pfb).




Figure 3.10 Expanded FAB mass spectrum of [RU3O(pfb)g(Et2O)3](pfb).


10 1203
"-4


78
%C = 21.20, %H = .73, and %F = 47.94, while values calculated for
[Ru30(pfb)6(Et20)1.5](pfb) give %C = 21.24, %H = .78 and %F =
48.46.
All attempts at growing crystals suitable for X-ray analysis
were unsuccessful, probably due to the highly unordered pfb ligands
as well as the ready substitution of water of the ether ligands.
Therefore, no definitive proof for the structure of this complex is
available. However, based on the evidence presented thus far, the
assumption of a complex having the formula [Ru30(pfb)5(Et2O)3](pfb)
and the "basic trinuclear acetate" structure is not unreasonable.
Based on this assumption, this complex was screened as a catalyst
for the oxidation of several organic substrates, as will be
discussed in the following chapters.
Experimental
All reagents used were reagent grade or better; the majority
were readily available from Aldrich Chemical Company. FTIR spectra
were collected on a Nicolet 5DXB FT spectrometer either as KBr
pellets or Nujol mulls. Proton and fluorine NMR spectra were
collected on either a Varian XL-200 (at 200MHz) or a Varian VXR-300
(at 300 MHz) FT spectrometers using TMS and CFC13, respectively, as
internal (or external where required) standards at 0 ppm.
Electronic spectra were performed on a PE 331 spectrophotometer and
were background corrected in all cases. Elemental analyses were
performed on a service basis by the University of Florida
microanalytical laboratory or by Galbraith Laboratories (Knoxville,
TN). Mass spectral determinations were carried out at the Middle


79
Atlantic Mass Spectrometry Laboratory at Johns Hopkins University, a
National Science Foundation Shared Instrument Facility.
Synthesis
Tri s(etherato)hexakis(heptafluorobutyrato)-M3-oxotriruthenium
(111,111,111) heptafluorobutyrate:
[Ru30(02CCF2CF2CF3)6(Et20)3](02CCF2CF2CF3).
Crude Ru30(prop)g(H20)3+ was prepared as described in the previous
chapter. The exchange was carried out by dissolving .5 g of the
crude propionate trimer in a mixture of 10 mL heptafluorobutyric
acid and 1 mL heptafl uorobutyric anhydride. The deep green solution
gradually changed to an olive-brown color upon refluxing under N2
for 90 minutes. The mixture was filtered warm and evaporated,
leaving a dark black, gummy solid. This solid was dissolved in 75
mL diethyl ether, filtered and evaporated; this process was repeated
twice. The final solid, obtained as a black powder, was dried i_n
vacuo at 50 C for 12 hours. Calculated for
[Ru30(02CCF2CF2CF3)6(Et20)i.5](02CCF2CF2CF3): % C = 21.24, % H =
0.78, % N = 0.0, % F = 48.46. Found (Galbraith Laboratories) % C =
21.20, % H = 0.73, % F = 47.94. Found (U. of F. Laboratories) % C =
20.91, % H = .12, % N = .11. This complex is highly hygroscopic, in
humid weather becoming quite gummy, and was stored in a desiccator.


CHAPTER IV
OLEFIN OXIDATIONS BY A NOVEL TRINUCLEAR
RUTHENIUM CARBOXYLATE COMPLEX
Introduction
The oxidation of hydrocarbons to a variety of oxygen-containing
organic chemicals is a highly useful industrial transformation as
outlined in Chapter I. Olefinic substrates were investigated
initially by the rubber industry as autoxidation substrates,4 which
led ultimately to the current process for the epoxidation of
ethylene and the Wacker process. Currently, the lack of a feasible
liquid-phase process for the epoxidation of propylene has generated
a great deal of interest in the selective oxidation of olefins.m
Ethylene and propylene are commercially inviting substrates for
study, due to the demand for their respective epoxides for plastics,
solvents, antifreeze, and other chemicals, The epoxidation of
ethylene over a silver-alumina catalyst is a unique, well-studied
heterogeneous system, and will not be further discussed.6,52,112
Propylene, on the other hand, cannot be oxidized to the epoxide
under similar conditions and is currently epoxidized using a
molybdenum-catalyzed process involving alkyl hydroperoxides.m
Asymmetric epoxidations have received a great deal of attention
in the literature due to Sharpless' discovery that chiral titanium-
80


81
isopropoxide complexes catalyze the epoxidation of allylic
alcohols.113-117 Epoxides formed in this fashion are generally
greater than 95% enantiomerically pure. This process has been
licensed by Aldrich Chemical Company and can be used to prepare
intermediates for a host of natural products of interest to the
pharmaceutical industry.H
Oxometal reagents containing most commonly the metals
molybdenum or vanadium, generally in combination with peroxides or
hydroperoxides, have been shown to actively epoxidize olefins. A
great deal of controversy concerning the mechanism of the
peroxomolybdenum-catalyzed epoxidation of olefins still exists.
Both Sharpless and Mimoun mechanisms are referred to for these and
similar oxidations.H*,*20 Ruthenium compounds, on the other
hand, have long been used to cleave double bond in organic
chemistry. Ruthenium tetroxide in combination with an oxygen
source, is a powerful reagent for cleaving carbon-carbon double
bonds to produce ketones or carboxylic acids. The wel1-studied,
widely-used RuCl2(PPh3)3 has also been shown to selectively oxidize
cyclohexene to the allylic ketone with O2 and styrene to styrene
oxide.119121 Selective epoxidation by ruthenium compounds is much
harder to achieve, however.
Reports of ruthenium complexes catalyzing a variety of olefin
oxidations using oxygen-atom transfer reagents such as iodosobenzene
instead of O2 abound. Commercially available ruthenium trichloride,
bipyridyl, and periodate in a biphasic solvent selectively oxidized
olefins to epoxides.I An electrogenerated compound thought to be


82
[Ru^(N4O)(0)]2+ (where N4O is bis[2-(2-pyridyl)ethyl][2-oxy-2-(2-
pyridyl)ethyl]amine) is reported to be the active species for the
epoxidation of olefins as well using oxygen-atom transfer
reagents. ^3
In contrast, only a few ruthenium compounds catalyze the
selective oxidation of olefins with molecular oxygen. The
epoxidation of norbornene was achieved with O2 using several Ru(II)
catalysts.^4 Jhis reaction was only about 10% selective to the
epoxide, generating oligomers of norbornene via a ring-opening
process as well as small amounts of norbornanone.
Metalloporphyrins have received a great deal of attention as
researchers try to mimic and understand the activities of biological
systems like cytochrome P-450.125 Iron and manganese prophyrins
have especially been used as probed for this system and will
epoxidize alkenes with oxygen atom transfer reagents.126-130
Recently a ruthenium porphyrin utilizing molecular oxygen as the
oxidant has been prepared by Groves and coworkers. 131132 /\
hindered trans-dioxo ruthenium(VI) porphyrin complex, at ambient
temperature and 1 atm O2, will react, albeit slowly, with a variety
of olefins to form epoxides. A similar compound,
[Ru(0)2(dmp)2](PFg)2, where dmp = 2,9-dimethyl-1,10-phenanthroline,
has been shown by Bailey and Drago to epoxidize olefins under
slightly more stringent conditions -- 55 C and 3 atm O2. 3 Another
dioxoruthenium(VI) complex containing acetate and pyridine ligands
will oxidize cyclohexene, hexene, and styrene slowly, presumably via
oxygen atom transfer.134


83
Meyer, in his extensive studies of ruthenium polypyridyl
complexes, has found a Ru(IV)-oxo complex will stoichiometrically or
electrocatalytically oxidize a variety of substrates including
olefins.95,135,136 ¡n general, high-valent ruthenium oxo complexes
have been studied intensively since the proposal that high valent
metal-oxo species are the active intermediates in metal 1oporphyrin
oxidations.
In light of the recent successes of ruthenium complexes as
olefin epoxidation catalysts, attempting to use the ruthenium
carboxylate trimers to oxidize olefins would be an intriguing
extension of the alcohol oxidation system encountered in Chapter II.
The use of such complexes as catalysts was indeed the primary
justification for the synthesis and characterization of the novel
perfluorobutyrate complex outlined in Chapter III.
Initial tests of the catalytic activity of the new ruthenium
perfluorobutyrate trimer were carried out using n-propanol as the
substrate. Under standard reaction conditions (see Chapter II), no
oxidation occurred after 12 hours of reaction time. The exchange of
alkyl for fluorinated carboxylate ligands was designed to increase
the ability of the metal centers to bind olefins, however.
Attempting to validate the assumption the perfluorobutyrate
complex would catalyze the epoxidation of olefins, a number of
olefins were tested with this complex (Table 4-1). Unfortunately,
the goal of selectively oxidizing olefins was not realized, since
the ruthenium perfluorobutyrate complex actively initiates the free
radical autoxidation of all of the substrates attempted.


84
Table 4.1
Olefin Substrates3
Substrate
Droduct
mmoles
turnovers^3
cyclohexene
2-cyclohexene-l-ol
1.37
67 /
3 hrs
2-cyclohexene-1-one
.764
38 /
3 hrs
cyclohexene oxide
.336
17 /
3 hrs
norbornene
norbornene oxide:
.055
4 /
24 hrs
(exo-2,3-epoxy-
norbornane)
.302
22 /
24 hrs
trans-#-
benzaldehyde
.325
3 /
48 hrs
methyl styrene
acetaldehyde
c
trans epoxide:
trace
(1R,2R-( + )-1-phenyl -
propylene oxide)
hexamethyl- hexamethylbenzene (major) c
Dewarbenzene hexamethylbenzene
oxide (minor) c
a) Reaction conditions are slightly different from reaction to
reaction. Specific details can be found under "Scope of Catalysis.
b) Turnovers = moles of product/moles of catalyst used in the
specified time.
c)Not quantified


85
Scope of Catalysis
The general setup and apparatus used are the same as described
in Chapter II under "Scope of Catalysis." For the olefin
oxidations, the temperature was held constant at 65 C; all
reactions were performed under 40 psi O2 initial pressure. Except
where otherwise noted, the solvent used was acetonitrile, and in all
cases a minimum 100-fold excess of substrate was used. All
reactions were monitored via GC or CG/MS. The procedure used varied
depending upon the state of the substrate. For norbornene, a solid,
the substrate was dissolved in 50 mL acetonitrile in the pressure
bottle and placed in the oil bath. Upon dissolution of the solid,
the catalyst was added and the apparatus assembled. No preliminary
preparation of norbornene was necessary; however, for some of the
liquid substrates used, pretreatment by washing through a neutral
alumina column to remove peroxides was required. In the case of
cyclohexene, a 20% by volume solution in acetonitrile (10 mL
substrate/40 mL solvent) was used; the other reactions were carried
out using 2 mL substrate in 50 mL solvent. No internal standard was
used in any of these reactions except 2-octanone in the cyclohexene
oxidation. In general, approximately 20 mg of catalyst was used,
corresponding to 10"^ moles. Products were determined by GC using a
DEGS column and FID detector. The amounts of products were
determined from a calibration curve relating moles of products to
relative area percents as described in Chapter II.


86
Results and Discussion
The oxidation of cyclohexene was attempted first, since it is
one of the easiest substrates to oxidize.^ After an induction
period of one hour, virtually all of the products typical of a free
radical autoxidation process were observed. The major products,
cyclohexene oxide, 2-cyclohexene-l-ol, and 2-cyclohexene-l-one, were
formed in roughly a 4:16:9 molar ratio after 3 hours. Significant
amounts of other products (approximately 10) were also observed but
not quantified. The addition of benzoquinone, a free radical trap,
inhibited the reaction for a finite period (between 6 and 9 hours,
depending on the amount of benzoquinone added), after which the
reaction resumed. Presumably, oxidation of the alkene resumes after
the oxidation of the quinone is complete.
At the other end of the spectrum in terms of oxidizability lies
norbornene. This substrate is widely used to prove the existence of
non-radical pathways in catalytic oxidation studies, since the kinds
of allylic hydrogen abstraction so prevalent in cyclohexene
oxidations are not possible in norbornene. Most of the norbornene
radicals produced are highly unstable, and would be expected to
decompose into alcohol and ketone products, as well as epoxide. An
induction period of 24 hours was observed in the oxidation of
norbornene also, after which primarily norbornene oxide was
produced.(Figure 4.1). This induction period is similar to that
seen in the oxidations by the high-valent oxo-rutheniurn complex
Ru(dmp).133 The Ru(pfb) complex was slightly less active than this
previously reported catalyst, producing 22 turnovers in 48 hours as


TURNOVERS
87
Figure 4.1 Activity curve for the oxidation of norbornene by
[Ru30(pfb)6(Et2O)3](pfb).


88
compared to 37 turnovers in the same period of time for the Ru(dmp)
catalyst.
Based on this result which seemed to indicate the selective
epoxidation of olefins by the Ru(pfb) catalyst, the oxidation of
trans-fl-methvlstyrene was attempted. This substrate had previously
been used by Groves to determine both the stereoselectivity and
possible mechanistic pathways for olefin epoxidation by his Ru(VI)
porphyrin.^2 jhe oxidation of trans-fl-methvlstyrene proceeded very
slowly. Only trace amounts of the trans epoxide were produced after
40 hours of reaction time. The major products of the reaction were
those due to the cleavage of the double bond--benzaldehyde and
acetaldehyde--indicative of a radical process.
Obviously, the perfluorinated ruthenium trimer acts as a potent
free radical initiator. Traylor, et. al, have shown that iron heme
complexes (cytochrome P450 analogs) also will catalyze radical-based
oxidations of alkenes.^6,127 jn orcjer to distinguish between a
free radical chain mechanism and a caged radical pair, the substrate
hexamethylDewarbenzene was chosen. Under Traylor's conditions,
autoxidation processes produce hexamethylbenzene as product, while a
caged radical pathway produces epoxide when m-chloroperbenzoic acid
is used as the oxidant. At 65 C and 3 atm O2 however, this
substrate is extremely reactive. No observable distinction in either
amount of type of product formed could be made between a blank
(using O2 and no catalyst present) and a typical catalytic run.
Another major disadvantage of this substrate is its sensitivity to
light.137-140 por these reasons, hexamethylDewarbenzene has limited


89
use as a substrate for mechanistic information in catalytic
oxidation studies under these stringent conditions. However, these
results do support the free radical nature of these oxidations.
These reactions show a marked solvent dependency as well (see
Table 4-2). Using norbornene as the substrate, a series of
reactions were run in a variety of solvents. In acetonitrile, 30
turnovers in a 48 hour period were achieved, while no reaction was
observed in benzonitrile, pyridine, or nitrobenzene. Approximately
5 turnovers in 48 hours were achieved in ethanol. These results can
be attributed to the increased solubility of O2 in acetonitrile
compared to the other solvents attempted. The decrease in activity
in ethanol is attributed primarily to its ability to act as a free
radical trap. These experiments, combined with the fact the addtion
of AIBN (azo-bis(isobutyronitrile)), a free radical initiator,
decreases the induction period and increases the number of turnovers
achieved (from 22 to 72 in 48 hours) all indicate a free radical
mechanism is involved in the oxidations (Figure 4.2). A caveat is
implicit in these results as wel1--norbornene is not as inert to
allylic hydrogen abstraction as has been previously assumed.
Observing changes in the catalyst during or after the reaction
would give some insight into the role the catalyst plays in these
reactions. The presence of the fluorinated ligands in the catalyst
enables the fate of the catalyst to be relatively easily monitored
via NMR. Variance in the structure of the compound, changes in
oxidation state, or complete degradation of the catalyst could be
discerned from changes in the resonances of the fluorinated atoms of


90
Table 4.2
Solvent Dependency
in Norbornene Oxidations
Turnovers3
sol vent
24 hrs
48 hrs
acetonitrile
3.5
29.8
pyridine
0.0
0.0
benzonitri1e
0.0
0.0
ethanol^
3.0
5.0
nitrobenzene
0.0
0.0
a) Turnovers =
moles of product/moles of
catalyst.
b) No oxidation
products from the solvent
were observed


TURNOVERS
91
Figure 4.2 Free radical experiments in norbornene oxidations.


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TRINUCLEAR RUTHENIUM CARBOXYLATE
COMPLEXES AS OXIDATION CATALYSTS
By
Leslie Shannon Davis
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
1988
ÃœTOF F
libraries

ACKNOWLEDGEMENTS
An undertaking of this size is rarely accomplished by a single
individual working entirely alone, and this is especially true of
this study. My advisor, Dr. Russell Drago, has been an inspiration,
mentor, and guide throughout this journey. I am indebted to him for
his "idears" and all his encouragement and advice during my sojourn
at Florida. Mrs. Ruth Drago, his kind, gracious wife, opened her
home and welcomed me as family, a gesture I certainly appreciated
and that eased my stay during the past four years. I would also
like to acknowledge Dr. Dave Richardson, Dr. Carl Stoufer, and the
remainder of my committee for their help and support.
The Drago Group as a whole has been an outstanding source of
hope, help and fun during our years together. For all the
camaraderie and aid, I thank each of them. I am especially
grateful, first of all, to my labmates, Alan Goldstein, Tom Cundari,
and Rich Riley, who endured all with happy faces, and were
consistent sources of good humor in the lab. I want to thank Ngai
Wong and Larry Chamusco for their computer and mechanical expertise,
without which much of this work would not have been possible. For
their assistance in various and sundry ways, I thank Jerry Grunwald,
Mark Barnes, and Cindy and Ed Getty. I am also grateful to former

group members Dr. Cindy Bailey and Dr. Iwona Bresinska for the
benefit of their wisdom. To Dr. Carl Bilgrien, the initiator of
this study, I owe a deep debt of gratitude. Very special thanks are
due Mrs. Maribel Lisk for her help, advice, and smiles.
Without the help of many others within the department many
"idears" could not be realized. I am grateful to Dr. Roy King for
his deep understanding of NMR and his willingness to share this
knowledge. The machine shop personnel, Chester, Vernon, and Daley,
were able to make anything I could describe, a talent I am most
grateful for. The creative talents of Rudy and Dick in the glass
shop in deciphering my sketches and still creating what I needed are
greatly appreciated. I also thank Chuck Christ and Paul Sharpe for
their assistance.
The experience of graduate school is not realized entirely in
the laboratory. I am grateful to Fran and Allan Goodman for
illustrating this lesson, and for many, many hours of plain old fun.
I also wish to thank Dr. Linda Lentz and Sasi Kalathoor for their
unswerving encouragement and support.
For first instilling in me an interest in chemistry, I thank
Mrs. Jackie Gay. My love of "things that turn pretty colors" is
entirely due to Dr. Alex Zozulin. I owe Alex an additional debt of
first showing me the joys of research.
My greatest debt is owed my family, without whose love and
support and encouragement I would not have accomplished this feat.
To them--Marcia, Larry and Debbie, and Drew--I dedicate this work.

TABLE OF CONTENTS
B.aqg
ACKNOWLEDGEMENTS ii
KEY TO ABBREVIATIONS vi
ABSTRACT vii
CHAPTERS
I. GENERAL INTRODUCTION 1
Catalytic Oxidations 1
Trinuclear Carboxylate Complexes 6
II. ALCOHOL OXIDATIONS BY TRINUCLEAR
RUTHENIUM CARBOXYLATE COMPLEXES 14
Introduction 14
Previous Work 18
Scope of Catalysis 19
Results and Discussion 24
Experimental 51
III. SYNTHESIS AND CHARACTERIZATION OF A
NOVEL TRINUCLEAR CARBOXYLATE COMPOUND 57
Background 57
Characterization 59
Experimental 78
i v

IV.OLEFIN OXIDATIONS BY A NOVEL TRINUCLEAR
RUTHENIUM CARBOXYLATE COMPLEX 80
Introduction 80
Scope of Catalysis 85
Results and Discussion 86
Experimental 96
V. ALKANE OXIDATIONS BY A NOVEL TRINUCLEAR
RUTHENIUM CARBOXYLATE COMPLEX 98
Introduction 98
Scope of Catalysis 105
Results and Discussion 105
Experimental 125
VI. CONCLUSIONS 126
REFERENCES 129
BIOGRAPHICAL SKETCH 139
v

KEY TO ABBREVIATIONS
Et20 = diethyl ether
OAc = CH3C02'
pfb = CF3CF2CF2CO2"
prop = CH3CH2C02'
PPh3 = triphenylphosphine
py = pyridine
tfa = trifluoroacetate
vi

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
TRINUCLEAR RUTHENIUM CARBOXYLATE
COMPLEXES AS OXIDATION CATALYSTS
By
Leslie Shannon Davis
December, 1988
Chairman: Russell S. Drago
Major Department: Chemistry
The family of trinuclear metal carboxylate complexes has been
known to chemists for over 100 years. General studies in terms of
their classical inorganic chemistry, such as ligand exchange or
electron transfer reactions and reactivity, are well documented in
the literature. However, little application of this knowledge has
been attempted.
The series of trinuclear ruthenium carboxylates is very
intriguing*in light of the extensive electrochemistry demonstrated
in these complexes. The ready accessibility to a variety of
oxidation states, combined with the facile exchange of ancillary
ligands, should make these complexes ideal subjects for catalytic
vi i
studies.

Previous work has shown a series of trinuclear ruthenium
carboxylates [Ru30(02CR)gL3]n to be active catalysts for the
oxidation of alcohols to carbonyl-containing products utilizing
dioxygen as the primary oxidant. Continuation of this study has
revealed a mechanism that utilizes the expected synergism between
the metals in this cluster to explain the unique features of this
oxidation.
A previously unknown member of this family, a complex
containing completely fluorinated ligands, has been synthesized and
characterized. Based on the accumulated evidence, this complex has
been formulated as [RU3O(02CCF2CF2CF3)g(Et20)3](O2CCF2CF2CF3). This
complex has been screened as a catalyst for a variety of organic
transformations and has excelled in initiating the free radical
autoxidation of several olefins again using dioxygen as the primary
oxidant. The oxidation studies were extended to alkane oxidations
as well, and were shown to occur by a slightly different mechanism
than that assumed to operate in the industrial, cobalt-catalyzed
oxidation of alkanes like cyclohexane.
VI 1 1

CHAPTER I
GENERAL INTRODUCTION
Catalytic Oxidations
The oxidation of organic substrates as a field of interest to
chemists has its origins in the beginnings of the history of
chemistry as a science. Lavoisier, the father of modern chemistry,
demolished the phlogiston theory when he explained the results of
Priestley and Sheele's air experiments.1 Air, he claimed, consisted
of two parts, one of which will support combustion (Priestley's
"fire gas") and one of which will not, and not "phlogiston." ^ He
named the "fire gas" oxygen (for acid former) and formulated the
theory of combustion in the late 1700s. In this origin the modern
field of oxidation chemistry has its roots.
Detailed studies of oxidation processes began in the 1800s.
The degradation of natural rubber was linked to oxygen absorption,
and a great deal of research was aimed at discovering anti-oxidants
for the rubber industry.3»^ The modern theories of autoxidation
processes (as the free radical oxidation of hydrocarbons by O2 is
known), were developed in the early 1900s. The effects of metal
ions on this process were studied during this period by Haber and
Weiss, who formulated the classical mechanism for metal-catalyzed
autoxidation in use today.^>4 (Figure 1.1)
1

2
In* +
R. +
R02* +
R« +
2 R02
In2
>
21 n •
RH —
>
R.
02
>
ro2*
RH —
>
ro2h +
ro2« —
>
ro2r
> RO4R > nonradical products
Figure 1.1 Basic autoxidation pathways.^>4

3
Autoxidation as a means for producing oxygenated compounds from
hydrocarbons is a highly desired process, although several serious
flaws exist in present processes. Controlling the selectivity of an
autoxidation process, a key element in terms of its usefulness, is
extremely difficult due to the radical nature of the chemistry. A
high activation energy, related to the spin-forbidden reaction
between dioxygen (a triplet state) and organic molecules (a singlet
state) is a barrier as well. Control of the process once initiated
is another disadvantage--the reaction is often hard to stop short of
CO2 and H2O.
Catalytic oxidations theoretically solve most of these problems
in that the addition of a catalyst should lower the energy barrier,
thus making the reaction easier to start. Product selectivity is
drastically affected by the presence of a catalyst as well. For
these reasons, the "Age of Petroleum" and the "Age of Catalysis" are
inescapably linked.^ Without catalysts to facilitate the conversion
of crude oil to useful products, a petroleum-based economy would not
be possible. Vice versa, without the widespread need for and use of
chemicals and products derived from oil, the study of catalysis
would be relegated to purely academic investigations. Sheldon and
Kochi estimate that today over 90% of the organic chemicals in use
are derived from petroleum, and the majority of petroleum and
petrochemical processes involve the use of catalysts.^ In terms of
the importance of catalytic oxidations, industrial organic
chemicals, including oxygenates from oxidative processes, made up
16.8% of the value of the total chemical industry in 1983.^ Seven

4
of the top fifty chemicals (by volume) were produced directly from
oxidation processes, and several others were produced from
oxidatively generated intermediates.®
Obviously, catalytic oxidations are industrially valuable.
Serious study and application of homogeneous, liquid-phase oxidation
began in the 1950s.® Before this time, the majority of industrial
processes used heterogeneous or supported catalysts. However,
homogeneous catalyst systems offer several advantages, especially to
the academician, over their heterogeneous counterparts. Generally
milder reaction conditions (i.e., lower temperature and pressure)
are used in homogeneous processes. Temperatures, mixing rates, and
catalyst concentration are more effectively controlled, and most
importantly, the reaction can feasibly be studied using standard
spectroscopic methods. New, improved surface science techniques
have made the study of heterogeneous catalysts easier, but the
relative perspicuity inherent in homogeneous systems still outweighs
these advances. The major disadvantages of homogeneous systems
industrially are the difficulty in separating products from the
reaction mixture and catalyst recovery. This last deterrent becomes
a major problem when dealing with catalysts containing noble metals
like rhodium or iridium due to their expense.
The advent of the Mid-Century Process (Equation 1-1) and the
discovery of the Wacker process (Equation 1-2) heralded a widespread
interest in homogeneous catalysis as well as organometallic
chemistry as fields of study.® Emphasis was placed on elucidation

5
p-CH3-(C5H4)-CH3 > C00H-(C6H4)-C00H Equation 1-1
cat = Co(0Ac)2 in HOAc, Br' promoter
200 °C
15 - 30 atm air
CH2=CH2 > CH3CHO Equation 1-2
cat = PdCI2/CuCl2
100 °C
10 atm air
of reaction mechanisms and the discovery of new compounds that would
catalyze transformations of organic compounds. Understanding the
chemistry of these processes eventually would lead to improvements
and enhancements of the catalysis. This understanding led to new
growth in both fields, and formed the basis for new expansions of
the chemistry and technology involved in catalytic oxidations.
The disciplines of homogeneous catalysis and organometallic
chemistry are closely related. So much so, in fact, that a major
justification for the study of organometal1ic complexes has been
their potential use as catalysts. Even though the overwhelming
majority of work in this area has dealt with mono-metallic systems,
the field of multi-metal 1ic catalysts is beginning to emerge as an
area rich in potential for catalytic research. Systems containing
more than one metal have several advantages over their mono-metal
counterparts. Enhanced stability as well as synergistic
interactions between the metals would give multi-metal 1ic systems a
range and versatility unknown in systems containing a single metal.
In theory, the judicious choice of the combination of metals should

6
lead to a "tunable" catalyst system - one where selectivity or
conversion is directly related to the metals involved.
The series of trinuclear metal carboxylate compounds is an
ideal choice for carrying out such studies. Their versatility,
combined with the wealth of knowledge available on the coordination
chemistry of these complexes, make them excellent choices as
subjects for the study of homogeneous catalysis.
Trinuclear Carboxylate Complexes
The family of trinuclear carboxylate complexes has been known
in the chemical literature for over 100 years. Only in the more
recent past have these complexes been extensively studied and
characterized. These studies are extremely interesting in light of
the versatility and uniqueness of multi-metallic systems in general.
The synergistic effects of the presence of two or more metals in
close proximity has been widely studied recently'7; several varied
applications of such systems are obvious in biochemistry and enzyme
studies (tryptophan 2,3-dioxygenase, for example, consists of both a-
Cu(II) and an iron porphyrin in the active site)® as well as
industrial processes involving transition metals on inorganic
supports (SMSI interactions between Ti and Ru and other platinum
metals in Fisher-Tropsch synthesis), and other commercial
applications (oxidations by Co(11) involving Mn(II) as a cocatalyst,
and the widely studied Ziegler-Natta polymerization system which
involves the combination of Zr or Ti and A1 as the catalytic
species).

7
All of the trinuclear metal carboxylate complexes or "basic
carboxylates"^ discovered to date have virtually the same basic
structure (Figure 1.2). (Although not trimers by the strictest
definition, these trinuclear ruthenium carboxylates will be referred
to as "trimers" for the sake of brevity.) The major differences in
these systems occur in metal-metal distances and the planarity of
the M3-O core. The equilateral triangle formed by the metals (as
the apices of the triangular M3-O core) is bridged above and below
the plane of the triangle by bidentate carboxylate ligands, and each
metal is connected via a central, three-coordinate oxygen atom.
Unlike their dimeric cousins,^ these complexes contain no formal
metal-metal bond. The remainder of the pseudo-octahedral
configuration around each metal atom is completed by the ancillary
ligand L. Obviously a great deal of versatility is inherent in
these complexes - not only can the metals used be widely varied, but
the carboxylate bridges and L also increase the permutations
possible. To date, almost all of the first-row transition metals^'
have been isolated as "basic trinuclear carboxylates" (V,15,16
Cr,17-20 Mn,21-24 pe,25-29 an(j qo30-31). others, like Ir,32 Ru,33-37
and Pd, Pt, and Rh38-41 have also been prepared. Titanium^ and
zirconium^ w-¡n also form a trinuclear complex slightly distorted
from the traditional basic carboxylate structure involving a central
hydroxy bridge between the metal centers. The carboxylate ligand
can vary from acetate to butyrate for all of these compounds;
partially chlorinated carboxylates as well as fluorinated ones have
also been used. The ligand L is most often a classical coordination

8
Figure 1.2 Generalized structure of basic trinuclear carboxylates
having the formula [M30(02CR)gL3]n.44

9
ligand such as pyridine, PPh3, or even H2O or diethyl ether. The
last forms of variation take place in terms of the metals involved.
As these complexes are most commonly isolated, the metals are found
in the +3 oxidation state, causing the cluster as a whole to have a
+1 charge. The other most commonly found form of these basic
trinuclear carboxylates is one with one metal in the +2 oxidation
state, rendering the complex neutral. In these systems, the
assignment of oxidation states is truly a formalism. The iron and
ruthenium complexes in particular can be classified as Robin and Day
Class III compounds, indicating complete delocalization of the metal
electrons. This classification is especially important in the
trinuclear ruthenium carboxylates, the subject of this work.
This versatility has made this family of complexes choice
candidates for a wide range of studies. The mixed-valence, neutral
species (primarily the iron complexes) have been extensively studied
in terms of intramolecular electron transfer reactions.45-48 j^g
manganese clusters, as well as similar dimeric systems, have been
studied in hopes of elucidating the role of Mn in photosynthesis as
well as for catalysis.49-51 j^e more classical inorganic chemistry
of these complexes has also been studied, including ligand-exchange
reactions, for example.
Due to the vast information available on these trimers, it is
not unreasonable to expect some studies in terms of their
usefulness. The synergism expected to occur between the metal
centers should manifest distinct differences from their monomeric
analogs. The variety of oxidation states available, combined with

10
the ready exchange of ligands, makes these complexes ideal choices
for catalysts, especially of homogeneous processes. Finally, the
carboxyl ate ligands have been shown to be relatively inert to
oxidation processes, as evidenced by the widespread use of metal
acetates (specifically Co(II) and Mn(II)) and acetic acid in
industrial oxidations.52 For these reasons, the family of
trinuclear metal carboxylate complexes would be expected to be good
catalysts for a variety of homogeneous processes.
Only in the last few years have widespread attempts been made
to utilize these complexes as catalysts. The cobalt acetate trimer
(Co30(0Ac)g(H20)3+ and others) has been proposed to be one of the
active catalytic species in the oxidation of p-xylene to
terephthalic acid (the Mid-Century/Amoco Process).55 It has also
been shown to oxidize toluene and other hydrocarbons under
relatively mild conditions.13,30,31,53 others have been used as
catalysts as well. The rhodium acetate complex uses t-butyl
hydroperoxide to selectively oxidize cyclohexene,5^ while the iron
acetate trimers have been proposed to catalyze a variety of organic
transformations.55-5^ By far the most widely studied (in catalytic
terms) of these trinuclear species is the ruthenium complex.
More literature is available on the ruthenium system in both
catalytic and chemical terms than most of the other trimers. These
complexes were first isolated and characterized55 in 1972; Spencer
and Wilkinson found these trimers to be unique in the family of
basic trinuclear acetates for several reasons. Both mixed-valence
and cationic trimers were readily isolable. These complexes

11
underwent a one-electron non-reversible wave electrochemically, and
readily underwent ligand exchange as well. The most unusual feature
of the ruthenium trimers was the reversible removal of the central
oxygen atom, a reaction unknown for the other trinuclear metal
systems. Later studies by Meyer et al., expanded^^'®^ the original
electrochemical studies, revealing for the Ru30(0Ac)g(pyz)3 (where
pyz = pyrazine) complexes a series of five one-electron reversible
waves. Four of these waves were attributed to the metal center,
corresponding to formal oxidation state changes from Ru(II1, 111,11)
to Ru(IV,IV,III). Linking these complexes into multinuclear
oligomers revealed systems that would undergo up to ten one- or two-
electron waves, justifiying the nickname "electron sponge" for these
ruthenium complexes. A generalized MO scheme,shown in Figure
1.3, shows several orbitals of the 7r system of the RU3O core in a
relatively small energy range. For the cationic, Ru(111,111,111)
complexes, all levels up to Ej" are filled; the A2' level is only
partially occupied. The orbitals containing the metal electrons are
virtually indistinguishable, the justification for the Class III
label of delocalization. This depiction also helps explain Meyer's
electrochemistry as well as other spectroscopic properties of these
complexes.
In all likelihood, the electrochemistry revealed for these
complexes prompted the widespread study of the ruthenium complexes
as catalysts. Olefin hydrogenations were first studied*^ by
Wilkinson; further studies were carried out both homogeneously and
heterogeneously supported on a carboxyl ate resin by Rempel and

"O X) TD
12
O
•> X
Ru
Figure 1.3 Qualitative molecular orbital description for [Ru30(02CR)6l-3]n
comgjexes, involving only the n system of the RU3O core. After Wilson et

13
others.64-66 Ziolkowski, et al,. have studied^,67-69 ^e Nineties
of cumene hydroperoxide decomposition and the exchange of DMF for
H2O using NMR techniques; they also have reported^ some catalytic
work in alkane oxidations. The ruthenium trimers have also been
involved in the Prins reaction,^ oxidative dehydrogenation of
saturated carbinols,^’^ and dimerization of acrylonitrile.^ ¡n
terms of oxidation catalysis, the ruthenium trimers, in the presence
of hydrogen peroxide, will oxidize substituted phenols to the
corresponding hydroquinone.^ ¡n a mixed solvent system containing
water, carbon tetrachloride, and acetonitrile, the Ru acetate trimer
with periodate will oxidatively cleave alkenes,^ similar to the
traditional chemistry observed for RUO4. They will also catalyze
the isomerization of allylic alcohols.^
A study of the use of the ruthenium trimers as oxidation
catalysts for a variety of organic transformations seemed
potentially interesting, based on their previous use as catalysts
and the large amount of electrochemical potential to be tapped in
these complexes. The use of molecular oxygen as the primary oxidant
has been an ongoing area of research, and the ruthenium trimers have
not previously been shown to be active as catalysts in such a
system. This work involved the continuation of the study of the
oxidation of alcohols by the ruthenium carboxylate trimers,^ as
well as an extension of these catalytic studies to the oxidations of
alkenes and alkanes by a new member of the trinuclear ruthenium
carboxylate family, [Ru30(pfb)g(Et20)3](pfb), which has been
synthesized and characterized.

CHAPTER II
ALCOHOL OXIDATIONS BY TRINUCLEAR
RUTHENIUM CARBOXYLATE COMPLEXES
Introduction
The oxidation of alcohols is a procedure long known and used in
organic chemistry for the production of aldehydes, ketones and
carboxylic acids. Mild reagents, such as Cr03/pyridine or Mn02,
react with alcohols to give primarily the carbonyl product (aldehyde
or ketone). Stronger oxidants, like RUO4, continue to oxidize
primary alcohols through an aldehyde intermediate to carboxylic
acids. Other high-valent ruthenium-oxo ions such as RuO^' or RUO4'
will oxidize primary alcohols to carboxylic acids, secondary
alcohols to ketones, and will oxidize unsaturated alcohols without
attacking the double bond.^ Autoxidation of alcohols tends to
produce ketone or acid along with hydrogen peroxide. Shell
commercialized a process for the production of hydrogen peroxide by
the autoxidation of 2-propanol (Eqn 2-1).^
(CH3)2C(H)0H > (CH3)2C=0 + H202 (98%) Equation 2-1
The reaction of alcohols with noble metals such as Pd or Pt to
give carbonyl products and a metal hydride species is well-
14

15
documented.a mechanism involving a ^-hydride elimination to
give metal hydrides is generally assumed. This mechanism is also
invoked for the Pd(II)-catalyzed oxidation of secondary alcohols to
ketones with oxygen at 25 °C.8
A great deal of literature has been published on ruthenium-
catalyzed oxidations of alcohols.^ Besides the general uses of
RUO4, low-valent Ru(II) complexes have been widely studied as
oxidation catalysts with both O2 and milder oxidants such as
iodosobenzene.8 The most widely studied compound of this type,
RuCl2(PPh3)3, has been used to oxidatively dehydrogenate alcohols
with oxygen.80 with iodosobenzene, RuCl2(PPh3)3 will selectively
oxidize primary alcohols to aldehydes.8 Sharpless et al. have
found N-oxides combined with RuCl2(PPh3)3 and other ruthenium
compounds will also oxidize alcohols to their respective carbonyl
products.81 In benzene solvent, this complex preferentially
oxidizes long-chain primary alcohols over the corresponding
secondary alcohol.82 Using O2 as the primary oxidant, RuCl2(PPh3)3
oxidizes allyl alcohols to a,/?-unsaturated carbonyl complexes in a
variety of relatively poorly coordinating solvents.83 in all of
these oxidations, several general trends arise. All of these
oxidations are shut down in the presence of strong donor solvents
like acetonitrile, indicating coordination of the substrate is
necessary for oxidation to occur. Replacement of a coordinated
nitrile by an alcohol is not highly likely. The mechanism
consistently invoked for these reactions involves the coordination
of alkoxides to a Ru(IV) species with subsequent ^-hydride

16
elimination to give carbonyl product and a Ru(II) hydride. The
hydridic species can be oxidized back to Ru(IV) by the available
oxidant, creating a catalytic cycle.
A wide variety of other ruthenium compounds have also been used
to catalyze the oxidation of alcohols. A ruthenium hydride,
RuH2(PPh3)4, catalyzes the condensation of alcohols to esters and
lactones at elevated temperatures.®^ Monomeric ruthenium complexes
containing fluorinated carboxylate ligands have also been shown to
dehydrogenate primary and secondary alcohols via a /3-hydride
elimination pathway.85,86 slightly more active compounds containing
diphosphine ligands have also been prepared and demonstrated to be
catalytic.®'7 The mixed-valence ruthenium carboxylate dimer
[Ru2(0CR)4C1] has been shown to dehydrogenate methanol to
formaldehyde under relatively mild conditions.®® Ruthenium
complexes as simple as commercially available ruthenium trichloride
have also been shown to be active for both the oxidation of
secondary alcohols and amines with oxygen.®® Ruthenium(111)
solutions will also oxidize allyl alcohol to acrolein.®®
Ruthenium complexes containing large, bulky ligands have also
been used for alcohol oxidations. Riley demonstrated a DMSO adduct
of Ru(II), RuX2(DMSO)3L, would catalyze the aerobic oxidation of
thioethers to sulfoxides. This reaction required a reducing
solvent, alcohol, to reduce the Ru(IV) species back to the active
Ru(II) complex, generating a carbonyl product.®*
Bidentate imines have also been used with Ru(II) to oxidize
coordinated alcohols in conjunction with 03.®^ In these systems, a

17
Ru(IV) to Ru(II) cycle is again proposed as the pathway of the
oxidation, and a disproportionation step enabling an escape from an
inactive Ru(III) species to active Ru(II) and Ru(IV) complexes is
also invoked. An unusual account of a Ru(III) complex containing
l,3-bis(2-pyridylimino)isoindoline (BPI) ligands is also involved in
the oxidation of alcohols.^ The use of Ru(III) is unusual in that
Ru(III) complexes, generally low-spin t2g^, tend to be
substitutionally inert.94,95 since the availability of open
coordination spaces is a requirement for a feasible, selective
homogeneous catalyst, Ru(III) complexes would not be expected to be
vary active catalytically. Gagne's system was active for alcohol
oxidations, producing around 60 turnovers (moles of product per
moles of catalyst used) in 24 hours when a strong, noncoordinating
base is present. Secondary alcohols formed ketones which were inert
to further oxidation. Primary alcohols were oxidized initially to
aldehydes (the primary product) which could react further giving
acetals and other products. Again a disproportionation of Ru(III)
to Ru(II) and Ru(IV) is proposed, with a Ru(IV)-coordinated alkoxide
species as the active intermediate. Hydridic ruthenium(II) may be
an intermediate in this reaction as well, arising from the /5-hydride
elimination of the Ru(IV)-alkoxide species.
T. J. Meyer has also contributed to this area with his well-
studied ruthenium polypyridyl complexes. Extensive kinetic and
mechanistic studies on alcohol oxidations by these high valent
ruthenium-oxo complexes have been carried out.96-98

18
In light of the extensive, ongoing research into oxidations by
ruthenium complexes in general, and the high potential for catalysis
demonstrated by the trinuclear ruthenium carboxylate complexes,
these particular complexes were chosen to screen as catalysts for
the oxidation of alcohols by molecular oxygen.
Previous Work
Bilgrien discovered that Ru30(prop)6(H20)3+ would catalyze the
selective oxidation of primary and secondary alcohols to the
corresponding carbonyl product using O2 as the primary oxidant.44 A
wide number of alcohols were active in this system, and in all cases
the only product formed was the aldehyde or ketone, with no traces
of carboxylic acid observed. Several different trimeric ruthenium
carboxylate complexes were found to be effective catalysts as well.
These oxidations exhibit a slight rate dependence upon acidity, as
demonstrated by the inhibition of the reaction upon the addition of
acids. On the other hand, bases had a curious effect on the
reaction. Sodium ethoxide enhanced the catalysis, 2,6-lutidine
inhibited the reaction, and NaOH caused precipitation of the
catalyst.
In mechanistic terms, Bilgrien found that for every mole of
carbonyl product produced, a mole of water was formed as well,
implying the four-electron reduction of oxygen to water.44 Hydrogen
peroxide, a likely intermediate in this process, was never detected
in the reaction mixture. These complexes would also oxidize
alcohols with H2O2 in place of O2 as the primary oxidant. A rough
calculation using the pressure drop of the pressure gauge for the O2

19
consumption showed that for each mole of O2 consumed, two moles of
product are produced. The rate of the reaction in terms of oxygen
pressure was found to be .25. The catalyst did not seem to
decompose during the reaction, as indicated by both IR and *H NMR
results. Bilgrien also found that the mixed-valence trimer was
readily oxidized by O2 in alcohol solution to the Ru(111,111,111)
complex, but the reduction of this species by alcohol did not occur.
Bilgrien's mechanism for the ruthenium trimer-catalyzed
oxidation of alcohols is shown in Figure 2.1.^ This scheme invokes
the Ru(III,III,II)-alcohol species as the active intermediate, which
undergoes intramolecular disproportionation to form a Ru(IV,II,II)
ruthenium species. The decomposition of this intermediate could
occur via a number of pathways, the most likely of which involves a
two-step reduction of the alcohol by the trimer. This reduction
would generate the Ru(II,II,II) species, without the central /j.3-
oxygen first observed by Spencer and Wilkinson,33 which would
readily be oxidized back to the Ru(III,III,II) intermediate in the
presence of O2.
Scope of Catalysis
Apparatus. All pressurized oxidations were carried out in
slightly modified Parr hydrogenation setups (Figure 2.2). This
apparatus has previously been described in detail by Zuzich and
Bilgrien.44,99 por these oxidations, stainless steel pressure
heads, constructed from Swagelok fittings and equipped with standard
sample valves, gas gauges, were also equipped with pressure relief
valves as a safety precaution. This apparatus was directly

20
Ruin
I
0
RÍ»I \u”l
/\
°2
RCH2OH
RuII
I
'\
Ru111 Ruin(RCH2OH)
Ru
II
Ru
II
Ru11(RCH2OH)
RuII
H • I
Ru1I RuiV(0CH2R)
RCHO + H20
rch2oh
Figure 2.1 Bilgrien's proposed mechanism for the oxidative
dehydrogenation of alcohols by [Ru30(02CR)g(L)3]n. Other ligands
have been omitted for clarity.^

21
Gas
Outlet
1/4" Silicone
Sceptum
1/4" Tu be to 1/8" NPT
Adapter
On/Off Ball Valve
1/8"NPT
Pressure
Gage
1/4"El bow.
1/4"Cross
d?—Wing Nut
Pressure
Bottle
Figure 2.2 Schematic diagram of a standard pressure head without
safety release valves.”

22
connected to an oxygen tank by copper tubing. The direction and
path of exit gases were controlled by a length of tygon tubing
attached to the exit valve which ran to the back of the hood. A
glass, 250ml, Parr hydrogenation bottle (the reactor vessel) was
attached to this apparatus by a #6 silicone gum rubber stopper and a
metal cage. The bottle is surrounded inside of this cage by an
aluminum sheath, designed to theoretically reduce the amount of
glass shards that would be produced in an explosion.
Sampling techniques have been previously described in detail by
Bilgrien.44 Briefly, a 1-mL gastight stringe, equipped with a Leur-
lok syringe valve and a 12-inch needle, is inserted through the
septum at the top of the pressure head with the valve closed. The
needle is guided through the ball valve into the reaction mixture.
The syringe valve is opened, a small aliquot withdrawn ( -.2 mL),
the valve closed, and the needle withdrawn. With practice, this
procedure can be accomplished quickly, safely, and with no
observable pressure loss. The aliquots are analyzed using GC,
GC/MS, and GC/IR.
Oxidation procedure. Typical oxidations involved 50 mL of
alcohol (as both solvent and substrate), 1 mL ketone standard, and
10~5 moles of catalyst. Reactions were carried out in a 65 °C
silicone oil bath monitored by an Omega 6100 temperature controller
and thermocouple under initial pressures of 40 psig of 02- Stirring
rates of the solutions were controlled by magnetic stirrers beneath
the oil bath; the oil bath was circulated by an overhead stirrer.

23
A slightly modified version of Bilgrien's technique44 was used
for the alcohol oxidations. The pressure bottle was charged with
all components of the reaction except catalyst (i.e., substrate,
standard, and a stirbar), covered with Parafilm, and placed in the
oil bath to equilibrate for 20 - 30 minutes. The catalyst was added
to the warm solution, the apparatus assembled and pressurized,
placed in the oil bath, and a sample withdrawn. This sample was
denoted as time zero and the start of the reaction. The reaction
was stirred as rapidly as possible to ensure the saturation of the
solution by 02-
Safety precautions. CAUTION! Combinations of warm organic
liquids and dioxygen are potential 1v explosive. Great care should
be taken, especially during setup and dismantling of oxidation
reactions, to avoid sparking the reaction mixture and causing a
violent explosion. General safety precautions to follow include (1)
let the reaction mixture cool to room temperature before
depressurizing; (2) be sure outlet gases are directed away from any
source of sparks; and (3) become aware of the explosion limits of
solvent, substrates, and the oxidant (whether air or O2) before
beginning an oxidation.
Calculations. Amounts of products formed were determined by GC
in all cases. Calibration curves relating moles of product to
relative peak areas were constructed for all products formed.
Standard procedure involved making up a series of solutions
containing a varying, known amount of product, and a constant amount
of standard in the solvent (alcohol) used. Repeated (at least five)

24
injections of each of these solutions gave a statistically valid
value for the area percent of the product peak. Knowing the number
of moles of product and standard in each solution gave a mole ratio
of product to standard, which can be plotted against the ratio of
the area percents of the product and standard. The area percents
are obtained electronically from the integration of the peak areas
of the GC chromatogram by an integrator. From the graph of mole
ratio to area percent ratio, the number of moles of product can be
obtained, if the amount of standard added is known.
Results and Discussion
Although Bilgrien's proposed mechanism adequately described the
experimental data,44 further work remained to be done to
substantiate this proposal. Areas to be addressed included the
differences in terms of activity between the mixed-valence and the
cationic trimers, and the fate of the catalyst during the reaction.
Other substrates should be tested, and more kinetic data should be
accumulated as mechanistic support. For these reasons, this
research project was continued.
Bilgrien found the activity of the ruthenium trimer catalysts
varied greatly depending upon the amount of purification of the
complex.44 Liquid chromatography on a four-foot Sephadex LH-20
size-exclusion gel gave the best results, with dramatic effects on
the catalysis, as shown in Figure 2.3. A general activity curve for
the ruthenium propionate trimer, the standard catalyst for most of
the remaining reactions, is shown in Figure 2.4.

TURNOVERS
25
Figure 2.3 Chromatographed vs. unchromatographed
[Ru30(prop)g(H20)3](prop) in isopropanol oxidations.

1600
1400
1200
1000
800
600
400
200
0
igure 2
sopropa
26
t—i—i—i—|—i—i—i—i—|—i—i—i—i—|—i—i—i—r
50 100 150 2Ó0
TIME (HRS.)
Activity curve for [Ru30(prop)g(H20)3]+ - catalyzed
oxidations.

27
Other alcohol substrates were screened to further test the
versatility of these trimers as catalysts (Table 2-1). Benzyl
alcohol, as expected, produced only benzaldehyde, and allyl alcohol
was exclusively oxidized to acrolein. Both of these substrates were
oxidized significantly slower than the isopropanol oxidation used as
the common standard for comparison in Bilgrien's work. While
isopropanol oxidations resulted in 147 turnovers in 12 hours, these
substrates only produced 40. Bilgrien also found the rate of
reaction slowed as the substrate varied from primary to secondary
alcohols. These substrates follow this general trend, as the rate
of reaction for both benzyl and allyl alcohol is slower than that
for either primary or secondary alcohols. This reduction in rate
for benzyl alcohol is most probably due to steric bulk and
subsequent hindrance in binding the substrate to the metal center.
The oxidation of allyl alcohol is slower due to a different mode of
substrate binding similar to that proposed by Taqui Khan.90 In the
RuCl3-catalyzed oxidation of allyl alcohol by O2, the substrate is
bound in two sites around the octahedral ruthenium center - once by
the double bond and once at the OH moiety. A /5-hydride transfer
creates a Ru(III) hydride-alcohol(+) species which is quickly
oxidized by O2 to give acrolein and the regenerated catalyst. No
hydride species is postulated for the trimers, but the relative
slowness of the reaction could be attributed to the inability of the
alcohol to bind at the olefinic site (vide infra), and a loss of
stability in the reduced ruthenium-alcohol intermediate.

28
Table 2-1
Alcohol Substrates3
substrate
T (°C)
product
to/12 hrs^
to/24 hrs
ethanolc
25
No reaction
65
acetaldehyde
198
313
isopropanolc
65
acetone
147
254
100
acetone
685
1015
n-propanolc
65
propanal
430
645
n-butanolc
65
butanal
d
d
cyclohexanolc
65
cyclohexanone
d
d
t-butanolc
65
No reaction
—
—
benzyl alcohol
65
benzaldehyde
40
e
allyl alcohol
65 '
acrolein
42
e
50%
65
acetone
75
e
isopropanolf
phenolS
65
No reaction
—
—
a) reaction conditions are as outlined under "Scope of Catalysis"
b) to = turnovers defined as moles of product/moles of catalyst used
c) from Bilgrien^
d) not quantified
e) reaction run for only 12 hours
f) auxiliary solvent used was acetonitrile as 50% by volume
g) solvent used was acetonitrile

29
Another congener of the ruthenium carboxylate trimer family,
[Ru30(prop)6(py)3](PFe), was synthesized. A bar graph comparing all
of the different trimers used in shown in Figure 2.5. The pyridine
adduct is completely unreactive in the oxidation of isopropanol,
indicating that coordination of the substrate in place of the
ancillary ligand L is necessary for catalysis to occur. When
graphed in terms of turnovers, the differences in the cationic and
the mixed-valence trimers becomes even more striking than Bilgrien
reported. The mixed-valence compounds are greater than three times
more active than their cationic counterparts. These differences are
made more enigmatic by the known reaction chemistry of these
complexes. The Ru(111,II1,11) trimers are readily oxidized to the
Ru(111,111,111) complexes by 03.^ However, the non-lability of
Ru(III) centers towards substitution is wel 1 -documented5^>9E>; ^he
Ru(111,111,111) trimer would be expected to exchange H2O (or L) for
alcohol ligands very slowly. Assuming coordination of substrate is
necessary for oxidation to occur, the Ru(111,111,111) system should
oxidize alcohols more slowly than the more labile Ru(III,III,II)
counterparts. The exchange of ligands in the Ru(III, III, II) trimer
would be faster, so that even if the oxidation of the trimer from
the Ru(III,III,II) to the Ru(111,111,111) did occur, a molecule of
alcohol would already be present in the coordination sphere of the
catalyst. This would explain some of the differences in the
catalytic activity of these complexes.
A comparison of the ruthenium trimers with other complexes
reported in the literature to oxidize alcohols would be informative

TURNOVERS/24. HOURS
30
CATALYSTS
Figure 2.5 Comparison of [Ru30(02CR)g(L)3]n catalysts for
isopropanol oxidations. (A) [Ru30(0Ac)g(H20)3]+ (B)
[Ru30(prop)g(H20)3]+ (C) [Ru30(prop)g(py)3j(PFg) (D)
[Ru30(prop)g(H20)3] (E) [Ru30(prop)g(PPh3)3].

31
in terms of gauging the activity of this system. A graphical
comparison is shown in Figure 2.6. As mentioned in the introduction
to this chapter, both RUCI3 and RuCl2(PPh3)3 have been shown to
oxidize alcohols using molecular oxygen as the primary oxidant.
Since the trimers also operate using oxygen, these systems should be
enlightening for comparing relative reactivities of the catalysts.
The trimers are approximately 10 times more active than the other
ruthenium complexes attempted, on the basis of turnovers in 12
hours. Even taking into account that the trimers contain 3 moles of
ruthenium per mole of catalyst, while the others only have one, the
trimers are still over three times more active.
The mechanism of these oxidations could safely be assumed to
not involve autoxidation pathways, due primarily to the selectivity
observed in the reaction. If free radicals were involved in these
oxidations, the further oxidation of aldehydes to carboxylic acids
would be expected. However, acid products are not observed under
our conditions, leading to the assumption the trimers are selective
oxidants. To further justify this claim, typical reactions designed
to prove or disprove free radical chain mechanisms were carried out
(Figure 2.7). The addition of benzoquinone, a free radical trap, to
a typical oxidation has no effect on the reaction. A free radical
initiator, AIBN (azobis(iso-butyronitrile)), was added to the
reaction in place of the catalyst and achieved approximately 10
turnovers in 1 hour and ceased to function. These experiments
emphasize the non-radical nature of these oxidations.

TURNOVERS
32
Figure 2.6 Comparison of various ruthenium catalysts in isopropanol
oxidations.

Figure 2.7 Free radical experiments in isopropanol oxidations.

34
Bilgrien noted that for every mole of product formed, one mole
of water was also produced.44 If the assumption the substrate must
coordinate in order for oxidation to occur is valid, the effects of
adding or removing water in the reaction should prove useful in
determining a mechanism (Figure 2.8). The addition of 5A activated
molecular sieves to the reaction greatly accelerated the rate, while
a reaction run in a 50/50 mixture of isopropanol and water showed a
drop in activity after about five hours. This curvature, indicative
of catalyst deactivation, is not observed in the activity curve
until after 180 hours of reaction time. Seemingly, the presence of
water slowly inactivates the catalyst.
The catalyst does not seem to decompose during catalysis,
according to NMR and IR.44 UV-Visible spectroscopy has been very
informative in determining the active species in solution. Since no
induction period is observed for these oxidations, either the
trinuclear carboxylate complexes is the active catalytic species, or
it is a precursor that converts rapidly to the active species in
solution. The lack of an induction period also indicates that the
two different versions of the trimer (Ru(III,111,111) and
Ru( III, II1, 11)) perhaps perform the oxidation by slightly different
pathways. Bilgrien noted that while the Ru(III,III,II) could be
oxidized to the Ru(111,111,111) in solution, alcohol was not a
strong enough reducing agent to perform the reverse reaction.44
However, a distinct color change is observed when an alcoholic
solution of the catalyst is heated to 43 °C under an inert
atmosphere. The changes were monitored via UV as shown in

TURNOVERS
Figure 2.8 Effects of H2O on isopropanol oxidations.

36
Figure 2.9. These changes correspond to the conversion of the
Ru(111,111,111) to the Ru(111,111,11) complex as reported by
Wilkinson.33 Jo effect this change, the solution had to be heated
for 18 hours. However, oxidations were performed at 65 °C, so this
conversion may well occur under typical oxidation conditions. If
this conversion were accompanied by production of ketone, the amount
produced (assuming either a stoichiometric conversion either per
mole of catalyst or per mole of ruthenium) was too small to be
detected by GC. This change is reversible; the addition of O2, 30%
H2O2, or air to the warm alcohol solution immediately oxidizes the
Ru( III,111, II) back to the Ru(111,111,111) with the corresponding
color change. The color change corresponding to this conversion is
not observed under our catalytic conditions; if present, the
Ru( 111, III, II) complex would be a transient species at best. These
UV-visible studies indicate a Ru(111,III,II) intermediate created
from a Ru(III,III,III) precursor would be a very slow but possible
process. They give little or no information about the pathway used
by a Ru(III,III,II) precursor, however.
The role of H2O2 in these oxidations was also pursued further.
Figure 2.10 shows the effects of adding H2O2 to typical alcohol
oxidations. Hydrogen peroxide is a potent oxidant by itself, as
demonstrated by the upper curves. However, the ruthenium trimer
catalyst will use peroxide in the absence of O2 to oxidize alcohols
to the same carbonyl products. If hydrogen peroxide is an
intermediate in the reduction of O2 as postulated by Bilgrien,^
these graphs indicate the peroxide would be consumed as a co-oxidant

Irt (T >
37
X (nm)
Figure 2.9 UV-Vis studies of [Ru30(prop)g(H20)3](prop) (A)catalyst
in ethanol under N2, 25 °C (B) catalyst in ethanol under N2, 43 °C
(C) solution (B) exposed to 02-

TURNOVERS
38
Figure 2.10 Role of H2O2 in isopropanol oxidations.

39
in the oxidation reactions. The amount of peroxide formed would, in
all probability, be small and would be consumed as rapidly as it
formed. A low steady-state concentration of peroxide would be one
explanation for the failure to identify peroxide in the reaction
mixture as wel1.
Determining kinetics in this system was based on the method of
initial rates from initial concentrations.-^ The rate expression
was assumed to take the form of Equation 2-2.
dx/dt = k0bs [cat]3 (P02)^ [substrate]0 Equation 2-2
In the alcohol oxidations, the substrate alcohol is present in much
higher volume and the conversion of alcohol to product is relatively
small. Therefore, the substrate concentration was assumed to be
relatively constant, giving the rate equation 2-3.
dx/dt = k'obs [cat]3 (P02)^ Equation 2-3
To obtain the order of the reaction with respect to each remaining
component, one variable was held constant while the other varied.
The rate of the reaction (dx/dt) was assumed to be the slope of the
straight line obtained from a plot of mole of product formed vs.
elapsed time. The appropriate mathematical manipulations gives a
ratio of the rate laws which will yield a value for the reaction
order.(Equations. 2-4,2-5,2-6).

40
(dx/dt)i = k'obsi [cat]ai (PO2)bi Equation 2-4
(dx/dt)1 k'obsi [cat]ai (P02)bi
- = --- Equation 2-5
(dx/dt)2 k'obs2 [cat]a2 (P02)b2
Holding one variable constant (for example, (P02)) gives Eqn 2-6.
log (dx/dt)1 - log (dx/dt)2
a = Equation 2-6
log [cat]i - log [cat]2
Several reactions were run where each of the variables was changed
in turn; this data is given in Table 2-2 and graphically in Figures
2.11 to 2.14.
Varying the concentration of the ruthenium catalyst (numbers 1,
4, and 5 in Table 2-2) lead to essentially first-order kinetics
(Figures 2-11 and 2-12). Varying the oxygen pressure was slightly
more demanding in that the total pressure had to be kept at 40 psig
for comparison purposes (numbers 1, 2, and 3 in Table 2-2). The
remainder of the pressure was made up of argon. The reaction order
was found to be approximately .2 in 02, very close to the value of
.25 reported by Bilgrien (Figures 2.13 and 2.14).44 For all
practical purposes, however, the reaction could be considered zero-
order in oxygen, considering the amounts of cumulative error in the
analysis, calibration curves, calculations, and the differences in
the values obtained mathematically and graphically.
A proposed mechanism for these oxidations is shown in Figure
2.15. This scheme differs significantly from that given by Bilgrien
in several areas. The mechanism, beginning with the more labile

41
Table 2-2
Kinetic Data for Alcohol Oxidations
by Ru30(prop)6(H20)3+
rate
law = dx/dt
= k'obs
[Ru]a (P02)
b
Exd.
[Ru]J
rx io'4i
Dsiqb
P02
n/vc
dx/dt^
x 10-3
ae
b*
1
8.89
44.0
.108
8.60
1.10
.185
2
1.01
16.3
.040
7.16
—
.259
3
9.24
27.5
.0675
7.62
—
.119
4
2.80
45.5
.112
2.42
1.19
5
218.2
45.0
.110
22.3
1.33
a) Concentration calculated in moles/liter using 50 mL as the total
volume.
b) Initial pressure of reaction in psig
c) N/v calculated from the ideal gas law (PV = nRT) assuming a
volume of 270 mL and 65 °C.
d) dx/dt has units of molarity/hour; calculated as explained in text
e) a = 1.20 ± .12
f) b = .187 ± .07

Figure 2.11 Kinetics: varying catalyst concentration.

43
1.00 -I
0.90 -
0.80 -
0.70 -
0.60 -
X
T>
0.50 -
o>
o
I 0.40 -
0.30 -
0.20 -
0.10 -
0.00
m = 1.04
“T i 1 1 1 1 1 1 1 1 r
2.5 3.0 3.5
- log Ru
t r
4.0
Figure 2.12 Kinetics: order in catalyst

44
Figure 2.13 Kinetics: varying initial O2 pressure.

45
Figure 2.14 Kinetics: order in O2.

46
L-RuIII
\
RUni
/
0
Ruin
L-Ru111
\
RuIII-L
/
Rum-L
/
0
RUH
I
L
Figure 2.15 Proposed mechanism for the [Ru30(02CR)6(L)3]n -
catalyzed alcohol oxidations. Carboxylate ligands have been omitted
for clarity.

47
Ru(III,III,II) species, involves first replacement of the ligand L
by a substrate molecule with concomitant loss of a proton. This
Ru(III,III,II)-alcohol species is postulated to be the active
intermediate in this cycle. Oxidation of this species by O2 (or
later in the cycle, H2O2) gives a species that can be formulated as
a Ru(III,III,III)-alkoxy radical or a Ru(II1,111,IV)-alkoxide
species, depending on the placement of the extra electron. In Robin
and Day Class III systems, this placement is more or less semantics.
Reductive elimination from this species gives carbonyl product and a
coordinatively unsaturated Ru(111,III,II) species. Solvation of
this species by another mole of alcohol regenerates the active
Ru (111,III,II) species.
The Ru(111,111,111) complex is slightly different in that to
reach the active species it must undergo a one-electron reduction
and replace L by a mole of alcohol. Ruthenium(111) species are, in
general, substitutionally inert,94,95 so replacement step would
be expected to be very slow. The UV-vis studies have demonstrated
the reduction process to be slow as well. These two reasons help
explain the differences between the two congeners.
The kinetics observed experimentally can be verified
mathematically using the mechanism proposed in Figure 2.15. Each
step in the mechanism can be written out and a rate expression
derived for each step (Equations 2-7 through 2-17) using standard
procedures and assuming the steady state approximation is valid for
Equations 2-14, 15, and 16.

48
Ru3»3,2 . 0H2 + ROH v Ru3’3’2 - OR + H2O Equation 2-7
A k_i B H
k2
ru3,3,2 . or
B H k_2
± Ru3’3’2 - OR + H+
C
Equation 2-8
Ru3’3’2 - OR + O2
C
k3
4 Ru3»3’3 - OR + 022'
D
Equation 2-9
Ru3’3»3 - OR *
D
Ru3’3’2 + ROH
E
k4
^5
Ru3’3»2 + R2C=0 Equation 2-10
E P
-> Ru3’3’2 - OR
B H
Equation 2-11
dP/dt = k4(D)
Equation 2-12
dA/dt = -ki(A) + k_1(B)
Equation 2-13
dB/dt = -k2(B) + k_2(C) - k.¡(B) + ki(A)
k-2(C) + ki(A) = k2(B) + k_1(B)
k2(B) = k_2(C) + ki(A) - k_1(B)
0 Equation 2-14
Equation 2-14a
Equation 2-14b
dC/dt = -k3(C)(02) + k2(B) - k_2(C)
k3(C)(02) = k2(B) - k_2(C)
Equation 2-15
Equation 2-15a
dD/dt = k3(C)(02) - k4(D) = 0
k3(C)(02) = k4(D)
Equation 2-16
Equation 2-16a
Rearrangement and subsequent substitution of Equations 2-14,
15, and 16 into the expression for dP/dt as shown below give the
rate expression for dP/dt in Equation 2-20. This expression can be
reduced to pseudo first-order in catalyst if k.¡(B) is assumed to be

49
small. Under our conditions, a large excess of alcohol, the reverse
reaction in Equation 2-7 should only occur to a small extent by Le
Chatelier's Principle, so the assumption seems to be valid.
dP/dt = k4(D)
= k3(C)(O2) (from 2-16a)
= k2(B) - k_2(C) (from 2-15a)
= [k_2(C) + ki(A) - k_i(B)] - k_2(C)
(from 2-14b)
Equation 2-12
Equation 2-17
Equation 2-18
Equation 2-19
= ki(A) - k_1(B)
Equation 2-20
Equation 2-21 is the rate expression for the oxygen dependence
obtained from the proposed mechanism. Through appropriate
substitution from Equation 2-16a, this expression takes the form of
Equation 2-22. This equation can be reduced to pseudo zero-order in
oxygen pressure by assuming the concentration of D is constant
throughout the reaction by the steady state approximation.
-d02/dt = k3(C)(02) Equation 2-21
= k4(D) (from 2-16a) Equation 2-22
= k4' Equation 2-23
This mechanism also accounts for the product/02 and
product/water ratios previously observed by Bilgrien. An entire
reaction, consisting of two complete cycles, will produce two moles
of product while reducing one mole of 02 to two moles of water.
Hydrogen peroxide is most probably an intermediate in this
reduction, although never positively identified because it is
consumed as rapidly as it is formed.

50
A major driving force in this reaction is the large excess of
alcohol available. Ordinarily, the replacement of ligands such as
H2O or PPh3 by the poorly coordinating alcohol would be highly
unlikely. However, with the large excess of alcohol available, the
substitution occurs to a small extent. The low conversion rates
observed in this oxidation (about 2%) are also explained by the
small amount of substitution occurring in these systems.
Interestingly enough, exchange of deuterated methanol for water in a
mixed-metal (Ru2Rh) acetate trimer has been observed in *H NMR.101
Few detailed NMR studies of these complexes have been
reported, 13,66,68 so this exchange may be more extensive than
previously expected. The complete failure of the pyridine adduct to
catalyze the oxidation, even after 24 hours, lends support to the
idea of slow substitution by the alcohol substrate.
None of these theories, however, explain the surprising
activity of the PPh3 adduct. Of all the trimers screened as
catalysts, the mixed-valence Ru30(02CCH2CH3)6(PPh3)3 complex
demonstrated the highest activity. Triphenyl phosphine is expected
to be a reasonably strong donor ligand toward Ru(II) (more than
H2O), so the substitution by alcohol should be significantly slower
than for the aquo adducts. However, triphenylphosphine is very
easily oxidized to the oxide, a very poor ligand. If all three
phosphine ligands are removed and subsequently oxidized to
triphenylphosphine oxidej the trimer would be essentially naked, and
alcohol coordination would occur rapidly. The presence of
triphenylphosphine oxide was never observed in the reaction mixture;

51
however, if this hypothesis is true, the quantities of the oxide
would be minute (10'^ to 10"^ moles) and difficult to detect.
The question of nuclearity of the catalyst has yet to be
addressed. The phosphine oxide hypothesis leads to the question of
the number of ruthenium atoms active in the oxidation. In the case
of the triphenylphosphine adduct, theoretically all three atoms
could be involved in the oxidation. A mechanism similar to the
proposal outlined in this chapter could be operating for each metal
center, by virtue of the extensive delocalization over the RU3-O
core. The synergism and interactions between the metals could
support such reactions, as evidenced by Meyer's oligomers.61 The
spectral data show the catalyst is essentially the same before and
after catalysis, and literature evidence is also available to
support the assumption that the complex remains intact. Considering
the volume of literature available on these complexes with no
reports of decomposition during reaction, it is reasonable to assume
that even under these stringent conditions the cluster retains its
nuclearity. The only physical evidence available is the differences
observed in the catalytic activity of the trimer compared to
monomeric ruthenium systems. The large difference indicates the
chemistry is somehow affected by three metals in close proximity, as
was expected from the outset.
The series of trinuclear ruthenium complexes has not failed in
its promise of producing highly intriguing chemistry. These
complexes have been shown to catalyze the selective oxidation of
alcohol to aldehydes and ketones by dioxygen. These oxidations are

52
presumed to occur via a standard Ru(II)-Ru(IV) cycle, but the cycle
involves the reductive elimination of a Ru(111)-alkoxy radical.
Based on these reactions, these complexes have upheld the potential
promised by their electrochemistry. The unique role of three metal
centers, intimately involved in a chemical transformation, has been
demonstrated, and these complexes manifest unusual catalytic
properties compared to monomeric species. Another enigma is their
catalytic activity, considering that Ru(III) centers are
traditionally inactive species in oxidations. The interactions
between the metals in the trimers can also be supposed to overcome
this trend, and in all probability, actually enhance the catalytic
activity of this system. However, the versatility of these trimers
has not been extensively tested.
Bilgrien found initially these complexes would not oxidize
olefins in alcohol solvent.^ However, changing the solvent to
acetonitrile, widely used in oxidation studies for its inertness,
drastically changed the chemistry. Under 40 psig of O2, cyclohexene
was oxidized to numerous products in the 12 hours. The volume of
products formed generally is indicative of free radical chemistry,
which is antipodally related to selectivity (vide infra). The
oxidation of a substrate inert to free radical process, norbornene,
was a complete failure. The lack of success in this area led to
branching out into other trimers containing different ligands.
Experimental
Reagents and equipment. All reagents used were reagent grade
or better and were, for the most part, readily available from

53
Aldrich Chemical Company. All alcohols were passed through a column
of neutral alumina, purity checked by GC, and stored over activated
molecular sieves. If necessary, the substrates were further
purified by standard techniques. Prior to use, the alcohols were
again passed through an alumina column.
GC analysis was performed on a Varian 3300 instrument utilizing
packed, 8-ft, stainless steel columns and both FID and TCD
detectors. Analyses and calibration curves were obtained using 15%
DEGS (diethylene glycol succinate) on Chromosorb W (80/100 mesh). A
Varian 4290 integrator automatically calculated peak areas and
retention times. GCMS was performed on a service basis by Dr. R. W.
King at the University of Florida. All IR spectra were recorded
either as Nujol mulls or KBr pellets on a Nicolet 5DXB spectrometer
and were background corrected. A Perkin-Elmer model 330 UV-visible
spectrometer equipped with a circulating thermal bath was used to
collect UV-vis spectra; all spectra were background corrected.
Elemental analyses were performed on a service basis by the
microanalytical laboratory at the University of Florida.
Synthesis. Trisaquohexakis(propionato) -/i3-oxotriruthenium-
(111,111,111) propionate, [Ru30(02CCH2CH3)5(H20)3](O2CCH2CH3), was
prepared as modified by Bilgrien.^ a mixture of 50 mL propionic
acid, 50 mL ethanol, and 1.2 g NaOH were warmed under N3 until the
NaOH dissolved. Two grams of "RuC13»x(H20)3" were added and the
solution refluxed under nitrogen for four hours until deep green-
black. The solution was cooled to -78 °C for 3-4 hours and filtered

54
to remove impurities including excess sodium propionate and NaCl.
The filtrate was evaporated on a rotary evaporator and vacuum dried
12 hours at 50 °C to give the crude catalyst.
For chromatography on the Sephadex column as described by
Bilgrien, 1 g of crude trimer was dissolved in 100 mL of methanol
and chromatographed in approximately 25 mL fractions. The middle,
blue green fraction was collected, discarding the first and third
"bands," stripped of solvent, and rechromatographed in smaller (5 -
10 mL) fractions . This treatment yielded a product whose spectra
matched the reported data. Again, the presence of trace nitrogen in
the elemental analyses of this complex is an enigma. Interestingly,
commercial RuCl3»x(H20) from Aldrich also analyzes for trace
nitrogen, while "pure" RuCl3»x(H20) from Johnson Matthey does not.
Using RuCl3 from Johnson Matthey eliminates the trace nitrogen in
the analyses as shown in the table below. It should be noted that
commercial RUCI3 is an ill-defined, heterogeneous mixture of mono-
and polymeric ruthenium complexes, including oxochloro,
hydroxochloro, and occasional nitrosyl complexes. The average
oxidation state is closer to Ru(IV) than Ru(III), and the main
constituent of RuCl3»xH20 is considered to be a Ru(0H)Cl3
species.102 Although this does not definitively isolate the source
of the nitrogen in the analyses, this data leads to the conclusion
the nitrogen is most probably inherent in the starting material and
is carried through the reaction.

55
Table 2-3
Elemental
Analyses
%C
%H
%N
theoretical for
[Ru30(prop)6(H20)3](prop)
28.49
4.68
0.00
Chromatographed
(Aldrich)
27.89
4.29
0.50
crude trimer
(Aldrich)
29.63
4.65
0.57
crude trimer
(Johnson Matthey)
28.33
4.34
0.00
theoretical for
RuCl3»x(H20)
0.00
2.29
0.00
Aldrich
RuC13*x(H20)
0.75
2.16
0.79
Johnson Matthey
RuCl3»x(H20)
0.24
1.62
0.00
Tri s(pyr i di ne)hexakis( prop i onato) -/í3-oxotr i ruthenium-
(111,111,111) hexafl uorophosphate, [Ru30(02CCH2CH3)5(05^)31 (PFg),
was prepared using a modification of Wilkinson's procedure.33 Crude
[Ru30(prop)g(H20)3](prop), (.79 g) was dissolved in 5 mL methanol,
2.5 mL of pyridine was added, and the solution was stirred for 1
hour. A solution of 1 g NaPFg in 1 mL methanol was added to the
mixture and the resulting solution stored at -40 °C for 48 hours.
Dark blue crystals were filtered from the cold solution, washed
three times with diethyl ether, and dried under vacuum at room

temperature for 12 hours. The IR and UV-visible of this complex
matched the reported values. Calculated for
[Ru30(02CCH2CH3)6(C5H5N)3](PF6): %C = 34.77, %H = 3.95, %N = 3.69;
Found %C = 33.86, %H = 3.86, %N = 3.51.

CHAPTER III
SYNTHESIS AND CHARACTERIZATION OF A NOVEL
TRINUCLEAR CARBOXYLATE COMPOUND
Background
Although trinuclear metal carboxylate complexes have been
widely studied (as mentioned in Chapter I), little variation in the
nature of the bridging carboxylate ligands has been attempted. The
literature reports only two examples where trinuclear carboxyl ates
have been synthesized using ligands other than alkyl carboxyl ates -
a ruthenium trimer having dichloroacetate ligands^ and more recent
reports detailing the synthesis of Fe, Cr, and V trifluoro-
acetates. ^ The iack 0f such reports, especially for the
ruthenium complexes, most probably stems from Wilkinson's failure to
prepare the trifluoroacetate derivative of the [Ru30(02CR)g(L)3]n
system.33
As reported in the last chapter, a variety of complexes having
the basic structure Ru30(02CR)6L3n are catalysts for the selective
oxidation of alcohols employing molecular oxygen as the oxidant.
However, Ru30(prop)6(H20)3+ did not catalyze the reaction of O2 with
norbornene, even after 48 hours of reaction time. This failure
prompted an investigation into routes to a selective olefin
epoxidation catalyst of this general type. Initial attempts
centered around creating a catalyst containing fluorinated
57

58
carboxylate ligands, thus increasing the acidity of the metal
centers and making the metals more likely to bind an olefin.
Binding the substrate directly to the metal would also provide a
method of selectively oxidizing the substrate to the desired
product.
Previous work in our laboratory!^ had shown Rh2(0Ac)4 to have
a much lower acidity than Rh2(tfa)4, primarily due to the
differences in the electronic nature of the carboxylate ligand.
Doyle and othersl^-109 extended these observations to the area of
olefin binding. These workers showed that Rh2(tfa)4 will bind
olefins while Rh2(0Ac)4 will not, and studied the stability
constants for these reactions. Also, the fluorinated ligands should
be harder to oxidize, making the Ru(III) center a better oxidant. In
light of these discoveries, the exchange of fluorinated for non-
fluorinated carboxylate ligands in the Ru30(prop)5(H20)3+ system
would be an interesting extension of the previous studies.
Heptafluorobutyric acid (pfb acid) was chosen as the exchange
medium due to Doyle's reports*^ that the perfluorobutyrate rhodium
dimer bound olefins three times better than the trifluoroacetate
complexes, as well as the fact Wilkinson was unsuccessful^3 in his
attempts to perform this exchange with trifluoroacetic acid. The
method used was an adaptation of a previously reported synthetic
route for the conversion of Rh2(0Ac)4 to the trifluoroacetate
analogue.!^ a typical synthesis involved refluxing crude
Ru30(prop)g(H20)3+ in a 10:1 mixture of pfb acid and pfb anhydride,
stripping away the solvent, and dissolving the residue in diethyl

59
ether. The solution was then filtered and evaporated, leaving dark
black crystals which were dried i_n vacuo for 12 hours at 50 °C.
Characterization
Several spectroscopic methods were used to identify the nature
of this complex. FTIR showed a decided difference between this
complex and the starting propionate trimer (Figures 3.1 and 3.2).
The water absorbance at 3400 cm'* is absent, and the uqq stretch has
shifted from 1567 to 1704 cm'*. Other significant differences occur
as well in the CH3 and CF3 regions. As a reference, the uqq stretch
for neat perfluorobutryic acid occurs at 1774 cnT*.
Proton NMR, shown in Figures 3.3 and 3.4, indicate the absence
of the distinctive peaks representative of the starting material.
The resonances observed are undoubtedly due to a slight impurity,
either in the complex or the solvent since they cannot be attributed
to either coordinated ether or residual starting material. Fluorine
NMR, on the other hand, gives the expected splitting pattern for a
trinuclear carboxylate containing both bound and ionic carboxylates
(Figure 3.5). The resonances are broadened slightly at the base,
indicative of the paramagnetism of the RU3O core. Again, as a
comparison, the free acid gives rise to three resonances at 82.5,
121, and 128 ppm. The relative insolubility of the complex,
combined with parameters inherent to the program used to transform
the data and the presence of fluorine-containing polymers in the
probe, make precise integration virtually impossible. The best
integrated ratios obtained were 16:3, 11:2, and 10:2, not

Figure 3.1 FTIR of [Ru30(pfb)g(Et20)3](pfb) in KBr.

000 l’’ ó I s p7 /to 7 7

Figure 3.2 FTIR overlay of [RU3O(prop)g(HpO)3](prop) (
[Ru30(pfb)g(Et20)3](pfb) (-•—•—•-) as Nujol mulls.
) and

^TRANSMITTANCE
30.528 38 695 46.861 55.027 63.194 71.360 79. 527
WAVENUMBERS (CM-l)
cr>
co

igure 3.3 lH NMR of crude [Ru30(prop)6(H20)3](prop) in CD30D. An
marks residual solvent peaks.

65
Figure 3.4 *H NMR of [Ru30(pfb)6(Et20)3](pfb) in CD3OD. An * marks
residual solvent peaks.

Figure 3.5 19F NMR of [Ru30(pfb)6(Et20)3](pfb) in CD3OD.

en
'v"4

68
significantly different from the 6:1 ratio that would be expected
for a complex having the formula [Ru30(pfb)g(Et2Ü)3](pfb).
The UV-visible spectrum of the perfluorobutyrate complex shows
a similar shift with the absorbances at 610 and 670 nm (of the
original complex) moving to 575 and 760 nm. A new absorbance
appears at 950 nm as well (Figure 3.6). As further evidence for the
existence of the trinuclear species, a titration with pyridine shows
distinct changes in the spectrum upon the addition of three
equivalents of base (Figures 3.7 and 3.8). The shoulder of the
charge-transfer band at 375 shifts to 415 nm with a subsequent
decrease in intensity (e = 2500). The peak at 575 becomes more
distinct as well. These peaks and shoulders, along with the
epsilon values, are given in Table 3.1, as are other spectral data
of interest from IR and NMR spectra.
Molecular weight determinations using the Signer method^0 were
quite unsuccessful. Even after 3 weeks of equilibration, a constant
volume for the complex solution was not obtained, indicating the
complex was probably not stable in solution over extended periods of
time. Hovever, FAB mass spectroscopy, a useful technique for
obtaining molecular weights of materials having a high molecular
weight, gave a parent ion peak at 1675 mass units, corresponding to
a protonated Ru30(pfb)g(Et20)3 species. Other significant peaks in
the mass spectrum correspond to the successive loss of coordinated
ether and pfb ligands. (Figures 3.9 and 3.10). Elemental analysis
data further supports the proposed structure of
[Ru30(pfb)g(Et20)3](pfb). Analysis by Galbraith Laboratories gave

ü) cr >
69
2-
X (ran)
Figure 3.6 UV-Vis overlay of (A) [Ru30(prop)6(H20)3](prop) and (B)
[Ru30(pfb)6(Et20)3](pfb) in methanol.

Figure 3.7 UV-Vis titration of [RU3O(pfb)g(Et2Ü)3](pfb) in
acetonitrile with pyridine. (A) [RU3O(pfb)5(Et20)3](pfb) (B) one
equivalent of pyridine (C) two equivalents of pyridine (D) three
equivalents of pyridine.

i/> cr >
X (nm)

<-0 10
72
Figure 3.8 Expansion of the UV-vis titrations with pyridine. (A)
[Ru30(pfb)g(Et20)3](pfb) (B) one equivalent of pyridine (C) two
equivalents of pyridine (C) three equivalents of pyridine.

73
Table 3.1
Spectral Data for
[Ru30(pfb)6(Et20)3](pfb)
NMR (referenced to internal CFCI3 at 0 ppm)
80.8
116.7
126.7
81.2 (triplet)
117.4 (quartet)
127.2 (singlet)
FUR (Nujol mull)
1704(s) 974(m)
1342(m) 936(m)
1224(s) 821(m)
1120(s)
UV-Vis
nm
e
375
(sh)
3383
575
(sh)
1574
760
(sh)
1312
950
1444

Figure 3.9 FAB positive ion mass spectrum of
[Ru30(pfb)6(Et20)3](pfb).


Figure 3.10 Expanded FAB mass spectrum of [RU3O(pfb)g(Et2O)3](pfb).

100
1203

78
%C = 21.20, %H = .73, and %F = 47.94, while values calculated for
[Ru30(pfb)6(Et20)1.5](pfb) give %C = 21.24, %H = .78 , and %F =
48.46.
All attempts at growing crystals suitable for X-ray analysis
were unsuccessful, probably due to the highly unordered pfb ligands
as well as the ready substitution of water of the ether ligands.
Therefore, no definitive proof for the structure of this complex is
available. However, based on the evidence presented thus far, the
assumption of a complex having the formula [Ru30(pfb)5(Et2O)3](pfb)
and the "basic trinuclear acetate" structure is not unreasonable.
Based on this assumption, this complex was screened as a catalyst
for the oxidation of several organic substrates, as will be
discussed in the following chapters.
Experimental
All reagents used were reagent grade or better; the majority
were readily available from Aldrich Chemical Company. FTIR spectra
were collected on a Nicolet 5DXB FT spectrometer either as KBr
pellets or Nujol mulls. Proton and fluorine NMR spectra were
collected on either a Varian XL-200 (at 200MHz) or a Varian VXR-300
(at 300 MHz) FT spectrometers using TMS and CFC13, respectively, as
internal (or external where required) standards at 0 ppm.
Electronic spectra were performed on a PE 331 spectrophotometer and
were background corrected in all cases. Elemental analyses were
performed on a service basis by the University of Florida
microanalytical laboratory or by Galbraith Laboratories (Knoxville,
TN). Mass spectral determinations were carried out at the Middle

79
Atlantic Mass Spectrometry Laboratory at Johns Hopkins University, a
National Science Foundation Shared Instrument Facility.
Synthesis
Tri s(etherato)hexakis(heptafluorobutyrato)-M3-oxotriruthenium
(111,111,111) heptafluorobutyrate:
[Ru30(02CCF2CF2CF3)6(Et20)3](02CCF2CF2CF3).
Crude Ru30(prop)g(H20)3+ was prepared as described in the previous
chapter. The exchange was carried out by dissolving .5 g of the
crude propionate trimer in a mixture of 10 mL heptafluorobutyric
acid and 1 mL heptafl uorobutyric anhydride. The deep green solution
gradually changed to an olive-brown color upon refluxing under N2
for 90 minutes. The mixture was filtered warm and evaporated,
leaving a dark black, gummy solid. This solid was dissolved in 75
mL diethyl ether, filtered and evaporated; this process was repeated
twice. The final solid, obtained as a black powder, was dried i_n
vacuo at 50 °C for 12 hours. Calculated for
[Ru30(02CCF2CF2CF3)6(Et20)i.5](02CCF2CF2CF3): % C = 21.24, % H =
0.78, % N = 0.0, % F = 48.46. Found (Galbraith Laboratories) % C =
21.20, % H = 0.73, % F = 47.94. Found (U. of F. Laboratories) % C =
20.91, % H = .12, % N = .11. This complex is highly hygroscopic, in
humid weather becoming quite gummy, and was stored in a desiccator.

CHAPTER IV
OLEFIN OXIDATIONS BY A NOVEL TRINUCLEAR
RUTHENIUM CARBOXYLATE COMPLEX
Introduction
The oxidation of hydrocarbons to a variety of oxygen-containing
organic chemicals is a highly useful industrial transformation as
outlined in Chapter I. Olefinic substrates were investigated
initially by the rubber industry as autoxidation substrates,4 which
led ultimately to the current process for the epoxidation of
ethylene and the Wacker process. Currently, the lack of a feasible
liquid-phase process for the epoxidation of propylene has generated
a great deal of interest in the selective oxidation of olefins.m
Ethylene and propylene are commercially inviting substrates for
study, due to the demand for their respective epoxides for plastics,
solvents, antifreeze, and other chemicals, The epoxidation of
ethylene over a silver-alumina catalyst is a unique, well-studied
heterogeneous system, and will not be further discussed.6,52,112
Propylene, on the other hand, cannot be oxidized to the epoxide
under similar conditions and is currently epoxidized using a
molybdenum-catalyzed process involving alkyl hydroperoxides.m
Asymmetric epoxidations have received a great deal of attention
in the literature due to Sharpless' discovery that chiral titanium-
80

81
isopropoxide complexes catalyze the epoxidation of allylic
alcohols.113-117 Epoxides formed in this fashion are generally
greater than 95% enantiomerically pure. This process has been
licensed by Aldrich Chemical Company and can be used to prepare
intermediates for a host of natural products of interest to the
pharmaceutical industry.H®
Oxometal reagents containing most commonly the metals
molybdenum or vanadium, generally in combination with peroxides or
hydroperoxides, have been shown to actively epoxidize olefins.® A
great deal of controversy concerning the mechanism of the
peroxomolybdenum-catalyzed epoxidation of olefins still exists.
Both Sharpless and Mimoun mechanisms are referred to for these and
similar oxidations.®’®’H*,*20 Ruthenium compounds, on the other
hand, have long been used to cleave double bond in organic
chemistry. Ruthenium tetroxide in combination with an oxygen
source, is a powerful reagent for cleaving carbon-carbon double
bonds to produce ketones or carboxylic acids.® The wel1-studied,
widely-used RuCl2(PPh3)3 has also been shown to selectively oxidize
cyclohexene to the allylic ketone with O2 and styrene to styrene
oxide.119,121 Selective epoxidation by ruthenium compounds is much
harder to achieve, however.
Reports of ruthenium complexes catalyzing a variety of olefin
oxidations using oxygen-atom transfer reagents such as iodosobenzene
instead of O2 abound. Commercially available ruthenium trichloride,
bipyridyl, and periodate in a biphasic solvent selectively oxidized
olefins to epoxides.I®® An electrogenerated compound thought to be

82
[Ru^(N4O)(0)]2+ (where N4O is bis[2-(2-pyridyl)ethyl][2-oxy-2-(2-
pyridyl)ethyl]amine) is reported to be the active species for the
epoxidation of olefins as well using oxygen-atom transfer
reagents. ^3
In contrast, only a few ruthenium compounds catalyze the
selective oxidation of olefins with molecular oxygen. The
epoxidation of norbornene was achieved with O2 using several Ru(II)
catalysts.^4 Jhis reaction was only about 10% selective to the
epoxide, generating oligomers of norbornene via a ring-opening
process as well as small amounts of norbornanone.
Metalloporphyrins have received a great deal of attention as
researchers try to mimic and understand the activities of biological
systems like cytochrome P-450.125 Iron and manganese prophyrins
have especially been used as probed for this system and will
epoxidize alkenes with oxygen atom transfer reagents.126-130
Recently a ruthenium porphyrin utilizing molecular oxygen as the
oxidant has been prepared by Groves and coworkers. 1^1»*32 /\
hindered trans-dioxo ruthenium(VI) porphyrin complex, at ambient
temperature and 1 atm O2, will react, albeit slowly, with a variety
of olefins to form epoxides. A similar compound,
[Ru(0)2(dmp)2](PFg)2, where dmp = 2,9-dimethyl-1,10-phenanthroline,
has been shown by Bailey and Drago to epoxidize olefins under
slightly more stringent conditions -- 55 °C and 3 atm O2. 3 Another
dioxoruthenium(VI) complex containing acetate and pyridine ligands
will oxidize cyclohexene, hexene, and styrene slowly, presumably via
oxygen atom transfer.^4

83
Meyer, in his extensive studies of ruthenium polypyridyl
complexes, has found a Ru(IV)-oxo complex will stoichiometrically or
electrocatalytically oxidize a variety of substrates including
olefins.95,135,136 ¡n general, high-valent ruthenium oxo complexes
have been studied intensively since the proposal that high valent
metal-oxo species are the active intermediates in metal 1oporphyrin
oxidations.
In light of the recent successes of ruthenium complexes as
olefin epoxidation catalysts, attempting to use the ruthenium
carboxylate trimers to oxidize olefins would be an intriguing
extension of the alcohol oxidation system encountered in Chapter II.
The use of such complexes as catalysts was indeed the primary
justification for the synthesis and characterization of the novel
perfluorobutyrate complex outlined in Chapter III.
Initial tests of the catalytic activity of the new ruthenium
perfluorobutyrate trimer were carried out using n-propanol as the
substrate. Under standard reaction conditions (see Chapter II), no
oxidation occurred after 12 hours of reaction time. The exchange of
alkyl for fluorinated carboxylate ligands was designed to increase
the ability of the metal centers to bind olefins, however.
Attempting to validate the assumption the perfluorobutyrate
complex would catalyze the epoxidation of olefins, a number of
olefins were tested with this complex (Table 4-1). Unfortunately,
the goal of selectively oxidizing olefins was not realized, since
the ruthenium perfluorobutyrate complex actively initiates the free
radical autoxidation of all of the substrates attempted.

84
Table 4.1
Olefin Substrates3
Substrate
Droduct
mmoles
turnovers^3
cyclohexene
2-cyclohexene-l-ol
1.37
67 /
3 hrs
2-cyclohexene-1-one
.764
38 /
3 hrs
cyclohexene oxide
.336
17 /
3 hrs
norbornene
norbornene oxide:
.055
4 /
24 hrs
(exo-2,3-epoxy-
norbornane)
.302
22 /
24 hrs
trans-#-
benzaldehyde
.325
3 /
48 hrs
methyl styrene
acetaldehyde
c
trans eooxide:
trace
(1R,2R-( + )-1-phenyl -
propylene oxide)
hexamethyl- hexamethylbenzene (major) c
Dewarbenzene hexamethylbenzene
oxide (minor) c
a) Reaction conditions are slightly different from reaction to
reaction. Specific details can be found under "Scope of Catalysis.
b) Turnovers = moles of product/moles of catalyst used in the
specified time.
c)Not quantified

85
Scope of Catalysis
The general setup and apparatus used are the same as described
in Chapter II under "Scope of Catalysis." For the olefin
oxidations, the temperature was held constant at 65 °C; all
reactions were performed under 40 psi O2 initial pressure. Except
where otherwise noted, the solvent used was acetonitrile, and in all
cases a minimum 100-fold excess of substrate was used. All
reactions were monitored via GC or CG/MS. The procedure used varied
depending upon the state of the substrate. For norbornene, a solid,
the substrate was dissolved in 50 mL acetonitrile in the pressure
bottle and placed in the oil bath. Upon dissolution of the solid,
the catalyst was added and the apparatus assembled. No preliminary
preparation of norbornene was necessary; however, for some of the
liquid substrates used, pretreatment by washing through a neutral
alumina column to remove peroxides was required. In the case of
cyclohexene, a 20% by volume solution in acetonitrile (10 mL
substrate/40 mL solvent) was used; the other reactions were carried
out using 2 mL substrate in 50 mL solvent. No internal standard was
used in any of these reactions except 2-octanone in the cyclohexene
oxidation. In general, approximately 20 mg of catalyst was used,
corresponding to 10"^ moles. Products were determined by GC using a
DEGS column and FID detector. The amounts of products were
determined from a calibration curve relating moles of products to
relative area percents as described in Chapter II.

86
Results and Discussion
The oxidation of cyclohexene was attempted first, since it is
one of the easiest substrates to oxidize.^ After an induction
period of one hour, virtually all of the products typical of a free
radical autoxidation process were observed. The major products,
cyclohexene oxide, 2-cyclohexene-l-ol, and 2-cyclohexene-l-one, were
formed in roughly a 4:16:9 molar ratio after 3 hours. Significant
amounts of other products (approximately 10) were also observed but
not quantified. The addition of benzoquinone, a free radical trap,
inhibited the reaction for a finite period (between 6 and 9 hours,
depending on the amount of benzoquinone added), after which the
reaction resumed. Presumably, oxidation of the alkene resumes after
the oxidation of the quinone is complete.
At the other end of the spectrum in terms of oxidizability lies
norbornene. This substrate is widely used to prove the existence of
non-radical pathways in catalytic oxidation studies, since the kinds
of allylic hydrogen abstraction so prevalent in cyclohexene
oxidations are not possible in norbornene. Most of the norbornene
radicals produced are highly unstable, and would be expected to
decompose into alcohol and ketone products, as well as epoxide. An
induction period of 24 hours was observed in the oxidation of
norbornene also, after which primarily norbornene oxide was
produced.(Figure 4.1). This induction period is similar to that
seen in the oxidations by the high-valent oxo-rutheniurn complex
Ru(dmp).133 The Ru(pfb) complex was slightly less active than this
previously reported catalyst, producing 22 turnovers in 48 hours as

TURNOVERS
87
Figure 4.1 Activity curve for the oxidation of norbornene by
[Ru30(pfb)6(Et2O)3](pfb).

88
compared to 37 turnovers in the same period of time for the Ru(dmp)
catalyst.
Based on this result which seemed to indicate the selective
epoxidation of olefins by the Ru(pfb) catalyst, the oxidation of
trans-fl-methvlstyrene was attempted. This substrate had previously
been used by Groves to determine both the stereoselectivity and
possible mechanistic pathways for olefin epoxidation by his Ru(VI)
porphyrin.^2 jhe oxidation of trans-fl-methvlstyrene proceeded very
slowly. Only trace amounts of the trans epoxide were produced after
40 hours of reaction time. The major products of the reaction were
those due to the cleavage of the double bond--benzaldehyde and
acetaldehyde--indicative of a radical process.
Obviously, the perfluorinated ruthenium trimer acts as a potent
free radical initiator. Traylor, et. al, have shown that iron heme
complexes (cytochrome P450 analogs) also will catalyze radical-based
oxidations of alkenes.^6,127 jn orcjer to distinguish between a
free radical chain mechanism and a caged radical pair, the substrate
hexamethylDewarbenzene was chosen. Under Traylor's conditions,
autoxidation processes produce hexamethylbenzene as product, while a
caged radical pathway produces epoxide when m-chloroperbenzoic acid
is used as the oxidant. At 65 °C and 3 atm O2» however, this
substrate is extremely reactive. No observable distinction in either
amount of type of product formed could be made between a blank
(using O2 and no catalyst present) and a typical catalytic run.
Another major disadvantage of this substrate is its sensitivity to
light.137-140 por these reasons, hexamethylDewarbenzene has limited

89
use as a substrate for mechanistic information in catalytic
oxidation studies under these stringent conditions. However, these
results do support the free radical nature of these oxidations.
These reactions show a marked solvent dependency as well (see
Table 4-2). Using norbornene as the substrate, a series of
reactions were run in a variety of solvents. In acetonitrile, 30
turnovers in a 48 hour period were achieved, while no reaction was
observed in benzonitrile, pyridine, or nitrobenzene. Approximately
5 turnovers in 48 hours were achieved in ethanol. These results can
be attributed to the increased solubility of O2 in acetonitrile
compared to the other solvents attempted. The decrease in activity
in ethanol is attributed primarily to its ability to act as a free
radical trap. These experiments, combined with the fact the addtion
of AIBN (azo-bis(isobutyronitrile)), a free radical initiator,
decreases the induction period and increases the number of turnovers
achieved (from 22 to 72 in 48 hours) all indicate a free radical
mechanism is involved in the oxidations (Figure 4.2). A caveat is
implicit in these results as wel1--norbornene is not as inert to
allylic hydrogen abstraction as has been previously assumed.
Observing changes in the catalyst during or after the reaction
would give some insight into the role the catalyst plays in these
reactions. The presence of the fluorinated ligands in the catalyst
enables the fate of the catalyst to be relatively easily monitored
via NMR. Variance in the structure of the compound, changes in
oxidation state, or complete degradation of the catalyst could be
discerned from changes in the resonances of the fluorinated atoms of

90
Table 4.2
Solvent Dependency
in Norbornene Oxidations
Turnovers3
sol vent
24 hrs
48 hrs
acetonitrile
3.5
29.8
pyridine
0.0
0.0
benzonitri1e
0.0
0.0
ethanol^
3.0
5.0
nitrobenzene
0.0
0.0
a) Turnovers =
moles of product/moles of
catalyst.
b) No oxidation
products from the solvent
were observed

TURNOVERS
91
Figure 4.2 Free radical experiments in norbornene oxidations.

92
the ligands. Any alteration of the catalyst would justify any
mechanistic considerations as well.
Several attempts were made to identify the nature of the
catalytic species during and after the reaction using norbornene
oxidations. The activity of the catalyst levels off after 150 hours
of reaction time in a typical norbornene oxidation. Analysis of
the spent catalyst indicate the perfluorobutyrate complex decomposes
during the reaction. Fluorine NMR of aliquots of actual reaction
mixtures taken before the reaction, at the end of the induction
period (24 hours), and at 48 hours indicate significant changes in
the catalyst are occurring (Figure 4.3) A new resonance at 119 ppm
appears after 24 hours of reaction time, while the resonance at 117
begins to disappear. The 117 ppm peak has completely vanished at 48
hours, leaving only the 119 ppm and the original 116 ppm peaks in
that region. No integration of these peaks is possible, due to the
low concentrations of catalyst (10'^ M), the low signal-to-noise
ratio, and the baseline roll, an inherent feature of the probe in
fluorine NMR. The vqq stretch in the FTIR shows a similar shift
towards the free acid, shifting approximately 20 wavenumbers to
higher frequency, from 1704 to 1720 cm'*.
The products observed in the oxidation of these olefins, and
the proposed decomposition of the catalyst are consistent with a
typical free radical autoxidation mechanism similar to the standard
Haber-Weiss scheme mentioned in Chapter I (Figure 4.4). This
proposed mechanism involves the one-electron reduction of the
catalyst in the presence of an olefin substrate (bound to a Ru

Figure 4.3
norbornene
hours (C)
NMR of the [Ru30(pfb)g(Et20)3]+ catalyst during
oxidations. (A) catalyst before (t = 0 hours) (B) t
: = 48 hours.
24

94
B
JjL-M*****^^
—i—i—i—i—i i i—¡—r—i—i—i ~ -i—!■■■ -i— i ¡—r
â–  | ' r â–  i i i
-93
-100
-113
-123 «**

95
ROOH
RO« + OH’
ROOH > R02* + H+
RÜ2* + R* > RO2R
Figure 4.4 Proposed mechanism for the autoxidation of olefins by
[Ru30(pfb)6(Et20)3](pfb).

96
center) to give an alkyl radical, R* . This radical rapidly reacts
with O2 to give an alkyl hydroperoxide, RÜ2* , which can then be
decomposed either thermally or by the perfluorobutyrate catalyst to
give alcohol and ketone products via a classical autoxidation
pathway. A possible route to norbornene epoxide is also outlined in
Figure 4.4. This particular route has been previously proposed by
Kochi, Lyons and others for the Co(acac)3-catalyzed epoxidation of
norbornene.Here hydroperoxy radicals formed thorugh the metal -
catalyzed decomposition of hydroperoxides attack the double bond,
forming a radical that rearranges to an epoxide.
This mechanism also explains the unreactivity of the alcohol
substrates as well. As free radical traps, alcohols would be
expected to be inert to free radical initiated oxidations, except
under highly stringent conditions. The perfluorobutyrate complex,
although an excellent initiator for olefin autoxidations, is not
potent enough to initiate the oxidation of alcohols under these
conditions.
From the spectral evidence, the catalyst decomposes, indicating
the reduction of the catalyst destabilizes the complex in some
fashion. The induction period corresponds to the formation of the
active catalytic species, which could be a mono- or dinuclear
ruthenium species.
Experimental
All reagents used were reagent grade or better. Olefinic
substrates were handled as described under "Scope of Catalysis."

97
The acetonitrile solvent used was HPLC grade and was stored over
molecular sieves.
GC analysis was performed as described in Chapter II, using the
same DEGS column. All instrumentation used - IR, *H NMR, and
NMR - are the same as described in Chapters II and III. The
catalyst was synthesized as described in Chapter III.

CHAPTER V
ALKANE OXIDATIONS BY A NOVEL
TRINUCLEAR RUTHENIUM CARBOXYLATE COMPLEX
Introduction
The field of catalysis has grown dramatically over the past 20
years. The challenges of activating H2, olefins, O2, CO, and N2
have all been pursued feverishly by inorganic and organometal1ic
chemists in hopes of discovering unique catalysts or processes to
more effectively and selectively utilize these materials. These
investigations are further evidence of the cooperativity and
interdependence of the fields of homogeneous catalysis and inorganic
coordination/organometallic chemistry mentioned in Chapter I.
Activation of alkanes, on the other hand, has only recently received
similar attention. In 1968, Jack Halpern proclaimed one of the most
important and challenging problems in the whole field of homogeneous
catalysis was the development of a successful approach to the
activation of C-H bonds.141,142
The largest scale industrialâ– application of homogeneous
catalysis is in the oxidation of hydrocarbons by O2.
Primarily, this predomination is due to the control of selectivity
afforded by homogeneous processes. As mentioned in Chapter I, the
reaction of O2 with hydrocarbons, although thermodynamically
favorable, is hard to initiate, and once initiated, very difficult
98

99
to control, so the advantages inherent in homogeneous catalysis are
highly desired in industrial alkane oxidations.
Oxidation of hydrocarbons generally involves the formation of
an alkyl hydroperoxide which is decomposed into products by a
transition metal catalyst (see Figure 1.1). In the majority of the
industrial processes to be discussed, the metal complex functions as
an agent of decomposition for the alkyl hydroperoxide, and does not
interact with O2 or the substrate.^ The metal complexes used as
catalysts enhance the production of desirable products while also
stimulating the production of more radical species to perpetuate the
chain.^2
Two of the most important industrial processes involving the
oxidation of hydrocarbons are the production of acetic acid (2.64
billion lbs. in 1984) and adipic acid (1.39 billion lbs. in 1984).^
Acetic acid is made via the oxidation of butane using a soluble
Co(11) salt under highly stringent conditions (Equation 5.1).,143
n-C4Hio > CH3CC2H5 > CH3COOH Equation 5.1
0
150 - 225 °C 93 % yield
800 psig O2 45 % selectivity
This process achieved a great deal of its popularity due to the
abundance of butane as a feedstock. Newer technology, the
carbonylation of methanol by a soluble rhodium carbonyl catalyst, is
rapidly taking over as the method of choice for producing acetic
acid, however.

100
The production of adipic acid has been carried out in
essentially the same fashion since the 1940's. The commercial
process involves the oxidation of cyclohexane to KA oil (a mixture
of cyclohexanone (K) and cyclohexanol(A)) by a soluble Co(11)
catalyst, again under stringent conditions.3,52,143,144 The
subsequent oxidation of the KA oil in the presence of Cu, O2, and
HNO3 leads to adipic acid with approximately a 90% yield (Equations
5-2, 5-3).52
C6HnOH
C5H12 > + Equation 5-2
C6HhO
C6HhOH
+ > C00H-(CH2)-C00H Equation 5-3
C6HnO
The yields of alcohol and ketone range between 60 and 70% in the
initial step, but with an incredibly low rate of conversion (10%)
for a commercial process.50,143,144 jhe process is deliberately
kept at such low conversion rates to avoid further oxidation of
either the ketone or the alcohol to undesirable products (such as
ring-opened acids, alcohols, ketones, and lactones).^
Figure 5.1 describes the proposed mechanism for the oxidation
of cyclohexane as carried out industrially. In this cycle, the
cobalt catalyst is not involved in the reaction of cyclohexane and
O2 to produce cyclohexyl hydroperoxide. The catalyst controls the
product distribution (in the absence of any metal ions, the amount

101
c6h12
c6h1 1°
In»
-V->
InH
<
02
Co
> products
Figure 5.1 Proposed mechanism for industrial cyclohexane
oxidations

102
of ketone formed is twice that of alcohol) and influences the rate
of the reaction in that the metal-catalyzed decomposition of the
hydroperoxide generates radicals to continue the reaction^
However, at high concentrations of catalyst, complexes between the
metal ion and the hydroperoxide may form, inhibiting the catalysis
to a small degree.52,146 j^e actual pathways involved in this
reaction are significantly more complex than depicted in this simple
diagram.^
The academic world has also been interested in the
organometal1ic chemistry involved in the activation of C-H bonds.
Initial studies were based on the assumption that the oxidative
addition of a C-H bond to a low-valent Group VIII metal should be
analogous to the well-studied addition of H2.^»142,147 In general,
however, saturated hydrocarbons are inert to soluble transition
metal complexes. Chloroplatinum complexes (particularly [PtCl4]^)
have been shown to effect the exchange of hydrogen for deuterium
atoms in a number of lower alkanes.142,147-150
Although this line of research was unsuccessful in producing
synthetically useful methods for the activation of alkanes, the
recent upsurge of activation mimicking enzymes which are capable of
reacting with saturated C-H bonds looks promising. The majority of
work in this area has centered around creating a porphyrin analog to
the monooxygenase cytochrome P450, as mentioned in Chapter IV.
Cytochrome P450 has long been known to hydroxyl ate alkanes; the
mechanism for this reaction involves the Fe(III)-heme center

103
undergoing a one-electon reduction, binding O2, and transferring
one oxygen atom to a bound substrate and producing one mole of
water.1®1 Synthetic analogs to this enzyme initially were i ron(III)
porphyrins, which have been shown to hydroxylate alkanes and
epoxidize olefins via an iron-oxo intermediate.152-155 Metals other
than iron as the active site in the porphyrin have also been used in
further attempts to create a selective oxidation catalyst. Groves
has reported a Cr(III) tetraphenylporphyrin that epoxidizes olefins;
the active intermediate in this reaction is presumed to be a Cr(V)-
oxo species.Manganese, another metal of biological significance
due to its involvement in photosynthesis, has also been studied and
shown to transfer oxygen to a variety of organic substrates
including al kanes.157-159 The active species here is again a high-
valent metal oxo species, similar to those proposed for the olefin
oxidations (see Chapter IV). However, all of these systems use
oxygen-atom transfer reagents like iodosobenzene or N-morpholine-N-
oxide as the oxygen sources in these reactions rather than molecular
oxygen.
Although ruthenium complexes are expected to activate C-H bonds
both through direct activation of alkanes and through intramolecular
activation of alkyl ligands, studies involving these complexes have
been significantly absent.149 Both Ru(II) and Ru(0) complexes have
been shown to catalyze the H/D exchange in alkanes in ligand side
chains.3’^»149 Reversible insertion of the metal into a ligand C-H
bond has also been accomplished by a Ru(II) phosphine complex.1®®

104
In sulfuric acid, both Ru(III) and Ru(IV) will oxidize
hydrocarbons, and a combination of Ru(IV)/Cr(VI) will oxidize lower
alkanes to alkyl chlorides in aqueous chloride solutions.^2
High-valent ruthenium-oxo complexes, widely used as catalysts
for olefin oxidations (vide supra), have not been successful as
hydrocarbon oxidation catalysts. Ruthenium porphyrins have been
used to oxidize alkanes with iodosobenzene, but the results are less
than spectacular. In methylene chloride solution, the oxidation of
cyclohexane produces mostly cyclohexanone and cyclohexyl bromide via
a radical process using Ru(OEP)(PPh3)Br as the catalyst. The
mechanism invokes a standard Ru(IV)-oxo species as the active
intermediate in this system.^jhe ubiquitous ruthenium
polypyridyl-oxo complexes will not oxidize alkanes, although these
complexes will oxidize water and aromatic hydrocarbons. >164
Taqui Khan has extensively studied the oxidation of cyclohexane
using Ru (III)- EDTA compl exes. 165- 167 These complexes will oxidize a
variety of organic substrates, and detailed kinetic and
thermodynamic studies of these oxidation have been published. On
the basis of these studies, a mechanism involving a /¿-peroxo-Ru-
cyclohexane adduct is proposed. Cyclohexanol is the primary
product; the ketone is formed through the subsequent oxidation of
the alcohol. These reactions have also been carried out in the
presence of a micelle in a mixed solvent system, and the kinetics of
this oxidation studied. The mechanism in this case also involves a
peroxo-Ru(IV)-cyclohexane adduct.

105
The lack of information on the ruthenium-catalyzed oxidation of
alkanes combined with the ability of the Ru(pfb) complex to initiate
free radical autoxidation prompted the further investigation of the
capabilities of this catalyst. The commercial oxidations of alkanes
are free radical reactions, so initial studies centered around
attempts to oxidize cyclohexane and other alkanes as another test of
the versatility of the perfluorobutyrate complex.
Scope of Catalvsis
Standard oxidation setups as described in Chapter II were used.
Generally, the oxidation reactions were comprised of 40 mL of
substrate, 10 mL of acetonitrile solvent, 1 mL benzene standard, and
10'5 moles of catalyst. The standard was omitted when it interfered
with the analysis of the products in some oxidations. The
temperature was maintained at 75 °C by an oil bath, and initial
pressures of 40 psig of compressed air were used in all cases.
Products were determined by GC using a FID detector and DEGS column.
Appropriate calibration curves were constructed and used to quantify
the products. In these oxidations, blanks using all reactants
except catalyst indicated some reaction occurred after an average of
20 hours of reaction time. For this reason, the length of all
alkane oxidations was limited to 12 hours.
Results and Discussion
Although the main substrate of interest was cyclohexane, due to
its commercial importance, other alkanes were attempted as
substrates as well (Table 5.1). These substrates show strikingly
different reactivities. At 75 °C, cyclohexane was oxidized to a

106
Table 5.1
Alkane Substrates3
Substrate
Product
Turnovers/12 hours*3
cyclohexane cyclohexanol
cyclohexanone
n-hexane
no reaction
n-hexanec
2-hexanol
2-hexanone
1-hexanol
15
3
16
16
trace
methyl cyclohexane
1-methylcyclohexanol major^
2-methylcyclohexanol rninor1^
3-methyl cyclohexanol minor^j
2-methyl cyclohexanone traced
3-methyl pentane
2-methyl pentane
toluene
cyclohexanol
no reaction
no reaction
no reaction
no reaction
a) All reactions, unless otherwise noted, used acetonitrile (10 mL)
as the solvent for the catalyst and 40 mL of substrate. For
specific detail concerning reaction conditions, see "Scope of
Catalysis."
b) Turnovers = mole of product/mole of catalyst used in the
specified time.
c) M-pyrol (or 1-methyl-2-pyrrol idinone) (10 mL) was used as
sol vent.
d) Not quantified.

107
mixture of cyclohexanol and cyclohexanone in acetonitrile. N-
hexane was unreactive in acetonitrile; in m-pyrol (vide infra) the
major products were 2-hexanol and 2-hexanone, with trace amounts of
1-hexanol formed after 12 hours of reaction time. Methyl cyclohexane
was oxidized to primarily 1-methylcyclohexanol. Smaller amounts of
2- and 3-methylcyclohexanol were also produced, along with trace
quantities of 2-methylcyclohexanone. Neither 3-methyl pentane nor 2-
methylpentane were reactive under these conditions. Cyclohexanol
was not oxidized to cyclohexanone, not surprising in light of the
results reported in Chapter IV. Under slightly different
conditions, refluxing acetic acid ( 140 °C) with a Br' promoter in a
ground glass apparatus, toluene was not oxidized after 1 hour under
a flow of 02- The change in reaction conditions was warranted by
the corrosive nature of the solution.
Of these substrates, cyclohexane was studied the most
extensively, in spite of the complexity of the reaction.145 ^ 55
°C, the reaction was very slow, forming only trace amounts of
products. Raising the temperature to 75 °C increased the rate of
the reaction significantly, forming 18 turnovers of total product in
12 hours. Thus, 75 °C was used as the reaction temperature for all
subsequent oxidations. Raising the temperature higher is not
recommended, since both cyclohexane and acetonitrile boil at 81 °C,
and the flammability limits of cyclohexane and air would be
exceeded.
The selectivity observed in the oxidation of cyclohexane by the
ruthenium perfluorobutyrate catalyst is very high. No traces of

108
other products are observed in either 6CIR or GCMS. Proton NMR of
the residue obtained by completely evaporating a standard oxidation
shown only small traces of products that could be acids or lactams,
leading to the assumption this oxidation occurs with a selectivity
for the alcohol and ketone products of 99% or greater.
Work in cyclohexane oxidations has involved the addition of
various additives and promoters to enhance product yields, the
alcohol/ketone ratio, or suppress undesirable side reactions. The
effects of several additives were tested with the pfb catalyst
(Table 5.2, Figure 5.2). Bromide ions have long been used as
promoters for free radical reactions; Br» species are involved in
both initiation and propagation steps in the chain reactions.52
Under these condition, the addition of bromide enhanced the
formation of product; the blanks were inactive. Other transition
metals have also been used as cocatalysts in cyclohexane oxidations
to influence the product ratios. Chromium(III) enhances the
production of ketone and increases the rate of oxidation of
cyclohexanol to the ketone.51 Iron(III) increases the yields of
lactone and alcohols and decreases production of acids, while
manganese(II) complexes generally increase yields of alcohols and
acids.143 Although the roles of each of these additives in the
catalysis is undoubtedly complex, high degrees of variance are
possible depending upon the choice of catalysts.
In the ruthenium-catalyzed oxidations, these metals follow
these general trends. The addition of MnCl2 increased alcohol
production, while FeCl3 increased overall production and increased

109
Table 5.2
Cyclohexane Oxidation Studies:
Additive3
Conditions
Turnovers/12 hours*3
Alcohol Ketone
Mole ratio
A1cohol/Ketone
standard0
14.8
3.2
4.6
Mn2+
37.0
9.0
4.1
Br‘
65.1
16.2
4.0
Fe3+
39.7
3.3
12.0
Cr3+
150.3
45.7
3.3
Co2+
--
--
--
Co(oct)2c*
--
--
--
10-5 moies 0f
additive were
used in combination with 10"5 moles
of catalyst. See text for details.
b) Turnovers = mole of product/mole of catalyst in the specified
time.
c) Standard conditions are 10'5 mole of catalyst. For more detail,
see "Scope of Catalysis."
d) Co(11) octoate, DuPont's commercial catalyst, trade name Cobalt
Hex-Cem.

TURNOVERS/1 2 HOURS
110
ADDITIVES
Figure 5.2 Effects of additives in cyclohexane oxidations catalyzed
by [Ru30(pfb)6(Et20)3]+. (A) [Ru30(pfb)6(Et20)3]+ (B) MnCl2.2H20
(C) NaBr (D) CuCl2»H20 (E) Co(oct)2 (F) FeCl3*6H20 (G) CrCl3*6H20.

Ill
the A/K ratio by a factor of 3. Chromium(III) chloride had a small
effect of the A/K ratio while drastically increasing the formation
of product by ten-fold. The addition of copper chloride to this
reaction completely shut down the reaction. Only a few of these
metals show any activity in the blanks. Iron(III) also reacts,
producing almost as much product as the iron-ruthenium combination
with a similar A/K ratio . Chromium(III) is much less active and
produces a 1:1 molar ratio of alcohol to ketone. Interestingly,
copper(II) alone is slightly active, producing 3.5 times as much
alcohol as ketone. Copper(I) will reduce alkyl hydroperoxides
catalytically,3 but is much less active that the traditional
catalysts Mn(II) and Co(II). Since the Cu(II) blank is active, the
combination of ruthenium and copper should not have been inactive
because both metals will decompose peroxides. Copper is used as the
catalyst in the second step in the oxidation of KA oil to adipic
acid, so the products formed by the bi-metallic system may have been
undetectable under the analysis conditions used.
A sample of the cobalt catalyst used in DuPont's commercial
oxidation of cyclohexane was tested under our reaction conditions.
This complex, cobalt(II) octoate (or 2-ethylhexanoate), was
unreactive at 75 °C, a very encouraging result. Acetonitrile and
cyclohexane are only slightly miscible, even at elevated
temperature. The biphasic nature of the solution implies that all
catalysis is taking place at the interfaces of the micelles that
form as the solution is vigorously stirred. In order to draw the

112
catalyst into the substrate phase, and hopefully enhance the
catalysis, sodium octoate [Na(oct)] was attempted as a phase-
transfer reagent (Table 5.3). Figure 5.3 illustrates the effects of
the addition of various amount of Na(oct) to a typical reaction.
More typical phase transfer reagents such as Aliquat 336 are
cationic, used for transferring anions from aqueous to organic
phases. In this case, transfer of the cationic [Ru30(pfb)g(Et20)3]+
species necessitated the use of an anionic phase transfer reagent.
Sodium octoate was chosen for its solubility in cyclohexane and for
the use of an anionic carboxyl ate to replace the perfluorobutyrate
counterion. Adding only one-half mole of phase transfer reagent per
mole of catalyst increases the reaction rate, but the differences in
the amount of product formed with one-half mole and three moles of
Na(oct) are insignificant. Also, the solutions are still biphasic,
even with four moles of Na(oct) present. Other studies have shown
that while Mn(II) stearate will catalyze the oxidation of alkanes,
the combination of Nn(II) and potassium stearates is five times more
effective.^ /\ similar type of synergistic effect may be occurring
in these experiments.
The solvent 1-methyl-2-pyrrolidinone (M-pyrol) has been shown
to be an excellent medium in which to carry out oxidations.
Solvent experiments similar to those in the olefin study were
performed for the cyclohexane oxidations (Table 5.4) Several tandem
experiments were performed using both m-pyrol and acetonitrile as
solvents and varying the ratio of solvent to substrate for each
solvent. Under standard conditions (40 mL substrate/10 mL

113
Table 5.3
Cyclohexane Oxidation Studies:
Phase Transfer Catalysis
by Sodium Octoate
Conditions
Turnovers/12 hours3
Alcohol Ketone
Mole ratio
Alcohol/Ketone
standard^
14.8
3.2
4.6
.5 molec
57.0
13.0
4.5
3 mole
66.0
12.0
5.5
4 mole
27.0
7.0
3.8
Turnovers =
mole of product/mole
of catalyst
in the specified
time.
b) Standard conditions are 10"^ moles of [Ru30(pfb)s(Et20)3](pfb)
catalyst.
c) Amount of Na(oct) per mole of catalyst added to a standard
reaction.

TURNOVERS/12 HOURS
114
\m Na(odDatB)/MOLES Ru(pfb)
Figure 5.3 Phase transfer in cyclohexane oxidations (A)
[Ru30(pTb)5(Et2O)3](pfb) (B) .5 mole Na(oct) (C) 3 moles Na(oct) (D)
4 moles Na(oct).

115
Table 5.4
Cyclohexane Oxidation Studies:
Solvent Effects
Turnovers/12 hours3
Mole ratio
Sol vent
Alcohol
Ketone
A1cohol/Ketone
10 mL MeCNb
14.8
3.2
4.6
25 mL MeCN
23.8
6.4
3.7
40 mL MeCN
--
--
--
10 mL m-pyrolb’c
20.5
8.9
2.3
25 mL m-pyrol
9.8
2.2
4.4
40 mL m-pyrol
1.6
0.5
3.6
cyclopentanone0
_ _
_ _
_ _
cyclopentanone/
MeCNCjd
5.2
trace
ethanolc>e
--
--
--
nitrobenzenec,e
--
--
--
DMFe
trace
--
--
benzonitrilee
4.7
1.6
3.0
pyridinec’e
a) Turnovers = mole of product/mole of catalyst in the specified
time.
b) The-total volume of solvent and substrate was held constant at 50
mL, with the amount of solvent varied as indicated.
• c) one phase
d) The mixed solvent system used 5 mL MeCN and 6.5 mL
cyclopentanone.
e) All reactions were run under standard conditions using 10 mL
solvent.

116
solvent), little difference is observed for either solvent, except
the m-pyrol solution is one phase. Inverting the ratio to 10 mL
substrate/40 mL solvent had a considerable effect on the reaction.
In 1-methyl-2-pyrrolidinone significantly less activity was
observed; in acetonitrile, the catalyst became completely inactive.
Changing the ratio to 25/25 increased the reaction in acetonitrile
but decreased the reaction in m-pyrol.
Cyclopentanone was attempted as a solvent to solve the phase
problem. Industrially, KA oil is used as a solvent to test
catalysts, but distinguishing between the solvent added and
productformed would be difficult under our conditions. The five-
carbon ketone was utilized as the solvent to circumvent this
difficulty. The solution was one phase, but no oxidation occurred
in 12 hours. A similar reaction was run using 5 mL of acetonitrile
as solvent, with enough cyclopentanone added to make the solution
one phase. In this mixed solvent system, oxidation of cyclohexane
did occur, but the amount of product decreased by one-third.
Substituting ethanol for the acetonitrile completely shut down the
reaction; interestingly, with alcohol as solvent the reaction
mixture was one phase. Catalysis did not occur in either
nitrobenzene or pyridine, but in DMF trace amounts of cycloh-exanol
were found after 12 hours, and the reaction in benzonitrile produced
measurable amounts of both alcohol and ketone. These results are
strikingly different from those found in the alkene oxidations,
where only acetonitrile and ethanol were active solvents.

117
A plausible explanation involves a major difference in the
types of electron transfer involved in these two oxidations. In the
olefin oxidations, the substrate is coordinated to one of the
ruthenium atoms in the trimer, and inner-sphere electron transfer
occurs, forming alkyl radicals and destabilizing the trimer. These
radicals initiate the autoxidation chain and the perfluorobutyrate
complex decomposes. The solvent variations in the oxidation of
cyclohexane indicate a slightly different mechanism is involved.
The reaction occurs in relatively polar, moderately coordinating
solvents, so that the catalyst should be more properly formulated as
[RU3O(pfb)6(S)3]+ in solution. These coordinating solvents help
stabilize the complex during the oxidation, which necessarily occurs
via an outer-sphere electron transfer. Alkanes are extremely poor
ligands, so it is doubtful that a molecule of cyclohexane is
contained in the coordination sphere of the trimer.
If this supposition is true, the possibility arises that the
catalyst could remain intact during the oxidation. A slight
induction period (approximately 2 hours) is seen in these
oxidations, after which the production of alcohol and ketone rises
exponentially (Figure 5.4). NMR studies similar to the ones
performed on the catalyst in. situ for the olefin oxidation were not
possible, since the solutions were biphasic. However, NMR of
the spent catalyst indicates little change has occurred during the
course of the reaction (Figure 5.5). FTIR of the catalyst before
and after also shows little change. The stability of the catalyst
during the oxidation is more support for a different mechanism than

TURNOVERS
118
Figure 5.4 Activity curve for the oxidation of cyclohexane by
[Ru30(pfb)6(Et2O)3](pfb).

Figure 5.5 NMR of [Ru30(pfb)6(Et2Ü)3](pfb) during cyclohexane
oxidations (A) before catalysis (B) after 12 hours of reaction time.

120

121
that involved in the olefin oxidations, where the catalyst
decomposes.
Standard free radical experiments were performed in these
oxidations as well (Figure 5.6). The addition of benzoquinone to
the reaction decreased the amount of product formed by one-third
(from 28 turnovers of total product to 16). The addition of AIBN to
a standard reaction had no effect at all. Using AIBN alone (with no
ruthenium catalyst) did not initiate the reaction, only achieving .8
turnovers of total product. These experiments indicate the reaction
involves free radicals, and the ruthenium perfluorobutyrate complex
is a more potent initiator than AIBN in this biphasic system.
Since free radicals are involved, more than likely peroxides or
hydroperoxides are formed as intermediates as well. Sharpless*^
has published a method for titrating peroxides in organic solvents
that has proven to be useful in other biphasic catalyst systems
studied.169 For these peroxide titrations, several modifi-cations
of the original procedure were necessary. A larger aliquot of
analyte (10 mL) was used, and the solutions were not diluted with
100 mL of water. The peroxide concentrations were very low, so
larger aliquots and no dilution was necessary; the water also formed
an emulsion with the cyclohexane, making the endpoint of the
titration very difficult to see. The titrant was diluted from .10 N
to .001 N to obtain reasonable values for the volume of titrant
required. Two different blanks were run, one with all components of
a standard oxidation without catalyst stirred at room temperature
for 30 minutes, and one under pressure for 12 hours. A standard

TURNOVERS/12 HOURS
122
Figure 5.6 Free radical experiments in cyclohexane oxidations. (A)
[Ru30(pfb)6(Et20)3](pfb) (B) catalyst and benzoquinone (C) catalyst
and AIBN (D) AIBN alone.

123
oxidation including catalyst was also set up for 12 hours. The
solvents already contain a small amount of peroxides (7.0 x 10"® N),
and the blank under pressure titrated for two orders of magnitude
more (1.2 x 10'^ N). Titrating the catalyst run was a challenge
since the greenish solution were to be titrated from yellow to
colorless. A larger amount of error is present in the titrations of
the catalyst, but this reaction contained the most peroxide (8.65 x
10'4 N), as would be expected. However, these small differences in
peroxide concentration between the catalyst and blanks do not
account for the large amount of product formed. The method is
suspect, especially at such low concentrations, but an additional
factor to consider is that the decomposition of peroxides in
solution continues in the presence of the catalyst, even after
removal from heat and pressure, making the amount of peroxide in the
standard catalyst solution lower than it should be. However, these
experiments do indicate more peroxides are formed in the presence of
the ruthenium perfluorobutyrate catalyst.
The mechanism proposed for these oxidations is slightly
different than that for the industrial-scale oxidations (Figure
5.7). The catalyst, solvated by a coordinating solvent, forms an
outer-sphere complex with the substrate. Electron-transfer from the
RU3O core to the alkane forms an alkyl radical, which rapidly reacts
with O2 beginning the chain reaction, and a reduced form of the
trimer. The catalyst is probably involved in the decomposition of
the alkyl hydroperoxides after they are formed as well. By analogy
with the alkyl carboxylate trimers, the reduced ruthenium trimer is

Rum
Ruin . l
Ruin
3 S
3 L
S
Ruin
\
Ruin .
y
o
s
Figure 5.7 Proposed mechanism for alkane oxidations by
Ru30(pfb)6(Et20)3+.

125
rapidly oxidized back to a Ru(111,111,111) species by O2,
regenerating the active catalyst.
Experimental
All reagents used were of the best quality available.
Cyclohexane (Spectranalyzed) and benzene (SpectrAR) were used as
received; HPLC grade acetonitrile was stored over activated
molecular sieves. Hexane was purified by stirring for several hours
with H2SO4, washing with water several times, and drying over
CaCl2* ^Before use, the hexane was passed through a column of
neutral alumina. No purification was necessary for any other
substrates used. Du Pont's catalyst, Co(octoate), trade name Cobalt
Hex-Cem at 12% cobalt was donated by Mooney Chemical Company.
All equipment was described in Chapter II. The catalyst was
prepared as described in Chapter III. All instruments used (FTIR,
NMR) are as described in Chapters II, III, or IV.

CHAPTER VI
CONCLUSIONS
The series of trinuclear ruthenium carboxylate complexes has
fulfilled its promise of highly intriguing chemistry. The catalytic
potential of these compounds, tapped essentially only in the area of
reductions, has been explored further, primarily in the area of
oxidations. The use of molecular oxygen as the primary oxidant, in
combination with these complexes, has made these inquiries even more
unique. A wide variety of ruthenium carboxylate complexes were
investigated as catalysts for the oxidation of essentially the
entire range of organic substrates.
Trinuclear ruthenium carboxylates of the formula
[Ru30(02CR)6(L)3]n+ where R = CH3, C2H5; L = H20, PPh3, py; n = +1,
0, selectively oxidize primary alcohols to aldehydes and secondary
alcohols to ketones in the presence of molecular oxygen. The
proposed mechanism for these oxidations involves a Ru(II)-Ru(IV)
cycle that assumes retention of the trimeric structure. However,
these complexes are unable to oxidize olefins.
A novel member of this series, a trinuclear ruthenium
carboxylate containing completely fluorinated carboxylate ligands,
has been synthesized and characterized. The exchange of alkyl for
fluorinated carboxylates was based on the assumption the electron-
126

127
withdrawing character of the fluorinated ligand would increase the
acidity of the metal, making it more likely to bind olefins, as well
as making the Ru(III) centers better oxidants. Without a crystal
structure, the exact nature of this complex cannot be definitively
identified; however, the experimental evidence leads to the
formulation of this novel complex as [Ru30(pfb)g(Et20)3](pfb)_
This novel complex was screened as a catalyst for the oxidation
of olefins by O2. The goal of synthesizing a catalyst, i.e., the
ruthenium perfluorobutyrate complex, that would selectively oxidize
olefins to the epoxide was not realized. This catalyst initiated
the free radical autoxidation of all of the alkene substrates
attempted. Spectral evidence indicated the decomposition of the
catalyst, probably to a mono- or dinuclear ruthenium species. The
mechanism for this reaction involves a standard Haber-Weiss
decomposition of alkyl hydroperoxides by the catalyst, forming
products, with subsequent decomposition of the trimeric catalyst.
Other substrates were screened with the perfluorobutyrate
complex as well. Alkanes, the most difficult substrates to oxidize,
were successfully converted to oxygenated products by this complex
and molecular oxygen. Interestingly, this complex will oxidize
cyclohexane to a mixture of cyclohexanol and cyclohexanone under
conditions significantly less strenuous than those used in industry.
The mechanism for this reaction is decidedly different from the
typical free radical mechanism proposed for the olefin oxidations.
No decomposition of the catalyst occurs, and the cluster seems to

128
remain intact during the production of cyclohexyl radicals and the
subsequent decomposition of the peroxides formed form this radical.
The successes of these complexes as catalysts lead to
speculation about the catalytic potential hidden in other trinuclear
carboxylate complexes. Most intriguing in light of current
interests in mixed-metal catalysis is the possibility of mixing
different metals in the M3O core. Such a catalyst theoretically
could be designed and "tuned" to achieve selective activation of
substrates by the judicious choice of metals placed in the trimer.
Further investigations of the catalytic potential of this family of
compounds is warranted based on these studies of the ruthenium
system.

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BIOGRAPHICAL SKETCH
Shannon Davis, born December 6, 1963, lived for most of her
formative years in the town of her birth, Statesboro, Georgia,
except for a sojourn of two years in Athens, Georgia. As a senior
in high school, part of a joint enrollment program, she entered
Georgia Southern College in 1980. She majored in chemistry, and
after a short stint as a lab technician at Braswell Food Company
(determining the pH of pickles and the sugar content of jams and
jellies) she became firmly committed to pursuing the science in
graduate school. She graduated from GSC in June of 1984 with the
degree B. S. in Chemistry, and the following August she moved to
Gainesville and began her graduate career at the University of
Florida. After a brief stay of four years, she expects to receive
her Ph.D. in inorganic chemistry officially in December 1988. In
the fall of 1988, she will move to Pensacola, Florida, where she
will join the Fiber Intermediates Division of Monsanto Chemical Company.
139

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.
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.
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.
R. Carl 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.
William M. Jones
Professor of Chemis
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.
Gar Hoflund (J
Professor of Chemical Engineering
This dissertation 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.
December, 1988
Dean, Graduate School

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