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Synergism in metal carboxylate clusters

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Title:
Synergism in metal carboxylate clusters
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Bilgrien, Carl Joseph, 1959-
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English
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ix, 231 leaves : ill. ; 28 cm.

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Subjects / Keywords:
Adducts ( jstor )
Alcohols ( jstor )
Carboxylates ( jstor )
Dimers ( jstor )
Enthalpy ( jstor )
Ligands ( jstor )
Orbitals ( jstor )
Oxidation ( jstor )
Oxygen ( jstor )
Trimers ( jstor )
Chemistry thesis Ph. D
Dissertations, Academic -- Chemistry -- UF
Metallic soap ( lcsh )
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bibliography ( marcgt )
non-fiction ( marcgt )

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Thesis:
Thesis (Ph. D.)--University of Florida, 1986.
Bibliography:
Bibliography: leaves 222-230.
General Note:
Typescript.
General Note:
Vita.
Statement of Responsibility:
by Carl Joseph Bilgrien.

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SYNERGISM IN METAL CARBOXYLATE CLUSTERS
By
CARL JOSEPH BILGRIEN
A DISSERTATION PRESENTED TO THE GRADUATE SCHOOL
OF THE UNIVERSITY OF FLORIDA IN
PARTIAL FULFILLMENT OF THE REQUIREMENTS
FOR THE DEGREE OF DOCTOR OF PHILOSOPHY
UNIVERSITY OF FLORIDA
1986


To Deanna,
who always knew.


ACKNOWLEDGEMENTS
No scholar or scientist works alone: each must rely on the
labors of past workers and the assistance of his contemporaries.
I have incurred many debts in this regard.
First and foremost I thank Professor Russell S. Drago for his
support, encouragement and advice. He has generously shared his
insight and perseverance and I am grateful for having had the
opportunity to work with him.
I thank my committee members, Professors Earl Muschlitz,
David Richardson, Harry Sisler and E. Dow Whitney for their
efforts.
For his continued friendship and enthusiasm I am especially
grateful to Dr. Barry B. Corden.
To all the group members who have shared their time, advice,
expertise, grousing and politics, I am grateful. These include
Kenneth Balkus, Iwona Bresinska, Jeffrey Clark, Richard Cosraano,
Shannon Davis, Peter Doan, Andrew Griffis, Karen Jongeward,
Ernest Stine, Joshua Telser, Keith Weiss and Ngai Wong.
For their ability to build anything I drew I thank the men
in the glass shop, Rudy and Dick. For their sweat and good
humor, I thank Vernon, Chester and Daly of the metal shop. For
not giving up on the calorimeter I thank Russell Pierce. For his
many suggestions and services I thank Dr. Roy King.
iii


Finally, I can never repay the time and sacrifices of ray
wife, Deanna Saint Souver. For her unabated
inspite of all too many hours spent in lab, I
always.
encouragement
wi11 love her
iv


TABLE OF CONTENTS
PAGE
ACKNOWLEDGEMENTS iii
KEY TO ABBREVIATIONS vii
ABSTRACT viii
CHAPTER I. GENERAL INFORMATION 1
CHAPTER II. REACTIVITY OF Cr2(02CR)4 14
A. Introduction 14
B. Results and Discussion 18
C. Conclusion 67
D. Experimental 68
CHAPTER III. TRANS INFLUENCE ACROSS A
METAL-METAL BOND 78
A. Introduction 78
B. Results and Discussion 81
C. Conclusion 96
CHAPTER IV. THE ELECTRONIC SPECTRA OF
M(III)2M(II)(02CR)6L3 SPECIES 98
A. Introduction 98
B. Results and Discussion 100
C. Conclusion 125
D. Experimental 125
CHAPTER V. THE OXIDATIVE DEHYDROGENATION OF
ALCOHOLS CATALYZED BY OXOTRIRUTHENIUM
CARBOXYLATES 129
A. Introduction 129
B. Results and Discussion 132
C. Conclusion 185
D. Experimental 188
CHAPTER VI. SUMMARY AND CONCLUSIONS 194
v


APPENDIX I. SPECTRAL AND CALORIMETRIC DATA
198
APPENDIX II. OPERATION OF CALORIMETER 215
APPENDIX III. DERIVATION OF EQUATION 2-17 219
REFERENCES 222
BIOGRAPHICAL SKETCH 231
vi


KEY TO ABBREVIATIONS
but
butyrate = O2CCH2CH2CH2
hept
heptanoate = 02C(CH2)5CH
hfb
heptafluorobutyrate = O21
OAc
acetate = O2CCH2
oct
octanoate = 02C(CH2)gCH2
prop
propionate = O2CCH2CH2
tfa
trifluoroacetate = 02CCF
vii


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
SYNERGISM IN METAL CARBOXYLATE CLUSTERS
By
Carl Joseph Bilgrien
August 1986
Chairman: Professor Russell S. Drago
Major Department: Chemistry
Synthetic, spectroscopic and reactivity studies on several
metal clusters with carboxylate ligands are described. These
complexes are of the general formula M2(02CR)jL2 or M'^CC^CR^L^
where M is Cr or Rh; M' is Co, Cr, Fe or Ru; C^CR is a bridging
carboxylate ligand; and L is a neutral donor ligand. These
studies were undertaken to examine the influence of reaction at
one metal center upon that at an adjacent metal atom and to help
understand the metal-metal bonding interactions which contribute
to the transmission of coordination effects. Ligand exchange
reactions of C^C C^CC F^) 4 [ (CH^ C^) 2 3 2 were monitored.
Equilibrium constants and enthalpys for exchange reactions with a
variety of donors were determined from calorimetry data. The
resulting enthalpys were used in a correlation analysis which
demonstrated that the Cr(II) centers are significant Lewis acids
and interact with axial ligands almost exclusively in an
viii


electrostatic sense. Despite a relatively weak metal-metal bond
the first exchange enthalpy is appreciably higher than the
second. Magnetic susceptibility measurements show increasing
paramagnetism for the Cr2^ + unit as stronger donors displace
coordinated ether. Strong donors promote oxidation and
rearrangement of the dinuclear unit.
Mixed ligand complexes of the form BRh2(02CCF2CF2CF2)|L where
B is a Lewis base were examined by electron paramagnetic
spectroscopy (L is the spin label 2,2,6,6,-tetramethylpyridine-N-
oxyl) and infrared spectroscopy (L is CO) and the spectral data
used to calculate acid parameters which describe the Rt^"1" unit.
The spin label g-value, the CO stretching frequency and the
calorimetric enthalpy, all of which describe the perturbation the
base makes on the trans-metal atom, show different relative
electrostatic/covalent responses as the base is varied,
demonstrating the method dependence of monitoring donor-acceptor
adduct formation. The electronic spectra of clusters of the form
M^OCOgCR^Lj (M = Co, Fe, Ru) exhibit a donor ligand (L)
dependence only for M = Ru.
Lastly, clusters of the form Ru20(02CR)^L2n+ (n = 0,1) are
shown to catalyze the selective oxidative dehydrogenation of
alcohols. A mechanistic proposal incorporates the demonstrated
nonradical behavior and the observed stoichiometry. Nonradical
chemistry and reduction of oxygen to water are demonstrated and a
mechanism proposed. High catalyst activity is suggested to arise
from the multiple metal centers acting in concert.
ix


CHAPTER I
GENERAL INFORMATION
An underlying theme in the current and recent intense
interest in the chemistry of molecules with multiple metal
centers is the way in which the metal centers influence each
other and generate unique reactivity. This synergistic interplay
is implicated in a host of chemical systems. Many enzymes rely
on multiple metal centers for substrate binding (e.g. heraocyanin,
lacease) while others employ proximal metal atoms for electron
transport (e.g. cytochrome c oxidase) or perturbation of the
substrate binding (e.g. nitrogenase). Examples of homogeneous and
heterogeneous catalyses with multiple metal centers abound
whereas interesting physical properties or reactivities of a
stoichiometric nature are often introduced by metal atoms in
close proximity.
The nature and extent of discrete metal-metal interactions
can vary from direct orbital overlap to long range electron
transfer. Studies along this continuum have proceeded via
several fronts: physical and theoretical studies of orbital
interactions; introduction of specific reactivity or binding;
modeling of metal loenzyraes; and studies of electron transfer.
One area in particular, the study of molecules with metal-metal
bonds, has received considerable attention following the
1


2
1 o
pioneering work in the laboratories of Wilkinson J and
Cotton.The latter area is the focus of this thesis; a
historical perspective is given here as introduction.
Indentification of single and multiple metal-metal bonds as
stereoelectronic elements of molecular structure has come about
only recently. Prejudiced by the conceptual framework of
coordination chemistry established by Alfred Werner, metal-metal
bonds were deemed inconsistent with one center coordination
chemistry. The advent of x-ray crystallography heralded
recognition of bonding between metal atoms. Although a
crystallographic report"^ as early as 1946 demonstrated Mo-Mo bond
distances shorter than those in metallic molybdenum, unequivocal
evidence for metal-metal bond formation came from the structure
determination of Mn2(C0)-|Q in 1957. Evidence for multiple
metal-metal bonding (in this case, a quadruple bond) came from
the initial structure determination^ of the RejClg^- anion and
its reinvestigation1 in 1964. These early discoveries and the
chemistry of metal-metal bonds have been throughly reviewed
through 1980 in "Multiple Bonds Between Metal Atoms" by F. A.
Cotton and R. A. Walton.11 Several other more specific reviews
have appeared recently.1^-1
The importance of the Re2Cli|i<- anion to the development of
this field lies in the electronic structure of the Re-Re bond.
The crystal structure1^ of K2(Re2Clg) (F^O^ demonstrated a Re-
Re distance of 2.24 A (2.75 A in metallic Re) and eclipsed
conformation of the two ReCl^ halves. Both of these features can
be qualitatively explained by considering d orbital overlap. The
in phase and out of phase combination of pairs of d orbitals on


3
the two metal atoms generates five bonding and the corresponding
antibonding molecular orbitals (MO) in Dooh symmetry. The in
phase bonding interactions are shown schematically in Fig 1-1.
Positive overlap of the dz2 orbitals generates a a bond (Fig 1-
1a) whereas negative overlap generates the corresponding a
antibonding orbital (not shown in Fig 1-1). The dxz and dyz
combinations give rise to two orthogonal but degenerate it bonds
and the corresponding tt antibonding orbitals. Lastly, in the
absence of any ligand interactions, cofacial overlap of the pairs
of dxy and dx2_y2 results in degenerate 6 bonds and their <5*
complements. Introduction of ligand orbitals lowers the symmetry
to D2 dx2_y2 orbitals point towards the ligands; the dxy orbitals point
between them). The dx2_y2 orbitals are thus utilized in ligand
bonding and play effectively no role in metal-metal bond
formation. The extent of orbital overlap dictates the MO
energies and the orbital diagram which results is shown in Fig 1-
2. In terms of the (Re2Clg)2 anion, each Re^ + center
contributes four electrons to give a net quadruple bond,
which explains the short Re-Re distance. The twist angle
dependence of the 6 interaction results in maximum overlap when
the Cl are eclipsed. Thus, the stabilization gained in 5 bond
formation is greater than the repulsion energy of the eclipsed
halides. The qualitative features of this MO description are
supported by quantitative calculations (SCF-Xa-SW) on the
(Mo2Cl8)i|" ion1-1^ and the (Re2Clg)2~ ion20-21 and is general
for dimers of the 2nd and 3rd row d2-^ transition metals with
octa-halo ligand sets


4
Lol,
a)
Co CO
ctCSo
b)
SS * SS
806
c)
<% d ft -
c5
d) ^ d*y ^
6}
Fig 1-1.
The five nonzero d-d overlaps between two metal atoms.


5
a
K
7T*
S*
S
TT
CT
Fig 1-2. Qualitative description of the primary metal-metal d
orbital interactions for a dinuclear metal carboxylate in
symmetry. The dx2- 2 orbitals are involved with carboxylate
bridge orbitals and do not contribute significantly to the metal-
metal bond.


6
The utility of this MO description extends to other ligand
sets as well, most notably the tetracarboxylates, M2(C>2CR)ijL2.
As with the octahalogenates, work with these complexes appeared
in the literature long before their general structural features
pp
were revealed by a crystallographic study'1*1 of
(which, incidentally, carries no metal-metal bond) in 1953. The
tetracarboxylate ligand framework is ubiquitous in transition
metal chemistry; examples exist for vanadium, chromium, cobalt,
copper, raolydbenum, technetium, ruthenium, rhodium, tungsten,
rhenium, and osmium where L represents a neutral donor molecule
or coordinated anion (naked clusters with no L also abound). The
general structure of M2(02CR)jjL2 type dimers* is shown in Fig 1-
3. Symmetry dictates the qualitative orbital interactions
displayed in Fig 1-2 apply to the tetracarboxylates also, as
borne out in extensive numerical calculations on dimers of the
2nd and 3rd row transition metals. Dinuclear carboxylates of the
first row, however, are not adequately described by this orbital
description. Representative c lusters,23-26
(^(quin^C^H^CC^)^ and (^(^O^CCH^CC^)^ all exhibit longer
metal-metal distances than found in the parent metal. As such,
no metal-metal bond is proposed. For neutral dinuclear
tetracarboxylates of these metals, the observed magnetic behavior
is rationalized in terms of antiferromagnetic exchange between
equivalent spin centers, presumably through the carboxylate
The term "dimer" is traditionally reserved for addition
complexes formed from two monomeric units. In the present
context, the term "dinuclear complex" would be more appropriate;
"dimer" is favored in the interest of brevity.


7
R
R
Fig 1-3. The general structure of MjCC^CR^CUg. Diffraction
studies support minor deviation from idealized symmetry.


8
bridges. The Cr21,+ center of the chromium tetracarboxylates is
the exception to the rule. Although there has been disagreement
as to the nature of the bonding in the systems,
that some bonding interaction is present has been reasonably
established. Results of SCF-Xa-SW^ and ab initio^-^
calculations support a weak quadruple bond as the generic MO
description of Fig 1-2 would suggest for two d^ centers.
As a family, the transition metal tetracarboxylate dimers
comprise a series of stable complexes, structurally well
characterized in most cases, whose electronic structures and many
physical properties can be interpreted in terras of a generic MO
description. As such they provide an excellent opportunity to
study the effect of reactivity at one metal center upon another.
This idea has been exploited by Drago and coworkers in the series
of coraplexes335 Ri^Cbut)^, Rh2(pfb)p Mo2(pfb)jj, Ru2(but)jjCl
and Cu2(hept)jj (but = butyrate, pfb = perfluorobutyrate, hept =
heptanoate). In these studies, stepwise adduct formation of the
Lewis acid dimer with first one, then two donor molecules was
monitored. Lower measured enthalpies for the second reaction
indicated a weakening of the second metal center acidity as a
result of base coordination to the first metal atom. By studying
a range of bases with each dimer, the enthalpy data were treated
in terms of the Drago-Wayland E and C equation.3639
empirical model3132 was pUt forwarci to describe the lowered
acidity in terms of the ability of the metal-metal bond to
transmit electrostatic (E) and covalent (C) effects. The shorter
quadruple Mo-Mo bond was found more able to transmit
electrostatic effects while the longer, more polarizable, single


9
Rh-Rh bond was found more able to transmit covalent effects.
Consistent with the generic MO scheme in Fig 1-2, the 14 electron
rhodium dimers, with filled tt orbitals, undergo tt-back bonding
to tt acids (e.g. CO, pyridine) while the 8 electron Mo2(pfb)/},
with no tt density, does not. The 11 electron Ru2(but)iiCl, with
half filled tt orbitals exhibited intermediate interactions.
As mentioned above, description of the metal-metal bond in
the chromium carboxylates has been the focus of recent
40
controversy. Ab initio calculations suggest that correlation
4 +
effects are very important in the description of the C^
complexes. These calculations involve description of the ground
state wavefunction of C^iC^CH)^ in terms of contributions from
PUP PUP
q tt 5 and excited states such as a tt 6 The contribution of
PUP
a tt 6 is 16% which contrasts markedly with similar calculations
for Mo2(02CH)jj in which the ct^tt**5^ configuration contributes 67%.
That is, the quadruple bond adequately describes the ground state
of Mo2(02CH)1j. It does not do so for C^CC^CH)^. deMello et al.
have suggested1* 1-1*^ that the dominant description of the bonding
in C^1*'1 is one of two Cr atoms antiferromagnetical ly coupled
with some contribution from multiple bonding. The chromium
carboxylate dimers are also unique in exhibiting a strong
dependence of the metal-metal bond length (and hence electronic
structure) upon axial ligation. Depending upon the nature of L
0
and R, molecules of the type C^CC^CR^L^ display a 0.57 A range
of bond lengths J from 1.97 A to 2.54 A. By contrast, adducts
of the Mo and Rh carboxylates display metal-metal bond length
O
ranges of 0.13 and 0.12 A, respectively. The sensitivity of the
metal-metal bond to donor molecule coordination may be manifested


10
in the way coordination at one chromium center affects the
second. A quantitative study of the coordination properties of
4 +
the Cr2 unit is the subject of chapter two. A literature
report of observed paramagnetism in the chromium dimer studied
here along with the antiferromagnetic description put forward by
4?
deMel lo et al. prompted a magnetic susceptibility investigation
which is also reported in chapter two.
Considerable effort has been expended upon the coordination
chemistry of the carboxylate diraer3 towards understanding the
nature of the metal-metal bondits electronic structure,
theoretical description and physical properties and reactivity.
The effects of varying both the axial ligand and the bridging
chelate on the metal-metal length bond have been extensively
explored.11 The perturbation of an axial ligand upon the metal-
metal bond has been termed a trans influence,11--*1 referring to
the influence of a ligand upon the bond directly trans to it. In
general, the axial ligand bond competes with the second metal as
a ligand, weakening the metal-metal bond; and conversely, the
stronger the metal-metal bond, the weaker is the metal ligand
interaction.*1 Another way of viewing the effects of axial
coordination is the influence of the ligand upon a second ligand
opposite the metal-metal bond. This secondary trans influence or
inductive effect can serve to identify the primary orbital
interactions in the metal ligand bond and the extent to which the
metal-metal bond transmits the ligand influence.
An infrared study of a series of L-Rl^Cpfb^-CO adducts and
an EPR study of a series of L-Rt^pfb^-TEMPO (TEMPO is the free


11
radical 2,2,6,6-tetraraethylpiper idine-N-oxyl) adducts has been
performed1^ and is further examined in chapter three.
A natural extension of these studies in the transmission of
bonding effects in dimers would be to consider trinuclear
complexes. Retaining the carboxylate ligand set still allows one
to choose from a diverse field of trinuclear complexes (trimers).
A desire to work in noncoordinating solvents and the need for
identical metal sites dictated the use of neutral, trigonal
complexes, exemplified by complexes of the basic iron acetate
structure.^ Q~-> 1 These complexes, of general formula
(M20(RC00)gL2)n+, where L is a neutral monodentate ligand,
contain a triply bridging oxide ion at the center of a (generally
equilateral) triangular array of metal ions; their structure is
illustrated in Fig 1-4. The electronic and structural details of
many of these compounds with M^^CnsI) have been studied. The
neutral, mixed valence compounds with n=0 have received special
attention as models for intramolecular electron transfer;
examples are known for M = Fe, Cr, Ru, Mn, V and perhaps Co.-^-^
A generic MO description for the mixed valence trinuclear
carboxylates is not available, and definitive MO calculations
remain prohibitive in light of their complexity. At least one
system, Ru20(02CCH2)5(PPh^)3, has been addressed from a LCAO
perspective, however.^ Understandably, assignment of the
electronic spectra of the trinuclear carboxylates remains
ambiguous. With regards to the neutral mixed valence trimers,
electronic spectra have been reported only for complexes of Co,
Fe, and Ru.-^--^ To gauge the effect of ligand substitution
reactions on their electronic structures, representative neutral


12
Fig 1-4. General trinuclear, oxo-centered, basic
carboxylate structure of formula [M20(RCC>2)6L2]n+.
metal


13
mixed valence trimers were examined in coordinating solvents and
the results are presented in chapter four.
The optical spectra of both neutral and cationic ruthenium
carboxylate trimers exhibit composite bands which originate from
a series of closely spaced molecular electronic transitions.^
The cluster system Ru2O(OAc)g(py)2^ + //2+,/+//0/,~ displays an
extensive reversible redox chemistry,^ prompting the name
"electron sponge". The stability of different redox states and
availability of substrate binding sites in the ruthenium trimers
are promising for homogeneous redox catalysis. The trirutheniura
acetate clusters have shown utility as homogeneous hydrogenation
catalysts for unsaturated substances. Attempts to employ
these same clusters as olefin oxidation catalysts in this
laboratory revealed the reversible reduction by alcohol solvent
at elevated temperatures. Subsequent specific catalytic
oxidative dehydrogenation of a range of alcohols and mechanistic
features were explored. These findings are given in chapter
five


CHAPTER II
REACTIVITY OF Cr^O^R)^
A. Introduction
The dimeric metal carboxylates, f^OgCR)^, are convenient
clusters for studying the effects of coordination at adjacent
metal sites. An extensive array of complexes has been isolated
and structurally characterized, synthesis methodology is
relatively straightforward, charge neutrality allows study in
non- or weakly coordinating solvents, and the metal centers
display open axial (trans to the metal-metal bond) coordination
sites to which Lewis bases readily bind.'1^ Work in this
research group has focused upon quantitative description of the
Lewis acid centers. It has been found that the enthalpy measured
for formation of the second metal-base bond is less than that for
the first metal-base bond in dimers of rhodium and molybdenum,
indicating that this may be a general phenomenon of this family
of compounds. By working with a range of characterized bases,
the enthalpy data could be treated with the Drago E and C
raodel.^~38 The empirical equation (2-1)
- AH + W = EaEb + CACB (2-1)
describes the enthalpy of adduct formation where EA and CA are
14


15
the acid parameters, and Eg and Cg are the base parameters
corresponding to the tendencies of the acid or base to undergo
electrostatic or covalent interactions. The W terra is included
when any constant contribution to the measured enthalpies
independent of acid or base variation accompanies adduct
formation. Drago, Long and Cosmano suggested an inductive
transfer model to describe the Lewis acidity of the second metal
center.3|t32 jn this model, the acid parameters of the 1:1
adduct are reduced from that of the free acid by an amount that
is proportional to the corresponding base parameter.
E
A
1:1
kEg
(2-2)
C
A
1:1
= C,
k'Cg
(2-3)
The k and kf have physical significance and represent the ability
of the metal-metal bond to transmit electrostatic and covalent
coordination effects. This description can be thought of as
parameterization of a trans effect. To date, four dinuclear
metal carboxylates have been studied and are summarized in Table
2-1. It should be emphasized that the intent of this methodology
lies not in determination of E^ and numbers per se (although
the experimentally determined numbers can be used to predict
unmeasured enthalpies), but their relative magnitudes serve to
illustrate the nature of the metal-ligand and metal-metal bonds.
For example, both Rl^tpfb^ and Mo2(pfb)jj exist in the +2
oxidation state and contain the same bridging ligand; a similar


16
Table 2-1.
dimers.
Acidity Parameters of various metal carboxylate
Acid
M2n+ dn BO Ea Ca CA/Ea k k'
Rh?(but)¡|
4 14 1 3.21 1.32 0.411 1.16 0.0364
Rh^Cpfb)^
4 14 1 5.06 1.74 0.344 a
MopCpfb)^
484 5.92 0.385 0.065 1.46 0.022
RupCbut^Cl
5 11 2.5 7.73 1.27 1.64 b
a.) Not determined though the butyrate and perfluorobutyrate
bridges were found similar in their transmission capability.
b.) Cl coordination precludes bonding of a second base.


17
partial positive charge exists at each metal center, and the two
dimers have similar EA numbers. The less electronegative bridge
in Rh2(but)2} results in a lower EA for this complex. The CA
numbers, on the other hand, reflect the polarizability of the
metal-metal bond. The quadruple metal-metal bond of Mo2(pfb)(j is
not as likely to redistribute electron density over the entire
molecule as the more flexible, single Rh-Rh bonds. The raolydenum
carboxylate has a lower CA.
Similar rationale lends physical significance to the
transmission coefficients, k and k'. The shorter metal-metal
bond in Mo2(pfb)^ allows for greater electrostatic interaction of
the second molybdenum with the first coordinated base and greater
interaction of the two base molecule dipoles. The greater
polarizability of the metal-metal bond in Rh2(pfb)1( allows for
better electron density redistribution as manifested by the
larger k' value.
The analogous chromium carboxylates have such a strong
tendency to coordinate electron pair donors in the axial
positions that they are only rarely seen without ligands.11 In
the two cases where unsolvated chromium dimers were studied
structural ly,^*^ axial coordination occurred by association of
the molecules to form infinite chains. The nature of the axial
44 64 6R
ligand has a pronounced effect upon the Cr-Cr distances * J
which range from 2.214 A in the (Cr-^CCO^^Ci^O)^11- ion to 2.541
A in Cr2(02CCF2)4(Et20)2. The experimentally observed range of
bond distances would suggest a shallow potential well for the
4.
Cr2 unit. The marked dependence of the electronic structure of


18
the Cr2i,+ unit upon the axial ligand provides a unique
opportunity to study transmission of bonding effects.
Reactivity studies of the dichroraium tetracarboxylates have
focused primarily on the acetate which has found wide utility as
a reducing agent and as a starting material in the preparation of
other compounds containing the unit. To this end, dimers
have been isolated and structurally characterized for a variety
of bridging ligands with C, N and 0 donor atoms and a range of
axial bases. Comparison of the structural parameters shows no
clear relationship between the nature of the bridging ligand and
metal-metal bond strength/length. Axial ligation, however, is
found to strongly influence the Cr-Cr bond with stronger donors
generally dictating longer bonds.11
Unlike the strongly bonded Mo-)1*'*' ion, there is no evidence
for the existence of the naked cluster Cr^*. Bridging ligands
may play a role in metal-metal bond formation other than keeping
the metal centers in close proximity; the (Cr^CH^g)*1"* and
(Cr2(C1jHg)i|)il ions exist without bridging ligands.1^6*67
The Cr2(tfa)iJ(Et20)2 adduct was first reported in 1966 to
exhibit weak paramagnetism. The structural report which followed
showed the dimer to contain the longest Cr-Cr bond known. Our
interest lay in probing the transmission of bonding effects
across such a weak, loosely interacting metal-metal bond, and
reactivity studies are reported here.
B. Results and Discussion
1. Qualitative Reactivity
Initial investigations were performed with the simple


19
carboxylates C^COAc^O^O^ and Cr2(but),(H20)2. The hydrates
are easily desolvated by heating in vacuo. An x-ray diffraction
study^3 0f anhydrous Cr2(0Ac)^ prepared by sublimation of the
hydrate demonstrated that bridging oxygens of neighboring
clusters satisfy the strong coordination requirements of the
Cr(II) centers. A portion of the polymeric compound which
results is shown in Fig 2-1; this compound is soluble only in
coordinating solvents. To minimize the nucleophilicity of the
carboxylate oxygens, the tr i fluoroacetate bridged dimer
Cr 2( t f a) i (E120) 2 was utilized. Again, the need for a
coordinating ligand attests to the Lewis acidity of the metal
atoms. Here, however, axial coordination is superceded by a weak
donor. Adduct formation in these studies proceeds via an
exchange reaction to displace diethylether.
Weak donors such as acetonitrile do not displace ether.
Intermediate donors such as dimethylacetamide cleanly displace
coordinated ether to give first 1:1, then 2:1, adduct formation.
Strong donors such as pyridine rapidly cleave the complex (vide
infra). A good measure of solvent donor capabilities come from
their E and C numbers; donors studied are summarized in Table 2-
2.
For the intermediate case, donor exchange is an equilibrium
process, readily monitored by changes in the electronic spectra
of the Cr2^+ chromophore. Analagous to earlier studies^ with
Rh2(pfb)jj and Rh2(but)jj, evidence for the formation of 1:1 and
2:1 adducts of Cr2(tfa)^ is provided by spectral studies.
Representative spectra of Cr2(tfa)j(Et20)2 and its adduct
exchange forms with DMA are shown in Fig 2-2 where the


20
y
/
\
o o
1/ 1/ /I
o Cr CrO
/ /I /i :c
Cr-\ /
/C
R
O \) Cr
i/ 1/ /
Cr Cr-. O
\
/
R
/I
Fig 2-1. The formation of infinite chains of 0^(0201?)^
molecules by oxygen bridge bonding.


21
Table 2-2. Donor parameters.
E
C
diethyl ether (Et20)
.963
3.25
do not displace Et20
acetone
.987
2.33
acetonitrile
.886
1.34
dimethyl sulfide
.343
7.46
methyl acetate
.903
1.61
tetrahydrothiophene
.341
7.90
triphenylphosphine
(a)
displace Et20
dimethyl acetamide (DMA)
1.32
2.58
dimethyl cyanaraide (DMCA)
1.10
1.81
dimethyl formamide (DMF)
1.23
2.48
dimethyl sulfoxide (DMSO)
1.34
2.85
dimethyl thioformamide (DMTF)
(a)
dioctyl ether (DOE)
1.10
3.40
p-dioxane
1.09
2.38
hexaraethylphosphoramide (HMPA)
1.52
3.55
triethylphosphate
(1.36)b
(1.81)
trimethylphosphine oxide
(1.53)b
(3.32)
trimethylphosphite
(1.03)b
(5.99)
tetrahydrofuran (THF)
.978
4.27
tetramethylurea
1.20
3.10
dissociate complex
diethylamine
1.17
8.51
N-raethyl imidazole
.934
8.96
pyridine
1.12
6.89
quinuclidine
.704
13.2
a) unknown
b) tentative parameters calculated from limited data sets


Fig 2-2. Spectrophotoraetric titration of chromium
trifluoroacetate, diethylether adduct, with DMA in The
free acid spectrum is labeled 0. The titration was ended with 50
molar excess added DMA, labeled X.


X (nm)


24
concentration of the dimer is kept constant and the donor
concentration gradually increased.
The presence of an isosbestic point at 572 nm is evidence for
only two species in solution at low base concentration: the free
acid, A, and the 1:1 adduct, AB. These spectral curves define
the equilibrium in Eqn 2-4.
(2-4)
A + B
AB
Further base addition results in spectral deviation from the
first isosbestic point as a third species is formed in solution,
Eqn 2-5.
(2-5)
AB + B
In these equilibria, the free acid, A, refers to the bis ether
adduct, Cr2(tfa)ii(Et20)2. Dissociation of coordinated Et20
accompanies both equilibria. At high base concentration, an
isosbestic point appears at 515 nm which upon cursory
examination would appear to correspond to the second equilibrium
process in Eqn 2-5. Quantitative analysis (vida infra), however,
reveals that the limiting spectrum centered at 585 nm
corresponds to the AB2 chromophore while the isosbestic point at
575 nm can be assigned to yet another equilibrium, Eqn 2-6.
(2-6)
The absence of a clearly defined isosbestic point for the second


25
equilibrium implies one of two things: 1) K^>> butthe
spectral change associated with Eqn 2-5 is slight, and an
isosbestic point cannot readily be discerned. 2) K1 is not
significantly greater than K2 and no appreciable amount of AB
forms in solution. The presence of a well defined isosbestic
point at low base concentration argues against the latter while
quantitative results (vide infra) indicate the former.
The observed spectral changes are understandable in terms of
the primary orbital interactions. Rice et al.^9 have examined
the single crystal polarized electronic spectrum of red chromous
acetate dihydrate, C^CC^CCH^C^O^ whose spectral features are
similar to that of Cr2(tfa)jj(Et20)2. Two bands are observed for
the former; the lower energy (465 nra, e = 120 M1 cm-1)
transition is associated with a metal centered 6 ->71* promotion;
the other (333 nra, e = 200) is attributed to charge transfer
from a nonbonding carboxylate tt orbital to metal centered tt *
#
(np^->Tr ). Violet Cr2(tfa)¡|(Et20)2 displays a similar spectrum.
The ->tt* (550 nm, e= 133) and nn ->ir* (328 nm, = 380)
^TT
transitions are assigned by analogy. An MO diagram of the metal
centered orbitals for C^itfa)^ and their changes upon adduct
formation are shown in Fig 2-3. The dimer functions as a Lewis
*
acid, accepting electron density in the antibonding a orbital.

Upon complexation of a stronger donor, the metal a and a
orbitals become closer in energy while adduct formation is
realized through stabilization of the donor lone pair orbital.
The dimer orbitals (6->tt*) involved in the transition are not
directly involved in adduct formation. Weakening of the metal-
metal bond through partial population of a* decreases the d


a *
IT* r
8*
8 t-l
77 44- -44-
o- 44--"
Cr2(tfa)4 2L Cr2(tfa)4 l_L L'
44-
44
44""
44-
N3
Fig 2-3. Effect of donor exchange upon metal centered MO's and corresponding change in optical transition.


27
orbital overlap, compressing the entire d orbital manifold in the
MO scheme. Replacing coordinated ether with stronger donors
gives a color change from violet to blue, consistent with the
expected red shift from the MO description.
Displacement of coordinated Et20 by stronger donors was
monitored by FTIR in order to verify the exchange processes.
Within the detection limits of the FTIR, coordinated Et20 was
completely exchanged for coordinated donor at 2:1 donortdiraer
molar ratio, indicating the exchange equilibria (Eqns 2-4 and
2-5) lie far to the right. Figure 2-4 shows an FTIR titration of
Cr2(tfa)¡|(Et20)2 with 0, 0.5, 1.0, 2.0, and 5.0 equivalents of
dimethylacetamide. Only above 2 equivalents is free DMA clearly
discernible. Figure 2-5 shows the ether vc_0_c region of the
same titration. Bound ether appears completely exchanged at 2
equivalents added donor. Some dissociation of diethylether may
occur even at 0 equivalents added donor as evidenced by a slight
shoulder absorbance at 1113 cm^. Only a slight frequency shift
is observed for vCQ of coordinated dimethylacetamide as the
titration proceeds. A similar result is observed for the
asymmetric v0_c_0 stretch of bridging trifluoracetate, vco2faSy:
equiv
DMA
VcoOMA)
" co2asy
0
1680.1 cm'
0.5
1610.0 cm-^
1680.3
1.0
1608.2
1682.1
2.0
1606.2
1684.0
5.0
1606.2
1683.9
free DMA
1651


28
(CM-I)
O
O 5
I O
20
5.0
Fig 2-4. FTIR titration of Cr-Ctfa^CEtgOjg (8.9 X 10"3M in
C H 2 C12) with 0, 0.5, 1.0, 2.0 and 5.0 equivalents of
dimethyl acetamide. Bound DMA: v co = 1160 cm-1; free: vco =
1639 era" .


29
Fig 2-5. FTIR titrationn of Cr
) with 0, 0.5, 1.0,
acetamide. Bound Et20:
= 1113 cm-1.
CH2C12
dimethyl
9(tfa)4(EtpO)2 (8
2.0 a nd 5.0 eq
v = 1053 cm"
v coc
.9 X 10"3 M in
uivalentsof
¡ free: vcoc


30
Coordinated DMA is relatively insensitive to the nature of the
trans coordinated ligand (Et2 or DMA). The slight shift which
occurs would indicate that DMA coordination is strengthened as a
stronger metal-ligand bond is formed on the opposite side. Such
a trans strengthening is surprising in light of results for the
Rl^pfb)^ dimer (chapter 3). Across the rhodium-rhodium bond, a
sigma donor on one side decreases the strength of a metal-ligand
sigma bond at the other side, as the two donors compete for the
same rhodium d orbitals.
The observed shifts for coordinated DMA are too slight to
resolve into their individual contributions from 1:1 and 2:1
adduct species. The FTIR experiment does provide the satisfying
result that the exchange process is indeed ocurring and that it
is essentially complete after two equivalents of donor has been
added at these chromium concentrations.
The trifluoroacetates are convenient spectroscopic labels and
display strong sharp absorbances for the asymmetric and symmetric
carboxylate stretches which have shown utility in differentiating
70
between unidentate, ionic, bidentate, or bridging coordination.1
The analagous Mo2(tfa)j displays the symmetric and asymmetric
stretches at 1592 and 1459 cm"'' (bridging). Upon adduct
71
formation, bulky phosphine donors occupy equatorial positions,'
and the resulting unidentate assymmetric stretch occurs at ca.
1680 cm -1. The IR spectrum of Cr2(tfa)j(Et20)2 (Fig 2-6) in
methylene chloride is assigned the bridging asymmetric and
symmetric stretches at 1680.1 and 1480.0 cm The similarity
of the molybdenum unidentate stretch (1680) and chromium


Fig 2-6. FTIR spectrum of Cr2(tfa)j(Et20)2, methylene chloride
solution.


9/o TRANSMITTANCE
800
2000.0 15000 1000.0
5000
WAVENUMBERS
LO


33
bridging stretch suggests the possibility of unidentate
coordination for the latter in methylene chloride solution.
Observation of only one asymmetric stretch (four unidentate
carboxylates would disociate the complex) and an identical Nujol
mull spectrum rule out this possibility. Though not detected at
five equivalents, a large excess (50 equiv.) of DMA results in an
FTIR spectrum in which another asymmetric carboxylate stretch
appears at 1717 cm"'. This, presumably, is the unidentate
carboxylate which arises from equatorial coordinationn of DMA and
is consistent with the formation of the AB^ species postulated by
Eqn 2-6.
Similar titrations were performed and spectra recorded for a
range of donors. Results for methylene chloride solutions of
p
Cr2(tfa)i|(Et20>2 (ca. 5x10 M) with two equivalents donor are
given in Table 2-3. Reported are the absorbances for the
appropriate functional groups of the free and complexed donor
molecules. Exchange adducts were readily isolated for several
donors by stoichiometric admixture and recrystallization from
benzene, and in all cases except that for DMTF, gave identical IR
spectra to those prepared from the same donor in situ. For those
donors whose functional groups showed a pronounced frequency
shift upon coordination, the FTIR spectra indicate the magnitude
of K1 and In all cases, no free donor could be detected at
the 1:1 level. With two added equivalents, the absorbance
spectra peak areas demonstrate 90% or greater complexation for
DMA, Et3P04, DMCA, DMSO, HMPA, and DMTF. The extent of exchange
could not be readily gauged for the remaining donors. A 1:1
equilibrium process (AB + BAB2) with initial concentrations


34
Table 2-3. FTIR data for trifluoroacetate bridges and donor
functional groups. v
CO2* asy CO2
functional
1 1
sym
group
free complexed
Rh2(tfa)4(EtOH)2a
1664
1467



Mo2(tfa)4b
1680 1592
1459



Cr2(tfa)4(Et20)2
1680
1480
o-c-o
1115
1053
(ch2ci2)
Cr,(tfa)4(Et,0),
1682
1484
o-c-o

1055
(mull)
(DMA)2
1717 1684
1476
c = 0
1651
1606
(Et^POjj) 2
1682
1478
P-O-C
1034,
,1037,
979
984
(DMCA),
1681
1477
C = N
2217
2239
(dmso)2
1680,
1478
s = 0
1057
1012,
1708
1022
(Me3PO)2
1681,
1473
P = 0
1179
1134
1712
(HMPA)2
1685
1473
P = 0
980
992
(dmtf)2
1679
1478
(d)
1538
1563
(dmf)2
1680
1477
c = 0
1676
1562
[(MeO)3P]2
1681
1477
(e)


(Me4Urea)2
1685
1477
C = 0
1640
1581
a) reference 70.
b) reference 71.
c) Slow oxidation follows complexation as evidenced by shift in
asymmetric stretch and color change from violet to green over 24
hour period.
d) assignment unknown.
e) Ligand vibrations shifted but specific shifts not assigned.


35
of 5x10-2M going to 90% completion yields K2 = 1800. Finally, as
with DMA, little or no frequency shift was observed for the bound
donor as the trans ligand was exchanged from diethylether to the
donor of interest.
An attempt was made to gauge the magnitude of the equilibrium
constants, and Kg, by calculating the mole fractions of bound
and coordinated diethylether from the fast exchange region 1H
NMR. The 100 MHz FT-NMR spectra of OgCtfa^EtgOJg and exactly
one equivalent added DMA and HMPA give the following diethylether
shifts in CgDg:
EtgO
5(-CHg-)
3.34(q)
6 (-CHO
1.07U)
Cn ^ ( tf3 )|(Et20) 2
5.22
1.40
+ 1 DMA
4.51
1.26
+ 1 HMPA
4.50
1.26
The chromium containing solutions all exhibit broadened
resonances, devoid of any spin-spin splitting. Treating the
coalesced resonance shifts as a mole fraction weighted average of
the free and bound species gives the same results for both HMPA
and DMA. Using the methylene resonances: 64% complexed, 36%
free; methyl, 58% complexed, 42% free. That different mole
fractions are calculated dependent on the resonance used is a
feature of the paramagnetic complex. Stronger donors raise the
paramagnetism of the Crgi*+ center (vide infra). The 1:1 adducts
then will display downfield shifted coalescence peaks.
Calculations thus favor a higher concentration of complexed ether
than is actually present. This effect should be more pronounced


36
for the methylene protons which experience a greater contact
shift. The only information these NMR 3pectra add to the
picture, then, is support for the paramagnetic nature of the
dinuclear metal center.
2. Quantitative Reactivity
Stepwise adduct formation between a metal center and first
one, then two donor ligands is described by two successive
equilibria (Eqns 2-4, 2-5). For the case where formation of the
first bond exerts no influence upon the second, A H-| = A H2 and
entropy considerations predict K-j = Inductive effects, as
in the case of the metal carboxylate dimers, perturb the second
equilibrium such that A H1 > A H2 and K1 > 41^. In this study,
adductformation is accompanied by a dissociation step,and the
reactions are described by exchange equilibria:
AL2 + B ABL + L (2-7)
ABL + B AB2 + L (2-8)
In subsequent discussion, the AL2 species is referred to a3 the
free acid and denoted by 0, the ABL as the 1:1 species (1) and
the AB2 as the 2:1 species (2).
a. Electronic spectra
UV-VIS titrations were performed by successive raicroliter
injections of a concentrated base solution into an inert
atmosphere cuvette containing the dimer solution. Despite the


37
presence of an isosbestic point for the initial spectral curves,
absorbance changes are too slight to allow satisfactory
definition of K.¡. In plots of K^ vs. Ae the resulting curves
intersect in a region of very small K_1 and small
absorbance/concentration errors can vary K by several orders of
magnitude. Consistent results are obtained, however, for Ae and
these values provide the initial estimates for the e at various
wavelengths.
Treatment of all the spectral curves using the program SPEC
allows the computer to calculate the best fit values of K-j, K2,
and £q, e .j and for the wavelengths analyzed (up to five).
The success of the fit relys upon good initial estimates of each
of these parameters. The Eg value comes from the free acid
spectrum. The e -j values can be calculated at any wavelength
from any curve that passes through the first isosbestic point,
once Ae has been determined from a K~1 vs. Ae plot at one
wavelength. The 2 values are estimated from the limiting
spectra which result from just above 2 added molar equivalents of
base. Specific details and precautions have been described. 0
The spectral absorbance data are listed in Appendix I.
Iteratively varying some of the spectral paramaters while
fixing the remainder allows the computer to uniquely define the
extinction coefficients and the picture which emerges in the same
absorption profile for the free acid, 1:1 and 2:1 species,
successively red-shifted. This is the expected behavior based
upon the MO arguments described earlier. Based upon the large
equilibrium constants (vide infra) these three species can be
observed in solution at the appropriate stoichiometry. The


38
UV-VIS spectra of the dimers with 0, 1, and 2 equivalents of
dimethylacetamide are shown in Fig 2-7, roughly illustrating the
spectral curves which the e 's define. Calculated values for the
equilibrium constants depend upon the initial guesses.
Consistently, allowing the computer to vary only K2 gives values
of 4K2 slightly less (ca. 50S) than K-j. Again, however,
absorbance changes are too slight to define both equilibrium
constants, and a wide range of K2 pairs satisfy the spectral
data with impossibly small standard deviations. If were
located, K2 could be uniquely defined, but the lack of a good
estimate for K-] prohibits their determination from the titration
curves. The spectral titrations indicate a relationship between
the equilibrium constants which would be consistent with weak
communication between the metal centers but do not adequately
define the magnitudes of the K's. The utility of these
experiments, then, lies in verifying that the exchange reactions
(Eqns 2-7 and 2-8) do occur, that their equilibria lie far to
the right and that the metal centers do not influence each others
coordination chemistry significantly.
Calorimetry
As Long has indicated,^ it is best to determine the
equilibrium constants which apportion the experimental heats and
the molar enthalpies from separate experiments, since the four
parameters, K^, K2, A H^.|, and A H2.1 are frequently highly
correlated. When this happens, a situation arises such as that
for the chromium spectral titrations above which a specific
solution relies on definition of (usually) one parameter.


39
Fig 2-7. Electronic spectrum of the chromium dimer (methylene
chloride solution) with 0, 1, and 2 equivalents DMA. Given the
large equilibrium constants for donor exchange, these curves
approximate the 0, 1 and 2 species.


40
Exceptions to the rule are instances in which the first or both
equilibrium constants are very large. Long showed that the
enthalpies for the rhodium systems, Rh^but)^ and Rh2(pfb)¡j,
could be reasonably determined from exclusively calorimetric data
Q
provided that was of the order of 10 or greater, as verified
by comparison of the enthalpies determined when the equilibrium
constants were solved for independently. A similar situation
occurs for the chromium calorimetric data, and the enthalpies are
extricable given the large equilibrium constants. Using the
program HEAT, values were fixed and the K2 varied to minimize
the conditional standard deviations associated with the
enthalpies. K-j was then varied by a factor of ten and the
minimization repeated. In this manner, the best K^t K2 pair
which minimizes the enthalpy deviationns was used to define
AH1#1 andAH2.^. The raw calorimetric data and best fit solution
values of each of the parameters are given in Appendix I. It is
not possible to assign deviations to the equilibrium constants
determined in this manner, and it was found that varying K^s of
this magnitude by a factor of 10^ had little effect (<1%) on the
enthalpies. Thus, though the enthalpies are uniquely determined,
is simply described as "large" and K2 represents the best fit
value. No inferences can be made for the relationship between K1
and K2 determined in this manner. Table 2-4 contains the results
of the calorimetric titrations.
Several aspects of the trends and some anamolies of the data
set deserve comment. First and foremost, the enthalpies of the
2nd exchange, AH2.^, are consistently lower than those for the
first, aH-|.-| in all cases except that for Me^PO. The lowering


41
Table 2-4. Thermodynamic data for the exchange reaction of
Cr2(tfa)jj(Et20)2 with various donors.
Base
K1
K2
- AH1
(kcal
; 1
mol"1)3
" A H2:1
(kcal
mol-1
DMTF
1 X
105
500
1.89
(0.06)
1.80
(0.12)
DMCA
1 X
108
3 X
103
3.75
(0.14)
2.66
(0.24)
DMF
1 X
108
1 X
104
4.52
(0.31)
3.01
(0.54)
(MeO)3Pb
1 X
108
2 X
103
0.79
(0.06)
0.73
(0.10)
Et^POjj
1 X
1010
1 X
106
4.37
(0.18)
3.05
(0.30)
DMA
1 X
109
1 X
105
5.26
(0.24)
3.59
(0.39)
DMSO
1 X
109
6 X
104
5.58
(0.27)
3.95
(0.45)
Me3PO
1 X
107
2 X
104
5.95
(0.26)
6.55
(0.43)
HMPA
1 X
1011
6 X
106
5.97
(0.23)
5.48
(0.37)
a. Values in parentheses are conditional standard
deviations.
b. Enthalpies are tentative; see text.


42
is slight but measurable, manifesting the inductive effect of the
first coordinated donor upon the second exchange reaction. As
expected, the inductive influence is small, in keeping with the
long, weak metal-metal bond in Cr2(tfa)2j(Et20)2* and is consistent
with the K^, K2 relationship inferred by the spectral data (vide
supra). Second, the trends in the enthalpy values is in keeping
with our intuitive knowledge of donor strengths, an aspect which
is taken advantage of in quantitative correlations to follow.
The higher >AH^.^ value for Me^PO is inconsistent
with the remainder of the data set and probably results from
subsequent further reactions. The spectral titration for this
system was well behaved only to a 2.0 molar ratio of Me^PO to
dimer at which point the solution began to turn green, indicating
oxidation of Cr(II). Heat evolution during calorimetric
titration for the other donors was essentially complete following
addition of two equivalents of donor. Heat continued to be
evolved in the Me^PO calorimetric titration which was carried out
to a 2.5 molar ratio. The stability of the dimer adducts toward
oxidation is decreased by donor complexation (vide infra) and the
C^ttfa^iMe^PO^ adduct may be readily oxidized via oxygen atom
transfer from the phosphine oxide.
Another inconsistency in the data set is the apparent
anomalously low heats observed for (MeCO^P despite the agreeable
AHi;i >AH2;i relation. Again, the spectral titration lends
insight. During intermediate stages of the titration (ca. 1:1)
greatly decreased (< 50%) absorbances were recorded in spite of
the presence of an isosbestic point at lower ratios. This
behavior would arise if 1) extinction coefficients for the 1:1


43
species were much lower than for the 0:1 or 2:1 species or 2)
solid formation is occuring in this region. The latter was
observed to be a common phenomenon in titrations with donors
containing two potential donor sites. Maintaining a strict 1:1
stoichiometry allowed preparation of polymeric species and
clearing of the solution for reactions with p-dioxane,
tetraraethyl urea and 1-methyl imidazole. Solid formation was not
obvious with (MeO)^P but could occur if both phosphorous and
oxygen donor sites are stronger Lewis bases than the ether
oxygen. A low value for AH-], for endothermic desolvation,
however, would be compensated by a high value of AH2 for
exothermic solvation, and this is not observed in the
calorimetric data. The low heats for (MeCO^P remain a surprise
and should be considered tentative.
Solution studies were conducted in methylene chloride solvent
which behaves as a Lewis acid towards donors. Solvation effects
are anticipated and can be corrected for since methylene chloride
has been shown to undergo primarily specific reaction (electron
pair sharing of a Lewis base with a Lewis acid) with donor
molecules.To assess their contribution to the experimental
heats, the enthalpy components are specified:
AH! AH2 AH3 AH4
(Et20)Cr2-(Et20) + B-S (Et20)Cr2-B + Et20-S (2-9)
ah5 ah2 ah6 ah4
(Et20)-Cr2B + 3-S^ B-Cr2B + Et20-S (2-10)
AH1;1 = AH4 + AH3 AH2 AH1 (2-11)
A H2.1 = A H4 + A Mg A H2 A Mg
(2-12)


44
Assuming nonspecific solvation differences between all the
various chromium adduct species to be negligible, the solvent
contributions, AH2 and A can be calculated from the Eg
and Cg values for the various donors and the E'A and C'A values
for methylene chloride, 1.66 and 0.01, determined by Drago et
al.^ The primes indicate that these are the best fit values for
methylene chloride adduct formation with a series of donors and
may contain small contributions from nonspecific solvation.
Solvation corrected enthalpies are given in Table 2-5. In three
instances, the donor Eg and Cg have been determined from a
limited data set and are not well defined. The corrected heats
for (MeO^P, Et^PC^ and Me^PO are thus considered tentative. To
alleviate propagation of their uncertainties, these three heats
will not be used in the following analysis.
Included in Table 2-5 are the experimental frequency shifts
from the spectral titrations. The observed red shifts indicate
the perturbation on the transition orbitals which results upon
donor exchange and provide another measure of the strength of the
interaction. The perturbations roughly follow the measured
enthalpies and include two donors not looked at by calorimetry.
Dimethyl sulfide titrations gave very small absorbance changes
and heats could not be accurately measured. Spectral changes are
apparent only at high donor concentration in which entropy
effects predominate and the weaker donor displaces ether to give
a blue shift. Me^Urea forms a polymer with the dimer at 1:1
molar ratios and redissolves as more donor is added, prohibiting
calorimetric determination of the enthalpies.


45
Table 2-5. Experimental frequency shifts and solvation corrected
enthalpies for the exchange reaction of Cr2(tfa)j(Et20)2 with
various donors.
-ah2;1
Base
(kcal mol^)
(kcal mol^)
Av.
Me2S
a

-480
Et20
0
0
0
DMTF
b

705
DMCA
3.96
2.87
710
DMF
4.96
3.45
770
(MeO)3P
C 0.13)G
(0.07)
1010
Et^POjj
(5.02)
(3.70)
1160
DMA
5.85
4.18
1160
Me^Urea
a

1270
DMSO
6.20
4.57
1285
Me3PO
(6.89)
(7.49)
1720
HMPA
6.90
6.41
2120
a. Nature of reaction prohibits enthalpy determination by
calorimetry.
b. Eg, Cg values unknown.
c. Calculated from tentative Eg, Cg numbers.


46
c. ECW model
We are now in a position to correlate the solvation corrected
enthalpies (corrected for AHg and AHj) whose contributions are
given below:
Ah.j..| = A A (solvent corrected) (2-13)
AH2.1 = AHg A(solvent corrected) (2-14)
For the observed 1:1 heats, AH^.-|, with a series of
different donors, the contribution A is independent of the
donor employed and must be treated as a constant, W. Using Eqn
2-1, where AH = AH^ and W = A Hp gives the equation:
-AH1;1 = -AH3 + Ah, = Ea1:1 Eb + CA1:1 CB (2-15)
Using the experimental solvent corrected heats for the six donors
whose Eg and Cg are well defined along with the Eg and Cg numbers
from Table 2-2 allows calculation of the constant contribution,
1 # 1 1*1
AH1fand the EA and CA associated with the acid which
defines Ah^. This acid is the 1:1 adduct (Et20)Cr2(tfa)j. A
plot of the simultaneous equations is shown in Fig 2-8. The best
fit values for EA, CA and W along with their standard deviations
are EA1:1 = 13.6 (0.75), CA1:1 = -1 .57 (0.24) and W = 8.00
(0.94), all in kcal mol The experimental enthalpies used in
the correlation along with those calculated from Eqn 2-15 are
given in Table 2-6 and demonstrate the quality of the fit.


Fig 2-8. Plot of Ea v3 CA for Cr2(tfa) jjCEtgO)


48
Table 2-6. Enthalpy data used to determine acid parameters for
Cr 2(tfa)jj.
Usedto determine Used to determine
E^:^ and C^:^ k and k'
Base
AH1:1,exp _AH1:1,calc aunexp ncalc
AAH,
AAH,
(kcal raol^)a (kcal mol-^)13 (kcal/mol)c (kcal/raol)d
Et20
0
0
0
0
DMCA
3.5
(0.1)
3.7
1.1
(0.3)
1.1
DMF
4.1
(0.3)
4.0
1.5
(0.6)
1.4
DMA
4.7
(0.2)
4.8
1.7
(0.5)
1.7
DMSO
5.0
(0.3)
4.5
1.6
(0.5)
1.7
HMPA
5.0
(0.2)
5.3
0.5
(0.4)
a.Solvation corrected values. Quantity in parentheses is the
conditional standard deviation for the measured uncorrected
heats.
b.Calculated from equation 2-15.
c.Difference between the first and second experimental
enthalpies. AAH = AH2>1 A H1.1 Quantity in parentheses is
the propagated conditional standard deviation, calculated from
a = a + a 0
B
d.Calculated from equation 2-17


49
Several points about these best fit parameters warrant
comment. The calculated W = AH1 = 8.0 kcal mol -1 refers to
breaking of one of the adduct bonds in (Et20)Cr2(tfa)j(Et20).
The relatively large EA^ and small CA1:^ demonstrate a
pronounced tendency for this Cr(II) acid center to interact
primarily in an electrostatic sense as might be expected for a
first row transition metal ion. The negative CA1:1 value does
not imply that the dimer interacts covalently in an antibonding
sense but simply that this is the best fit parameter to a data
set in which implicit assumptions have been made in defining the
magnitudes of the original parameters.3 For all intents and
purposes, the CA1:' value merely suggests little or no covalent
interaction between the (Et20)Cr2(tfa)j adduct and donors.
The real quantities of interest for comparison to other dimer
systems are the acid parameters associated with the naked dimer,
C^Ctfa)^ and are extricable from an analysis of the 2:1
enthalpies. The EA1:1 and CA 1:1 values are defined by Eqns 2-2
and 2-3.
EA1:1 = EA k EB (2-2)
CA1:1 = ca k' cb (2-3)
Thus, determination of the inductive transfer parameters, k and
k', and using the base parameters for Et20 allows calculation of
Ea and CA for C^Ctfa)^. By substituting Eqns 2-2 and 2-3 into
the E and C equation, 2-1, the first and second enthalpies are
32
related to each other by


50
- AH2;1 = -AH1;1 -kEB2 k'CB2 (2-16)
In the present study, where an exchange reaction is being
studied, Eqn 2-16 takes on a slightly different form. Defining
all the enthalpy components AH-|, AH^, AH^ and AHg in terras of
their E and C components results in many cross terras which define
the perturbation that coordinated ether makes on the first
exchange reaction and that coordinated base, B, makes on the
second. The derivation is given in Appendix III and the
simplified solution has the familiar form
- AH2;i = -AH1s1 k(Eg EEt20)2 k?^CB CEt20^2 t2-1?)
This equation has the same form as Eqn 2-1 and the ECW
program is used to solve for k and k'. The input data are the
heat differences as AAH, the squares of the quantities in
parentheses as the base parameters and W is, of course, zero.
The experimental heat differences, AAH = AH7.^ AH^.^, used
to determine k and k, along with those calculated from the best
fit parameters are given in Table 2-6. The attendant conditional
standard deviations are given in parentheses. The value for HMPA
was rejected in light of the large standard deviation and k and
k' were determined from the remaining five data points. The
best fit values (standard deviation) are k = 1.54 (0.13) and k* =
0.0079(0.0043). These are the parameters associated with the Cr-
Cr bond and demonstrate the ability to transmit electrostatic (k)
and covalent (k') effects. Equations 2-2 and 2-3 may now be used


51
to calculate the EA and CA values for the naked chromium dimer,
Cr2(tfa)4. The derived values are given in Table 2-7 along with
those for the molybdenum and rhodium dimers for comparison.
Again, the large EA and very small (negative) CA demonstrate
the pronounced Lewis acidity of the chromium dimer and its
tendency to interact with Lewis bases in primarily an ionic
fashion. This finding is consistent with the lack of
experimental success towards generating a dinuclear chromium
carboxylate free of axial coordination.11 In the absence of
donor molecules, the chromium carboxylates exist as polymers
where the very ionic carboxylate oxygens of neighboring dimers
serve as donors.
The low k* value which corresponds to transmission of the
base covalent parameters indicates that base binding does not
serve to polarize the bonding density of the chromium-chromium
bond, consistent with the poor orbital overlap between the two
metal centers. The rather large k value which corresponds to
transmission of the base electrostatic parameter is rather
surprising. Electrostatic interaction of the base dipole on one
side with the second metal center and with the second base dipole
p
should vary as 1/r Of the three systems studied, the chromium
dimer exhibits the longest metal-metal bond (2.54 A for1121
Cr2(tfa)4(Et20)2, 2.39 l for73 Rh2Ac4(H20)2 and 2.09 for73
Mo2(tfa)4) and should exhibit the smallest value for k in the
series. For comparison purposes, it is probably more appropriate
to consider the relative values of k for each metal system. The
k values quantify the change in EA at the second metal center as
a consequence of bonding a donor at the first, irrespective of


52
Table 2-7. Acid parameters for various dinuclear carboxylates.
ea
CA
k
k'
Rl^Cbut)^
5.06
1.74
1.16
0.0364
Mo2(pfb)|(
5.92
0.385
1.46
0.022
Cr 2 (t-f 3) i
15.1
-1.54
1.54
0.0079


53
the magnitude of EA. The roughly twice as large EA for the
chromium system would display relatively half as large
perturbations on the magnitude of EA compared to the other two
dimers. This insight, then, is consistent with the longer metal-
metal bonds in the chromium dimers. Thus, more appropriate
measures of the ability of a metal-metal bond to transmit
coordination effects would come from k/EA and k'/CA instead of
comparing the absolute k and k* values.
d. Magnetic susceptibility
Theoretical investigations on dichromium tetraformate at the
SCF-level have predicted both the quadruple bond, 0^71^6^, and
P p # p # p
no bond, a 6 a configurations, neither of which
correspond to realistic descriptions.^ Incorporation of
correlation effects which allow mixing in of excited states in
P ft
the ground state description give more satisfying results^0 in
P P
which the term is important, but the ground state bond
order is closer to 1.5. Dichroraium tetraacetate is diamagnetic^
and deMello et al. have proposed two antiferromagnetical ly
coupled chromium centers and no net covalency between the
4 +
chromium atoms as the dominant description for Crg centers.
This latter description, too, suffers from inconsistencies by
predicting unrealistically long Cr-Cr distances. At shorter
(1.8-2.5 t) experimental bond lengths, d orbital overlap is
expected, and Zerner qualifies his calculations to suggest that
some degree of covalency accompany the antiferromagnet ic
description.
The trifluoroacetate bridged dimer studied here may represent


54
just such a borderline example with a chromium-chromium
separation of 2.54 A. An early magnetic study' reported
magnetic moments of u 0.74 BM for "CrCF^CCOO^" and 0.85 BM
for "CrCF^CCOO^.Ather". These species are presumably dimeric,
and the increase in magnetic moment upon ether complexation is
consistent with both an antiferroraagnetic description and the
proposal that ligand donors weaken the metal-metal bond through
partial population of the a* orbital.46,77 oniy other first
row transition metal for which a range of metal carboxylate
dimers has been prepared is copper. The cupric carboxylate
dimers exhibit long (2.6-2.9 A) metal-metal distances and
incomplete spin pairing between the two d^ centers. An
impressive array of these complexes^ has been prepared and
characterized magnetically and structurally, with a general goal
of determining the factors (structural parameters; nature of
bridge; axial ligands) and mechanism which contribute to spin
exchange. The following conclusions appear to be general for the
cupric carboxylate dimers:
1) the unpaired electrons reside in the dx2_y2 orbitals;
overlap is minimal at these metal-metal distances, ruling out any
direct exchange or covalency;
2) overlap with carboxylate tt orbitals allows a super
exchange pathway via the carboxylate bridges, resulting in
antiferromagnetic interactions with -2J ranging from 217 to
555 cm-1;
3) the singlet-triplet separation (-2J) is relatively
insensitive to the Cu-Cu distance;
4) -2J is sensitive to the Cu-0-C-0-Cu bridge distance and


55
angle, generally decreasing at longer distances;
5) -2J is sensitive to axial ligation, generally increasing
as the terminal ligands become stronger electron pair donors.
The chromium dimers, on the other hand, exhibit some degree
of orbital overlap, suggested to be incomplete in the
Cr2(tfa)Jj(Et20)2 adduct. The strong dependence of the Cr-Cr
distance upon axial ligation suggested magnetic studies to
complement the copper studies in treating a metal-metal bond with
some degree of covalency.
The room temperature magnetic susceptibilities of various
adducts of the form C^Ctfa^I^ were investigated by the solution
Evans method^ by generating the complexes in situ. This method
allows calculation of experimental magnetic susceptibilities by
measuring the proton chemical shift of an inert substance in the
presence of a paramagnetic complex. A coaxial tube arrangement
was used which contained a solvent system of 2% v/v CgHg and 2%
v/v TMS in CgDg in both the outer 5mm tube and the inner 1mm
capillary. The inner capillary also contained the paramagnetic
species at a concentration of ca. 5x10 M. For an inert
substance (in this case, CgHg and TMS) the shifts caused by the
paramagnetic substance when one employs a nonsuperconducting NMR
instrument are given by
AH = 2 7T (2-18)
I 3 4K
where Ak is the change in volume susceptibility. With the
cocentric tube arrangement, two resonance lines will be obtained


56
for each standard, with the line from the more paramagnetic
solution lying at higher field. The mass susceptibility, x of
the dissolved substance is given by Evans as
X
3Av
2irvm
Xo
+ Xo(d0-d )
m
(2-19)
where Av is the frequency shift separation in Hz, v is the
spectrometer frequency (99.55 MHz), m is the mass of substance
contained in 1 ml solution, xo is the mass susceptibility of
the solvent, dQ is the density of the solvent and ds is the
density of the solution. Brault and Rougee have presented0 a
modified form of Eqn 2-19 to calculate the molar magnetic
susceptibility
XM
3Av 1000 +
2ttv c
XoM
X
D
(2-20)
where C is the molar concentration of the paramagnetic complex, M
is the molecular weight of the complex and xd i s the
diamagnetic susceptibility of the paramagnetic complex. As
Desmond points out,^ the original Evans equation required noxn
term since complex dismagnetic susceptibility was compensated for
by including an equal concentration of ligand in the reference
solution. The Brault and Rougee equation is more general but
fails to account for density difference contributions. The
density term can often be neglected but becomes important when
measuring the small shifts for weakly paramagnetic substances.
Desmond includes the density term to obtain the equation


57
= 5Av 1000. + y m y + Xo(dp-d ) 1000, (2-21)
XM 2ttv c Xo xD c
The Xd term is usually approximated from Pascals constants^
and may be quite erroneous for large molecules. To alleviate
this hazard and the need for a density determination, the
diamagnetic correction may be obtained by NMR. Setting = 0
and moving the first two right-hand terras of Eqn 2-21 to the
other side gives an expression for the combined diamagnetic
susceptibility and density terms. An Evans experiment performed
on the diamagnetic analogue Rl^Ctfa^CTHF^ gave no peak shift
nor was any asymmetry evident in the TMS or peaks. The sum
of the diamagnetic term and density correction then is -5.63 x
10-i< (obtained by setting both X^ and Avequal to zero). The
calculated diamagnetic term is -3.0 x 1 0_i* (from Pascals
constants, where each Rh(II) is given a value of -20 x 10^ ml
raol-^), and the difference approximates the density correction.
An upper error limit on the shift determination is about 0.3 Hz
which corresponds to a 0.43 x 1Q~^ contribution to the
diamagnetic/density term. Substitution of the value for the
combined complex diamagnetism plus density correction into Eqn 2-
21 gives
xm = : Av T- + XoM + 5-63 x 10"4 (2-22)
At the concentrations used, complex precipitation warranted shift
determination in CD2CI2 in several instances. Equation 2-22 was
used to calculate molar magnetic susceptibilities from both the


58
C^Dg and CD2CI2 solution data. The single greatest source of
error in these measurements probably comes from determining the
shift magnitude. If the error calculated for the diamagnetic
correction is doubled to account for dual shift measurements (one
for the rhodium complex and one for the paramagnetic complex) the
attendant error is estimated to be 0.86 or about 1 x 10*"^.
Use of the "spin only" formula for the molar susceptibility
allows determination of the effective magnetic moment
P eff = 2.84(x mT)1/2 (2-23)
where the product is expressed in Bohr magnetons. X m is a molar
quantity, calculated per mole of dimer, while yis per metal
and must be determined from 1/2 X M.
The molar magnetic susceptibilities and moments calculated
from the observed frequency shifts are reported in Table 2-8.
Benzene solution was the method of choice (economy), but some of
the dimer adducts were too unstable towards oxidation in this
solvent, necessitating determination in methylene chloride.
Immediately obvious is a very satisfying trend towards larger
susceptibilities and moments with increasing donor strength.
This relationship provides the first conclusive evidence for
genuine paramagnetism in a series of chromium carboxylates since
this consistency would not be observed with random trace
contamination by Cr(III) impurities.11
The measured susceptibilities roughly parallel the
experimental enthalpies except in one instance. The DMCA adduct
displays a larger moment than might be expected from the heat


59
Table 2-8. Magnetic susceptibilities and moments for C^tfa^L^
species.
conc(mM)
Av(Hz)
XMax104
yeff(B.M.)
Rh2(tfa)4(THF)2
26.2
0
0
0
Cr2(tfa)4(Et20)2
28.4
3.54
6.66
0.89
[(MeO)3P]2
46.9
6.77
6.91
0.91
c
50.4
7.32
6.95
0.91
(Et3P04)2d
61.3
13.24
10.94
1.15
(DMTF)2d
62.1
13.98
12.40
1.22
(Me4Urea)2
51.1
15.38
14.53
1.32
(DMF)2d
21.3
17.64
18.18
1.48
(DMS0)2d
62.2
24.66
20.73
1.58
(DMA)2d
62.1
24.84
20.80
1.58
(DMCA)2
27.19
20.50
36.90
2.11
c
88.31
62.20
34.52
2.04
(HMPA)2
17.7
17.64
47.01
2.38
(MenPO)o
20.9
33.20
77.86
3.06
a) per dimer, + 1 X 10\
b) per Cr(II) atom, + 0.35 B.M.
c) multiple determinations.
d)determined in CD2CI2


60
data. Subsequent redetermination gave essentially the same
value, however, and both are reported.
Besides the chromium and widely studied copper systems
reports of two other discrete first row transition metal
carboxylate dimers have appeared, both of which display weak
antiferromagnetic exchange interactions. Dicobalt (II)
tetrabenzoate bis quinuclidine,-^ Co2(CgH^C02)j(quin)2 shows a
O
long Co-Co distance of 2.83 A and the high spin Co(II) centers
display weak anti ferromagnetism of -2J = 38 cm^. At this long
metal-metal distance, spin coupling is expected to occur via the
bridge super exchange pathway. A vanadium complex,^
V2(tfa)i|(C5H^)2 exhibits a V-V distance of 3.7 A and weak
antiferromagnetism, suggested also to proceed through a super
exchange path involving the carboxylate bridges.
Treatment of the exchange mechanism operative in the chromium
dimers is necessarily more complex than for the copper
Q|l
carboxylates, the former involving S = 2 metal centers. MO
calculations support direct orbital overlap and theoretical
treatment would include both direct and super exchange pathways.
Calculations showing low lying excited states do not rule out the
possibility of thermal equilibrium between the populations of
ground level and the first excited level with other multiplicity
(singlet-triplet equilibrium).^ The latter ha3 not been
demonstrated in the binuclear carboxylates but occurs in
monomeric complexes when the ligand field splitting is close in
energy to the electron pairing energy. The trend towards larger
moments with increased ligand field, however, is not consistent
with a singlet ground state.


61
Determination of the singlet-triplet splitting (-2J) can only
come from measurement of the temperature dependence of the
magnetic susceptibility, and these data are not yet available for
the chromium trifluoroacetate dimers. The data in Table 2-8 do,
however, suggest a direct exchange mechanism. In modeling the
magnetic susceptibility data for some 140 compounds of the
binuclear copper carboxylates, Jothara and Kettle observed a
general trend for J to increase as either the terminal or the
bridging ligands become better electron donors.At a given
temperature, larger values of -2J would give smaller experimental
moments. Thus in these systems in which the super exchange
pathway predominates, better donors give smaller magnetic
moments. The reverse is true for the chromium series studied
here. Donor lone pairs interact primarily with the chromium
dimer a orbitals. Stronger donors would serve to destabilize
the Cr-Cr a orbital while stabilizing the Cr-Cr cr* orbital, and
partial population would weaken the covalent bonding, consistent
with the observed trend for larger moments with better donors.
e. EPR, Electrochemistry
The cyclic voltametry of Cr2(tfa)j(Et20)2 was attempted with
the aim of finding the proper conditions for controlled oxidative
electrolysis. An EPR investigation of the cationic product might
add insight into the nature of the chromium-chromium bond. The
EPR prarmeters would be descriptive of the HOMO orbital which
surrendered the electron. In methylene chloride, the dimer was
found to be incompatible with the electrolytes Bu¡jNBFjjf Bu^PF^


62
and BujjNI, all of which reacted to give blue or green solutions.
No reaction was observed with Bu^NClO^. No oxidation wave could
be detected in solution, however, with this electrolyte though a
slow irreversible reduction occurred at -0.45 V (Ag/AgCl). No
further attempts were made to oxidize the complex.
The magnetic studies suggested unpaired electron density at
the Cr(II) centers, prompting investigation by X-band EPR. The
solid state adduct Cr2(tfa)2|(Et20)2 displayed no EPR down to 10K.
In CH2C12 solution, the ether and HMPA adducts gave no signal
down to 85K. One attempt was made to record the solution
spectrum of Cr2(tfa)ij(Et20)2 at liquid helium temperatures. A
1x10 M CH2CI2 solution glass gave the spectrum in Fig 2-9. The
observed asymmetric signal displays no ^Cr (9.5%) hyperfine
splitting and would appear to be that for a system of axial
symmetry with g^ = 1.98 and g1 1 = 1.87. A forbidden S = 2
transition may be present at H/2 (1600 G) but this is speculative
in lieu of experimental clarification. Lack of any observed
signal for a non-dilute powder sample at the same temperature is
consistent with significant spin lattice relaxation while lack of
any EPR at 77K (methylene chloride glass) would indicate
extensive spin-orbit coupling. EPR in d1* systems is very rare
due to short spin lattice relaxation times and a large zero
field splitting. A comprehensive review of the literature,^
albeit in 1972, yielded only five d1* systems for which an EPR had
been reported; three of which were Cr(II) as dilute ionic salts.
The estimated g values reported here do not, however, resemble
any of those reported for Cr(II) or Cr(III) systems. This


Fig 2-9. The X-band EPR spectrum (CH2CI2 glass, 9K) of
Cr2(tfa)i|(Et20)2. The field ranges from 1-5 kG.


64


65
finding warrants further investigation and is reported here only
as a matter of record.
3. Further Reactivity
Chromium (II) salts have been extensively employed as
O O
reducing agents in preparative organic chemistry.00 Chromium
(II) chloride, sulphate and perchlorate are similar in their
scope; chromium (II) ethylenendiamine cation functions as a more
efficient reducing agent while chroraous acetate is a milder
reducing agent, reacting under relatively neutral conditions. By
contrast, the utility of chromous trifluoroacetate as a reducing
agent remains relatively unexplored and the ligand exchange
reactions explored here offer some insight into its redox
chemistry.
Reduction of alkylhalides by chromium (II) salts to generate
olefins, carbenes, or alkyl chromium species via haloatom
abstraction represents one of the best known applications of
these reagents. Methylene cloride solutions of chroraous
trifluoroacete were found to be stable for several months under
dinitrogen as evidenced by retention of the robust purple color.
g
Irradiation (550 nm, 6-- tt ) of solutions for 30 minutes in the
UV-VIS beam path also showed no appreciable spectral changes.
As mentioned for the case with DMA, UV-VIS titrations were
well behaved, defining a two step exchange process. Deviation
from the 2:1 limiting spectrum with excess base (> 5 equiv.) was
apparent in all cases and was accompanied by an immediate blue to
green color change which became more pronounced with additional
added base. FTIR titrations revealed a new vCo2,asy at 1?05


66
to 1710 cm-1 with excess (> 2 equiv.) donor which was assigned to
raonodentate trifluoroacetate. Chromium (II) oxidation would thus
appear to be facilitated by bridge removal and equatorial donor
coordination. Additionally, except for C^Ctfa^l^ species
where L = EtgO, DMTF or Et^POjj, the final blue solutions from
both spectral and calorimetric titrations turned green within 2
to 24 hours despite rigorous exclusion of dioxygen, indicating
oxidation when excess donor was present in solution. These
observations are consistent with electrochemical measurements on
a series of tetracarboxylato-dirhodiura (II) complexes.^ Das et
al. found that electron withdrawing substituents on the
carboxylate bridges produced less easily oxidized dimers. In
their studies, the lower oxidation state of Rl^itfa)^ was
stabilized to such an extent as to see no oxidation step within
the solvent limits. A similar trend in solvent effect was
observed with coordinating solvents such as pyridine and DMSO
giving the most negative oxidation potentials. The chromium
dimers are stable indefinitely in solution if stoichiometric
amounts of base are present, allowing ready isolation of the
exchange species.
A crystal structure has been reported-^ for the reaction
product of the strong donors pyridine and 4-cyanopyridine with
Cr2(02CCF2H)|j(Et20)2. The isolated crystals are trichromium (II,
III, III) compounds with the basic iron acetate trimer structure.
No spectral results were reported. The same procedure was
followed using the trifluoroacetate and pyridine to give long (up
to 2 cm) pale green needles which turn dark olive upon exposure
to air. Assuming the same structure, the bridging


67
trifluoroacetates show a vco2>a3y at 1671 cm-1 in the mixed
valence trimer. Attempts to induce triraer formation with the
donors DMA and DMCA were unsuccessful, surrendering no solid
products after three weeks. Both the blue DMA-containing and
dark green DMCA-containing solutions gave VC02 aSy peaks at
1717 and 1713 cm-^, respectively, indicating monodentate
trifluoroacetate. Triraer formation does not appear to be
general, with most donors simply inducing slow oxidation of
chroraous trifluoroacetate, and the effect is more pronounced in
benzene than in methylene chloride.
C. Conclusion
A thorough investigation of the coordination chemistry of
chroraous trifluoroacetate has been performed with the aim of
understanding the transmission of bonding effects through the
weak Cr-Cr bond and the perturbation of donor ligands on the
metal-metal bond. Spectral titrations indicate that donor
ligands decrease the d orbital overlap between the metal centers.
Calorimetric studies show a lower enthalpy for second base
coordination resulting from transmission of donor effects from
one metal center to the next. The effect is less pronounced than
in the previously studied Rh2(but), and Mo2(pfb)j systems. A
correlation analysis of the calorimetric data allows description
of the Cr(II) Lewis acidity as a relatively strong acid,
interacting in primarily an electrostatic fashion with Lewis
bases. Excellent data fits support an inductive transfer model
used to describe communication between the metal centers.


68
Magnetic susceptibility measurements on a range of Crjitfa^L^
adduct species reveal a pronounced influence by donors upon
dimer paramagnetism and support a direct exchange pathway for
spin pairing. UV-VIS, NMR and IR data support the donor exchange
reactions proposed and reveal a destabilization towards
1 .
oxidation when stronger donors coordinate to the center.
D. Experimental
1. Data Analysis
Programs. The following computer programs were used for
analysis of raw spectral and calorimetric data.
AB
£ O
Written by J. R. Long, 0 this program utilizes a standard non
linear least squares routine to provide the best Ae(ei e0) arKl
K which describe a 1:1 equilibrium, A + B i=?AB. Raw data needed
are concentrations and absorbance changes for a spectral
titration. Raw heats can be used to solve for AH and K for a
calorimetric titration.
SPEC
Written by T. Kuechler,^ SPEC utilizes concentrations and
spectral changes to solve for the best K-j and K2 via least
squares for a two step equilibrium. The large number of unknowns
solved for (K^, K2, and Eg, e^, at each wavelength used)
requires good initial estimates of each of these parameters to
avoid false minima. Useful algorithms for initial estimates are


69
provided by J. R. Long.6 In practice, best results are obtained
by fixing most of the parameters (usually the e's) and allowing
the computer to vary those that are least well known (usually the
K's).
HEAT
ft
Written by J. R. Long, 0 this least squares program solves for
the best fit molar enthalpies, AH^ and AH2, for a two step
equilibrium. Input includes concentrations, raw heats and
equilibrium constants for a calorimetric titration.
Error analysis. Output parameters from the above programs
are provided with marginal and conditional standard deviations,
MSD and CSD, which demonstrate how well the model fits the data.
The conditional standard deviation defines the magnitude of error
while the MSD/CSD ratio indicates how well defined the parameter
iS>68,90 For a rea3ona5ie CSD, results are considered meaningful
if the ratio is less than 4, tentative if the ratio is between 4
and 12, and not meaningful if the ratio is larger than 12.
ECW Program. Revised by M. K. Kroeger, this program uses
experimentally determined enthalpies to calculate the best fit
values of the unknown parameters EA, CA (and W if necessary) to
the following equation.
-AH + W = EAEB CACB


70
The Eg and Cg are found in references 38 and 39. Alternatively,
measured heats and and known EA, pairs can be used to solve
for Eg, Cg. A description of the computer program has been
Q1
previously reported.^
2. Materials
Metal complexes. All syntheses and manipulations were
performed under dinitrogen using Schlenk techniques or an inert
atmosphere box. Chroraous complexes are generally quite oxygen
sensitive and the compounds used here are oxidized within seconds
upon air exposure.
Dichromium (II) tetrakistrifluoroacetate bisdiethylether, _I.
Chroraous acetate was generated from Zn/Hg reduction of Cr(III)
(aq) and sodium acetate.^ Yellow chromous carbonate was
synthesized through exchange of the acetate bridges by reaction
with potassium carbonate in water.^3 Chroraous trifluoroacetate
was prepared by a modified literature procedure.Under
dinitrogen, 5.5 g (9.9 mmol) KjjC^CO^ij^O)^ and 8.0 ml (100
mmol) trifluoroacetic acid were refluxed in 80 ml deaerated
diethylether for six hours. The Schlenk flask was plunged
briefly into a dry ice bath, the purple ether layer decanted from
the frozen blue aqueous layer and the ether removed by vacuum.
Extraction with benzene followed by two recrystallizations from
benzene and vacuum drying (30 rain, 25) produced purple blocks of
C^itfaJjjiEtgO^. Prolonged evacuation 06 hr) was found to
strip off coordinated ether. Final yields were typically about


71
2O. Attempts to purify by sublimation decomposed the complex.
Repeated elemental analyses typically showed loss of 510%
coordinated diethylether. Calculated for ci6H20Cr2F1210:
C, 27.29; H, 2.86; Cr, 14.77. Found: C, 25.91; H, 2.82;
Cr, 15.01.
Dichromium(II) tetrakistrifluoroacetate bisdimethylthioform-
amide. I, 34 rag (0.048 mmol) was dissolved in 1 ml benzene.
Addition of 9.0 y1 (0.11 mmol) DMTF produced a clear solution and
a fine blue powder within ca. 5 min. The solid was filtered,
washed once with cyclohexane and vacuum dried (25) one hour.
Dichromium(II) tetrakistrifluoroacetae bis dimethylformamide.
I_, 37 nig (.053 mmol) was dissolved in 1 ml benzene. Addition of
9.0 y 1 (0.12 mmol) DMF produced a clear solution and a fine blue
powder within ca. 5 rain. Solid was filtered, washed once with
cyclohexane and vacuum dried (25) one hour.
Dichromium(II) tetrakistrifluoroacetate bis trimethylphosphite
I_, 33 mg (0.047 mmol), was dissolved in 1 ml benzene. (MeO^P, 12
y1 (0.10 mmol), was added to give a slightly bluer solution.
Stripping off the solvents by vacuum produced a blue violet
powder which was not treated further.
Dichromium(II) tetrakistrifluoroacetate bis tetramethylurea.
I, 222 mg (0.32 mmol), was dissolved in 20 ml methylene chloride
and 76 yl (0.64 mmol) tetramethylurea added to give a


72
blue solution. Solvents were stripped off, the solid extracted
with benzene and the volume reduced to 5 ml. The solution was
suspended in a dewar above ice. The rectangular dark blue
crystals which formed after three days were collected and vacuum
dried (25) one hour. This compound slowly decomposes in the
solid state, turning green after about one month under
dinitrogen.
Dichromium(II) tetrakistrifluoroaectate bis Triethylphosphate.
_1, 490 mg (0.70 mmol), was dissolved in 20 ml benzene and 0.35
ml (2.1 mmol) triethyl phosphate added to give a deep blue
solution. Volume was reduced to 5 ml and cooled to 5 for 24
hours to produce a mass of dark blue cubic crystals. Solid was
collected, washed with cyclohexane and vacuum dried (25) 30 min.
Calculated for C2oH20Cr2rl 216P2: c 26.10; H, 3.29; F, 24.77.
Found: C, 26.11; H, 3.40; F, 24.44.
Dichromium(II) tetrakistrifluoroacetate bis hexamethylphosphor-
amide. _I, 525 mg (0.74 mmol), was dissolved in 20 ml benzene and
0.30 ml (1.7 mmol) hexamethylphosphoramide added to give a deep
blue solution. Solution was cooled to 5 and bright blue cubic
crystals began to form after 15 rain. After 24 hours, crystals
were filtered and dried under a stream of dinitrogen. Calculated
for C20H36Cr2F12N6O10P2: C, 26.27; H, 3.97; F, 24.93; N, 9.19.
Found: C, 26.34; H, 3.99; F, 24.38; N, 9.21


73
Attempts to isolate solid adducts of dimethylacetamide and
diraethy 1su1foxide by similar procedures in benzene were
unsuccessful, producing a pale green gel in both instances.
Cr30(02CCF,)6(C5H5N),. I, 140 mg (0.20 mmol), in 5 ml EtgO was
placed in a test tube. 4 ml of hexane was layered onto the ether
and 0.8 ml (9.9 mmol) pyridine in 8 ml hexane added to the hexane
layer. After four days, solutions were almost clear. Large (1 X
1 X 20ram) pale olive needles had grown down from the solvent
interface. A much smaller yield of small dark olive crystals
were sparsely formed on tube walls near the bottom. Crystals
were collected and dried (vac., 25) one hour. Only the pale
green complex changes color (dark olive) upon exposure to air.
Bases, solvents. Methylene chloride was dried over CaC^
and vacuum distil led from P2O5 at 25. Benzene was dried over
CaCl2 and vacuum distilled from CaC^ at 25. Pyridine was dried
over Na and fractionally vacuum distilled from BaO. Diethyl
ether was vacuum distilled from Na/benzophenone at 25. DMSO was
dried over NaOH and fractionally vacuum distilled from NaOH. DMF
was dried over KOH and fractionally vacuum distilled from BaO.
DMA was dried over molecular sieves and fractionally vacuum
distilled from BaO. DMTF was dried over BaO and fractionally
vacuum distilled. DMCA was fractionally vacuum distilled. HMPA
was fractionally vacuum distilled from BaO. Triraethyl phosphite
was fractionally vacuum distilled. Triraethyl phosphine oxide was
used as received. Tetramethyl urea was dried over BaO and
fractionally vacuum distilled. Triethyl phosphate was
fractionally vacuum distilled from BaO. DMTF was dried over BaO,


74
fractionally vacuum distilled and stored in the dark. All
distillations were performed under dinitrogen or vacuum and only
the heart cut saved.
3. Data Collection
All manipulations were performed under dinitrogen in an inert
atmosphere box or employing syringe techniques. All glassware,
and cells were stored in dessicators over CaSOjj or in the dry
box.
Electronic spectroscopy. UV-VIS titrations were performed by
repeated microliter injections of a concentrated stock base
solution into the standardized 5 ml chromium solution. The
quartz inert atmosphere cell has been described.^ Methylene
chloride solutions were prepared in volumetric flasks inside the
inert atmosphere box. Spectra were recorded on a Perkin Elmer
330 spectrophotometer.
Calorimetry. Design and operation of the calorimeter has
been previously described.To correct stability problems, a
new unit was constructed in the Electronic Shop of the University
of Florida chemistry department. Schematics and operating
procedure are given in Appendix II. Typically, a methylene
p
chloride solution (ca. 10 M) of the dimer was prepared in a 50
ml volumetric flask and decanted into the inert atmosphere 55 ml
adiabatic cell. Five milliliters of solvent were syringed into
the volumetric flask and washed into the cell. A 5 ml aliquot of
solution was removed from the cell and transferred to the UV-VIS


75
cell. Concentrations were calculated from the absorbance at 550
nm (^550 = 132.5). The base solution (ca. 0.5 M) was prepared in
a volumetric cell and used to charge a 1 ml Hamilton gas-tight
syringe. After expelling any dead volume and air bubbles, the
syringe was emptied to the zero calibration stop and secured into
the cell. While stirring, the cell and contents were allowed to
equilibrate to room temperature (25) for several hours. A
series of thirteen calibrated syringe stops allowed incremental
injection of known volumes of base solution. Recorder
deflections which accompanied the exothermic base additions were
ratioed to a calibration deflection of known heat. Identical
titrations were performed in the absence of metal complex to
measure the base solution dilution heats (if any). Measurable
dilution heats were not observed with any of the donor solutions.
FTIR. Methylene chloride solutions were scanned in airtight
0.2 mra solution cells, after preparation in the glove box.
Solvent absorbances were subtracted. Spectra were recorded on a
Nicolet 5DX-B instrument.
FT-NMR. CD2CI2 and CgDg solutions were prepared in the inert
atmosphere box and 100 MHz spectra recorded on a JEOL XL-100 FT-
NMR. Determinations of magnetic susceptibilities with the EVANS
method were performed by modified procedures'^ with a newly
designed coaxial tube arrangement, shown in Fig 2-10. Chromium
solutions of 1-5 X 10^M were prepared by addition of 2.2-3
donor of interest to CgDg solutions of
equivalents of the


teflon spacer
paramagnetic solution
in I mm i d. capillary
Fig 2-10. Modified NMR tube for determination of solution
magnetic susceptibilities of air-sensitive complexes.


77
Cr2(tfa)lt(Et20)2 Based on the large calculated equilibrium
constants, the added donor effectively displaces coordinated
ether. The same solvent system, CgD^ containing 2% v/v TMS and
2 v/v CgHg was used for both tubes. Paramagnetic solutions were
syringed into clean dry melting point capillary tubes (1 ram i.d.)
and sealed by fitting with a machined teflon spacer. A tight
fitting cap provides a second oxygen barrier and the NMR Spectra
were collected within 15 minutes of sample preparation. The dual
standards provide a check on locating the paramagnetic shifted
resonance and signal integration provides another. This tube
arrangement allows facile manipulation of the air sensitive
solutions and avoids the expense of coaxial tubes.
Elemental analyses. Analyses were performed by Galbraith
Laboratories, Knoxville, TN. The expense of routine analysis of
the air sensitive complexes prohibited characterization of each
adduct species. The HMPA and Et^PO^ ( a strong and intermediate)
donor complexes were characterized as representative complexes.


CHAPTER III
TRANS INFLUENCE ACROSS A METAL-METAL BOND
A. Introduction
Trans influencethe ground state influence on a metal-ligand
bond strength by another trans coordinated ligandplays a
fundamental role in coordination chemistry and transition metal
complex reactivity. The ability of coordinated ligands to
labilize or stabilize a trans coordination site serves to define
the metal coordination sphere. A general feature of trans ligand
influence in mononuclear chemistry occurs when both metal-ligand
bonds are in competition for the same metal d orbital.^
The idea of a trans influence in dinuclear complexes has been
vigorously pursued, and, in a general sense, the metal-ligand
bond strength has an inverse relationship with the metal-metal
bond strength. As the studies in the last chapter illustrate,
the effect is pronounced in the chromium carboxylates as both
ligand and adjacent metal atom compete for the same chromium
orbital; stronger chromium-ligand bonds result in weaker metal-
metal bonds. Conversely, the strong quadruple Mo-Mo bonds in
the molybdenum carboxylates dictate weak metal-ligand
interactions. The effect is schematically illustrated in Fig 3-1.
While structural studies of dimers with stronger metal
78


79
L r=^> M L
L^>M< >MC=^L
L L
Fig 3-1. Schematic representation of the trans influence in
metal carboxylate dimers.


80
interactions illustrate slight lengthening of the raetal-raetal
bond with ligand donor strength, the overall effect is a
weakening of the second metal-ligand interaction as demonstrated
by calorimetric measurements.^ Redistribution of electron
density donated at one metal over both metal atoms lowers the
Lewis acidity of the second, and the effect has been
parameterized in terms of communication between the two metal
centers. As a further test of this model, the studies reported
here examine the effect of varying one ligand on the spectral
properties of a second, fixed, trans ligand in a rhodium dimer.
Examples of mixed ligand and corresponding fixed ligand complexes
for comparison are scarce. The results reported here provide the
first systematic investigation of the trans influence across a
metal-metal bond. First the EPR parameters of a coordinated spin
label (TEMPO) in complexes of the type (DRI^Cpfb^CTEMPO) are
examined, followed by the stretching frequency of coordinated CO
in complexes of the type (L) Rl^Cp f b) ^ (CO). The original
experimental results have been communicated^ by James Stahlbush;
a reinvestigation of the data is offered here.
Relationships between spectral and bonding properties of 1:1
adducts have been reported,^-10"' such as the linear correlation
between Avqh and AH for a series of 1:1 phenol-base adducts.
Breakdown of this correlation upon extension to a larger donor
set is attributed to fundamental differences in AvQH and AH for
gauging E and C effects. A more appropriate treatment of
spectral changes which accompany adduct formation involves an
adaptation of Eqn 2-1. Replacing AH by Ax, the spectral shift,
gives^ Eqn 3-1 for the case where a base is held constant and a


81
series of acids studied.
AX + W = EaEb* + CACB* (3-1)
The asterisks imply that conversion units for converting EA from
(kcal raol-^)^^ are included in Eg* along with the response to
the quantity being measured induced in the base by the acid.
The analysis of spectral data via Eqn 3-1 serves several
purposes. Besides establishing a correlation for a particular
acid or base which allows prediction of bond strengths from
spectral shifts, this treatment extends the E and C basis set
beyond only enthalpy data. And as with the enthalpy data,
correlation of the easier to obtain spectral shifts lends insight
into the nature of the metal-metal and metal-ligand bonds.
B. Results and Discussion
TEMPO
a. EPR spectra
Species of the type (B) Rl^Cpfb^CTEMPO), where B is a
coordinated Lewis base and TEMPO is 2,2,6,6-tetraraethy 1 -
piperidine-N-oxy1, were investigated using EPR in order to
measure the influence of B on the EPR spectrum of the coordinated
nitroxyl radical. TEMPO is a donor of moderate strength that
does not bind to RhjCOAc)^ or Rh2(but)^ but forms adducts with
the corresponding fluorinated derivatives, Rl^Ctfa)^ and
Rl^Cpfb)^ Typically, solutions were prepared in a 9:10:1 molar
ratio of base: rhodium complex: TEMPO in oxygen-free methylene
chloride. Most of the nitroxide existed in the unbound state,
while most of that coordinated existed as the WRl^Cpfb^CTEMPO)


82
species. A representative spectrum for diraethylacetamide with
Rl^Cpfb)^ and TEMPO is shown in Fig 3-2. In all cases, the
signal for the (BjRt^Cpfb^CTEMPO) species appears between those
for the free nitroxide and Rl^Cpfb^dEMPO). The nitrogen
hyperfine for these 2:1 adducts was usually either equal to or
slightly less than that observed for the Rl^Cpfb^CTEMPO) adduct.
Both of the above effects would be expected from an inductive
weakening of the rhodium nitroxide bond by the coordinated base.
That is, a coordinated donor weakens the Rh-TEMPO bond, and the
EPR parameters of the spin label move towards those of free
TEMPO. A wide range of g-values was observed for the mixed donor
2:1 adducts while only minor or no changes were observed for the
nitrogen hyperfine, A^. A tabulation of the EPR parameters is
given in Table 3-1.
Bonding in Rl^Ctfa^CTEMPO) has been described in terms of o-
donation from a nitroxide oxygen lone pair into the Rh-Rh a
orbital with concomitant orbital mixing of the Rh-Rh tt* orbital
with the nitroxide it* containing the unpaired spin. When B is a
sigma donor, the metal nitroxide bond is weakened, and the g
value moves toward that of free nitroxide. When B is a pi
acceptor the metal nitroxide bond can be weakened in two ways.

Directly, competition for metal-metal tt electron density
decreases the iT-backbonding to TEMPO. Indirectly, a pi acceptor
exhibits enhanced sigma donation, causing weakening of the metal
nitroxide bond in the same way as a pure sigma donor. The net
result is that both a-donors and ir-acceptors serve to weaken the
metal nitroxide bond and cause a lowering of the g-value back
towards the free solution value of TEMPO.


83
a
Fig 3-2. Representative spectrum observed for a CH2C12 solution
of Rh2(pfb)4, TEMPO and dimethyl acetamide. Species are a) free
TEMPO, b) (DMA) Rh2(pfb)4(TEMPO), c) Rh2(pfb)4(TEMPO), d)
precipitated Rh2(pfb)4(TEMPO).


84
Table 3-1. EPR Parameters of (¡DRI^Cpfb^CTEMPO) adducts.3
-AH 1:1
B
no base
methyl acetate
ethyl acetate
acetone
p-dioxane
dimethyl acetamide
bridged ethere
tetrahydroduran
dimethylsulfoxide
hexamethylphos-
phoramide
dimethyl formamide
acetonitrile
pyridine-N-oxide
f
cage phosphite
diethylsulfide
4-picoline
pyridine
1-methylimidazole
piperidine
triethylamine
An( 10-:>cm ) g'
1.57
2.0152
1.55
2.0128
1.55
2.0127
1.55
2.0122
1.55
2.0120
1.55
2.0119
1.55
2.0119
1.55
2.0118
1.55
2.0116
1.55
2.0115
1.55
2.0114
1.56
2.0108
1.55
2.0101
1.56
2.0095
1.55
2.0093
1.55
2.0083
1.56
2.0081
1.56
2.0079
h
2.0074
h
2.0069
Scale (kcal mo1
2.0152
0
2.0130
7.37
2.0129
7.96
2.0125
9.05
2.0122
9.66
2.0119
11.17
2.0114
11.64
2.0111
12.38
2.0117
11.74
2.0109
13.87
2.0120
10.54
(2.0132)
6.81
2.0106
14.65
(2.0104)
13.93
(2.0100)
14.59
(2.0092)
17.75
(2.0096)
17.06
2.0081
20.32
2.0078
21.28
2.0068
24.31


85
Table 3-1 (continued)
a) Compare to solution parameters of free TEMPO: g = 2.0047, A
= 1.47 X 10"3cm1.
b) g of (B)Rh2(pfb)4(TEMPO).
c) g calc from equation 3-4 using the C^:^ and E^1:1 from
equations 3-2 and 3-3 and values reported in Table 3-2.
d) Calculated enthalpy for adding B to Rhp(pfb)4 using values in
Table 3-2.
e) 7-0xabicyclo [2.2.1] heptane.
f) 1-Phospha-4-ethyl-2,5,7-trioxabicylo [2.2.1] octane.
> H
g) Systems in which metal tt to ligand tt back-bonding
occurs.
h) Unresolved.


86
b. Quantitative correlations
The general trends discussed in the previous section suggest
that g-values from the EPR spectra may be used to provide
quantitative data about the strength of binding. This encouraged
us to investigate the quantitative relationship between the
enthalpys of adduct formation and changes in the g-values. A
model has been proposed and tested for predicting the enthalpy of
coordination of a second donor, B, to an f^CC^ClOjjCB) adduct to
form a 2:1 adduct.^ ^33 jn this model, the EA parameter for the
1:1 adduct behaving as an acid to form a 2:1 adduct, EA1:\ is
given by
ea':1 h k%
and C.1:^ is given by
V = CA -X,CB
(3-2)
(3-3)
Where k and k* reflect (for a-donors) the effectiveness of the
metal-metal bond at transmitting the inductive influence of base
1*1 1*1
coordination to the second metal center. EA and CA values
for(B)Rl^ipfb)^ can be calculated from equations 3-2 and 3-3.
The reported32*38,39 £ ancj q values used in this analysis are
given in Table 3-2. The EA1:1 and CA1:1 of the various
(B^t^ipfb)^ adducts are calculated and also listed in Table 3-2.
Note how the acid parameters decrease as the inductive effect
increases. That is, the greater the donor strength, the lower is


87
Table 3-2. E and C Parameters for
Species
Used in
This Study
B
eb
CB
c 1:1
ea
r 1:1
methyl acetate
0.903
1.61
4.01
1.68
ethyl acetate
0.975
1.74
3.93
1.68
acetone
0.987
2.33
3.92
1.66
p-dioxane
1.09
2.38
3.80
1.65
dimethyl acetamide
1.32
2.58
3.53
1.65
bridged ether
0.887
4.11
4.03
1.59
tetrahydrofuran
0.978
4.27
3.93
1.58
dimethylsulfoxide
1.34
2.85
3.51
1.64
hexaraethylphosphoramide
1.52
3.55
3.30
1.61
dimethyl formamide
1.23
2.48
3.63
1.65
acetonitrile
0.886
1.34
4.03
1.69
pyridine N-oxide
1.34
4.52
3.51
1.58
cage phosphite
0.548
6.41
4.42
1.51
diethylsulfide
0.339
7.40
4.67
1.47
4-picoline
1.17
6.80
3.70
1.49
pyridine
1.17
6.40
3.70
1.51
1-raethylimidazole
0.934
8.96
3.98
1.41
piperidine
1.01
9.29
3.89
1.40
triethylamine
0.991
11.09
3.91
1.34
TEMPO
0.915
6.21
4.00
1.51
A
ea
CA
k
k'
Rh2(pfb)4
5.06
1.74
1.16
0.0364


88
the Lewis acidity of the second metal center. With these E and C
values for the various 1:1 adducts, we are now in a position to
attempt a correlation of the g-values obtained when TEMPO is
coordinated to the second coordination site to form a series of
2:1 adducts of general formula (BJRhgtpfb^CTEMPO). Substitution
of Eqns 3-2 and 3-3 into 3-1 gives
g + W = Ea1:1Eb* + CA1:1CB* (3-4)
where g has been substituted for Ax The simultaneous equations
are solved for Eg* and Cg* which are the spectroscopic parameters
for TEMPO needed to predict g. The quantity W includes the g
value for free TEMPO (2.0047) as well as any nonzero enthalpy
components of the spectroscopic relation.^ 02-1 03 The best fit
results yield
Eg* = 1.16 x 10~3 (0.29 x 10-3)
Cg* = 1.78 x 10"2 (0.10 x 10-2)
W = -1.9784 (.0018)
with standard deviations in parentheses. These parameters for
TEMPO allow calculation of the g-value for any
(B)Rh2(pfb)¡j(TEMP0) complex when the base is a o-donor whose Eg
and Cg parameters are known. Data for adducts with donors known
to act as T:-acceptors (acetonitrile, cage phosphite,
diethysulfide, 4-picoline and pyridine) are not included in the


89
calculation of the TEMPO parameters. Attempted fits which
include the ^-acceptors give larger standard deviations, as
expected. Table 3-1 contains the g-values calculated from this
fit (scaic) The columns of g and gcaic show excellent
agreement, generally within the accuracy of the measured numbers,
except in those cases where metal to base u-backbonding occurs
(data in parentheses). For the latter, these are the g-values
expected if the Lewis bases, B, utilize only their o-bonding
capabilities in forming the (B^l^ipfb^ adducts. The close
agreement between g and gca^c demonstrates that the inductive
model (equations 3-2 and 3-3) adequately describes the
transmission of coordination effects through the metal-metal
bond, for it is this model which describes the varying acidity of
the second metal center. For the five donors which also behave
as it-acids, the lower observed g-values than calculated by Eqn 3-
4 manifest the metal to B iT-backbonding contribution in these
adducts. Thus this analysis would suggest that bothoand tt
interactions in the B-Rh bond serve to weaken the Rh-TEMPO bond.
To see if a relationship exists between g and the strength of
B binding to Rh2(pfb)^, enthalpies for binding bases were
calculated from the E and C equation (2-1); these 1:1 adduct
heats, AHg1:1, are given in Table 3-1. The E and C parameters
are derived from a-only interactions and hence the calculated Ah
in Table 3-1 reflect only the o component of the adduct bond.
The experimental g-values, however, reflect the sum of o-donor
and any fT-acceptor interactions. In Fig 3-3 the enthalpies of
1:1 adduct formation for (B)Rh2(pfb)j are plotted as a function
of the experimental g values for the 2:1 adducts (B^l^pfb^tTEMPO).


90
Both calculated (o,A) and experimental (X) enthalpies are
included. (It would perhaps be more direct to compare the g-
values of the 2:1 adducts with the enthalpy of 2:1 adduct
formation, that is, the enthalpy for B + Ri^Cpfb^CTEMPO) or even
TEMPO + Rt^Cpfb^B). The qualitative conclusions are, however,
the same, and experimental heats are available for B +
Rl^pfb)^.) In Fig 3-3, the calculated (cr-only) enthalpies are
all lower than the g-values would suggest for the five donors
which act as n-acceptors: acetonitrile, cage phosphite,
diethylsulfide, 4-picoline and pyridine. Additional
stabilization in the (B^hgipfb)^ adduct bond is consistent with
Rh to B 7T-backbonding. In the two cases where experimental
AHg^:1 are available (acetonitrile and pyridine), the measured
heats lie much closer to the correlation line. Thus, the g-
values manifest both a and tt effects across the metal-metal bond,
both serve to lower the g-value of coordinated TEMPO, and o-
donation appears to exert a stronger influence than tt-acceptance.
(Inclusion of tt-effects brings Ahb^:1 closer to the correlation
line but not all the way.)
2. Carbon Monoxide
a. FTIR spectra
The CO ligand is ubiquitous in organoraetal lie chemistry, and
considerable effort has been put forth to understand the nature
of M-C-0 bonding and the influence of various ligands upon the
reactivity and spectroscopic properties of carbonyls. In probing


Full Text
mm

SYNERGISM IN METAL CARBOXYLATE CLUSTERS
By
CARL JOSEPH BILGRIEN
A DISSERTATION PRESENTED TO THE GRADUATE SCHOOL
OF THE UNIVERSITY OF FLORIDA IN
PARTIAL FULFILLMENT OF THE REQUIREMENTS
FOR THE DEGREE OF DOCTOR OF PHILOSOPHY
UNIVERSITY OF FLORIDA
1986

To Deanna,
who always knew.

ACKNOWLEDGEMENTS
No scholar or scientist works alone: each must rely on the
labors of past workers and the assistance of his contemporaries.
I have incurred many debts in this regard.
First and foremost I thank Professor Russell S. Drago for his
support, encouragement and advice. He has generously shared his
insight and perseverance and I am grateful for having had the
opportunity to work with him.
I thank my committee members, Professors Earl Muschlitz,
David Richardson, Harry Sisler and E. Dow Whitney for their
efforts.
For his continued friendship and enthusiasm I am especially
grateful to Dr. Barry B. Corden.
To all the group members who have shared their time, advice,
expertise, grousing and politics, I am grateful. These include
Kenneth Balkus, Iwona Bresinska, Jeffrey Clark, Richard Cosraano,
Shannon Davis, Peter Doan, Andrew Griffis, Karen Jongeward,
Ernest Stine, Joshua Telser, Keith Weiss and Ngai Wong.
For their ability to build anything I drew I thank the men
in the glass shop, Rudy and Dick. For their sweat and good
humor, I thank Vernon, Chester and Daly of the metal shop. For
not giving up on the calorimeter I thank Russell Pierce. For his
many suggestions and services I thank Dr. Roy King.
iii

Finally, I can never repay the time and sacrifices of ray
wife, Deanna Saint Souver. For her unabated
inspite of all too many hours spent in lab, I
always.
encouragement
wi11 love her
iv

TABLE OF CONTENTS
PAGE
ACKNOWLEDGEMENTS iii
KEY TO ABBREVIATIONS vii
ABSTRACT viii
CHAPTER I. GENERAL INFORMATION 1
CHAPTER II. REACTIVITY OF Cr2(02CR)4 14
A. Introduction 14
B. Results and Discussion 18
C. Conclusion 67
D. Experimental 68
CHAPTER III. TRANS INFLUENCE ACROSS A
METAL-METAL BOND 78
A. Introduction 78
B. Results and Discussion 81
C. Conclusion 96
CHAPTER IV. THE ELECTRONIC SPECTRA OF
M(III)2M(II)(02CR)6L3 SPECIES 98
A. Introduction 98
B. Results and Discussion 100
C. Conclusion 125
D. Experimental 125
CHAPTER V. THE OXIDATIVE DEHYDROGENATION OF
ALCOHOLS CATALYZED BY OXOTRIRUTHENIUM
CARBOXYLATES 129
A. Introduction 129
B. Results and Discussion 132
C. Conclusion 185
D. Experimental 188
CHAPTER VI. SUMMARY AND CONCLUSIONS 194
v

APPENDIX I. SPECTRAL AND CALORIMETRIC DATA
198
APPENDIX II. OPERATION OF CALORIMETER 215
APPENDIX III. DERIVATION OF EQUATION 2-17 219
REFERENCES 222
BIOGRAPHICAL SKETCH 231
vi

KEY TO ABBREVIATIONS
but
butyrate = O2CCH2CH2CH2
hept
heptanoate = 02C(CH2)5CH
hfb
heptafluorobutyrate = O21
OAc
acetate = O2CCH2
oct
octanoate = 02C(CH2)gCH2
prop
propionate = O2CCH2CH2
tfa
trifluoroacetate = O2CCF
vii

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
SYNERGISM IN METAL CARBOXYLATE CLUSTERS
By
Carl Joseph Bilgrien
August 1986
Chairman: Professor Russell S. Drago
Major Department: Chemistry
Synthetic, spectroscopic and reactivity studies on several
metal clusters with carboxylate ligands are described. These
complexes are of the general formula M2(02CR)¿jL2 or M'^CC^CR^L^
where M is Cr or Rh; M' is Co, Cr, Fe or Ru; C^CR is a bridging
carboxylate ligand; and L is a neutral donor ligand. These
studies were undertaken to examine the influence of reaction at
one metal center upon that at an adjacent metal atom and to help
understand the metal-metal bonding interactions which contribute
to the transmission of coordination effects. Ligand exchange
reactions of C^C C^CC F^) 4 [ (CH^ C^) 2® ^ 2 were monitored.
Equilibrium constants and enthalpys for exchange reactions with a
variety of donors were determined from calorimetry data. The
resulting enthalpys were used in a correlation analysis which
demonstrated that the Cr(II) centers are significant Lewis acids
and interact with axial ligands almost exclusively in an
viii

electrostatic sense. Despite a relatively weak metal-metal bond
the first exchange enthalpy is appreciably higher than the
second. Magnetic susceptibility measurements show increasing
paramagnetism for the Cr2^ + unit as stronger donors displace
coordinated ether. Strong donors promote oxidation and
rearrangement of the dinuclear unit.
Mixed ligand complexes of the form BRh2(02CCF2CF2CF2)¿|L where
B is a Lewis base were examined by electron paramagnetic
spectroscopy (L is the spin label 2,2,6,6,-tetramethylpyridine-N-
oxyl) and infrared spectroscopy (L is CO) and the spectral data
used to calculate acid parameters which describe the Rt^"1" unit.
The spin label g-value, the CO stretching frequency and the
calorimetric enthalpy, all of which describe the perturbation the
base makes on the trans-metal atom, show different relative
electrostatic/covalent responses as the base is varied,
demonstrating the method dependence of monitoring donor-acceptor
adduct formation. The electronic spectra of clusters of the form
M^OCOgCR^Lj (M = Co, Fe, Ru) exhibit a donor ligand (L)
dependence only for M = Ru.
Lastly, clusters of the form Ru20(02CR)^L2n+ (n = 0,1) are
shown to catalyze the selective oxidative dehydrogenation of
alcohols. A mechanistic proposal incorporates the demonstrated
nonradical behavior and the observed stoichiometry. Nonradical
chemistry and reduction of oxygen to water are demonstrated and a
mechanism proposed. High catalyst activity is suggested to arise
from the multiple metal centers acting in concert.
ix

CHAPTER I
GENERAL INFORMATION
An underlying theme in the current and recent intense
interest in the chemistry of molecules with multiple metal
centers is the way in which the metal centers influence each
other and generate unique reactivity. This synergistic interplay
is implicated in a host of chemical systems. Many enzymes rely
on multiple metal centers for substrate binding (e.g. hemocyanin,
lacease) while others employ proximal metal atoms for electron
transport (e.g. cytochrome c oxidase) or perturbation of the
substrate binding (e.g. nitrogenase). Examples of homogeneous and
heterogeneous catalyses with multiple metal centers abound
whereas interesting physical properties or reactivities of a
stoichiometric nature are often introduced by metal atoms in
close proximity.
The nature and extent of discrete metal-metal interactions
can vary from direct orbital overlap to long range electron
transfer. Studies along this continuum have proceeded via
several fronts: physical and theoretical studies of orbital
interactions; introduction of specific reactivity or binding;
modeling of metal loenzyraes; and studies of electron transfer.
One area in particular, the study of molecules with metal-metal
bonds, has received considerable attention following the
1

2
1 o
pioneering work in the laboratories of Wilkinson J and
Cotton.The latter area is the focus of this thesis; a
historical perspective is given here as introduction.
Indentification of single and multiple metal-metal bonds as
stereoelectronic elements of molecular structure has come about
only recently. Prejudiced by the conceptual framework of
coordination chemistry established by Alfred Werner®, metal-metal
bonds were deemed inconsistent with one center coordination
chemistry. The advent of x-ray crystallography heralded
recognition of bonding between metal atoms. Although a
crystallographic report"^ as early as 1946 demonstrated Mo-Mo bond
distances shorter than those in metallic molybdenum, unequivocal
evidence for metal-metal bond formation came from the structure
determination® of Mn2(C0)-|Q in 1957. Evidence for multiple
metal-metal bonding (in this case, a quadruple bond) came from
the initial structure determination^ of the RejClg^- anion and
its reinvestigation10 in 1964. These early discoveries and the
chemistry of metal-metal bonds have been throughly reviewed
through 1980 in "Multiple Bonds Between Metal Atoms" by F. A.
Cotton and R. A. Walton.11 Several other more specific reviews
have appeared recently.1^-1®
The importance of the Re2Cli|i<- anion to the development of
this field lies in the electronic structure of the Re-Re bond.
The crystal structure1^ of K2(Re2Clg) (F^O^ demonstrated a Re-
Re distance of 2.24 A (2.75 A in metallic Re) and eclipsed
conformation of the two ReCl^ halves. Both of these features can
be qualitatively explained by considering d orbital overlap. The
in phase and out of phase combination of pairs of d orbitals on

3
the two metal atoms generates five bonding and the corresponding
antibonding molecular orbitals (MO) in Dooh symmetry. The in
phase bonding interactions are shown schematically in Fig 1-1.
Positive overlap of the dz2 orbitals generates a a bond (Fig 1-
1a) whereas negative overlap generates the corresponding a
antibonding orbital (not shown in Fig 1-1). The dxz and dyz
combinations give rise to two orthogonal but degenerate it bonds
and the corresponding tt * antibonding orbitals. Lastly, in the
absence of any ligand interactions, cofacial overlap of the pairs
of dxy and dx2_y2 results in degenerate 6 bonds and their <5*
complements. Introduction of ligand orbitals lowers the symmetry
to D2 dx2_y2 orbitals point towards the ligands; the dxy orbitals point
between them). The dx2_y2 orbitals are thus utilized in ligand
bonding and play effectively no role in metal-metal bond
formation. The extent of orbital overlap dictates the MO
energies and the orbital diagram which results is shown in Fig 1-
2. In terms of the (Re2Clg)2“ anion, each Re^+ center
contributes four electrons to give a net quadruple bond,
which explains the short Re-Re distance. The twist angle
dependence of the 6 interaction results in maximum overlap when
the Cl“ are eclipsed. Thus, the stabilization gained in 5 bond
formation is greater than the repulsion energy of the eclipsed
halides. The qualitative features of this MO description are
supported by quantitative calculations (SCF-Xa-SW) on the
(Mo2Cl8)i|~ ion1®-1^ and the (Re2Clg)2~ ion20-21 and is general
for dimers of the 2nd and 3rd row d2-^ transition metals with
octa-halo ligand sets

4
Lol,
a)
Cío CÍO
ctCSo
b)
SS *â–  SS
— 806
c)
<% d« ¿ft -
— c«5
d) ^ d*y ^
6}
Fig 1-1.
The five nonzero d-d overlaps between two metal atoms.

5
a
K
7T*
S*
S
TT
Fig 1-2. Qualitative description of the primary metal-metal d
orbital interactions for a dinuclear metal carboxylate in
symmetry. The dx2- 2 orbitals are involved with carboxylate
bridge orbitals and do not contribute significantly to the metal-
metal bond.

6
The utility of this MO description extends to other ligand
sets as well, most notably the tetracarboxylates, M2(C>2CR)ijL2.
As with the octahalogenates, work with these complexes appeared
in the literature long before their general structural features
pp
were revealed by a crystallographic study*1*1 of (^((^CCH^jjC^O^
(which, incidentally, carries no metal-metal bond) in 1953. The
tetracarboxylate ligand framework is ubiquitous in transition
metal chemistry; examples exist for vanadium, chromium, cobalt,
copper, raolydbenum, technetium, ruthenium, rhodium, tungsten,
rhenium, and osmium where L represents a neutral donor molecule
or coordinated anion (naked clusters with no L also abound). The
general structure of M2(02CR)jjL2 type dimers* is shown in Fig 1-
3. Symmetry dictates the qualitative orbital interactions
displayed in Fig 1-2 apply to the tetracarboxylates also, as
borne out in extensive numerical calculations on dimers of the
2nd and 3rd row transition metals. Dinuclear carboxylates of the
first row, however, are not adequately described by this orbital
description. Representative c lusters,23-26
(^(quin^C^H^CC^)^ and Cv^Ct^O^CCH^CC^);; all exhibit longer
metal-metal distances than found in the parent metal. As such,
no metal-metal bond is proposed. For neutral dinuclear
tetracarboxylates of these metals, the observed magnetic behavior
is rationalized in terms of antiferromagnetic exchange between
equivalent spin centers, presumably through the carboxylate
The term "dimer" is traditionally reserved for addition
complexes formed from two monomeric units. In the present
context, the term "dinuclear complex" would be more appropriate;
"dimer" is favored in the interest of brevity.

7
R
R
Fig 1-3. The general structure of MjCC^CR^CUg. Diffraction
studies support minor deviation from idealized symmetry.

8
bridges. The center of the chromium tetracarboxylates is
the exception to the rule. Although there has been disagreement
as to the nature of the bonding in the systems,
that some bonding interaction is present has been reasonably
established. Results of SCF-Xa-SW^ and ab initio^®-^
calculations support a weak quadruple bond as the generic MO
description of Fig 1-2 would suggest for two d^ centers.
As a family, the transition metal tetracarboxylate dimers
comprise a series of stable complexes, structurally well
characterized in most cases, whose electronic structures and many
physical properties can be interpreted in terms of a generic MO
description. As such they provide an excellent opportunity to
study the effect of reactivity at one metal center upon another.
This idea has been exploited by Drago and coworkers in the series
of coraplexes3°“35 Ri^Cbut)^, Rl^Cpfb)^ Mo2(pfb)¡j, Ru2(but)jjCl
and Cu2(hept)jj (but = butyrate, pfb = perfluorobutyrate, hept =
heptanoate). In these studies, stepwise adduct formation of the
Lewis acid dimer with first one, then two donor molecules was
monitored. Lower measured enthalpies for the second reaction
indicated a weakening of the second metal center acidity as a
result of base coordination to the first metal atom. By studying
a range of bases with each dimer, the enthalpy data were treated
in terms of the Drago-Wayland E and C equation.36—39
empirical model31“32 was pU£ forward to describe the lowered
acidity in terms of the ability of the metal-metal bond to
transmit electrostatic (E) and covalent (C) effects. The shorter
quadruple Mo-Mo bond was found more able to transmit
electrostatic effects while the longer, more polarizable, single

9
Rh-Rh bond was found more able to transmit covalent effects.
Consistent with the generic MO scheme in Fig 1-2, the 14 electron
rhodium dimers, with filled tt orbitals, undergo tt-back bonding
to it acids (e.g. CO, pyridine) while the 8 electron Mo2(pfb)/},
with no tt * density, does not. The 11 electron Ru2(but)iJCl, with
half filled ir orbitals exhibited intermediate interactions.
As mentioned above, description of the metal-metal bond in
the chromium carboxylates has been the focus of recent
40
controversy. Ab initio calculations suggest that correlation
4 +
effects are very important in the description of the C^
complexes. These calculations involve description of the ground
state wavefunction of C^iC^CH)^ in terms of contributions from
p n p PUP
q% 5 and excited states such as a tt 6 . The contribution of
PUP
a tt 6 is 16% which contrasts markedly with similar calculations
for Mo2(02CH)jj in which the configuration contributes 67%.
That is, the quadruple bond adequately describes the ground state
of Mo2(02CH)ij. It does not do so for C^CC^CH)^. deMello et al.
have suggested1* 1-1*^ that the dominant description of the bonding
in C^1*'1’ is one of two Cr atoms antiferromagnetical ly coupled
with some contribution from multiple bonding. The chromium
carboxylate dimers are also unique in exhibiting a strong
dependence of the metal-metal bond length (and hence electronic
structure) upon axial ligation. Depending upon the nature of L
0
and R, molecules of the type C^CC^CR^L^ display a 0.57 A range
of bond lengths J from 1.97 A to 2.54 A. By contrast, adducts
of the Mo and Rh carboxylates display metal-metal bond length
O
ranges of 0.13 and 0.12 A, respectively. The sensitivity of the
metal-metal bond to donor molecule coordination may be manifested

10
in the way coordination at one chromium center affects the
second. A quantitative study of the coordination properties of
4 +
the Cr2 unit is the subject of chapter two. A literature
report of observed paramagnetism in the chromium dimer studied
here along with the antiferromagnetic description put forward by
U?
deMel lo et al. prompted a magnetic susceptibility investigation
which is also reported in chapter two.
Considerable effort has been expended upon the coordination
chemistry of the carboxylate diraer3 towards understanding the
nature of the metal-metal bond—its electronic structure,
theoretical description and physical properties and reactivity.
The effects of varying both the axial ligand and the bridging
chelate on the metal-metal length bond have been extensively
explored.11 The perturbation of an axial ligand upon the metal-
metal bond has been termed a trans influence,11-’-*1® referring to
the influence of a ligand upon the bond directly trans to it. In
general, the axial ligand bond competes with the second metal as
a ligand, weakening the metal-metal bond; and conversely, the
stronger the metal-metal bond, the weaker is the metal ligand
interaction.*1® Another way of viewing the effects of axial
coordination is the influence of the ligand upon a second ligand
opposite the metal-metal bond. This secondary trans influence or
inductive effect can serve to identify the primary orbital
interactions in the metal ligand bond and the extent to which the
metal-metal bond transmits the ligand influence.
An infrared study of a series of L-Rl^Cpfb^-CO adducts and
an EPR study of a series of L-Rt^Cpfb^-TEMPO (TEMPO is the free

11
radical 2,2,6,6-tetraraethy lpiper idine-N-oxyl) adducts has been
performed1^ and is further examined in chapter three.
A natural extension of these studies in the transmission of
bonding effects in dimers would be to consider trinuclear
complexes. Retaining the carboxylate ligand set still allows one
to choose from a diverse field of trinuclear complexes (trimers).
A desire to work in noncoordinating solvents and the need for
identical metal sites dictated the use of neutral, trigonal
complexes, exemplified by complexes of the basic iron acetate
structure.^ Q~-> 1 These complexes, of general formula
(M20(RC00)gL2)n+, where L is a neutral monodentate ligand,
contain a triply bridging oxide ion at the center of a (generally
equilateral) triangular array of metal ions; their structure is
illustrated in Fig 1-4. The electronic and structural details of
many of these compounds with M^^CnsI) have been studied. The
neutral, mixed valence compounds with n=0 have received special
attention as models for intramolecular electron transfer;
examples are known for M = Fe, Cr, Ru, Mn, V and perhaps Co.-^”-^
A generic MO description for the mixed valence trinuclear
carboxylates is not available, and definitive MO calculations
remain prohibitive in light of their complexity. At least one
system, Ru20(02CCH2)5(PPh^)3, has been addressed from a LCAO
perspective, however.^ Understandably, assignment of the
electronic spectra of the trinuclear carboxylates remains
ambiguous. With regards to the neutral mixed valence trimers,
electronic spectra have been reported only for complexes of Co,
Fe, and Ru.-^--^ To gauge the effect of ligand substitution
reactions on their electronic structures, representative neutral

12
Fig 1-4. General trinuclear, oxo-centered, basic
carboxylate structure of formula [M2O(RCO2)0L2]n+.
metal

13
mixed valence trimers were examined in coordinating solvents and
the results are presented in chapter four.
The optical spectra of both neutral and cationic ruthenium
carboxylate trimers exhibit composite bands which originate from
a series of closely spaced molecular electronic transitions.^
The cluster system Ru2O(OAc)g(py)2^ + //2+,/+//0/,~ displays an
extensive reversible redox chemistry,^ prompting the name
"electron sponge". The stability of different redox states and
availability of substrate binding sites in the ruthenium trimers
are promising for homogeneous redox catalysis. The trirutheniura
acetate clusters have shown utility as homogeneous hydrogenation
catalysts for unsaturated substances. Attempts to employ
these same clusters as olefin oxidation catalysts in this
laboratory revealed the reversible reduction by alcohol solvent
at elevated temperatures. Subsequent specific catalytic
oxidative dehydrogenation of a range of alcohols and mechanistic
features were explored. These findings are given in chapter
five

CHAPTER II
REACTIVITY OF Cr^O^R)^
A. Introduction
The dimeric metal carboxylates, f^OgCR)^, are convenient
clusters for studying the effects of coordination at adjacent
metal sites. An extensive array of complexes has been isolated
and structurally characterized, synthesis methodology is
relatively straightforward, charge neutrality allows study in
non- or weakly coordinating solvents, and the metal centers
display open axial (trans to the metal-metal bond) coordination
sites to which Lewis bases readily bind.0'1 Work in this
research group has focused upon quantitative description of the
Lewis acid centers. It has been found that the enthalpy measured
for formation of the second metal-base bond is less than that for
the first metal-base bond in dimers of rhodium and molybdenum,
indicating that this may be a general phenomenon of this family
of compounds. By working with a range of characterized bases,
the enthalpy data could be treated with the Drago E and C
raodel.^~38 The empirical equation (2-1)
- AH + W = EaEb + CACB (2-1)
describes the enthalpy of adduct formation where EA and CA are
14

15
the acid parameters, and Eg and Cg are the base parameters
corresponding to the tendencies of the acid or base to undergo
electrostatic or covalent interactions. The W terra is included
when any constant contribution to the measured enthalpies
independent of acid or base variation accompanies adduct
formation. Drago, Long and Cosmano suggested an inductive
transfer model to describe the Lewis acidity of the second metal
center.31»32 jn model, the acid parameters of the 1:1
adduct are reduced from that of the free acid by an amount that
is proportional to the corresponding base parameter.
E
A
1:1
kEg
(2-2)
C
A
1:1
= C,
k'Cg
(2-3)
The k and kf have physical significance and represent the ability
of the metal-metal bond to transmit electrostatic and covalent
coordination effects. This description can be thought of as
parameterization of a trans effect. To date, four dinuclear
metal carboxylates have been studied and are summarized in Table
2-1. It should be emphasized that the intent of this methodology
lies not in determination of E^ and numbers per se (although
the experimentally determined numbers can be used to predict
unmeasured enthalpies), but their relative magnitudes serve to
illustrate the nature of the metal-ligand and metal-metal bonds.
For example, both Rl^tpfb^ and Mo2(pfb)jj exist in the +2
oxidation state and contain the same bridging ligand; a similar

16
Table 2-1.
dimers.
Acidity Parameters of various metal carboxylate
Acid
M2n+ dn BO Ea Ca CA/Ea k k'
RhpCbut)^
4 14 1 3.21 1.32 0.411 1.16 0.0364
RhgCpfb)^
4 14 1 5.06 1.74 0.344 a
MopCpfb)^
484 5.92 0.385 0.065 1.46 0.022
RupCbut^Cl
5 11 2.5 7.73 1.27 1.64 b
a.) Not determined though the butyrate and perfluorobutyrate
bridges were found similar in their transmission capability.
b.) Cl“ coordination precludes bonding of a second base.

17
partial positive charge exists at each metal center, and the two
dimers have similar EA numbers. The less electronegative bridge
in Rh2(but)2} results in a lower EA for this complex. The CA
numbers, on the other hand, reflect the polarizability of the
metal-metal bond. The quadruple metal-metal bond of Mo2(pfb)(j is
not as likely to redistribute electron density over the entire
molecule as the more flexible, single Rh-Rh bonds. The raolydenum
carboxylate has a lower CA.
Similar rationale lends physical significance to the
transmission coefficients, k and k'. The shorter metal-metal
bond in Mo2(pfb)^ allows for greater electrostatic interaction of
the second molybdenum with the first coordinated base and greater
interaction of the two base molecule dipoles. The greater
polarizability of the metal-metal bond in Rh2(pfb)1( allows for
better electron density redistribution as manifested by the
larger k' value.
The analogous chromium carboxylates have such a strong
tendency to coordinate electron pair donors in the axial
positions that they are only rarely seen without ligands.11 In
the two cases where unsolvated chromium dimers were studied
structural ly,^*^ axial coordination occurred by association of
the molecules to form infinite chains. The nature of the axial
44 64 6R
ligand has a pronounced effect upon the Cr-Cr distances * * J
which range from 2.214 A in the (Cr-jCCO^^Cf^O)^11- ion to 2.541
A in Cr2(02CCF2)4(Et20)2. The experimentally observed range of
bond distances would suggest a shallow potential well for the
4.
Cr2 unit. The marked dependence of the electronic structure of

18
the Cr2i,+ unit upon the axial ligand provides a unique
opportunity to study transmission of bonding effects.
Reactivity studies of the dichroraium tetracarboxylates have
focused primarily on the acetate which has found wide utility as
a reducing agent and as a starting material in the preparation of
other compounds containing the unit. To this end, dimers
have been isolated and structurally characterized for a variety
of bridging ligands with C, N and 0 donor atoms and a range of
axial bases. Comparison of the structural parameters shows no
clear relationship between the nature of the bridging ligand and
metal-metal bond strength/length. Axial ligation, however, is
found to strongly influence the Cr-Cr bond with stronger donors
generally dictating longer bonds.11
Unlike the strongly bonded Mo-)1*'*' ion, there is no evidence
for the existence of the naked cluster Cr^*. Bridging ligands
may play a role in metal-metal bond formation other than keeping
the metal centers in close proximity; the (Cr^CH^g)*1"’ and
(Cr2(C1jHg)i|)il” ions exist without bridging ligands.^*67
The Cr2(tfa)it(Et20)2 adduct was first reported in 1966 to
exhibit weak paramagnetism. The structural report which followed
showed the dimer to contain the longest Cr-Cr bond known. Our
interest lay in probing the transmission of bonding effects
across such a weak, loosely interacting metal-metal bond, and
reactivity studies are reported here.
B. Results and Discussion
1. Qualitative Reactivity
Initial investigations were performed with the simple

19
carboxylates C^COAc^O^O^ and Cr2(but)¿,(H20)2. The hydrates
are easily desolvated by heating in vacuo. An x-ray diffraction
study^3 0f anhydrous Cr2(0Ac)^ prepared by sublimation of the
hydrate demonstrated that bridging oxygens of neighboring
clusters satisfy the strong coordination requirements of the
Cr(II) centers. A portion of the polymeric compound which
results is shown in Fig 2-1; this compound is soluble only in
coordinating solvents. To minimize the nucleophilicity of the
carboxylate oxygens, the tr i fluoroacetate bridged dimer
Cr 2( t f a) ¿i (E120) 2 was utilized. Again, the need for a
coordinating ligand attests to the Lewis acidity of the metal
atoms. Here, however, axial coordination is superceded by a weak
donor. Adduct formation in these studies proceeds via an
exchange reaction to displace diethylether.
Weak donors such as acetonitrile do not displace ether.
Intermediate donors such as dimethylacetamide cleanly displace
coordinated ether to give first 1:1, then 2:1, adduct formation.
Strong donors such as pyridine rapidly cleave the complex (vide
infra). A good measure of solvent donor capabilities come from
their E and C numbers; donors studied are summarized in Table 2-
2.
For the intermediate case, donor exchange is an equilibrium
process, readily monitored by changes in the electronic spectra
of the Cr^'1' chromophore. Analagous to earlier studies^® with
Rh2(pfb)jj and Rh2(but)jj, evidence for the formation of 1:1 and
2:1 adducts of Cr2(tfa)^ is provided by spectral studies.
Representative spectra of Cr2(tfa)¿j(Et20)2 and its adduct
exchange forms with DMA are shown in Fig 2-2 where the

20
y
/
\
o o
1/ 1/ /I
o — Cr — Cr—O
/ /I /i :c
Cr-°N /°
/C
R
O \) " Cr
i/ 1/ /
Cr Cr-.— O
\
/
R
/I
Fig 2-1. The formation of infinite chains of 0^(0201?)^
molecules by oxygen bridge bonding.

21
Table 2-2. Donor parameters.
E
C
diethyl ether (Et20)
.963
3.25
do not displace Et20
acetone
.987
2.33
acetonitrile
.886
1.34
dimethyl sulfide
.343
7.46
methyl acetate
.903
1.61
tetrahydrothiophene
.341
7.90
triphenylphosphine
(a)
displace Et20
dimethyl acetamide (DMA)
1.32
2.58
dimethyl cyanaraide (DMCA)
1.10
1.81
dimethyl formamide (DMF)
1.23
2.48
dimethyl sulfoxide (DMSO)
1.34
2.85
dimethyl thioformamide (DMTF)
(a)
dioctyl ether (DOE)
1.10
3.40
p-dioxane
1.09
2.38
hexaraethylphosphoramide (HMPA)
1.52
3.55
triethylphosphate
(1.36)b
(1.81)
trimethylphosphine oxide
(1.53)b
(3.32)
trimethylphosphite
(1.03)b
(5.99)
tetrahydrofuran (THF)
.978
4.27
tetramethylurea
1.20
3.10
dissociate complex
diethylamine
1.17
8.51
N-raethyl imidazole
.934
8.96
pyridine
1.12
6.89
quinuclidine
.704
13.2
a) unknown
b) tentative parameters calculated from limited data sets

Fig 2-2. Spectrop hotoraetric titration of chromium
trifluoroacetate, diethylether adduct, with DMA in Ct^C^. The
free acid spectrum is labeled 0. The titration was ended with 50
molar excess added DMA, labeled X.

X (nm)

24
concentration of the dimer is kept constant and the donor
concentration gradually increased.
The presence of an isosbestic point at 572 nm is evidence for
only two species in solution at low base concentration: the free
acid, A, and the 1:1 adduct, AB. These spectral curves define
the equilibrium in Eqn 2-4.
(2-4)
A + B
AB
Further base addition results in spectral deviation from the
first isosbestic point as a third species is formed in solution,
Eqn 2-5.
(2-5)
AB + B
In these equilibria, the free acid, A, refers to the bis ether
adduct, Cr2(tfa)ii(Et20)2. Dissociation of coordinated Et20
accompanies both equilibria. At high base concentration, an
isosbestic point appears at 515 nm which upon cursory
examination would appear to correspond to the second equilibrium
process in Eqn 2-5. Quantitative analysis (vida infra), however,
reveals that the limiting spectrum centered at 585 nm
corresponds to the AB2 chromophore while the isosbestic point at
575 nm can be assigned to yet another equilibrium, Eqn 2-6.
(2-6)
The absence of a clearly defined isosbestic point for the second

25
equilibrium implies one of two things: 1) K^>> butthe
spectral change associated with Eqn 2-5 is slight, and an
isosbestic point cannot readily be discerned. 2) K1 is not
significantly greater than and no appreciable amount of AB
forms in solution. The presence of a well defined isosbestic
point at low base concentration argues against the latter while
quantitative results (vide infra) indicate the former.
The observed spectral changes are understandable in terms of
the primary orbital interactions. Rice et al.^9 have examined
the single crystal polarized electronic spectrum of red chromous
acetate dihydrate, C^CC^CCH^C^O^ whose spectral features are
similar to that of Cr2(tfa)jj(Et20)2. Two bands are observed for
the former; the lower energy (465 nra, e = 120 M“1 cm-1)
transition is associated with a metal centered 6 ->71* promotion;
the other (333 nra, e = 200) is attributed to charge transfer
from a nonbonding carboxylate tt orbital to metal centered tt *
#
(np^->-rr ). Violet Cr2(tfa)¡|(Et20)2 displays a similar spectrum.
The ó ->tt* (550 nm, e= 133) and nn ->ir* (328 nm, = 380)
^TT
transitions are assigned by analogy. An MO diagram of the metal
centered orbitals for C^itfa)^ and their changes upon adduct
formation are shown in Fig 2-3. The dimer functions as a Lewis
*
acid, accepting electron density in the antibonding a orbital.
»
Upon complexation of a stronger donor, the metal a and a
orbitals become closer in energy while adduct formation is
realized through stabilization of the donor lone pair orbital.
The dimer orbitals (6->tt*) involved in the transition are not
directly involved in adduct formation. Weakening of the metal-
metal bond through partial population of a* decreases the d

cr*
IT* -r-
8*
8 44—
^ 44- 44-
Cr2(tfa)4 2L Cr2(tfa)4 LL L'
44-
^-44-
44-
44-"
—H-
ho
Fig 2-3. Effect of donor exchange upon metal centered MO's and corresponding change in optical transition.

27
orbital overlap, compressing the entire d orbital manifold in the
MO scheme. Replacing coordinated ether with stronger donors
gives a color change from violet to blue, consistent with the
expected red shift from the MO description.
Displacement of coordinated Et20 by stronger donors was
monitored by FTIR in order to verify the exchange processes.
Within the detection limits of the FTIR, coordinated Et20 was
completely exchanged for coordinated donor at 2:1 donortdiraer
molar ratio, indicating the exchange equilibria (Eqns 2-4 and
2-5) lie far to the right. Figure 2-4 shows an FTIR titration of
Cr2(tfa)¡|(Et20)2 with 0, 0.5, 1.0, 2.0, and 5.0 equivalents of
dimethylacetamide. Only above 2 equivalents is free DMA clearly
discernible. Figure 2-5 shows the ether vc_0_c region of the
same titration. Bound ether appears completely exchanged at 2
equivalents added donor. Some dissociation of diethylether may
occur even at 0 equivalents added donor as evidenced by a slight
shoulder absorbance at 1113 cm“^. Only a slight frequency shift
is observed for vCQ of coordinated dimethylacetamide as the
titration proceeds. A similar result is observed for the
asymmetric v0_c_0 stretch of bridging trifluoracetate, vc02>aSy:
equiv
DMA
VcoOMA)
“ co2»asy
0
1680.1 cm'
0.5
1610.0 cm-^
1680.3
1.0
1608.2
1682.1
2.0
1606.2
1684.0
5.0
1606.2
1683.9
free DMA
1651

28
(CM -1)
O
O 5
I O
20
5.0
Fig 2-4. FTIR titration of Cr-Ctfa^CEtgOjg (8.9 X 10“3M in
C H 2 C12) with 0, 0.5, 1.0, 2.0 and 5.0 equivalents of
dimethyl acetamide. Bound DMA: v co = 1160 cm-1; free: vco =
1639 era” .

29
Fig 2-5. FTIR titrationn of Cr
) with 0, 0.5, 1.0,
acetamide. Bound Et20:
= 1113 cm-1.
CH2C12
dimethyl
9(tfa)4(EtpO)2 (8
2.0 a nd 5.0 eq
v „ = 1053 cm"
v coc
.9 X 10"3 M in
uivalentsof
¡ free: vcoc

30
Coordinated DMA is relatively insensitive to the nature of the
trans coordinated ligand (Et2Ü or DMA). The slight shift which
occurs would indicate that DMA coordination is strengthened as a
stronger metal-ligand bond is formed on the opposite side. Such
a ’’trans strengthening” is surprising in light of results for the
Rl^ípfb)^ dimer (chapter 3). Across the rhodium-rhodium bond, a
sigma donor on one side decreases the strength of a metal-ligand
sigma bond at the other side, as the two donors compete for the
same rhodium d orbitals.
The observed shifts for coordinated DMA are too slight to
resolve into their individual contributions from 1:1 and 2:1
adduct species. The FTIR experiment does provide the satisfying
result that the exchange process is indeed ocurring and that it
is essentially complete after two equivalents of donor has been
added at these chromium concentrations.
The trifluoroacetates are convenient spectroscopic labels and
display strong sharp absorbances for the asymmetric and symmetric
carboxylate stretches which have shown utility in differentiating
70
between unidentate, ionic, bidentate, or bridging coordination.1
The analagous Mo2(tfa)jj displays the symmetric and asymmetric
stretches at 1592 and 1459 cm"'' (bridging). Upon adduct
71
formation, bulky phosphine donors occupy equatorial positions,'
and the resulting unidentate assymmetric stretch occurs at ca.
1680 cm -1. The IR spectrum of Cr2(tfa)¿j(Et20)2 (Fig 2-6) in
methylene chloride is assigned the bridging asymmetric and
symmetric stretches at 1680.1 and 1480.0 cm The similarity
of the molybdenum unidentate stretch (1680) and chromium

Fig 2-6. FTIR spectrum of Cr2(tfa)¿j(Et20)2, methylene chloride
solution.

9/o TRANSMITTANCE
800
2000.0 15000 1000.0
5000
WAVENUMBERS
LO
PO

33
bridging stretch suggests the possibility of unidentate
coordination for the latter in methylene chloride solution.
Observation of only one asymmetric stretch (four unidentate
carboxylates would disociate the complex) and an identical Nujol
mull spectrum rule out this possibility. Though not detected at
five equivalents, a large excess (50 equiv.) of DMA results in an
FTIR spectrum in which another asymmetric carboxylate stretch
appears at 1717 cm"'. This, presumably, is the unidentate
carboxylate which arises from equatorial coordinationn of DMA and
is consistent with the formation of the AB^ species postulated by
Eqn 2-6.
Similar titrations were performed and spectra recorded for a
range of donors. Results for methylene chloride solutions of
Cr2(tfa)i|(Et20>2 (ca. 5x10~^M) with two equivalents donor are
given in Table 2-3. Reported are the absorbances for the
appropriate functional groups of the free and complexed donor
molecules. Exchange adducts were readily isolated for several
donors by stoichiometric admixture and recrystallization from
benzene, and in all cases except that for DMTF, gave identical IR
spectra to those prepared from the same donor in situ. For those
donors whose functional groups showed a pronounced frequency
shift upon coordination, the FTIR spectra indicate the magnitude
of K1 and In all cases, no free donor could be detected at
the 1:1 level. With two added equivalents, the absorbance
spectra peak areas demonstrate 90% or greater complexation for
DMA, Et3P04, DMCA, DMSO, HMPA, and DMTF. The extent of exchange
could not be readily gauged for the remaining donors. A 1:1
equilibrium process (AB + B AB2) with initial concentrations

3^
Table 2-3. FTIR data for trifluoroacetate bridges and donor
functional groups. v
CO2* asy CO2»
functional
1 1
sym
group
free complexed
Rh2(tfa)4(EtOH)2a
1664
1467
—
—
—
Mo2(tfa)4b
1680 1592
1459
—
—
—
Cr2(tfa)4(Et20)2
— 1680
1480
o-c-o
1115
1053
(ch2ci2)
Cr,(tfa)4(Et,0),
1682
1484
o-c-o
—
1055
(mull)
(DMA)2
1717 1684
1476
c = 0
1651
1606
(Et^POjj) 2
— 1682
1478
P-O-C
1034,
,1037,
979
984
(DMCA),
— 1681
1477
C = N
2217
2239
(dmso)2
1680,
1478
s = 0
1057
1012,
1708
1022
(Me3PO)2
— 1681,
1473
P = 0
1179
1134
1712
(HMPA)2
1685
1473
P = 0
980
992
(dmtf)2
— 1679
1478
(d)
1538
1563
(dmf)2
1680
1477
c = 0
1676
1562
[(MeO)3P]2
— 1681
1477
(e)
—
—
(Me4Urea)2
1685
1477
C = 0
1640
1581
a) reference 70.
b) reference 71.
c) Slow oxidation follows coraplexation as evidenced by shift in
asymmetric stretch and color change from violet to green over 24
hour period.
d) assignment unknown.
e) Ligand vibrations shifted but specific shifts not assigned.

35
of 5x10-2M going to 90% completion yields K2 = 1800. Finally, as
with DMA, little or no frequency shift was observed for the bound
donor as the trans ligand was exchanged from diethylether to the
donor of interest.
An attempt was made to gauge the magnitude of the equilibrium
constants, and Kg, by calculating the mole fractions of bound
and coordinated diethylether from the fast exchange region 1H
NMR. The 100 MHz FT-NMR spectra of OgCtfa^EtgOJg and exactly
one equivalent added DMA and HMPA give the following diethylether
shifts in CgD^:
EtgO
5(-CHg-)
3.34(q)
6 (-CHO
1.07U)
Cn ^ ( tf3 )¿|(Et20) 2
5.22
1.40
+ 1 DMA
4.51
1.26
+ 1 HMPA
4.50
1.26
The chromium containing solutions all exhibit broadened
resonances, devoid of any spin-spin splitting. Treating the
coalesced resonance shifts as a mole fraction weighted average of
the free and bound species gives the same results for both HMPA
and DMA. Using the methylene resonances: 64% complexed, 36%
free; methyl, 58% complexed, 42% free. That different mole
fractions are calculated dependent on the resonance used is a
feature of the paramagnetic complex. Stronger donors raise the
paramagnetism of the Crgi,+ center (vide infra). The 1:1 adducts
then will display downfield shifted coalescence peaks.
Calculations thus favor a higher concentration of complexed ether
than is actually present. This effect should be more pronounced

36
for the methylene protons which experience a greater contact
shift. The only information these NMR 3pectra add to the
picture, then, is support for the paramagnetic nature of the
dinuclear metal center.
2. Quantitative Reactivity
Stepwise adduct formation between a metal center and first
one, then two donor ligands is described by two successive
equilibria (Eqns 2-4, 2-5). For the case where formation of the
first bond exerts no influence upon the second, A H-| = A H2 and
entropy considerations predict K-j = Inductive effects, as
in the case of the metal carboxylate dimers, perturb the second
equilibrium such that A H1 > A H2 and K1 > 41^. In this study,
adductformation is accompanied by a dissociation step,and the
reactions are described by exchange equilibria:
AL2 + B ABL + L (2-7)
ABL + B AB2 + L (2-8)
In subsequent discussion, the AL2 species is referred to a3 the
free acid and denoted by 0, the ABL as the 1:1 species (1) and
the AB2 as the 2:1 species (2).
a. Electronic spectra
UV-VIS titrations were performed by successive raicroliter
injections of a concentrated base solution into an inert
atmosphere cuvette containing the dimer solution. Despite the

37
presence of an isosbestic point for the initial spectral curves,
absorbance changes are too slight to allow satisfactory
definition of K.¡. In plots of K“^ vs. Ae , the resulting curves
intersect in a region of very small K_1 and small
absorbance/concentration errors can vary K by several orders of
magnitude. Consistent results are obtained, however, for Ae and
these values provide the initial estimates for the e at various
wavelengths.
Treatment of all the spectral curves using the program SPEC
allows the computer to calculate the best fit values of K-j, K2,
and £g, e .j and e:^ for the wavelengths analyzed (up to five).
The success of the fit relys upon good initial estimates of each
of these parameters. The £q value comes from the free acid
spectrum. The e -j values can be calculated at any wavelength
from any curve that passes through the first isosbestic point,
once Ae has been determined from a K~1 vs. Ae plot at one
wavelength. The 2 values are estimated from the limiting
spectra which result from just above 2 added molar equivalents of
base. Specific details and precautions have been described. 0
The spectral absorbance data are listed in Appendix I.
Iteratively varying some of the spectral paramaters while
fixing the remainder allows the computer to uniquely define the
extinction coefficients and the picture which emerges in the same
absorption profile for the free acid, 1:1 and 2:1 species,
successively red-shifted. This is the expected behavior based
upon the MO arguments described earlier. Based upon the large
equilibrium constants (vide infra) these three species can be
observed in solution at the appropriate stoichiometry. The

38
UV-VIS spectra of the dimers with 0, 1, and 2 equivalents of
dimethylacetamide are shown in Fig 2-7, roughly illustrating the
spectral curves which the e 's define. Calculated values for the
equilibrium constants depend upon the initial guesses.
Consistently, allowing the computer to vary only K2 gives values
of 4K2 slightly less (ca. 50S) than K-j. Again, however,
absorbance changes are too slight to define both equilibrium
constants, and a wide range of K2 pairs satisfy the spectral
data with impossibly small standard deviations. If were
located, K2 could be uniquely defined, but the lack of a good
estimate for prohibits their determination from the titration
curves. The spectral titrations indicate a relationship between
the equilibrium constants which would be consistent with weak
communication between the metal centers but do not adequately
define the magnitudes of the K's. The utility of these
experiments, then, lies in verifying that the exchange reactions
(Eqns 2-7 and 2-8) do occur, that their equilibria lie far to
the right and that the metal centers do not influence each others
coordination chemistry significantly.
Calorimetry
As Long has indicated,^® it is best to determine the
equilibrium constants which apportion the experimental heats and
the molar enthalpies from separate experiments, since the four
parameters, K^, K2, A H^.|, and A H2.1 are frequently highly
correlated. When this happens, a situation arises such as that
for the chromium spectral titrations above which a specific
solution relies on definition of (usually) one parameter.

39
Fig 2-7. Electronic spectrum of the chromium dimer (methylene
chloride solution) with 0, 1, and 2 equivalents DMA. Given the
large equilibrium constants for donor exchange, these curves
approximate the 0, 1 and 2 species.

40
Exceptions to the rule are instances in which the first or both
equilibrium constants are very large. Long showed that the
enthalpies for the rhodium systems, Rh^but)^ and Rh2(pfb)¡j,
could be reasonably determined from exclusively calorimetric data
Q
provided that was of the order of 10° or greater, as verified
by comparison of the enthalpies determined when the equilibrium
constants were solved for independently. A similar situation
occurs for the chromium calorimetric data, and the enthalpies are
extricable given the large equilibrium constants. Using the
program HEAT, values were fixed and the K2 varied to minimize
the conditional standard deviations associated with the
enthalpies. K-j was then varied by a factor of ten and the
minimization repeated. In this manner, the best K^, K2 pair
which minimizes the enthalpy deviationns was used to define
AH1#1 andAH2.^. The raw calorimetric data and best fit solution
values of each of the parameters are given in Appendix I. It is
not possible to assign deviations to the equilibrium constants
determined in this manner, and it was found that varying K^s of
this magnitude by a factor of 10^ had little effect (<1í) on the
enthalpies. Thus, though the enthalpies are uniquely determined,
is simply described as "large" and K2 represents the best fit
value. No inferences can be made for the relationship between K1
and K2 determined in this manner. Table 2-4 contains the results
of the calorimetric titrations.
Several aspects of the trends and some anamolies of the data
set deserve comment. First and foremost, the enthalpies of the
2nd exchange, AH2.^, are consistently lower than those for the
first, aH-|.i» in all cases except that for Me^PO. The lowering

41
Table 2-4. Thermodynamic data for the exchange reaction of
Cr2(tfa)¿j(Et20)2 with various donors.
Base
K1
K2
- AH1
(kcal
; 1
mol"1)3
" A H2:1
(kcal
mol-1
DMTF
1 X
105
500
1.89
(0.06)
1.80
(0.12)
DMCA
1 X
108
3 X
103
3.75
(0.14)
2.66
(0.24)
DMF
1 X
108
1 X
104
4.52
(0.31)
3.01
(0.54)
(MeO)3Pb
1 X
108
2 X
103
0.79
(0.06)
0.73
(0.10)
Et^POjj
1 X
1010
1 X
106
4.37
(0.18)
3.05
(0.30)
DMA
1 X
109
1 X
105
5.26
(0.24)
3.59
(0.39)
DMSO
1 X
109
6 X
104
5.58
(0.27)
3.95
(0.45)
Me3PO
1 X
107
2 X
104
5.95
(0.26)
6.55
(0.43)
HMPA
1 X
1011
6 X
106
5.97
(0.23)
5.48
(0.37)
a. Values in parentheses are conditional standard
deviations.
b. Enthalpies are tentative; see text.

42
is slight but measurable, manifesting the inductive effect of the
first coordinated donor upon the second exchange reaction. As
expected, the inductive influence is small, in keeping with the
long, weak metal-metal bond in Cr2(tfa)2j(Et20)2* and is consistent
with the K^, K2 relationship inferred by the spectral data (vide
supra). Second, the trends in the enthalpy values is in keeping
with our intuitive knowledge of donor strengths, an aspect which
is taken advantage of in quantitative correlations to follow.
The higher >AH^.^ value for Me^PO is inconsistent
with the remainder of the data set and probably results from
subsequent further reactions. The spectral titration for this
system was well behaved only to a 2.0 molar ratio of Me^PO to
dimer at which point the solution began to turn green, indicating
oxidation of Cr(II). Heat evolution during calorimetric
titration for the other donors was essentially complete following
addition of two equivalents of donor. Heat continued to be
evolved in the Me^PO calorimetric titration which was carried out
to a 2.5 molar ratio. The stability of the dimer adducts toward
oxidation is decreased by donor complexation (vide infra) and the
C^ttfa^iMe^PO^ adduct may be readily oxidized via oxygen atom
transfer from the phosphine oxide.
Another inconsistency in the data set is the apparent
anomalously low heats observed for (MeO^P despite the agreeable
AHi;i >AH2;i relation. Again, the spectral titration lends
insight. During intermediate stages of the titration (ca. 1:1)
greatly decreased (< 50%) absorbances were recorded in spite of
the presence of an isosbestic point at lower ratios. This
behavior would arise if 1) extinction coefficients for the 1:1

43
species were much lower than for the 0:1 or 2:1 species or 2)
solid formation is occuring in this region. The latter was
observed to be a common phenomenon in titrations with donors
containing two potential donor sites. Maintaining a strict 1:1
stoichiometry allowed preparation of polymeric species and
clearing of the solution for reactions with p-dioxane,
tetramethyl urea and 1-methyl imidazole. Solid formation was not
obvious with (MeCO^P but could occur if both phosphorous and
oxygen donor sites are stronger Lewis bases than the ether
oxygen. A low value for AH-j, for endothermic desolvation,
however, would be compensated by a high value of AH2 for
exothermic solvation, and this is not observed in the
calorimetric data. The low heats for (MeCO^P remain a surprise
and should be considered tentative.
Solution studies were conducted in methylene chloride solvent
which behaves as a Lewis acid towards donors. Solvation effects
are anticipated and can be corrected for since methylene chloride
has been shown to undergo primarily specific reaction (electron
pair sharing of a Lewis base with a Lewis acid) with donor
molecules.To assess their contribution to the experimental
heats, the enthalpy components are specified:
AH! AH2 AH3 AH4
(Et20)Cr2-(Et20) + B-S (Et20)Cr2-B + Et20-S (2-9)
ah5 ah2 ah6 ah4
(Et20)-Cr2B + 3-S B-Cr2B + Et20-S (2-10)
AH1.1 = AH4 + AH3 - AH2 - AH1 (2-11)
A H2.1 = A H4 + A — A H2 — A Mg
(2-12)

44
Assuming nonspecific solvation differences between all the
various chromium adduct species to be negligible, the solvent
contributions, AH2 and A , can be calculated from the Eg
and Cg values for the various donors and the E'A and C'A values
for methylene chloride, 1.66 and 0.01, determined by Drago et
al.^ The primes indicate that these are the best fit values for
methylene chloride adduct formation with a series of donors and
may contain small contributions from nonspecific solvation.
Solvation corrected enthalpies are given in Table 2-5. In three
instances, the donor Eg and Cg have been determined from a
limited data set and are not well defined. The corrected heats
for (MeO^P, Et^PC^ and Me^PO are thus considered tentative. To
alleviate propagation of their uncertainties, these three heats
will not be used in the following analysis.
Included in Table 2-5 are the experimental frequency shifts
from the spectral titrations. The observed red shifts indicate
the perturbation on the transition orbitals which results upon
donor exchange and provide another measure of the strength of the
interaction. The perturbations roughly follow the measured
enthalpies and include two donors not looked at by calorimetry.
Dimethyl sulfide titrations gave very small absorbance changes
and heats could not be accurately measured. Spectral changes are
apparent only at high donor concentration in which entropy
effects predominate and the weaker donor displaces ether to give
a blue shift. Me^Urea forms a polymer with the dimer at 1:1
molar ratios and redissolves as more donor is added, prohibiting
calorimetric determination of the enthalpies.

45
Table 2-5. Experimental frequency shifts and solvation corrected
enthalpies for the exchange reaction of C^Ctfa^tEtgO^ with
various donors.
-ah2;1
Base
(kcal mol“^)
(kcal mol~^)
Av.
Me2S
a
—
-480
o
OJ
-p
w
0
0
0
DMTF
b
—
705
DMCA
3.96
2.87
710
DMF
4.96
3.45
770
(MeO)3P
C 0.13)G
(0.07)
1010
Et^POjj
(5.02)°
(3.70)
1160
DMA
5.85
4.18
1160
Me^Urea
a
—
1270
DMSO
6.20
4.57
1285
Me3PO
(6.89)c
(7.49)
1720
HMPA
6.90
6.41
2120
a. Nature of reaction prohibits enthalpy determination by
calorimetry.
b. Eg, Cg values unknown.
c. Calculated from tentative Eg, Cg numbers.

46
c. ECW model
We are now in a position to correlate the solvation corrected
enthalpies (corrected for AHg and AH¿j) whose contributions are
given below:
AH^.^ = A- A (solvent corrected) (2-13)
AH2.1 = AHg - A(solvent corrected) (2-14)
For the observed 1:1 heats, AH^.-|, with a series of
different donors, the contribution A is independent of the
donor employed and must be treated as a constant, W. Using Eqn
2-1, where AH = AH^ and W = A Hp gives the equation:
-AH1;1 = -AH3 + Ah, = Ea1:1 Eb + CA1:1 CB (2-15)
Using the experimental solvent corrected heats for the six donors
whose Eg and Cg are well defined along with the Eg and Cg numbers
from Table 2-2 allows calculation of the constant contribution,
1 # 1 1*1
AH1fand the EA * and CA * associated with the acid which
defines Ah^. This acid is the 1:1 adduct (Et20)Cr2(tfa)¿j. A
plot of the simultaneous equations is shown in Fig 2-8. The best
fit values for EA, CA and W along with their standard deviations
are EA1:1 = 13.6 (0.75), CA1:1 = -1 .57 (0.24) and W = 8.00
(0.94), all in kcal mol The experimental enthalpies used in
the correlation along with those calculated from Eqn 2-15 are
given in Table 2-6 and demonstrate the quality of the fit.

Fig 2-8. Plot of Ea v3 CA for Cr2(tfa) jjCEtgO)

48
Table 2-6. Enthalpy data used to determine acid parameters for
Cr 2(tfa)¿j.
Usedto determine Used to determine
E^:^ and C^:^ k and k'
Base
~AH1:1,exp _AH1:1,calc aunexp “ ncalc
AAH,
AAH
(kcal raol“^)a (kcal mol-^)'5 (kcal/mol)c (kcal/raol)^
Et20
0
0
0
0
DMCA
3.5
(0.1)
3.7
1.1
(0.3)
1.1
DMF
4.1
(0.3)
4.0
1.5
(0.6)
1.4
DMA
4.7
(0.2)
4.8
1.7
(0.5)
1.7
DMSO
5.0
(0.3)
4.5
1.6
(0.5)
1.7
HMPA
5.0
(0.2)
5.3
0.5
(0.4)
a.Solvation corrected values. Quantity in parentheses is the
conditional standard deviation for the measured uncorrected
heats.
b.Calculated from equation 2-15.
c.Difference between the first and second experimental
enthalpies. AAH = AH2>1 - A H1.1 Quantity in parentheses is
the propagated conditional standard deviation, calculated from
a * = a « + a 0
B
d.Calculated from equation 2-17

49
Several points about these best fit parameters warrant
comment. The calculated W = AH1 = 8.0 kcal mol -1 refers to
breaking of one of the adduct bonds in (Et20)Cr2(tfa)¿j(Et20).
The relatively large EA^:"* and small CA1:^ demonstrate a
pronounced tendency for this Cr(II) acid center to interact
primarily in an electrostatic sense as might be expected for a
first row transition metal ion. The negative CA1:1 value does
not imply that the dimer interacts covalently in an antibonding
sense but simply that this is the best fit parameter to a data
set in which implicit assumptions have been made in defining the
magnitudes of the original parameters.®® For all intents and
purposes, the CA1:”' value merely suggests little or no covalent
interaction between the (Et20)Cr2(tfa)¿j adduct and donors.
The real quantities of interest for comparison to other dimer
systems are the acid parameters associated with the naked dimer,
C^Ctfa)^ and are extricable from an analysis of the 2:1
enthalpies. The EA1:1 and CA 1:1 values are defined by Eqns 2-2
and 2-3.
EA1:1 = EA - k EB (2-2)
CA1:1 = CA - k' CB (2-3)
Thus, determination of the inductive transfer parameters, k and
k'f and using the base parameters for Et20 allows calculation of
Ea and CA for C^itfa^. By substituting Eqns 2-2 and 2-3 into
the E and C equation, 2-1, the first and second enthalpies are
32
related to each other by

50
- AH2;1 = -AH1;1 -kEB2 - k'CB2 (2-16)
In the present study, where an exchange reaction is being
studied, Eqn 2-16 takes on a slightly different form. Defining
all the enthalpy components AH-|, AH^, AH^ and AHg in terras of
their E and C components results in many cross terms which define
the perturbation that coordinated ether makes on the first
exchange reaction and that coordinated base, B, makes on the
second. The derivation is given in Appendix III and the
simplified solution has the familiar form
- A H2; i = -AH1:1 - k(Eg - EEt20)2 “ k?^CB " CEt20^2 t2-1?)
This equation has the same form as Eqn 2-1 and the ECW
program is used to solve for k and k'. The input data are the
heat differences as AAH, the squares of the quantities in
parentheses as the base parameters and W is, of course, zero.
The experimental heat differences, AAH = AH7.^ - AH1#^, used
to determine k and k’, along with those calculated from the best
fit parameters are given in Table 2-6. The attendant conditional
standard deviations are given in parentheses. The value for HMPA
was rejected in light of the large standard deviation and k and
k' were determined from the remaining five data points. The
best fit values (standard deviation) are k = 1.54 (0.13) and k* =
0.0079(0.0043). These are the parameters associated with the Cr-
Cr bond and demonstrate the ability to transmit electrostatic (k)
and covalent (k') effects. Equations 2-2 and 2-3 may now be used

51
to calculate the EA and CA values for the naked chromium dimer,
Cr2(tfa)4. The derived values are given in Table 2-7 along with
those for the molybdenum and rhodium dimers for comparison.
Again, the large EA and very small (negative) CA demonstrate
the pronounced Lewis acidity of the chromium dimer and its
tendency to interact with Lewis bases in primarily an ionic
fashion. This finding is consistent with the lack of
experimental success towards generating a dinuclear chromium
carboxylate free of axial coordination.11 In the absence of
donor molecules, the chromium carboxylates exist as polymers
where the very ionic carboxylate oxygens of neighboring dimers
serve as donors.
The low k* value which corresponds to transmission of the
base covalent parameters indicates that base binding does not
serve to polarize the bonding density of the chromium-chromium
bond, consistent with the poor orbital overlap between the two
metal centers. The rather large k value which corresponds to
transmission of the base electrostatic parameter is rather
surprising. Electrostatic interaction of the base dipole on one
side with the second metal center and with the second base dipole
p
should vary as 1/r . Of the three systems studied, the chromium
dimer exhibits the longest metal-metal bond (2.54 A for1111
Cr2(tfa)4(Et20)2, 2.39 A for73 Rh2Ac4(H20)2 and 2.09 Á for73
Mo2(tfa)4) and should exhibit the smallest value for k in the
series. For comparison purposes, it is probably more appropriate
to consider the relative values of k for each metal system. The
k values quantify the change in EA at the second metal center as
a consequence of bonding a donor at the first, irrespective of

52
Table 2-7. Acid parameters for various dinuclear carboxylates.
ea
CA
k
k'
Rl^Cbut)^
5.06
1.74
1.16
0.0364
Mo2(pfb)|j
5.92
0.385
1.46
0.022
Cr ^ (t Ts)
15.1
-1.54
1.54
0.0079

53
the magnitude of EA. The roughly twice as large EA for the
chromium system would display relatively half as large
perturbations on the magnitude of EA compared to the other two
dimers. This insight, then, is consistent with the longer metal-
metal bonds in the chromium dimers. Thus, more appropriate
measures of the ability of a metal-metal bond to transmit
coordination effects would come from k/EA and k'/CA instead of
comparing the absolute k and k* values.
d. Magnetic susceptibility
Theoretical investigations on dichromium tetraformate at the
SCF-level have predicted both the quadruple bond, and
P p # p # p
no bond, a 66 a , configurations, neither of which
correspond to realistic descriptions.^ Incorporation of
correlation effects which allow mixing in of excited states in
P ft
the ground state description give more satisfying results^0 in
P Ü P
which the term is important, but the ground state bond
order is closer to 1.5. Dichroraium tetraacetate is diamagnetic^
and deMello et al. have proposed two antiferromagnetical ly
coupled chromium centers and no net covalency between the
¿1.
chromium atoms as the dominant description for Crg centers.
This latter description, too, suffers from inconsistencies by
predicting unrealistically long Cr-Cr distances. At shorter
(1.8-2.5 t) experimental bond lengths, d orbital overlap is
expected, and Zerner qualifies his calculations to suggest that
some degree of covalency accompany the antiferromagnet ic
description.
The trifluoroacetate bridged dimer studied here may represent

54
just such a borderline example with a chromium-chromium
separation of 2.54 A. An early magnetic study' reported
magnetic moments of u 0.74 BM for "CrCF^CCOO^" and 0.85 BM
for "CrCF^CCOO^.Ather". These species are presumably dimeric,
and the increase in magnetic moment upon ether complexation is
consistent with both an antiferroraagnetic description and the
proposal that ligand donors weaken the metal-metal bond through
partial population of the a* orbital.46,77 xhe onpy other first
row transition metal for which a range of metal carboxylate
dimers has been prepared is copper. The cupric carboxylate
dimers exhibit long (2.6-2.9 A) metal-metal distances and
incomplete spin pairing between the two d^ centers. An
impressive array of these complexes^ has been prepared and
characterized magnetically and structurally, with a general goal
of determining the factors (structural parameters; nature of
bridge; axial ligands) and mechanism which contribute to spin
exchange. The following conclusions appear to be general for the
cupric carboxylate dimers:
1) the unpaired electrons reside in the dx2_y2 orbitals;
overlap is minimal at these metal-metal distances, ruling out any
direct exchange or covalency;
2) overlap with carboxylate tt orbitals allows a super¬
exchange pathway via the carboxylate bridges, resulting in
antiferromagnetic interactions with -2J ranging from 217 to
555 cm-1;
3) the singlet-triplet separation (-2J) is relatively
insensitive to the Cu-Cu distance;
4) -2J is sensitive to the Cu-0-C-0-Cu bridge distance and

55
angle, generally decreasing at longer distances;
5) -2J is sensitive to axial ligation, generally increasing
as the terminal ligands become stronger electron pair donors.
The chromium dimers, on the other hand, exhibit some degree
of orbital overlap, suggested to be incomplete in the
Cr2(tfa)¿j(Et20)2 adduct. The strong dependence of the Cr-Cr
distance upon axial ligation suggested magnetic studies to
complement the copper studies in treating a metal-metal bond with
some degree of covalency.
The room temperature magnetic susceptibilities of various
adducts of the form C^Ctfa^I^ were investigated by the solution
Evans method^ by generating the complexes in situ. This method
allows calculation of experimental magnetic susceptibilities by
measuring the proton chemical shift of an inert substance in the
presence of a paramagnetic complex. A coaxial tube arrangement
was used which contained a solvent system of 2% v/v CgHg and 2%
v/v TMS in CgDg in both the outer 5mm tube and the inner 1mm
capillary. The inner capillary also contained the paramagnetic
species at a concentration of ca. 5x10 M. For an inert
substance (in this case, CgHg and TMS) the shifts caused by the
paramagnetic substance when one employs a nonsuperconducting NMR
instrument are given by
AH = 2 7T . (2-18)
I 3 4K
where Ak is the change in volume susceptibility. With the
cocentric tube arrangement, two resonance lines will be obtained

56
for each standard, with the line from the more paramagnetic
solution lying at higher field. The mass susceptibility, x » of
the dissolved substance is given by Evans as
X
5Av
2irvm
Xo
+ Xo(d0-d )
m
(2-19)
where Av is the frequency shift separation in Hz, v is the
spectrometer frequency (99.55 MHz), m is the mass of substance
contained in 1 ml solution, xo is the mass susceptibility of
the solvent, dQ is the density of the solvent and ds is the
density of the solution. Brault and Rougee have presented®0 a
modified form of Eqn 2-19 to calculate the molar magnetic
susceptibility
XM
5Av 1000 +
2ttv c
XoM
X
D
(2-20)
where C is the molar concentration of the paramagnetic complex, M
is the molecular weight of the complex and xd i s the
diamagnetic susceptibility of the paramagnetic complex. As
Desmond points out,®^ the original Evans equation required noxn
term since complex dismagnetic susceptibility was compensated for
by including an equal concentration of ligand in the reference
solution. The Brault and Rougee equation is more general but
fails to account for density difference contributions. The
density term can often be neglected but becomes important when
measuring the small shifts for weakly paramagnetic substances.
Desmond includes the density term to obtain the equation

57
= 5Av 1000. + y m . y + Xo(d0-d ) 1000, (2-21)
XM 2ttv c Xo xD c
The Xd term is usually approximated from Pascal’s constants®^
and may be quite erroneous for large molecules. To alleviate
this hazard and the need for a density determination, the
diamagnetic correction may be obtained by NMR. Setting = 0
and moving the first two right-hand terras of Eqn 2-21 to the
other side gives an expression for the combined diamagnetic
susceptibility and density terms. An Evans experiment performed
on the diamagnetic analogue Rl^Ctfa^CTHF^ gave no peak shift
nor was any asymmetry evident in the TMS or peaks. The sum
of the diamagnetic term and density correction then is -5.63 x
10-i< (obtained by setting both X^ and Avequal to zero). The
calculated diamagnetic term is -3.0 x 1 0_i* (from Pascals
constants, where each Rh(II) is given a value of -20 x 10“^ ml
raol-^), and the difference approximates the density correction.
An upper error limit on the shift determination is about 0.3 Hz
which corresponds to a 0.43 x 1Q~^ contribution to the
diamagnetic/density term. Substitution of the value for the
combined complex diamagnetism plus density correction into Eqn 2-
21 gives
xm = : Av 1c°Q~ + XoM + 5*63 x 10"4 (2-22)
At the concentrations used, complex precipitation warranted shift
determination in CD2CI2 in several instances. Equation 2-22 was
used to calculate molar magnetic susceptibilities from both the

58
and CD2CI2 solution data. The single greatest source of
error in these measurements probably comes from determining the
shift magnitude. If the error calculated for the diamagnetic
correction is doubled to account for dual shift measurements (one
for the rhodium complex and one for the paramagnetic complex) the
attendant error is estimated to be 0.86 or about 1 x 10*"^.
Use of the "spin only" formula for the molar susceptibility
allows determination of the effective magnetic moment
P eff = 2.84(x mT)1/2 (2-23)
where the product is expressed in Bohr magnetons. X m is a molar
quantity, calculated per mole of dimer, while yis per metal
and must be determined from 1/2 X M.
The molar magnetic susceptibilities and moments calculated
from the observed frequency shifts are reported in Table 2-8.
Benzene solution was the method of choice (economy), but some of
the dimer adducts were too unstable towards oxidation in this
solvent, necessitating determination in methylene chloride.
Immediately obvious is a very satisfying trend towards larger
susceptibilities and moments with increasing donor strength.
This relationship provides the first conclusive evidence for
genuine paramagnetism in a series of chromium carboxylates since
this consistency would not be observed with random trace
contamination by Cr(III) impurities.11
The measured susceptibilities roughly parallel the
experimental enthalpies except in one instance. The DMCA adduct
displays a larger moment than might be expected from the heat

59
Table 2-8. Magnetic susceptibilities and moments for C^tfa^L^
species.
conc(mM)
Av(Hz)
XMax104
yeff(B.M.)
RMtfa)4(THF)2
26.2
0
0
0
Cr2(tfa)4(Et20)2
28.4
3.54
6.66
0.89
[(MeO)3P]2
46.9
6.77
6.91
0.91
c
50.4
7.32
6.95
0.91
(Et3P04)2d
61.3
13.24
10.94
1.15
(DMTF)2d
62.1
13.98
12.40
1.22
(Me4Urea)2
51.1
15.38
14.53
1.32
(DMF)2d
21.3
17.64
18.18
1.48
(DMS0)2d
62.2
24.66
20.73
1.58
(DMA)2d
62.1
24.84
20.80
1.58
(DMCA)2
27.19
20.50
36.90
2.11
c
88.31
62.20
34.52
2.04
(HMPA)2
17.7
17.64
47.01
2.38
(MenPO)o
20.9
33.20
77.86
3.06
a) per dimer, + 1 X 10~\
b) per Cr(II) atom, + 0.35 B.M.
c) multiple determinations.
d)determined in CD2CI2

60
data. Subsequent redetermination gave essentially the same
value, however, and both are reported.
Besides the chromium and widely studied copper systems
reports of two other discrete first row transition metal
carboxylate dimers have appeared, both of which display weak
antiferromagnetic exchange interactions. Dicobalt (II)
tetrabenzoate bis quinuclidine,®-^ C^iCgHcjC^^Cquin^» shows a
O
long Co-Co distance of 2.83 A and the high spin Co(II) centers
display weak anti ferromagnetism of -2J = 38 cm“\ At this long
metal-metal distance, spin coupling is expected to occur via the
bridge super exchange pathway. A vanadium complex,^
V2^fa)n(.Cc)ti^)2t exhibits a V-V distance of 3.7 A and weak
antiferromagnetism, suggested also to proceed through a super
exchange path involving the carboxylate bridges.
Treatment of the exchange mechanism operative in the chromium
dimers is necessarily more complex than for the copper
carboxy lates,®** the former involving S = 2 metal centers. MO
calculations support direct orbital overlap and theoretical
treatment would include both direct and super exchange pathways.
Calculations showing low lying excited states do not rule out the
possibility of thermal equilibrium between the populations of
ground level and the first excited level with other multiplicity
(singlet-triplet equilibrium).^ The latter ha3 not been
demonstrated in the binuclear carboxylates but occurs in
monomeric complexes when the ligand field splitting is close in
energy to the electron pairing energy. The trend towards larger
moments with increased ligand field, however, is not consistent
with a singlet ground state.

61
Determination of the singlet-triplet splitting (-2J) can only
come from measurement of the temperature dependence of the
magnetic susceptibility, and these data are not yet available for
the chromium trifluoroacetate dimers. The data in Table 2-8 do,
however, suggest a direct exchange mechanism. In modeling the
magnetic susceptibility data for some 140 compounds of the
binuclear copper carboxylates, Jothara and Kettle observed a
general trend for J to increase as either the terminal or the
bridging ligands become better electron donors.At a given
temperature, larger values of -2J would give smaller experimental
moments. Thus in these systems in which the super exchange
pathway predominates, better donors give smaller magnetic
moments. The reverse is true for the chromium series studied
here. Donor lone pairs interact primarily with the chromium
dimer a orbitals. Stronger donors would serve to destabilize
the Cr-Cr a orbital while stabilizing the Cr-Cr cr* orbital, and
partial population would weaken the covalent bonding, consistent
with the observed trend for larger moments with better donors.
e. EPR, Electrochemistry
The cyclic voltametry of Cr2(tfa)¿j(Et20)2 was attempted with
the aim of finding the proper conditions for controlled oxidative
electrolysis. An EPR investigation of the cationic product might
add insight into the nature of the chromium-chromium bond. The
EPR prarmeters would be descriptive of the HOMO orbital which
surrendered the electron. In methylene chloride, the dimer was
found to be incompatible with the electrolytes Bu¡jNBFjjf Bu^PF^

62
and BujjNI, all of which reacted to give blue or green solutions.
No reaction was observed with Bu^NClO^. No oxidation wave could
be detected in solution, however, with this electrolyte though a
slow irreversible reduction occurred at -0.45 V (Ag/AgCl). No
further attempts were made to oxidize the complex.
The magnetic studies suggested unpaired electron density at
the Cr(II) centers, prompting investigation by X-band EPR. The
solid state adduct Cr2(tfa)2|(Et20)2 displayed no EPR down to 10K.
In CH2C12 solution, the ether and HMPA adducts gave no signal
down to 85K. One attempt was made to record the solution
spectrum of Cr2(tfa)ij(Et20)2 at liquid helium temperatures. A
1x10 M CH2CI2 solution glass gave the spectrum in Fig 2-9. The
observed asymmetric signal displays no -‘^Cr (9.5%) hyperfine
splitting and would appear to be that for a system of axial
symmetry with g^ = 1.98 and g1 1 = 1.87. A forbidden S = 2
transition may be present at H/2 (1600 G) but this is speculative
in lieu of experimental clarification. Lack of any observed
signal for a non-dilute powder sample at the same temperature is
consistent with significant spin lattice relaxation while lack of
any EPR at 77K (methylene chloride glass) would indicate
extensive spin-orbit coupling. EPR in d1* systems is very rare
due to short spin lattice relaxation times and a large zero
field splitting. A comprehensive review of the literature,^
albeit in 1972, yielded only five d1* systems for which an EPR had
been reported; three of which were Cr(II) as dilute ionic salts.
The estimated g values reported here do not, however, resemble
any of those reported for Cr(II) or Cr(III) systems. This

Fig 2-9. The X-band EPR spectrum (CH2CI2 glass, 9K) of
Cr2(tfa)i|(Et20)2. The field ranges from 1-5 kG.

64

65
finding warrants further investigation and is reported here only
as a matter of record.
3. Further Reactivity
Chromium (II) salts have been extensively employed as
O O
reducing agents in preparative organic chemistry.00 Chromium
(II) chloride, sulphate and perchlorate are similar in their
scope; chromium (II) ethylenendiamine cation functions as a more
efficient reducing agent while chroraous acetate is a milder
reducing agent, reacting under relatively neutral conditions. By
contrast, the utility of chromous trifluoroacetate as a reducing
agent remains relatively unexplored and the ligand exchange
reactions explored here offer some insight into its redox
chemistry.
Reduction of alkylhalides by chromium (II) salts to generate
olefins, carbenes, or alkyl chromium species via haloatom
abstraction represents one of the best known applications of
these reagents. Methylene cloride solutions of chroraous
trifluoroacete were found to be stable for several months under
dinitrogen as evidenced by retention of the robust purple color.
g
Irradiation (550 nm, 6-»- tt ) of solutions for 30 minutes in the
UV-VIS beam path also showed no appreciable spectral changes.
As mentioned for the case with DMA, UV-VIS titrations were
well behaved, defining a two step exchange process. Deviation
from the 2:1 limiting spectrum with excess base (> 5 equiv.) was
apparent in all cases and was accompanied by an immediate blue to
green color change which became more pronounced with additional
added base. FTIR titrations revealed a new vCo2,asy at 1?05

66
to 1710 cm-1 with excess (> 2 equiv.) donor which was assigned to
raonodentate trifluoroacetate. Chromium (II) oxidation would thus
appear to be facilitated by bridge removal and equatorial donor
coordination. Additionally, except for C^Ctfa^l^ species
where L = EtgO, DMTF or Et^POjj, the final blue solutions from
both spectral and calorimetric titrations turned green within 2
to 24 hours despite rigorous exclusion of dioxygen, indicating
oxidation when excess donor was present in solution. These
observations are consistent with electrochemical measurements on
a series of tetracarboxylato-dirhodiura (II) complexes.^ Das et
al. found that electron withdrawing substituents on the
carboxylate bridges produced less easily oxidized dimers. In
their studies, the lower oxidation state of Rl^itfa)^ was
stabilized to such an extent as to see no oxidation step within
the solvent limits. A similar trend in solvent effect was
observed with coordinating solvents such as pyridine and DMSO
giving the most negative oxidation potentials. The chromium
dimers are stable indefinitely in solution if stoichiometric
amounts of base are present, allowing ready isolation of the
exchange species.
A crystal structure has been reported-^ for the reaction
product of the strong donors pyridine and 4-cyanopyridine with
Cr2(02CCF2H)|j(Et20)2. The isolated crystals are trichromium (II,
III, III) compounds with the basic iron acetate trimer structure.
No spectral results were reported. The same procedure was
followed using the trifluoroacetate and pyridine to give long (up
to 2 cm) pale green needles which turn dark olive upon exposure
to air. Assuming the same structure, the bridging

67
trifluoroacetates show a vco2>asy at 1671 cm-1 in the mixed
valence trimer. Attempts to induce triraer formation with the
donors DMA and DMCA were unsuccessful, surrendering no solid
products after three weeks. Both the blue DMA-containing and
dark green DMCA-containing solutions gave VC02 aSy peaks at
1717 and 1713 cm-^, respectively, indicating monodentate
trifluoroacetate. Triraer formation does not appear to be
general, with most donors simply inducing slow oxidation of
chroraous trifluoroacetate, and the effect is more pronounced in
benzene than in methylene chloride.
C. Conclusion
A thorough investigation of the coordination chemistry of
chroraous trifluoroacetate has been performed with the aim of
understanding the transmission of bonding effects through the
weak Cr-Cr bond and the perturbation of donor ligands on the
metal-metal bond. Spectral titrations indicate that donor
ligands decrease the d orbital overlap between the metal centers.
Calorimetric studies show a lower enthalpy for second base
coordination resulting from transmission of donor effects from
one metal center to the next. The effect is less pronounced than
in the previously studied Rh2(but)¿, and Mo2(pfb)¿j systems. A
correlation analysis of the calorimetric data allows description
of the Cr(II) Lewis acidity as a relatively strong acid,
interacting in primarily an electrostatic fashion with Lewis
bases. Excellent data fits support an inductive transfer model
used to describe communication between the metal centers.

68
Magnetic susceptibility measurements on a range of Crjitfa^L^
adduct species reveal a pronounced influence by donors upon
dimer paramagnetism and support a direct exchange pathway for
spin pairing. UV-VIS, NMR and IR data support the donor exchange
reactions proposed and reveal a destabilization towards
Í1 .
oxidation when stronger donors coordinate to the center.
D. Experimental
1. Data Analysis
Programs. The following computer programs were used for
analysis of raw spectral and calorimetric data.
AB
C' o
Written by J. R. Long, 0 this program utilizes a standard non
linear least squares routine to provide the best AgCe^ - e0) arKl
K which describe a 1:1 equilibrium, A + B <=?AB. Raw data needed
are concentrations and absorbance changes for a spectral
titration. Raw heats can be used to solve for AH and K for a
calorimetric titration.
SPEC
Written by T. Kuechler,^® SPEC utilizes concentrations and
spectral changes to solve for the best K-j and K2 via least
squares for a two step equilibrium. The large number of unknowns
solved for (K^, K2, and Eg, e^, at each wavelength used)
requires good initial estimates of each of these parameters to
avoid false minima. Useful algorithms for initial estimates are

69
provided by J. R. Long.6® In practice, best results are obtained
by fixing most of the parameters (usually the e's) and allowing
the computer to vary those that are least well known (usually the
K's).
HEAT
ft
Written by J. R. Long, 0 this least squares program solves for
the best fit molar enthalpies, AH^ and AH2, for a two step
equilibrium. Input includes concentrations, raw heats and
equilibrium constants for a calorimetric titration.
Error analysis. Output parameters from the above programs
are provided with marginal and conditional standard deviations,
MSD and CSD, which demonstrate how well the model fits the data.
The conditional standard deviation defines the magnitude of error
while the MSD/CSD ratio indicates how well defined the parameter
is.68,90 For a rea3ona5ie CSD, results are considered meaningful
if the ratio is less than 4, tentative if the ratio is between 4
and 12, and not meaningful if the ratio is larger than 12.
ECW Program. Revised by M. K. Kroeger, this program uses
experimentally determined enthalpies to calculate the best fit
values of the unknown parameters EA, CA (and W if necessary) to
the following equation.
-AH + W = EAEB ♦ CACB

70
The Eg and Cg are found in references 38 and 39. Alternatively,
measured heats and and known EA, pairs can be used to solve
for Eg, Cg. A description of the computer program has been
Q1
previously reported.^
2. Materials
Metal complexes. All syntheses and manipulations were
performed under dinitrogen using Schlenk techniques or an inert
atmosphere box. Chroraous complexes are generally quite oxygen
sensitive and the compounds used here are oxidized within seconds
upon air exposure.
Dichromium (II) tetrakistrifluoroacetate bisdiethylether, I.
Chroraous acetate was generated from Zn/Hg reduction of Cr(III)
(aq) and sodium acetate.^ Yellow chromous carbonate was
synthesized through exchange of the acetate bridges by reaction
with potassium carbonate in water.^3 Chroraous trifluoroacetate
was prepared by a modified literature procedure.Under
dinitrogen, 5.5 g (9.9 mmol) KjjC^ÍCO^ijíí^O)^ and 8.0 ml (100
mmol) trifluoroacetic acid were refluxed in 80 ml deaerated
diethylether for six hours. The Schlenk flask was plunged
briefly into a dry ice bath, the purple ether layer decanted from
the frozen blue aqueous layer and the ether removed by vacuum.
Extraction with benzene followed by two recrystallizations from
benzene and vacuum drying (30 rain, 25°) produced purple blocks of
C^itfaJjjiEtgO^. Prolonged evacuation 06 hr) was found to
strip off coordinated ether. Final yields were typically about

71
20Í. Attempts to purify by sublimation decomposed the complex.
Repeated elemental analyses typically showed loss of 5—10%
coordinated diethylether. Calculated for ci6H20Cr2F12°10:
C, 27.29; H, 2.86; Cr, 14.77. Found: C, 25.91; H, 2.82;
Cr, 15.01.
Dichromium(II) tetrakistrifluoroacetate bisdimethylthioform-
amide. I, 34 rag (0.048 mmol) was dissolved in 1 ml benzene.
Addition of 9.0 y1 (0.11 mmol) DMTF produced a clear solution and
a fine blue powder within ca. 5 min. The solid was filtered,
washed once with cyclohexane and vacuum dried (25°) one hour.
Dichromium(II) tetrakistrifluoroacetae bis dimethylformamide.
I_, 37 nig (.053 mmol) was dissolved in 1 ml benzene. Addition of
9.0 y 1 (0.12 mmol) DMF produced a clear solution and a fine blue
powder within ca. 5 rain. Solid was filtered, washed once with
cyclohexane and vacuum dried (25°) one hour.
Dichromium(II) tetrakistrifluoroacetate bis trimethylphosphite
I_, 33 mg (0.047 mmol), was dissolved in 1 ml benzene. (MeO^P, 12
y1 (0.10 mmol), was added to give a slightly bluer solution.
Stripping off the solvents by vacuum produced a blue violet
powder which was not treated further.
Dichromium(II) tetrakistrifluoroacetate bis tetramethylurea.
I, 222 mg (0.32 mmol), was dissolved in 20 ml methylene chloride
and 76 yl (0.64 mmol) tetramethylurea added to give a

72
blue solution. Solvents were stripped off, the solid extracted
with benzene and the volume reduced to 5 ml. The solution was
suspended in a dewar above ice. The rectangular dark blue
crystals which formed after three days were collected and vacuum
dried (25°) one hour. This compound slowly decomposes in the
solid state, turning green after about one month under
dinitrogen.
Dichromium(II) tetrakistrifluoroaectate bis Triethylphosphate.
_1, 490 mg (0.70 mmol), was dissolved in 20 ml benzene and 0.35
ml (2.1 mmol) triethyl phosphate added to give a deep blue
solution. Volume was reduced to 5 ml and cooled to 5° for 24
hours to produce a mass of dark blue cubic crystals. Solid was
collected, washed with cyclohexane and vacuum dried (25°) 30 min.
Calculated for c20tt20('r2p^ 2°16P2: c» 26.10; H, 3.29; F, 24.77.
Found: C, 26.11; H, 3.40; F, 24.44.
Dichromium(II) tetrakistrifluoroacetate bis hexamethylphosphor-
amide. _I, 525 mg (0.74 mmol), was dissolved in 20 ml benzene and
0.30 ml (1.7 mmol) hexamethylphosphoramide added to give a deep
blue solution. Solution was cooled to 5° and bright blue cubic
crystals began to form after 15 rain. After 24 hours, crystals
were filtered and dried under a stream of dinitrogen. Calculated
for C20H36Cr2F12N6O10P2: C, 26.27; H, 3.97; F, 24.93; N, 9.19.
Found: C, 26.34; H, 3.99; F, 24.38; N, 9.21

73
Attempts to isolate solid adducts of dimethylacetamide and
diraethy 1su1foxide by similar procedures in benzene were
unsuccessful, producing a pale green gel in both instances.
Cr30(02CCF,)6(C5H5N),. I, 140 mg (0.20 mmol), in 5 ml EtgO was
placed in a test tube. 4 ml of hexane was layered onto the ether
and 0.8 ml (9.9 mmol) pyridine in 8 ml hexane added to the hexane
layer. After four days, solutions were almost clear. Large (1 X
1 X 20ram) pale olive needles had grown down from the solvent
interface. A much smaller yield of small dark olive crystals
were sparsely formed on tube walls near the bottom. Crystals
were collected and dried (vac., 25°) one hour. Only the pale
green complex changes color (dark olive) upon exposure to air.
Bases, solvents. Methylene chloride was dried over CaC^
and vacuum distil led from P2O5 at 25°. Benzene was dried over
CaCl2 and vacuum distilled from CaClj at 25°. Pyridine was dried
over Na and fractionally vacuum distilled from BaO. Diethyl
ether was vacuum distilled from Na/benzophenone at 25°. DMSO was
dried over NaOH and fractionally vacuum distilled from NaOH. DMF
was dried over KOH and fractionally vacuum distilled from BaO.
DMA was dried over molecular sieves and fractionally vacuum
distilled from BaO. DMTF was dried over BaO and fractionally
vacuum distilled. DMCA was fractionally vacuum distilled. HMPA
was fractionally vacuum distilled from BaO. Triraethyl phosphite
was fractionally vacuum distilled. Triraethyl phosphine oxide was
used as received. Tetramethyl urea was dried over BaO and
fractionally vacuum distilled. Triethyl phosphate was
fractionally vacuum distilled from BaO. DMTF was dried over BaO,

74
fractionally vacuum distilled and stored in the dark. All
distillations were performed under dinitrogen or vacuum and only
the heart cut saved.
3. Data Collection
All manipulations were performed under dinitrogen in an inert
atmosphere box or employing syringe techniques. All glassware,
and cells were stored in dessicators over CaSOjj or in the dry
box.
Electronic spectroscopy. UV-VIS titrations were performed by
repeated microliter injections of a concentrated stock base
solution into the standardized 5 ml chromium solution. The
quartz inert atmosphere cell has been described.^ Methylene
chloride solutions were prepared in volumetric flasks inside the
inert atmosphere box. Spectra were recorded on a Perkin Elmer
330 spectrophotometer.
Calorimetry. Design and operation of the calorimeter has
been previously described.To correct stability problems, a
new unit was constructed in the Electronic Shop of the University
of Florida chemistry department. Schematics and operating
procedure are given in Appendix II. Typically, a methylene
p
chloride solution (ca. 10 M) of the dimer was prepared in a 50
ml volumetric flask and decanted into the inert atmosphere 55 ml
adiabatic cell. Five milliliters of solvent were syringed into
the volumetric flask and washed into the cell. A 5 ml aliquot of
solution was removed from the cell and transferred to the UV-VIS

75
cell. Concentrations were calculated from the absorbance at 550
nm (^550 = 132.5). The base solution (ca. 0.5 M) was prepared in
a volumetric cell and used to charge a 1 ml Hamilton gas-tight
syringe. After expelling any dead volume and air bubbles, the
syringe was emptied to the zero calibration stop and secured into
the cell. While stirring, the cell and contents were allowed to
equilibrate to room temperature (25°) for several hours. A
series of thirteen calibrated syringe stops allowed incremental
injection of known volumes of base solution. Recorder
deflections which accompanied the exothermic base additions were
ratioed to a calibration deflection of known heat. Identical
titrations were performed in the absence of metal complex to
measure the base solution dilution heats (if any). Measurable
dilution heats were not observed with any of the donor solutions.
FTIR. Methylene chloride solutions were scanned in airtight
0.2 mra solution cells, after preparation in the glove box.
Solvent absorbances were subtracted. Spectra were recorded on a
Nicolet 5DX-B instrument.
FT-NMR. CD2CI2 and CgDg solutions were prepared in the inert
atmosphere box and 100 MHz spectra recorded on a JEOL XL-100 FT-
NMR. Determinations of magnetic susceptibilities with the EVANS
method were performed by modified procedures'^ with a newly
designed coaxial tube arrangement, shown in Fig 2-10. Chromium
solutions of 1-5 X 10“^M were prepared by addition of 2.2-3
donor of interest to CgDg solutions of
equivalents of the

teflon spacer
paramagnetic solution
in I mm i d. capillary
Fig 2-10. Modified NMR tube for determination of solution
magnetic susceptibilities of air-sensitive complexes.

77
Cr2(tfa)lt(Et20)2» Based on the large calculated equilibrium
constants, the added donor effectively displaces coordinated
ether. The same solvent system, CgD^ containing 2% v/v TMS and
2% v/v CgHg was used for both tubes. Paramagnetic solutions were
syringed into clean dry melting point capillary tubes (1 mm i.d.)
and sealed by fitting with a machined teflon spacer. A tight
fitting cap provides a second oxygen barrier and the NMR Spectra
were collected within 15 minutes of sample preparation. The dual
standards provide a check on locating the paramagnetic shifted
resonance and signal integration provides another. This tube
arrangement allows facile manipulation of the air sensitive
solutions and avoids the expense of coaxial tubes.
Elemental analyses. Analyses were performed by Galbraith
Laboratories, Knoxville, TN. The expense of routine analysis of
the air sensitive complexes prohibited characterization of each
adduct species. The HMPA and Et^PO^ ( a strong and intermediate)
donor complexes were characterized as representative complexes.

CHAPTER III
TRANS INFLUENCE ACROSS A METAL-METAL BOND
A. Introduction
Trans influence—the ground state influence on a metal-ligand
bond strength by another trans coordinated ligand—plays a
fundamental role in coordination chemistry and transition metal
complex reactivity. The ability of coordinated ligands to
labilize or stabilize a trans coordination site serves to define
the metal coordination sphere. A general feature of trans ligand
influence in mononuclear chemistry occurs when both metal-ligand
bonds are in competition for the same metal d orbital.^®
The idea of a trans influence in dinuclear complexes has been
vigorously pursued, and, in a general sense, the metal-ligand
bond strength has an inverse relationship with the metal-metal
bond strength. As the studies in the last chapter illustrate,
the effect is pronounced in the chromium carboxylates as both
ligand and adjacent metal atom compete for the same chromium
orbital; stronger chromium-ligand bonds result in weaker metal-
metal bonds. Conversely, the strong quadruple Mo-Mo bonds in
the molybdenum carboxylates dictate weak metal-ligand
interactions. The effect is schematically illustrated in Fig 3-1.
While structural studies of dimers with stronger metal
78

79
L r=^> M « L
L^>M< >M<^L
L L
Fig 3-1. Schematic representation of the trans influence in
metal carboxylate dimers.

80
interactions illustrate slight lengthening of the raetal-raetal
bond with ligand donor strength, the overall effect is a
weakening of the second metal-ligand interaction as demonstrated
by calorimetric measurements.^ Redistribution of electron
density donated at one metal over both metal atoms lowers the
Lewis acidity of the second, and the effect has been
parameterized in terms of communication between the two metal
centers. As a further test of this model, the studies reported
here examine the effect of varying one ligand on the spectral
properties of a second, fixed, trans ligand in a rhodium dimer.
Examples of mixed ligand and corresponding fixed ligand complexes
for comparison are scarce. The results reported here provide the
first systematic investigation of the trans influence across a
metal-metal bond. First the EPR parameters of a coordinated spin
label (TEMPO) in complexes of the type (DRI^Cpfb^CTEMPO) are
examined, followed by the stretching frequency of coordinated CO
in complexes of the type (L) Rl^Cp f b) ^ (CO). The original
experimental results have been communicated^^ by James Stahlbush;
a reinvestigation of the data is offered here.
Relationships between spectral and bonding properties of 1:1
adducts have been reported,^-’10"' such as the linear correlation
between Avqh anc* AH for a 3et*ies of 1:1 phenol-base adducts.
Breakdown of this correlation upon extension to a larger donor
set is attributed to fundamental differences in AvQH and AH for
gauging E and C effects. A more appropriate treatment of
spectral changes which accompany adduct formation involves an
adaptation of Eqn 2-1. Replacing AH by Ax, the spectral shift,
gives^ Eqn 3-1 for the case where a base is held constant and a

81
series of acids studied.
AX + W = EaEb* + CACB* (3-1)
The asterisks imply that conversion units for converting EA from
(kcal raol-^)^^ are included in Eg* along with the response to
the quantity being measured induced in the base by the acid.
The analysis of spectral data via Eqn 3-1 serves several
purposes. Besides establishing a correlation for a particular
acid or base which allows prediction of bond strengths from
spectral shifts, this treatment extends the E and C basis set
beyond only enthalpy data. And as with the enthalpy data,
correlation of the easier to obtain spectral shifts lends insight
into the nature of the metal-metal and metal-ligand bonds.
B. Results and Discussion
TEMPO
a. EPR spectra
Species of the type (B) Rl^Cpfb^CTEMPO), where B is a
coordinated Lewis base and TEMPO is 2,2,6,6-tetraraethy 1 -
piperidine-N-oxy1, were investigated using EPR in order to
measure the influence of B on the EPR spectrum of the coordinated
nitroxyl radical. TEMPO is a donor of moderate strength that
does not bind to RhjCOAc)^ or Rh2(but)^ but forms adducts with
the corresponding fluorinated derivatives, Rl^Ctfa)^ and
Rl^Cpfb)^ Typically, solutions were prepared in a 9:10:1 molar
ratio of base: rhodium complex: TEMPO in oxygen-free methylene
chloride. Most of the nitroxide existed in the unbound state,
while most of that coordinated existed as the WRl^Cpfb^CTEMPO)

82
species. A representative spectrum for diraethylacetamide with
Rl^Cpfb)^ and TEMPO is shown in Fig 3-2. In all cases, the
signal for the (BJRt^Cpfb^CTEMPO) species appears between those
for the free nitroxide and Rl^Cpfb^dEMPO). The nitrogen
hyperfine for these 2:1 adducts was usually either equal to or
slightly less than that observed for the Rl^Cpfb^CTEMPO) adduct.
Both of the above effects would be expected from an inductive
weakening of the rhodium nitroxide bond by the coordinated base.
That is, a coordinated donor weakens the Rh-TEMPO bond, and the
EPR parameters of the spin label move towards those of free
TEMPO. A wide range of g-values was observed for the mixed donor
2:1 adducts while only minor or no changes were observed for the
nitrogen hyperfine, A^. A tabulation of the EPR parameters is
given in Table 3-1.
Bonding in Rl^Ctfa^CTEMPO) has been described in terms of o-
donation from a nitroxide oxygen lone pair into the Rh-Rh a
orbital with concomitant orbital mixing of the Rh-Rh tt* orbital
with the nitroxide it* containing the unpaired spin. When B is a
sigma donor, the metal nitroxide bond is weakened, and the g
value moves toward that of free nitroxide. When B is a pi
acceptor the metal nitroxide bond can be weakened in two ways.
»
Directly, competition for metal-metal tt electron density
decreases the iT-backbonding to TEMPO. Indirectly, a pi acceptor
exhibits enhanced sigma donation, causing weakening of the metal
nitroxide bond in the same way as a pure sigma donor. The net
result is that both a-donors and ir-acceptors serve to weaken the
metal nitroxide bond and cause a lowering of the g-value back
towards the free solution value of TEMPO.

83
a
Fig 3-2. Representative spectrum observed for a CH2CI2 solution
of RhgCpfb)^, TEMPO and dimethyl acetamide. Species are a) free
TEMPO, b) (DMA) Rh2(pfb)4(TEMPO), c) Rh2(pfb)4(TEMPO), d)
precipitated Rh2(pfb)4(TEMP0).

84
Table 3-1. EPR Parameters of (¡DRl^Cpfb^CTEMPO) adducts.3
-AH 1:1
B
no base
methyl acetate
ethyl acetate
acetone
p-dioxane
dimethyl acetamide
bridged ethere
tetrahydroduran
dimethylsulfoxide
hexamethylphos-
phoramide
dimethyl formamide
acetonitrile
pyridine-N-oxide
f
cage phosphite
diethylsulfide
4-picoline
pyridine
1-methylimidazole
piperidine
triethylamine
An( 10-;>cm“ ) g'
1.57
2.0152
1.55
2.0128
1.55
2.0127
1.55
2.0122
1.55
2.0120
1.55
2.0119
1.55
2.0119
1.55
2.0118
1.55
2.0116
1.55
2.0115
1.55
2.0114
1.56
2.0108
1.55
2.0101
1.56
2.0095
1.55
2.0093
1.55
2.0083
1.56
2.0081
1.56
2.0079
h
2.0074
h
2.0069
Scale (kcal mo1
2.0152
0
2.0130
7.37
2.0129
7.96
2.0125
9.05
2.0122
9.66
2.0119
11.17
2.0114
11.64
2.0111
12.38
2.0117
11.74
2.0109
13.87
2.0120
10.54
(2.0132)®
6.81
2.0106
14.65
(2.0104)®
13.93
(2.0100)®
14.59
(2.0092)®
17.75
(2.0096)®
17.06
2.0081
20.32
2.0078
21.28
2.0068
24.31

85
Table 3-1 (continued)
a) Compare to solution parameters of free TEMPO: g = 2.0047, AN
= 1.47 X 10"3cm“1.
b) g of (B)Rh2(pfb)4(TEMPO).
c) g calc from equation 3-4 using the C^:^ and E^1:1 from
equations 3-2 and 3-3 and values reported in Table 3-2.
d) Calculated enthalpy for adding B to Rhp(pfb)4 using values in
Table 3-2.
e) 7-0xabicyclo [2.2.1] heptane.
f) 1-Phospha-4-ethyl-2,5,7-trioxabicylo [2.2.1] octane.
> H
g) Systems in which metal tt to ligand tt back-bonding
occurs.
h) Unresolved.

86
b. Quantitative correlations
The general trends discussed in the previous section suggest
that g-values from the EPR spectra may be used to provide
quantitative data about the strength of binding. This encouraged
us to investigate the quantitative relationship between the
enthalpys of adduct formation and changes in the g-values. A
model has been proposed and tested for predicting the enthalpy of
coordination of a second donor, B, to an f^CC^ClOjjCB) adduct to
form a 2:1 adduct.^ ^”33 jn this model, the parameter for the
1:1 adduct behaving as an acid to form a 2:1 adduct, E^1:\ is
given by
ea':1 * h - k%
and C.1:^ is given by
V = CA -X,CB
(3-2)
(3-3)
Where k and k* reflect (for o-donors) the effectiveness of the
metal-metal bond at transmitting the inductive influence of base
1*1 1*1
coordination to the second metal center. E^ and values
for(B)Rl^ipfb)^ can be calculated from equations 3-2 and 3-3.
The reported32*38,39 £ ancj q values used in this analysis are
given in Table 3-2. The and cA1i1 of the vari°us
(B)Rh2(pfb)¡j adducts are calculated and also listed in Table 3-2.
Note how the acid parameters decrease as the inductive effect
increases. That is, the greater the donor strength, the lower is

87
Table 3-2. E and C Parameters for
Species
Used in
This Study
B
eb
CB
c 1:1
ea
r 1:1
methyl acetate
0.903
1.61
4.01
1.68
ethyl acetate
0.975
1.74
3.93
1.68
acetone
0.987
2.33
3.92
1.66
p-dioxane
1.09
2.38
3.80
1.65
dimethyl acetamide
1.32
2.58
3.53
1.65
bridged ether
0.887
4.11
4.03
1.59
tetrahydrofuran
0.978
4.27
3.93
1.58
dimethylsulfoxide
1.34
2.85
3.51
1.64
hexaraethylphosphoramide
1.52
3.55
3.30
1.61
dimethyl formamide
1.23
2.48
3.63
1.65
acetonitrile
0.886
1.34
4.03
1.69
pyridine N-oxide
1.34
4.52
3.51
1.58
cage phosphite
0.548
6.41
4.42
1.51
diethylsulfide
0.339
7.40
4.67
1.47
4-picoline
1.17
6.80
3.70
1.49
pyridine
1.17
6.40
3.70
1.51
1-raethylimidazole
0.934
8.96
3.98
1.41
piperidine
1.01
9.29
3.89
1.40
triethylamine
0.991
11.09
3.91
1.34
TEMPO
0.915
6.21
4.00
1.51
A
ea
CA
k
k'
Rh2(pfb)4
5.06
1.74
1.16
0.0364

88
the Lewis acidity of the second metal center. With these E and C
values for the various 1:1 adducts, we are now in a position to
attempt a correlation of the g-values obtained when TEMPO is
coordinated to the second coordination site to form a series of
2:1 adducts of general formula (BJRhgtpfb^CTEMPO). Substitution
of Eqns 3-2 and 3-3 into 3-1 gives
g + W = Ea1:1Eb* + CA1:1CB* (3-4)
where g has been substituted for Ax . The simultaneous equations
are solved for Eg* and Cg* which are the spectroscopic parameters
for TEMPO needed to predict g. The quantity W includes the g
value for free TEMPO (2.0047) as well as any nonzero enthalpy
components of the spectroscopic relation.^ 02-1 03 The best fit
results yield
Eg* = 1.16 x 10~3 (0.29 x 10-3)
Cg* = 1.78 x 10"2 (0.10 x 10-2)
W = -1.9784 (.0018)
with standard deviations in parentheses. These parameters for
TEMPO allow calculation of the g-value for any
(B)Rh2(pfb)¡j(TEMP0) complex when the base is a o-donor whose Eg
and Cg parameters are known. Data for adducts with donors known
to act as T:-acceptors (acetonitrile, cage phosphite,
diethysulfide, 4-picoline and pyridine) are not included in the

89
calculation of the TEMPO parameters. Attempted fits which
include the Tr_acceptors give larger standard deviations, as
expected. Table 3-1 contains the g-values calculated from this
fit (scaic)« The columns of g and gcaic show excellent
agreement, generally within the accuracy of the measured numbers,
except in those cases where metal to base u-backbonding occurs
(data in parentheses). For the latter, these are the g-values
expected if the Lewis bases, B, utilize only their o-bonding
capabilities in forming the (B)Rhp(pfb)jj adducts. The close
agreement between g and gca^c demonstrates that the inductive
model (equations 3-2 and 3-3) adequately describes the
transmission of coordination effects through the metal-metal
bond, for it is this model which describes the varying acidity of
the second metal center. For the five donors which also behave
as it-acids, the lower observed g-values than calculated by Eqn 3-
4 manifest the metal to B iT-backbonding contribution in these
adducts. Thus this analysis would suggest that bothoand tt
interactions in the B-Rh bond serve to weaken the Rh-TEMPO bond.
To see if a relationship exists between g and the strength of
B binding to Rt^tpfb)^, enthalpies for binding bases were
calculated from the E and C equation (2-1); these 1:1 adduct
heats, AHg1:1, are given in Table 3-1. The E and C parameters
are derived from a-only interactions and hence the calculated Ah
in Table 3-1 reflect only the ocomponent of the adduct bond.
The experimental g-values, however, reflect the sum of o-donor
and any TT-acceptor interactions. In Fig 3-3, the enthalpies of
1:1 adduct formation for (B)Rh2(pfb)¿j are plotted as a function
of the experimental g values for the 2:1 adducts (B^l^pfb^tTEMPO).

90
Both calculated (o,A) and experimental (X) enthalpies are
included. (It would perhaps be more direct to compare the g-
values of the 2:1 adducts with the enthalpy of 2:1 adduct
formation, that is, the enthalpy for B + RhjCpfb^CTEMPO) or even
TEMPO + Rt^Cpfb^CB). The qualitative conclusions are, however,
the same, and experimental heats are available for B +
Rl^pfb)^.) In Fig 3-3, the calculated (cr-only) enthalpies are
all lower than the g-values would suggest for the five donors
which act as iT-acceptors: acetonitrile, cage phosphite,
diethylsulfide, 4-picoline and pyridine. Additional
stabilization in the (B^hgipfb)^ adduct bond is consistent with
Rh to B 7T-backbonding. In the two cases where experimental
AHg^:1 are available (acetonitrile and pyridine), the measured
heats lie much closer to the correlation line. Thus, the g-
values manifest both a and tt effects across the metal-metal bond,
both serve to lower the g-value of coordinated TEMPO, and o-
donation appears to exert a stronger influence than tt-acceptance.
(Inclusion of tt-effects brings Ahb^:1 closer to the correlation
line but not all the way.)
2. Carbon Monoxide
a. FTIR spectra
The CO ligand is ubiquitous in organoraetal lie chemistry, and
considerable effort has been put forth to understand the nature
of M-C-0 bonding and the influence of various ligands upon the
reactivity and spectroscopic properties of carbonyls. In probing

AH o ( kcal/mol)
91
Fig 3-3. Correlation of predicted enthalpy of adduct formation,
AHg , for Rh2(pfb)4 and donor, B, to the observed g value of
(B)Rh2(pfb)i|(TEMP0).> o,o-donors; A,iT-acceptors. For those
systems where AH3^:”' has been experimentally determined, the
corresponding experimental heats are given by X. Character size
indicates range of experimental error. The best fit line is for
a -donors (see equation S-^).

92
the influence of cis and trans coordinated ligands, NMR10i<“
and ir109-112 studies of the coordinated carbonyl have
investigated the relative donor and acceptor properties of
coordinated ligands in transition metal mononuclear carbonyl
complexes. Here, ligand effects on coordinated CO in a dinuclear
complex are investigated. The experimental vCQ values obtained
when CO is bubbled through methylene chloride solutions of
Rl^Cpfb)^ containing excess donor are given in Table 3-3- A
gradual but consistent variation in the coordinated carbonyl
stretching frequency occurs with varying donor strength. The
carbonyl stretching frequency is dependent upon the extent of Rh
to CO ff-backbonding. The lower the value of vco, the stronger is
the metal-carbon bond. The data indicate that stronger donors
induce greater ^-backbonding to CO. The trend, however, is the
reverse of that seen for coordinated TEMPO. Where as the Rh-
TEMPO bond strength decreases with donor strength, the Rh-CO bond
strength increases with donor strength. Thi3 apparent anomaly is
readily understood in terras of the ligand (TEMPO or CO) orbital
•
interactions. Sigma donors partially populate the metal-metal a
orbital, raising the electronegativity of the dinuclear metal
â– 3 p
center and lowering the Lewis acidity of both metal atoms.-1
TEMPO, which acts primarily as a a-donor, is less able to donate
electron density when another a-donor lies trans. CO interacts
significantly as a n-acid and is best able to do so when another
a-donor lies trans.
The effect of sigma donation of electron density from the
Lewis base, B, to the metal raises the energy of the d orbital
manifold. When B is a 7T-acceptor ligand, two interactions occur.

93
Table 3-3. CO stretching frequencies of (B)Rh2(pfb)^(CO)
adducts. vc0 vco
B
F a
lb
c a
F k
ea
O*
exptc
calcd
—
0
0
5.06
1.74
2135.8
2136.0
ethyl acetate
.975
1.74
3.93
1.68
2130.8
2129.1
acetonitrile
.886
1.34
4.03
1.69
2129.1
(2129.8)
acetone
.987
2.33
3.92
1.66
2128.6
2128.7
bridged ethere
.887
4.11
4.03
1.59
2127.0
2127.7
dimethylsulfoxide
1.34
2.85
3.51
1.64
2126.2
2126.2
triethylphosphate
1.36
1.81
3.48
1.67
2126.2
2126.7
diethylsulfide
.339
7.40
4.67
1.47
2126.1
(2128.4)
dimethylacetamide
1.32
2.58
3.53
1.65
2125.4
2126.5
f
cage phosphite1
.548
6.41
4.42
1.51
2125.1
(2128.0)
1-methylimidazole
.934
8.96
3.98
1.41
2124.7
2123.7
piperidine
1.01
9.29
3.89
1 .40
2124.2
2123.0
pyridine
1.12
6.89
3.76
1.49
2122.8
(2124.3)
triethylamine
.991
11.09
3.91
1.34
2121.6
2121.9
4-picoline
1.17
6.80
3.70
1.49
2121.5
(2124.0)
quinuclidine
.704
13.2
4.24
1.26
2120.8
2120.8
a) References 38, 39.
b) Calculated from equations 3-2 and 3-3 for (B)Rh2(pfb)^.
c) Experimental value for (BjRhgipfb^CCO).
d) Calculated from Eqn 3-5 using E^ (1:1) and C^(1:1) values.
Values in parenthesis are for TT-acceptors whose experimental
v £Q were not used in the correlation.
e) 7-oxabicyclo [2.2.1] heptane.
f) 1-phospha-4-ethyl-1,5,7-trioxabicyclo [2.2.1] octane.

94
1) The d-manifold is stabilized by 7T-backbonding. 2) Enhanced
a-donation destabilizes the d-raanifold. Which ever predominates
will dictate the effect of ligand, B, tt bonding on the trans
carbonyl stretching frequency. An E and C analysis reveals the
relative importance of these two interactions.
b. Quantitative correlations
The donor Eg and C0 values are given in Table 3-3 along with
the EA^:^ and CA1:1 values for the (B) Rl^ (p f b) ^ adducts
calculated from equations 3-2 and 3-3 using the reported
parameters for Rl^Cpfb)^ from Table 3-2. Again, the EA1:^ and
CA quantify the lowered acidity for the second metal atom as a
result of base coordination at the first rhodium atom. Because
the E and C parameters are rigorously derived from reactions in
which only o interactions are important, the E^:1 and CA1:^ do
not reflect the effects of tt -backbond ing in the (B^t^pftOjj
adduct bond (the metal to ligand ir-backbonding also includes a
proportionately enhanced a bonding).
The form of the E and C equation used to describe a
nonenthalpic change for a given base interacting with a series of
acids is given by Eqn 3-1. When the experimental carbonyl
frequencys are used for Ax and the acid is the 1:1 adduct,
(B)Rh2(pfb)jj, Eqn 3-5 results.
v00 . W = ES1:V
(3-5)

95
The least squared minimized solution for those systems in which B
is not a n-acceptor ligand gives
Eg* = 4.99 (.74)
CB* = 21.1 (2.0)
W = -2074.1 (4.2)
where standard deviations are given in parentheses. These
parameters for CO allow calculation of CO stretching frequencies
for any (B)Rh2(pfb)ii(C0) complex when the base is a sigma donor
whose Eg and Cg parameters are known. The calculated vcQ for the
systems studied here are given in Table 3-3. Data for adducts
with donors known to act as TT-acceptors (acetonitrile, cage
phosphite, diethyl sulfide, pyridine and 4-picoline) are not
included in the calculation of the CO parameters. As expected,
larger standard deviations result when the five ir-acceptors are
included.
Examination of the magnitude of vQ0 when o and effects are
both operative lends insight into the mechansira of inductive
transfer of base binding effects. The calculated vco for the
five '"-acceptors are also given (in parentheses) in Table 3-3;
these are the frequencies expected if the Lewis bases, B, utilize
on ly their o-bonding capabilities in forming the (B^hgtpfb^
adduct. The lower observed vCQ values than calculated from Eqn
3-5 manifest the metal to B ‘"'-backbonding contribution for

96
adducts with these five bases. Thus both o-donation and 7T -
acceptance in the B-Rh bond serve to strengthen the Rh-CO bond.
n-Backbonding stabilizes the d-manifold, and thus the
concoramitant a-donation which accompanies the it-backbond appears
to exert a greater influence than the tt backbond itself upon the
Lewis acidity of the Rh^* unit. The changes in the vco
frequencies are small, however, and can result from subtle
differences in the many perturbations made on the system by base
coordination. The key point is that sigma donors decrease vcq in
a predictable fashion, and the it acceptor bases cause a decrease
larger than expected from a o-only effect.
C. Conclusion
Trans influence across a metal-metal bond has been examined
in two ways. Various Lewis base, B, forms of (BiRl^Cpfb^CTEMPO)
were examined with regards to the EPR parameters of the
coordinated nitroxyl. The corresponding analogs with coordinated
carbon monoxide, (BJRhoipfb^CCO) were monitored by FTIR with
regards to carbonyl stretching frequency. While both methods
provide facile monitoring of the trans influence, the two series
of complexes reveal opposite trends. Base coordination weakens
the Rh-TEMPO bond but strengthens the Rh-CO bond. The apparent
anomaly is readily attributed to the different bonding
interactions the two spectral probes undertake with the metal
center. In all cases, Lewis base (B) coordination lowers the
Lewis acidity of the trans rhodium center. Carbon monoxide
interacts significantly as a tt acid; the Rh->CO bond is enhanced

97
by base coordination. TEMPO interacts primarily as a a donor;
Rh<-TEMP0 a bond is weakened by base coordination.
An E and C analysis of the spectral data demonstrated its
utility in correlating and predicting both EPR g-values and
carbonyl frequencies. This treatment reveals several features of
the electronic structure of the (BÍRl^ípfb^CL) adducts (L=CO,
TEMPO). Separate analysis of those systems which involve o
donors and those with potential it acceptors provides further
support for the backbonding capabilities of the Rh2^ + unit.
Second, in both the CO and TEMPO adducts, tt stabilization in the
B-Rh bond has the same effect and serves to enhance the influence
of a donation. That is, the concommitant strengthening of the o _
bond which occurs in the backbond appears to exert a stronger
influence on the electronic structure of the Rl^4"1" center than
does the tt bond itself. Lastly, the satisfying correlations
which result provide further support for the inductive transfer
model (Eqns. 3-2 and 3-3) for it is these equations which
quantify the inductive influence of coordinated donors.

CHAPTER IV
THE ELECTRONIC SPECTRA OF M(III)2M(II)(02CR)gL3 SPECIES
A. Introduction
The success in monitoring and gauging the transmission of
bonding effects across a range of metal-metal bonds in various
carboxylate dimers encourages extension of these studies to more
complex metal clusters. For their ease of synthesis, stability
and charge neutrality, the likely candidates are the mixed-
valence carboxylate trimers having the basic iron acetate
structure (Fig 1-4). While an extensive range of homo- and
hetero- metal cationic complexes have been prepared- usually with
equivalent oxidation states—solvation and ion pairing
preclude their utility for studying coordination effects. A much
smaller family of neutral M(II)M(III)20 centered carboxylate
complexes have been shown to possess the familiar equilateral
triangle of metal atoms for complexes of V, Cr, Mn, Fe, Ru and
perhaps Co.-^“^ These have received attention as structural
curiosities especially in terms of ascribing the metal oxidation
states. Two of these neutral trimers, Ru^OCOAcigCPPh^)^ and
Mn^OCOAcJgCpyridine)^, have been examined by single-crystal x-ray
crystallography, 54,113 ancj both were found to contain equivalent
metal ions suggesting formal oxidation states of 2.67 for each of
the metal atoms and delocalization of the odd electron about the
98

99
M^O framework on the very long x-ray time scale. Metal-metal bond
distances preclude direct bonding interactions in both the
o o
ruthenium (3.33 A) and manganese (3.36 A) mixed valence trimers.
However, magnetic susceptibility measurements on both cationic
and neutral trimers gave moments greatly reduced from the sura of
the spin-only values of the metal ion centers, indicating
significant spin pairing via a mechanism other than metal-metal
bonding.^ The neutral ruthenium triraer Ru20(0Ac)g(H20)2 is, in
114
fact, diamagnetic in the solid state.
The neutral mixed valence trimers have also been the focus of
a range of physical studies aimed at determining rates of
intramolecular electron transfer. To this end, the
triiron acetates, Fe^OÍOAc^L^ (L = H2O or pyridine) have been
the most vigorously pursued. Although absolute electron transfer
rates measured for these complexes in the solid state show a
11C
strong dependence upon lattice effects, J physical studies
generally place these iron trimers on the high end of the Robin
1 1
and Day class II (appreciable electron delocalization). ESCA
studies on both cationic and neutral mixed-valence ruthenium
trimers gave single binding energies, demonstrating complete
delocalization on the ESCA time scale.
Despite the absence of discrete metal-metal bonding
interactions in the neutral mixed valence trimers, the obvious
electronic communication between the metal centers suggests these
complexes may be good candidates for the monitoring of
transmitted coordination effects. Demonstrated Lewis acidity of
the metal atoms accounts for the ability to prepare a range of
symmetrically substituted clusters, M^CXC^CR^L^. Towards

100
developing a trimeric system for monitoring of the transmission
of bonding effects, neutral trimers of Ru, Fe and Co were
prepared and their electronic spectra examined with regards to
ligand substitution. Ideally, spectral changes accompanying
stepwise ligand binding or substitution would allow determination
of sequential equilibrium constants for adduct formation. The
results of this survey are the subject of this chapter.
B. Results and Discussion
1. Iron
Air sensitive Fe^OCOAc^^O)^ was prepared by the method of
c Q
Dziobkowski et al., from hydrated calcium acetate and ferrous
chloride. Though FT accumulation allows slightly better
resolution, the IR spectrum (Fig 4-1) agrees with that
reported,^® complete^with the asymmetric stretch of the M^O
frame, va3y 560 cm,-1 and the M-OH2 stretch at 530 cm“^.
Satisfactory elemental analysis indicated this compound had been
isolated as the dihydrate, Fe^OCOAc)^
This complex is insoluble in nonpolar nonbasic solvents but
dissolves readily in acetic acid, DMSO and pyridine. Solubility
is slight in methylene chloride, ca. lo”1* M. Solutions for UV-
VIS spectra were prepared by suspending the trimer in CH2CI2 and
adding a 100 molar excess (based on M^) of various donors.
Complex solubility was not effected with acetonitrile,
dimethyl sulfide or triphenyl phosphine, and spectra of filtered
solutions with each of these donors displayed only the trimer
charge transfer (CT) band. Solubility was effected by addition

Fig 4-1. FTIR Spectrum of Fe^CKOAc^CHpO^.HpO in KBr.
peaks are due to incomplete removal of HpO and COp for
scan.
Negative
background

^TRANSMITTANCE
8386 13 455 25 072 36.689 48 306 59 923
4600 0 3800 0 3000 0 2200.0 1800 0 1400.0 1000 0 800 00 600 00 350 00
WAVENUMBERS (CM-l)
102

103
of diraethylsulfoxide triethylphosphate or pyridine, and the
resulting spectra are shown in Fig 4-2. Presuming that solvation
is accomplished by partial or complete removal of coordinated
water, these scans represent the base coordinated complexes.
Despite almost complete obscuring of the ligand field transitions
due to a pronounced CT band, the apparent shoulder absorbtions
still display essentially no changes with ligand substitution.
Dubicki and Martin achieved some success in describing the
cationic FeCIID^ trimers in terras of a monomeric d^
configuration in an octahedral field, but spectral support for
their conclusions is clouded by poor resolution due to CT
11ft
bands. No attempts have been made to assign the electronic
spectra of the neutra 1 FedIDpdl) trimers apart from
speculation on the near IR bands in the solid state.1 1(^'1 1 ^
Should any ligand field absorbances manifest donor coordination
effects, their monitoring is difficult given low complex
solubility and obscuring CT bands.
2, Cobalt
While the dark green product formed from oxidation of pink
cobaltous acetate has been known for many years, conclusive
structural identification has not been realized. The reaction
product may depend upon the oxidation procedure used, and several
structural proposals have been suggested. Peracetic acid
oxidation of aqueous solutions of cobalt (II) acetate gave
diacetato-y-hydroxocobalt (III) as the dimer. Alternatively,
electrochemical oxidation of a mixture of cobalt (II) and
potassium acetates in aqueous solution gave a binuclear cation
described 121 as (HOAc)i
Fig 4-2. Electronic spectra of FeoCKOAc^CF^OK (3.4 X 10_^M)
with 100 equivalents (0.34 M) a) pyridine, tO DHSO and c)
triethylphosphate in methylene chloride using 2 mm quartz cells.
The aquo complex is insoluble in methylene chloride in the
absence of donor solvents.


106
1 O p
oxidation gave similar cationic products while ozonation of
cobalt (II) acetate in glacial acetic acid gave the former
neutral dimer.
On the other hand, peracetic acid oxidation of an acetic acid
solution of cobaltous acetate, pyridine and NHjjPFg gave a mixture
of interconvertible trimers of the basic iron acetate structure
formulated as [Py^Co^OCOAc^OH] PFg and [ Py^Co^OCOAc)^ ] PFg (py =
pyridine). The former was verified by an x-ray crystal structure
of the PyCoBr^ salt. Only one group has reported isolation of
a mixed valence trimer. Ziólkowski et al. reported-^ the
isolation of Co^OCOAc^CHOAC)^ from an acetic acid solution of
cobaltous acetate oxidized by a mixture of ozone and oxygen (ca.
2% ozone v/v). More recently, the same workers reported1^
formation of Co^OCC^CR^Hi^O)^ from acid solutions of Co(II)
carboxylate, oxidized by the appropriate aldehyde and C>2.
Variations of this latter method were attempted here.
Anhydrous cobaltous propionate turns green immediately upon
contact with propionaldehyde, producing cobalt (III) and,
presumably, propanol. Addition of either methanol or ethanol to
the resulting green slurry, however, gave the original rose color
and no oxidation occurred if either alcohol were used as solvent.
Both water and propionic acid were suitable solvents for the
oxidation of cobaltous propionate though the former is immiscible
with propionaldehyde. Oxidation of cobaltous propionate in
propionic acid was accomplished by addition of 1-5 equivalents of
propionaldehyde. Neat propionaldehyde gave only an intractable
green solid. Oxygen was bubbled through the solution for two
hours. The dark green brown solution which resulted was

107
filtered, passed through a Dowex (ion exchange resin) column to
remove unreacted Co(II), stripped to give a glassy khaki solid,
and dried under vacuum. Heating of solutions of this solid or
the previous solutions to 100° invariably gave a rose colored
product. Cobaltic propionate prepared in this manner was soluble
in alcohol, acetic acid, DMSO and acetonitrile and decomposes in
water, slowly turning red. Attempts to recrystallize from
ethanol or acetic acid were unsuccessful, and further
characterization was performed on the vacuum-dried solid.
The FTIR spectrum of cobaltic propionate (Fig 4-3) resembles
that of cobaltous propionate and hence only serves to confirm the
presence of propionate. The best fit of the elemental analysis
is for partially dehydrated trimer, Co20(prop)g(H20), relying on
the C and H combustion analysis only. The electronic spectrum
(Fig 4-4) is similar to that reported for both the product
prepared aldehyde/02 oxidation-^ and that prepared by
ozonization.12^ Given the spectral similarities of the possible
structures, the real question at hand is the nuclearity of the
complex.
Vacuum-dried cobaltic propionate was extracted with ethanol,
precipitated with acetone and vacuum dried. The precipitated
complex was used in the following experiments. The molecular
weight of cobaltic propionate was determined by isothermal
1
distillation in methanol using the Signer method. The
measured value of 568 + 24 is intermediate between the trimer,
Co^Oiprop^^O) (MW = 649) and the dimer, Ck^prop^iOH^ipropH)
(MW = 518). In this experiment the molecular weight was
monitored over a period of 40 days and significant fluctuations

Fig ^-3. FTIR spectrum of cobaltic propionate in KBr.

^TRANSMITTANCE
1655 7.967° 15.810 23 633 31 455
WAVENUMBERS (CM—1)

Fig 4-4. The electronic spectrum of cobaltic propionate in
acetic acid (1.43 mM based on Co^CKprop^C^O)^). a) 2 mm
quartz cell b) 1 cm cell.

111
X(NM)

112
observed near equilibration (Fig 4-5). The measured value of 568
+ 24 is the mean of the last nine points at the 99% confidence
level. The observed fluctuation may arise from temperature
variations as the apparatus was simply left on the bench.
The mass spectrum (70 eV electron impact) of the precipitated
cobaltic propionate is shown in Fig 4-6 and is consistent with
the trimer formulation, given a molecular weight of 631 for the
Co^Oiprop)^ core. Almost complete fragmentation has ocurred with
a base peak at m/e = 43, but an expansion of the higher m/e
region shows M + 1, M + 15 and M + 29 peaks. No parent peak at
631 is observed. The significant 573 peak (M-58) corresponds to
loss of two CH^C^ units. Catterick and Thornton, however, offer
this caveat^: "Mass spectrometry is the most misleading
technique available to the inorganic chemist seeking to
characterize a noncrystalline carboxylate." These workers refer
to the preponderance of rearrangement reactions and the wealth of
parent and fragment peaks which occur. The experiment was
repeated at a lower ionizing energy (50 eV El) and served to
confirm the formation of ions in the trimer m/e region (Fig 4-7)
though rearrangements may still monopolize the upper mass end.
The mass spectra would support the trimer cluster formulation but
without a clear fragmentation pattern at lower ionization energy,
definitive characterization is difficult.
The magnetic susceptibility measured by the Evans method in
methanol at 25° was 1.51 + 0.1 B.M. per Co atom. This is
similar to the solution value of 1.4 + 0.1 B.M. per Co
reported1^ for Co2(0Ac)i<(0H)2(H0Ac), but different from the
solid state value of 3.58 B.M. at 27° reported for

MW
Fig 4-5. Molecular weight determination of cobaltic propionate
by isothermal distillation using the Signer method. The
reference compound is azobenzene in methanol.

Fig 4-6. 70 eV electron impact mass spectrum of cobaltic propionate.
Above m/e = 200 the ordinate is expanded 10X.

100
43
*10
90
60
70
60
50
40
30
20
10
0
60
303
573
300
632
100
200
400
500
600

Fig 4-7. 50 eV electron impact mass spectrum of cobaltic
propionate.


118
Co^CKOAc^C^O)^ (This report-^ does not indicate if this is per
complex or per metal). The similar magnetic susceptibility
determination does not, however, confirm a structure but only
indicates agreement with another workers measurement.
Without an x-ray structure determination, the true nature of
cobaltic acetate (propionate) continues to be an open debate
given compelling evidence favoring both the dimeric and triraeric
formulations. Here, various methods of characterization showed
similarities to both. All the same, the electronic response to
donors was examined by scanning the visible spectrum of the
vacuum-dried solid in pyridine, acetonitrile and
tetrahydrothiophene. Only for pyridine solution were any
spectral changes obvious and in this base the electronic spectrum
of cobaltic propionate consists of only a featureless CT band
which tails well into the visible region. Thus, regardless of
its structure, cobaltic propionate does not appear amenable to
monitoring of base complexation in the electronic spectrum.
3. Ruthenium
The neutral, mixed valence Ru20(prop)g(H20)2 was synthesized
by PtC>2 catalyzed hydrogen reduction of the cationic
[Ru20(prop)g(H20)2] (prop) in water and is mildly air sensitive.
Neutral Ru^CXprop^tPPh^ was also prepared from the cation by
the addition of triphenylphosphine and recrystallization of the
precipitated complex from benzene. This adduct is stable to air.
The electronic spectrum of RujCKprop^^O)^ in acetone (Fig
5-15) displays several shoulder absorbances along with a distinct
transition (unassigned) at 865 nm. In acetonitrile solution this
band is shifted to 890 nm (Fig 4-8) and for the PPhj adduct in

119
cyclohexane appears at 980 nm (Fig 4-9). Wilkinson has reported
the pyridine adduct, Ru^OCOAc^Cpy)^ for which this transition
appears at 890 nm in methanol.11^ Baumann et al. have detailed^
C 11
an earlier proposal J to describe the uorbitals in the Ru^O
unit and their proposal is shown in Fig 4-10. In this scheme the
dz2 orbitals are involved with a bonding to the central oxygen
atom and the dxtyt with a bonding to the carboxylate oxygen
atoms and neither are shown. The primes on xT and y’ indicate
that the molecular orbitals chosen do not lie along the
coordinate axes system used, the x' and y' axes point between the
carboxylate oxygens and not at them. Only the dv,_ orbitals
combine with the oxygen lone p orbital, giving rise to the A"2
levels. The dx,2_yi2 orbitals can interact by symmetry with p
orbitals localized on the oxygen atoms of the bridging
carboxylates, which would provide a mechansira for Ru-Ru
interaction. Given the relative distances involved, overlap is
suggested to be slight though enough to remove the degeneracy of
the three dxl2_y,2 orbitals to give the E’(2) and A*1 orbitals.
C Jl O
Despite the long Ru-Ru distance-^ (3.33 A) observed for
Ru^0(OAc)g(PPh^)3, Baumann et al. propose direct Ru-Ru
combinations by mixing of the dx,z orbitals to generate metal-
metal bonding, E'^), and antibonding, A'2, levels. In the
diamagnetic III2II complex, the highest occupied molecular
orbital is either the A'2 or E", both of which are completely
filled.
From the spectral deconvolution of the electronic spectrum of
[Ru^0(0Ac)g(py)2pyr]+ (py = pyridine, pyz = pyrazine), Baumann et
al. demonstrated contributions from three bands to the low

Fig 4-8. The electronic spectrum of Ru^OCprop^C^OK in
acetonitrile.


Fig 4-9. The electronic spectrum of RuoO(prop)6(PPhOo in
cyclohexane (1.62 X

ro
LO

124
D
3h
XZ
xa-r£ ee
YZ
/
/
/
/
/
/
/
\
\
\
\
\
A;
\
\h2
\
-E" \
F1 x
A l(2) \
- p1 \
b(|) /V
/
/
/
/ -»
A-
Ru
0
Fig 4-10. Qualitative molecular orbital scheme for the Ru^O
system in symmetry.^

125
energy transition. Similarly in the neutral cluster the low
energy transition is suggested to comprise a three-band envelope
assigned to transitions from essentially nonbonding (EM, E'(2)
and A'p to antibonding (a"2) orbitals. These intracluster
transitions may indicate a pathway for electron transfer (via the
central oxygen atom) and appear to respond to donor effects in
concert. Stronger donors give rise to a red shift for this
absorbtion envelope. The ruthenium analogue is the only mixed
valence trimer examined which exhibited this behavior but was not
examined further in favor of examining its catalytic potential
(Chapter V).
C. Conclusion
In an effort to develop a three centered metal cluster
system in which to monitor stepwise adduct formation, the
electronic spectral response of neutral mixed-valence clusters,
M^OO^ClOgL^, to various donors was examined. In the case of
iron, any spectral changes were obscured by a predominant charge
transfer band. The reported cobalt analogue could not be
reconciled from a similar reported dimeric complex. The
ruthenium analogue was unique in demonstrating a donor spectral
dependence in the transitions associated with intracluster charge
transfer.
D. Experimental
Air sensitive complexes were synthesized using standard
Schlenk techniques and solutions prepared in a VAC dry box.
Electronic spectra were recorded on a PE 330 spectrophotometer
and are background corrected. FTIR Spectra were collected (ca.

126
50 scans) on a Nicolet 5DXB Spectrometer. NMR spectra were
collected on a JEOL FX-100 FT instrument. Mass spectra and
elemental analyses were performed on a service basis at the
University of Florida.
[ Fe^OC OAcJgCí^O)^], yT-Oxotriaquohexakis(acetato)iron( II)di-
iron(III), was prepared by a modification of the method of
Brown.Ferrous chloride was prepared according to the method
of Winter.To a solution of 4.0 g (28 mmol) of FeClg.^O in
13 ml water was added a slurry of 7.4 g (42 mmol) CaiCH^C^^^O
in 14 ml water and 25 ml acetic acid. The mixture was heated to
70° and aerated with air for 6 hours, the gas feed was switched
to nitrogen,and the flask contents flushed for 10 min, stoppered
and placed in the refrigerator. The dark orange brown solid
which precipitated after 24 hr was filtered off in a glove bag,
the mother liquor reduced to 15 ml in a glove bag and returned to
the refrigerator. After three days, the rust colored
precipitate was similarly collected, washed with five volumes of
0.083 M CH^C^H and dried under vacuum over KOH at 25°. The
product was stored under nitrogen. Elemental analysis
characterized the product as the dihydrate yield: 0.9 g (15%)
Caled, for C-^^s Fe3 °18: C* 22*95; H, 4.49; N, 0.00. Found:
C, 22.84; H, 4.40; N, 0.00.
Ca(02CCH2CH2CH2)2» Calcium butyrate. A suspension of 3.1 g (54
mmol) Ca(0H)2 in 25 ml butyric acid was heated to dissolve,
stripped and vacuum dried (100°). The product was recrystallized
from methanol/acetone and vacuum dried (60°). Combustion

127
analysis indicates a partial hydrate, Ca(02CCH2CH2CH3)2, 0.5 H20.
Anal, caled.: C, 43.03; H, 6.77; N, 0.00. Found: C, 43.36; H,
6.85; N,0.00.
[Fe30(but)g(H20)3] , y3-Oxotriaquohexakis(butyrato)iron(II)
dilron(III) was prepared in butyric acid by the above procedure.
After cooling the reaction mixture 4 hr at 0°, the acid layer was
decanted, stripped, dried under vacuum (100°, 36 hr) and stored
under nitrogen. Harsh drying removed some of the coordinated
water; the best formulation for the combustion analysis is for
Fe30(02CCH2CH2CH3)6(H20). Anal, caled.: C, 39.81; H, 6.12; N,
0.00. Found: C, 39.66; H, 6.14; N, 0.02.
Co(prop)2, Cobaltous propionate. A slurry of 25 g cobalt metal
as a fine powder in 100 ml propionic acid was refluxed for 24 hr
after which time about half of the metal had oxidized. The thick
purple solution was filtered, stripped and vacuum dried (110°).
After recrystallization from ethanol/acetone, the rose powder was
vacuum dried (70°). "Anhydrous" cobaltous propionate contains
some water of hydration and fails to deliver a good combustion
analysis though the analysis agrees with that of a sample later
supplied as a
gift by Shepard
Chemical
O
O
•
Anal.
caled.
CóHi2C0O5: C,
32.3; H, 5.42;
N, 0.00.
Found: C
, 31.39;
4.97, N, 0.00.
Found (Shepard
sample):
c,
31.65;
H, 4.96,
0.00.
"Cobaltic propionate" was prepared by a modification of the
procedure of Ziólkowski.125 To 1.0 g of cobaltous propionate was

128
added 1.8 ml propanal followed by 20 ml propionic acid to give a
murky dark green suspension. Oxygen was bubbled (20 ga. needle)
through the suspension for 100 min. An ion-exchange column was
prepared from 40 g Dowex 50X8-400 resin and glass wool plugs. An
initial wash with 250 ml 2N HC1 was followed by conductance water
until neutral, 50 ml acetic acid and 50 ml propionic acid. The
reaction slurry was passed through the column, depositing any
precipitated solids on the glass wool and reddening the resin
with Co^+. The column was washed down with 50 ml propionic acid
and the clear green solution stripped at 40°. The glassy
green/black product was vacuum dried at 25°. The complex appears
air stable though decomposes in water or upon heating > 100°,
giving a purple product. The best formulation for cabaltic
propionate prepared in this manner, suggested by combustion
analysis, is for the dehydrated triraer, Co^OCprop^i^O), though
characterization failed to confirm the complex nuclearity. Anal,
caled, for C1 g H35 Co^ 0^: C» 33.30; H, 4.97; N, 0.00. Found:
C, 32.82; H, 4.83; N, 0.00.
Ru^OCprop^Ct^O)^, yg-Oxotrlaquohexakis(propionato)ruthenium(II)
diruthenium(III), was prepared from the RuClID^ cation as
described in the next chapter.
Ru^0(pr op) g( PPh^) j» U2-0^otrls(triphenylph°3phlne)hexakis
(propionato)ruthenium(II) diruthenium(III), was also prepared
from the RuCIII)^ cation as described in the next chapter.

CHAPTER V
OXIDATIVE DEHYDROGENATION OF ALCOHOLS
CATALYZED BY OXOTRIRUTHENIUM CARBOXYLATES
A. Introduction
The design and use of metal clusters to effect unique
homogeneous catalytic transformations relies on two important
differences from their mononuclear counterparts. 1) In addition
to new available substrate/ligand binding modes (e.g., bridging,
capping), cluster reactivity predicated on contiguous multiple
metal atom sites may be realized.^® 2) The electronic
properties of metal clusters may approach those of the bulk metal
where ancillary metal atoms serve to open up and stabilize
additional redox states.a fine demonstration of the latter is
the ruthenium trimer, Ru^CXC^CRjgL^, encountered in the last
chapter. Wilson et al. have revealed an extensive
electrochemistry for the ruthenium trimersj^ The cyclic
voltammograra for the cation system Ru^CXCH^CC^^Cpy^PY2)* (py is
pyridine, pyz is pyrazine) displays four electrochemically
reversible one electron waves indicating that the cluster system
remains intact in five discrete molecular oxidation states: +3,
+2, +1, 0, -1. Cluster linking via pyrazine bridges to prepare
higher oligomers afforded a trimer of trimers for which a series
129

130
of ten one or two electron waves were recorded, prompting the
epithet "electron sponge."
Facile ligand substitution and exchange has been established
for both the neutral and cationic trimers, [Ru^OCC^ClOgL^]n+.
Spencer and Wilkinson exchanged coordinated 1^0 and MeOH for PPh^,
MeOH, pyridine and various ir-acid ligands such as CO, CH^NC,
SOg» NO and PCOMe)^ 1 3° Fouda et al.131 demonstrated facile
displacement of DMF for whereas Trzeciak and Ziolkowski1^
monitored the exchange kinetics of 1^0 and DMF for cumene
hydroperoxide by NMR line broadening, concluding rapid exchange at
253K.
The fundamental requirements for a successful homogeneous
catalyst would appear to have been satisfied—reversible access
to multiple oxidation states and coordinative unsaturation/facile
ligand exchange. Indeed, the cationic ruthenium carboxylate
trimer ha3 been successful 1 ly utilized as a homogeneous catalyst
for a variety of transformations. In early work Wilkinson et al.,
employed the cationic trimer as an olefin hydrogenation catalyst
precursor; the active catalyst is believed to be
mononuclear.^1 *133»13^ Subsequent olefin hydrogenation studies in
DMF demonstrated retention of the cluster framework with
coordination of both the olefin and dihydrogen at the same metal
center.131 jn the absence of dihydrogen, cationic ruthenium
trimer isomerizes terminal olefins.131 The cationic trimer has
also been used to catalyze disproportionative coupling of
acrylonitrile,1allylic alcohol isomerization,1 3^ cumene
peroxide decomposition, 13^ periodate oxidation of olefin,13?
peroxide oxidation of olefins,1^® the Prins reaction,139 and

131
peroxide oxidation of phenols.140 Spencer and Wilkinson reported
central y-oxo atom abstraction and insertion for the ruthenium
11 4
trimer. Under dihydrogen pressure, the cationic trimer
undergoes first a one electron reduction to the mixed valence
Ru(III2II) oxobridged complex followed more slowly by reduction
to the Rudl)^ complex from which the y-oxo atom has been
removed. Oxygen atom insertion can be accomplished by 02 or
pyridine-N-oxide at room temperature:
[Ru30(C02R)L3]+ Ru30(C02R)L3 Ru3(C02R)L3 (5-1)
-e 0
With respect to the other metal carboxylate trimers, this
demonstration of reversible central oxygen atom abstraction is
unique to the ruthenium cluster. Our interest in this metal
cluster centered on this reactivity and its possible incorporation
into a substrate oxidation scheme. Provided that a reducing
substrate could effect the two electron step, the trimer offers
intriguing possibilities as a non radical autoxidation catalyst.
Towards this end, initial experiments explored the reactivity of
[Ru30(prop)g(H20)3]prop with olefins under dioxygen. No olefin
oxidation products were observed; instead the metal complex was
engaged in catalyzing the oxidative dehydrogenation of solvent
alcohol. As Gagne has pointed out,141 few transition metal
complexes are known that catalyze the homogeneous oxidative
dehydrogenation of alcohols, and in most cases, the reaction
142—154 *
mechanisms are not well understood. ' Investigations with
* Provesta Corp.introduced an alcohol oxidase enzyme in 1985
which catalyzes the reaction: Alcohol + 02 -> Aldehyde + H202.
The enzyme is believed to contain catalytically active copper.

132
the ruthenium trimer were undertaken to define the scope and
possible mechanism of these oxidations. The previous homogeneous
studies have utilized solely monomeric noble metal complexes. The
present work offers an additional incentive for study: to
disceren the synergistic role of three metal atoms in this
reaction, provided the cluster remains intact.
B. Results & Discussion
Scope of Catalysis
At elevated temperatures, [Ru30(02CR)gL3]n+ (n=0, R=CH3,
L=H20, PPh3; n= 1, R=CH3, CH2CH3, L=H20, MeOH) catalyzes the
homogeneous oxidative dehydrogenation of primary alcohols to
aldehydes and secondary alcohols to ketones. Typically,
oxidations were carried out in a glass pressure bottle fitted with
a brass pressure head, employing 50 ml solvent (usually substrate
alcohol) and 40 rag catalyst (ca. 1 raM) under 40 psi dioxygen
pressure. Product formation was followed as a function of time
after the catalyst had been added to the incubated solution.
Regular sampling of the reaction solution allowed monitoring of
products by GC. Identification was confirmed by GCMS in most
cases. Internal ketone standards were used for quantitative
analysis; product formation was measured by GC peak area. A
calibration curve of mole ratio vs peak area ratio was
constructed for a range of product/standard ratios in alcohol
solution. In some instances, trace amounts of carbonyl product
were present in starting alcohol solution. Where detected,
product yields were corrected for this initial concentration.
Oxygen uptake was estimated from the pressure drop registered on

133
the pressure head gauge. In control experiments, a pressure
bottle charged with i-propanol, heated to 65° and pressurized to
40 psi O2 registered no (<0.5 psi) pressure drop over 43 hours.
In the absence of gas formation then, this method provides a
rough measure of oxygen consumption.
For initial experiments, ethanol solutions of 1-hexene or
cyclohexene (10% v/v) and ruthenium trimer cation gave no products
derived from the olefin. Similarly, an ethanol solution
containing 10% 1-hexene and the mixed valence catalyst Ru^O
(prop)g(PPh^)^ delivered no olefin oxidation products.
Acetaldehyde was the only product detected when the solution was
heated to 65° under dioxygen pressure. Olefins were omitted in
all subsequent experiments.
The propionate bridged cation, Ru^CKprop^tHgO^"1" , was
preferred to the acetate bridged for its better solubility; the
former was used for most of the experiments though the latter
demonstrated comparable activity. (Other trimer and one dimer
form of ruthenium carboxylates were also investigated, these
results are deferred to a later section. The following discussion
refers to catalysis with the propionate bridged cation.). Shown
in Fig 5-1 are the production profiles for acetaldehyde, acetone
and propanal from ethanol, i-propanol and n-propanol at various
temperatures. In no case was an induction period observed.
Ru^CKpropJgíí^O)^ was inactive for ethanol oxidation at 25° but
did show slight activity for n-propanol oxidation at this
temperature. The temperature dependence of the initial rates is
consistent with a doubling of rate for each 10° (i-propanol: 65°,

13^
Fig 5-1. Oxidative dehydrogenation catalyzed by
Ru^OCprop^CF^O)^. Numbers refer to Table 5-1.

135
Table 5-1. Products of alcohol oxidation catalyzed by
[ru3o(o2cch2ch3)6(h2o)3](o2cch2ch3).
no.*5
solvent/substrate
TC°
C) product (s)
TO/12°
TO/24'
1
ethanol
25
no reaction
—
—
2
ethanol
65
acetaldehyde
198
313
3
i-propanol
65
acetone
147
254
4
i-propanol
100
acetone
685
1015
5
n-propanol
25
propanal
16
25
6
n-propanol
65
propanal
430
645
n-butanol
65
butanal
(e)
cylohexanol
65
cyclohexanone
(e)
t-butanol
65
no reaction
_,
a) Reaction conditions: 50 ml alcohol, 40 mg catalyst, 1 ml 2-
hexanone standard and 40 psi 02 initial pressure.
b) See Fig 5-1.
c) Catalyst turnovers defined as mole product/mole catalyst in
12 hours.
d) Catalyst turnovers in 24 hours.
e)not quantified

136
0.6 mmol/hr; 100°, 5.7 mmol/hr.) Rates of reaction using this
catalyst follow the order n-propanol > ethanol > i-propanol.
Table 5-1 demonstrates that catalytic oxidation of primary
and secondary alcohols appears to be general. t-Butanol, which
contains no a -hydrogen atoms, was unreactive, as expected.
Primary and secondary alcohols selectively gave only the two
electron oxidation products under these conditions, and further
oxidation of aldehydes to acid was not detected for any of the
primary alcohol solutions.
Product concentration determination involved syringe sampling
of the pressure bottle solution under reaction conditions. This
method assumes the solution concentrations accurately reflect the
true product/standard concentration ratios. The high volatility
of acetaldehyde (bp 21°) made this method suspect for ethanol
oxidations. A pressure bottle experiment with ethanol was run
under identical conditions as in Table 5-1 (no. 2) and sampled
only before pressurizing. After 24 hours the reaction was
quenched by immersing the entire pressure bottle assembly—under
pressure—into an ice bath for 20 minutes and sampling as usual.
Analysis gave 339 turnovers compared to 313 for the reaction
sampled at 65°. The difference is in the expected direction but
not significantly large for our purposes. Acetaldehyde volatility
does not present a sampling problem.
Product identification was afforded by spiking the reaction
solutions to verify GC retention times, and GCMS characterization
of ethanol, n-propanol and i-propanol solutions revealed the
characteristic fragmentation patterns for acetaldehyde, propanol
and acetone. Primary alcohols and their aldehyde products might

137
be expected to form acetals especially in alcohol solution; ketal
formation from secondary alcohol and ketone would be less
favorable. These species were looked for in oxidations employing
ruthenium trimers. Isopropanol gave only acetone in all
experiments with no other significant products detected by GC or
GCMS. n-Butanol gave only butyraldehyde with no other significant
products detected by GC. n-Propanol gave only propanal with no
other products detected by GC (both DEGA and Carbowax columns) or
GCMS. Ethanol oxidations gave only acetaldehyde by GC. The
acetal was, however, detected by GCMS after 24 hours and found to
elute with ethanol making quantification difficult. Conversion of
some of the product acetaldehyde to acetal would imply higher
activity than reported here by monitoring of only the
acetaldehyde peak area. Detectable acetal or ketal products with
substrates other than ethanol might be expected in that solvent
alcohol should exhibit a mass action influence on the ketal/acetal
equilibria. Equilibration is driven by acid catalysis, however;
pH readings of ethanol and isopropanol oxidation solutions gave
values of 7.2-7.7. Thus, the ruthenium trimers examined here are
selective to only the two electron alcohol oxidation products and
acetal/ketal conversion was only detected for ethanol.
GCMS analysis also identified some unanticipated products.
n-Propanol solution showed the ketal of 2-hexanone and n-propanol.
Ethanol solution showed the heraiketal formed from 2-hexanone and
ethanol. i-Propanol solution gave the hemiketal of 2-hexanone and
i-propanol (though not that of acetone and i-propanol). Although
none of these addition products were produced in large enough
concentration to be detected by flame ionization (and hence alter

138
the standard concentration), MS detection dictates a caveat for
future studies to avoid use of ketones as standards in alcohol
solutions.
While the above discussion pertains to reactions run with
Ru^OCprop^Ct^O)^, several other ruthenium carboxylates were
examined with regard to catalyzing alcohol dehydrogenation as
shown in Table 5-2. The acetate bridged cation appears to be a
slightly better catalyst than the propionate bridged congener
though the difference is not great. The most striking result,
however, is that choice of catalyst oxidation state is flexible -
both the cationic (III, III, III) and neutral mixed valence (III,
III, II) trimers are active. And except for the third entry in
Table 5-2 (albeit under different reaction conditions), the mixed
valence complexes afford slightly higher product yields. This
latter point is surprising in that the mixed valence trimers are
synthesized via reduction of the cationic trimer, the aquo adduct
by H2/PtÜ2 and the PPh^ adduct by reaction with PPh^ (Fig 5-2).
Spencer and Wilkinson have shown the aquo adduct to revert to the
cation upon standing in air, while the PPh^ adduct is stable to
air. Under oxidizing (C^) conditions, either mixed valence
complex might be expected to shuttle through the (III)^ oxidation
state. On the other hand, the (III)^ complex may form the
(III)2(H) complex under reducing (ROH) conditions. In either
case, catalytic activity would be similar for Ru^Oiprop^O^O)^
and Ru^Oiprop^G^O)^ unless an induction period were evident for
the formation of the active catalyst species. One is not. Fig 5-
3 demonstrates the catalytic activity of the various trimers for

139
Table 5-2. Product yield dependence upon ruthenium catalyst.3
substrate
catalyst T
RC(0)R'/24
(mmol)b TO
ethanol
0.035 mmol
Ru30(prop)6(H20)3+ 65°
12.0
339
ethanol
0.032 mmol
Ru30(prop)g( ^0)3 52
13.9
433
ethanol0
0.019 mmol
Ru30(prop)g(P<|i3)3 60
4.3
228
n-propanol
0.051 mmol
Ru30(0Ac)6(H20)3+ 65
34.6
672
n-propanol
0.045 mmol
Ru30(prop)g(H20)3+ 65
29.1
645
(1) i-propanol
0.041 mmol
Ru30(0Ac)6(H20)3+ 65
11.0
270
(2) i-propanol
0.036 mmol
Ru30(prop)6(H20)3+ 65
9.0
254
(3) i-propanol
0.018 mmol
Ru30(prop)g(H20)3 65
14.1
764
(4) i-propanol(
^ 0.015 mmol Ru^OCprop)^P 1^)3 65
13.5
904
i-propanol
0.052 mmol
Ru2(but)4Cl 65
0.7
13
a) Reaction conditions: 50 ml alcohol, 1 ml 2-hexanone or 2-
octanone standard, 40 psi initial 02 pressure, catalyst.
b) Product carbonyl yield after 24 hours.
c) 50 ml ethanol, 10 ml benzene, 5 ml 1-hexene
d) 40 ml i-propanol, 10 ml benzene

140
ru3o(co2r)6(h2o)3 +
+e
V
°2
Ru30(C02R)é(H20)3
+2e
V
A
°2
ru3(go2r)5(h2o)3
114
Fig 5-2. Reactions of oxotrirutheniura complexes.

141
Fig 5-3. i-Propanol oxidation as a function of catalyst.
Numbers refer to Table 5-2.

142
i-propanol oxidation. No induction periods are evident; the
catalysts differ only by their relative activities.
That the two mixed valence trimers gave consistent results
under similar conditions (Fig 5-3) was unexpected.
Triphenylphosphine does not, apparently, coordinatively saturate
the triraer in alcohol solution. Other donors—sodium ethoxide and
2,6-lutidine—had accelerating and inhibiting effects on i-
propanol oxidation catalyzed by the cation (vide infra). To
reaction solutions run with Ru^CXpropJgiPPh^)^ it was necessary
to add 20% v/v benzene to dissolve the catalyst in alcohol.
To discern if the catalysis studied here is general for
ruthenium carboxyl ates, the Ru2(but)¿jCl dimer was also screened
for i-propanol oxidation. While several turnovers were achieved,
the dimer does not rival the trimer activity (Table 5-2)
suggesting unique electronic or coordination properties associated
with the trimers.
During initial experiments employing the cationic ruthenium
trimers, it became clear that catalyst activity was a function of
catalyst preparation, with significant variation depending upon
extent of purification. Spencer and Wilkinson reported that crude
[Ru20(0Ac)g(H20)2]0Ac contained a considerable excess of sodium
acetate. The pure compound could be obtained via dialysis in
water, by chromatography on Sephadex or, best, by
recrystallization from raethanol/acetonej ^ Fouda et al. found
this latter method inadequate and reported exhaustive ethanol
extraction followed by impurity precipitation gave best
1 O 1
results. J In light of these reports and the observed
inconsistencies in trimer cation activity for catalysts prepared

143
by Wilkinson's method, chromatographic separation was undertaken
to isolate the catalyst. Crude [Ru^CKOAc^íf^O^ÜOAc was
prepared by refluxing RuCl^HgO^ with excess sodium acetate and
acetic acid in ethanol, repeated cooling to -78° to precipitate
sodium acetate and sodium chloride followed by solvent stripping.
Cation prepared in this manner repeatedly gave low analysis
for C and H while showing the presence of trace N. Activated and
neutral aluminas, silica and activated Florasil were inadequate
while Sephadex (a size exclusion) gel gave adequate separation.
Fractions were rechromatographed to elute three components which
were stripped of solvent methanol. The first (highest mol. wt.)
to elute, I, gave a grey green solid (olive in solution), the
second, II, a black green solid (blue green in solution) and the
third, III, a light green solid (lime in solution) in a mass
ratio of about 10/100/1. Except for the NMR, spectral analysis—
Table 5-3—indicates the major fraction, II, agrees most
favorably with reported parameters for the cation. The reason
for different chemical shifts than that reported for the cation
accompanied by C10jj“ anion is not clear; a satisfactory
integration ratio of ca. six to one is observed , however.
Similar chromatagraphic preparation of [Ru^Oiprop^i^O)^]-
(prop) yielded similar results. The middle, major fraction gave
the expected NMR: 6= 2.18(q, 2H, anion), 1.42 (q, 12H, bridge),
1.08 (t, 3H, anion), 0.80 (t, 18H, bridge), with paramagnetic
broadening of bridge propionate protons (Fig 5-4). The FTIR
spectrum of this fraction is shown in Fig 5-5. Despite this
treatment, both the acetate and propionate bridged cations

144
Table 5-3
I
II
III
reported
(114,156)
Analysis of chromatographed fractions.
coo,asy vcoo,sym UV-VIS
1H NMR
elem.
anal.
1563 1418 580
31 Osh
1.87(s) 24.53ÍC
1,02(t) 4.12 H
2.14(q)
1560 1426
680(1200) 1 .87(s) 20.02 C
610(1900) 0.54(s) 3.36 H
390sh(ca.3400) 0.41 N
310sh(ca.5800)
several
singlets
1557
1430 686(1100) 21.38 C
629(1000) 2.30 3.46 H
391sh(ca .1250) (bridge) 0 N

Fig 5-4. 1H NMR spectrum of [RuoCKprop^f^O)^] (prop) (fraction
II) in MeOH-d . The methanol methyl resonance ( = 3.30) is used
as the reference.

146
- iD

Fig 5-5. FT IR spectrum of [Ruo0(prop)g(H20)?](prop) (fraction
II) in KBr.

^TRANSMITTANCE
WAVENUMBERS (CM-l)

149
continue to give trace nitrogen analysis (see exptl. sect.). The
nitrogen source is unclear. Twelve different preparations of
acetate and propionate bridged cation all analyzed for nitrogen
present, ranging from 0.35 - 0.635. Ru^CKprop^iPPhj)^, formed
by addition of PPh^ to [Ru^OCprop)^ (i^O)^](prop), analyzed for
05 N.
To assess the catalytic properties of the chromatographed
species, I and II (acetate bridged cation) were used in tandem
standard oxidations of i-propanol at 65°. After 24 hours, II had
generated 270 turnovers for acetone, based upon
[Ru^OiOAc^iHpO^KOAc) while I gave only 32 "turnovers" based on
the same catalyst formula for comparison. Not enough of III had
been generated for a coraparitive experiment. Clearly then, the
active catalyst precursor is the cationic (or neutral) ruthenium
trimer which can be isolated by size exclusion chromatography.
Catalyst prepared in this manner consistently analyzes for small
amounts of nitrogen, the source of which is unclear.
Finally, in control experiments, omission of either alcohol
(cyclohexane solvent) or catalyst gave no detectable quantities of
carbonyl product after 24 hours.
2. Mechanistic Considerations
Having shown the catalytic nature of the ruthenium trimer for
alcohol oxidative dehydrogenation, the objective became one of
establishing a rational mechanism. Towards this end, the
following experiments address the mechanistic details of this
reaction

150
Ruthenium complexes have previously been employed for
homogeneous catalytic oxidative dehydrogenations. Tang, et al.,
communicated oxidation of amines and alcohols with ruthenium
trichloride at ca. 100° and 2-3 atm O2 to form both
dehydrogenation and dehydration products.1Rileyutilized
ruthenium complexes of the form RuX2(Me2S0)^ for the tandem
oxidation of thioether and alcohol to sulfoxide and carbonyl
products by oxygen at 100° and 100 psi O2. Gagne and Marks11*1
found ruthenium complexes of 1,3-bis(2-pyridyliraino)isoindolines
BPI, (Fig 5-6) active for alcohol oxidations even at ambient
temperature and pressure. Only in the latter were conclusive
mechanistic details reported. The activity of the
RudllKBPDCl^/base system approaches that of the trimers but is
not quite as high. In terras of mole product per mole ruthenium,
ethanol oxidation achieves 25 turnovers at 70° with the
Ru(III)(BPI)C13 catalyst while [ Ru(III>3 (prop) g ( H20) 3 ] (pr op)
delivers 104 turnovers at 65°. Additionally, strong base is
required in the former system and can lead to product aldehyde
disproportionation.1¡*1
Gagne and Marks found base addition necessary (Fig 5-7). In
their proposed mechanism, strong but non-complexing bases (such as
2,6-lutidine) were necessary for HC1 abstraction, H+ from the
ligand (4'-MeLH) and Cl“ from ruthenium. Removal of Cl“ allows
alcohol coordination and solubilizes the catalyst. A complexing
base (such as pyridine) hinders the oxidation. Catalysis is
faster with the stronger base sodium ethoxide than with 2,6-
lutidine, which is attributed to the base dependence of the
disproportionation step, favored in basic solution. The

151
R
R
LH, R = H
4'MeLH, R = CH3
¿J'-s-BuLH, R = sec-butyl
Fig 5-6. The 1 f3-bis-(2-pyridylimino)isoindoline (BPI) ligands
utilized by Gagné and Marks.

152
2(4'-MeLH)RuCl3
base
v
2[(4'-MeL)RuCl3]“
-2 Cl-
+2 C2H5OH
h2o
c2h4o c2h5oh
Fig 5-7. Mechanism for ethanol oxidation proposed
Gagne and Marks. (4'-MeLH) is the ligand (1,3-bis (4-methyl
2-pyridylimino)isoindoline) or 4'-MeBPI.
by

153
disproportionation of Ru(III) amines was shown to be base
dependent by Rudd and Taube156 (Eqn 5-2).
2(NH3)5Ru(III)py3+
(NH3)5Ru(II)py2+ + (NH3)i|(NH2)Ru(IV)py3+ + H+ (5-2)
Ru(III) disproportionation followed by stoichiometric
oxidation of bound alcohol has been reported by Tovrog et
al.^^ (Fig 5-8). This reaction too was favored in basic
solution.
A similar role for Bronsted base might be expected utilizing
the ruthenium triraer for alcohol oxidations. Disproportionation,
presumably intramolecular (Eq 5-3)»
[III, III, III (ROH)]+ [II, III, IV(OR)] + H+ (5-3)
would similarly be base enhanced. As in the BPI system,
coordinated ligand (OAc“ or even u3-02“) may act as the
alternative proton acceptor.
The effect of sodium ethoxide and 2,6-lutidine on isopropanol
oxidations is shown in Fig 5-9. As with the BPI system higher
activity is observed for added NaOEt than 2,6-lutidine, but, more
importantly, 2,6-lutidine has an inhibitatory effect, the reason
for which is unclear as coordinative saturation is unlikely with
this base. If base perturbation of the disproportionation
equilibrium (Eq 5-3) is indeed important, the basicity order
OEt” > prop” > 2,6-lutidine
would predict rate enhancement for added sodium ethoxide with
little or no effect due to added 2,6-lutidine since the propionate

154
(m^1
'U
O-CH
/ \
H ch3
8
Fig 5-8. Oxidation of coordinated alcohol at high pH reported by
Tovrog et al.158

155
Fig 5-9. Effect of added base upon isopropanol oxidations with
Ru^CKprop^Cf^CO^prop (0.046 mmol).

156
bridges are always present. Consistent with this reasoning,
tandem isopropanol oxidations run with and without NaBr (1.3 fold
excess) showed no effect from added NaBr (with, 111 TO/12 hr;
without, 110 TO/12 hr). Still, the effect of 2,6-lutidine is not
readily explained in this context.
The use of hydroxide ion gave a unique result in isopropanol
oxidation. Following accelerated acetone production, catalyst
activity quickly fell as the catalyst was precipitated to yield a
clear orange solution (originally deep blue-green) after 24 hours
(Fig 5-10). Using a 73-fold excess of NaOH (untreated, 2.6 mmol)
produced the same molar quantity (2.6 mmol, 73 TO) of acetone,
indicating stoichiometric consumption of hydroxide. Catalyst
precipitation warrants this an unattractive option for tailoring
activity.
Bronsted acid would have an inhibiting effect on the
disproportionation (Eqn 5-3) and, if this step was rate
determining, an inhibiting effect upon rate of oxidation.
Conversely, strong acid would protonate carboxylate bridges,
displacing strongly coordinated carboxylate for weakly coordinated
conjugate base. This "exposed edge" strategy has been used to
synthesize a number of metal-metal bond complexes with
coordinative unsaturation^® and to generate hydrogenation
catalysts from ruthenium and rhodium carboxylates.^ If acid were
to open up additional reaction sites, activity enhancement would
not be unexpected. The effects of perchloric (as 60% aq. soln.)
and trifluororaethyl sulfonic (neat) acids upon isopropanol
oxidations is shown in Fig 5-11. Both serve to seriously inhibit
the oxidation reaction with the latter shutting down the reaction

157
Fig 5-10. Effect of added base upon isopropanol oxidation
with Ru^OCprop)^ F^O^prop (0.035 mmol).

158
Fig 5-11. Effect of acid upon isopropanol oxidation with
iU^OCprop^Cf^CD^prop (0.035 mmol).

159
almost completely. Conceivably the perchlorate ion may act as a
stoichiometric oxidant (this was not separately tested for) and
hence the higher activity observed with this acid. Regardless,
these strong acids only serve to inhibit oxidation catalyzed by
ruthenium trimer which is consistent with disproportionation
(Eqn 5-3) as the rate limiting step.
The effects of some other metal salts were also looked at in
isopropanol oxidations. ZnCOAc^O^O^ (2.3-fold excess) was
employed as an arbitrary Lewis acid. CudlO^^Q^O)^ (3.4-fold
excess) was introduced since Cu(II) was shown to play an important
role in mediating reoxidation of Pd(II) salts for alcohol
oxidations (Wacker chemistry).11*^ Apparently Cu(II) is not needed
or does not function in this manner, nor does Zn(II) enhance
isopropanol oxidation as shown in Fig 5-11. Both metals inhibit
this reaction, the Cu(II) precipitating slowly as an insoluble
light solid. Blackburn and Schwarz reported a similar result in
PdC^-NaOAc catalyzed oxidations of secondary alcohols. These
workers reported11*1* that alcohol oxidation is retarded by the
presence of Cu(II) salts, the Cu(II) slowly precipitating as a
blue-green material.
Alcohol oxidation may proceed via a free radical mechanism.
The autoxidation of primary or secondary alcohols affords
hydrogen peroxide and aldehydes or ketones, respectively, as the
primary products.1-^ Catalysis by metal salts can result from
chain initiation, which stems from the reaction of the metal
species with H2O2 or with alcohol substrate:

160
Mn+ + H202 -—> M(n-1)+ + H02 + H+ (5-4)
M(n-1)+ + h2o2 * Mn+ + HQ* + H0- (5-5)
Mn+ + ^CHOH > M(n-1)+ + ^COH + H+ (5-6)
Under the usual autoxidation conditions (mild heat, 02) the
aldehydes derived from primary alcohols undergo further oxidation
to the corresponding carboxylic acids. Alcohol oxidations
catalyzed by oxoruthenium trimer were examined with regard to
radical participation, and these results are shown in Fig 5-12.
In a standard run (65°, 40 psi initial 02 pressure) with
isopropanol, no ruthenium catalyst and 16 mg (0.097 mmol) 2,2-
azobis(2-raethylpropionitrile) (ABMP), a free radical initiator, no
acetone production was detected in the first six hours though
0.385 mmol acetone (4.0 TO w.r.t. ABMP) had formed after 12 hours.
Incorporation of benzoquinone (0.26 mmol, 6.8-fold excess) had no
effect upon acetone production. A radical or radical chain
process does not appear to be involved from these findings.
Further support comes from the reaction selectivity, with no acid
products detected due to further oxidation of aldehydes formed
from primary alcohols.
While alcohol oxidation clearly produces carbonyl products,
the fate of reduced oxygen has not been discussed. Oxygen
reduction in this system can proceed along several routes with an
initial two electron step most likely. Product peroxide may be
decomposed by a variety of species (oxotriruthenium cation

mmol acetone
161
Fig 5-12. Effects of free radical trap and radical initiator (no
catalyst) upon isopropanol oxidations with Ru^CXprop^C^CO^prop.

162
effectively decomposes cumene hydroperoxide); ^2 or catalyzed
disproportionation may occur. Hydrogen peroxide may also undergo
another two electron reduction with concomitant production of a
second mole of carbonyl product:
2 HO*
H+
->H2C>2-
1/2 02 + H20
H+ 2 H20
(5-7a)
(5—7b)
(5-7c)
The hydroxyl radical is rather indiscriminate in its reactions
with organic substrates, and the observed selectivity argues
against its formation. Specifically, hydroxyl radical production
would generate products derived from a - and ¡3-hydrogen
abstraction of alcohol and, if formed, aldehyde.1^
Water is indeed a product in these reactions and can be
monitored by GC using a thermal conductivity detector (TCD). This
analysis reveals two aspects of water participation in these
alcohol oxidations: 1) H20 wa3 present in the starting alcohol
solutions, of the order for that which was produced during the
course of the reaction in the case of i-propanol, ca. 0.1 M. 2) 1
mole of water was produced for every mole of carbonyl product
formed. Dual analysis of an isopropanol oxidation—FID for
acetone and TCD for water—allowed comparison of their yields as
shown in Fig 5-13. Water initially present was compensated for
(7.0 mmol in this experiment) and the two profiles represent least
square lines for the data points, forced through the origin. This
treatment was used given the large scatter in water concentration
data points determined by the less sensitive TCD method. After 24
hours, this translates to 5.9 mmol H20 and 5.3 mmol acetone

mmol
163
Fig 5-13. Product yields for isopropanol oxidation catalyzed by
Ru^OCprop^Cf^CO^prop (0.050 mmol).

164
produced in 50 ml isopropanol for a molar ratio of 1.1,
demonstrating equimolar production of water and acetone. Product
water does not greatly inhibit the oxidation of isopropanol. A
solution of 45 ml HgO/IO ml ispropanol generated 1.67 mmol acetone
after 24 hours despite incomplete dissolution of the catalyst
while a 50 ml ispropanol solution run under identical conditions
resulted in 3.65 mmol acetone after 24 hours. Because catalyst
dissolution was a problem in the former, this result only serves
to illustrate that Ru^CKprop^C^O^prop continues to function in
the presence of water.
Though unlikely, hydrogen peroxide might not undergo
decomposition, disproportionation or reduction to water,
remaining instead in solution unreacted. Catalyst solutions were
tested for the presence of hydrogen peroxide using the
ferrithiocyanate test.^1 This qualitative indicator gives a
blood red color if peroxide is present due to oxidation of ferrous
iron, but does not quantitate the peroxide concentration. Testing
of reagent ethanol stored over molecular sieves and used for
oxidation experiments gave a slight red coloration. Passage
through neutral alumina gave ethanol which gave a less pronounced
color indicating incomplete scrubbing of peroxide contaminant. An
ethanol oxidation that had run for 72 hours to produce 18 mmol
acetaldehyde was tested after the solution had been allowed to
cool, ca. one hour. The red tint produced was even more subtle
than for the alumina treated ethanol. Identical testing was
performed on ispropanol solutions stored over molecular sieves,
passed over alumina and ran in an oxidation run for 72 hours to
give 5 mmol acetone. Again the untreated alcohol showed some

165
peroxide contamination, alumina removed most (but not all) of this
contamination, and an actual catalyst solution, tested about an
hour after acetone production was quenched by cooling the
reaction, gave the lowest peroxide concentration. Referring back
to the ethanol solutions, the acetaldehyde concentration after 72
hours was 0.35 M. An aqueous hydrogen peroxide solution of about
105& of this (0.031 M) was tested and gave a deep crimson color.
These results demonstrate that, if produced, hydrogen peroxide is
rapidly consumed or decomposed in the catalyst solution and is not
a stable product of oxygen reduction under these conditions. The
demonstration of peroxide contamination in reagent alcohol and
(incomplete) removal by alumina prompted regular scrubbing of the
alcohols. Though significant differences between chromatographed
and untreated alcohol solutions were not observed, alcohols were
passed through neutral alumina just prior to using for oxidations.
If hydrogen peroxide is the intermediate reduction product,
the above experiments demonstrate only its disappearance, but not
how. Disproportionation (Eq 5-7b) vs reduction (Eq 5-7c) can be
demonstrated by the manner in which the ruthenium catalyst
utilizes hydrogen peroxide under catalysis conditions. Addition
of 1.0 ml of 30% ^(^(aq) to a standard isopropanol oxidation
using a reduced amount of catalyst resulted in a significant
acetone concentration by the time of the first sampling (0.5 hr)
as shown in Fig 5-14. With or without the help of the ruthenium
catalyst, hydrogen peroxide rapidly oxidizes alcohol under the
standard reaction conditions. That the ^2^2 consumption is not
quantitative (8.8 mmol H2O2 added, 3.9 mmol acetone produced
initially) indicates that disproportionation to H2O and O2 may

mmol acetone
166

167
also occur. Under normal reaction conditions, consumption of HgO^,
may be different in that its production occurs much more slowly
than its appearance in this experiment, with some steady-state
concentration likely. Regardless, these results demonstrate that
the ruthenium trimers are capable of two two-electron reductions
of dioxygen to give 2 moles of carbonyl product and two moles of
water.
Oxygen uptake in these alcohol oxidations was apparent by the
pressure drop recorded from the pressure head gauge in all
experiments. While this is not the preferred method for strict
determination of oxygen uptake due to the insensitivity of the
pressure gauge, it does allow estimation of the oxygen
stoichiometry in these reactions. Continuous sampling of the
pressure bottle setup during a run undoubtedly exaggerates the
oxygen consumption despite the fact that a pressure drop could not
be detected during a given sampling. Oxygen uptake was therefore
estimated from an experiment in which the solution was not sampled
during the 24 hour duration. The previously described experiment
used to verify acetaldehyde product concentration from ethanol
oxidation was used for this purpose. Ethanol and standard were
incubated for 30 min., catalyst added, the solution sampled, and
the system flushed and pressurized as usual. The pressure was
immediately recorded and this was considered the start of the
reaction. After 24 hours, the pressure was recorded and the
system plunged into an ice bath before final sampling. In this as
in other experiments the pressure drop was initially accelerated
before stabilizing to a constant rate as shown in Fig 5-15.
Attributing this initial rate to solvent uptake, the equilibrium

168
Fig 5-15. Oxygen consumption in the oxidation of ethanol
catalyzed by Ru^OCprop^Ci^O^prop (0.035 mmol).

169
uptake was extrapolated back to time zero to estimate the oxygen
consumed in this reaction. Recorded pressure drops were converted
to molar consumption using the ideal gas law. Gauge pressure can
be read to about 0.3 psi accuracy which converts to +/- 0.2 mmol
oxygen. This treatment gives an oxygen consumption of 5.8 +/- 0.2
mmol with corresponding production of 12.0 mmol acetaldehyde over
24 hours for a ratio of 2.1. Within the error of this method
then, the reduction of one mole of dioxygen by ruthenium trimer
nets two moles of alcohol oxidation product.
A dependence of reaction rate upon oxygen pressure was
observed, and the reaction order in oxygen was determined.
Isopropanol oxidations using Ru^OCprop^Cf^O^prop (0.046 mmol)
were ran at 65° under (initially) 69 psi O2, 42 psi O2 and
ambient air. Acetone concentrations determined by GC provide the
best measure of extent of reaction, and initial reaction rates
were determined from tangents to the acetone production curves
through their origins. The van't Hoff differential method allows
determination of reaction order for a given species from a plot of
In rate vs In concentration:
d[A] = k[A]n (5-8)
dt
In d[A] = In k + n In [A] (5-9)
dt
A plot of In initial rate (TO/hr) vs In initial PQ2 (psi) is
shown in Fig 5-16. The slope, n, gives a reaction order of 0.25
for dioxygen

In initial rate
170
2.0 -\
1.5 H
1.0 -\
o.5
0
T
In P,
4
02
Fig 5-16. Determination of isopropanol oxidation rate order in
dioxygen catalyzed by Ru^OCprop)^I^Oj^prop.

171
This rate dependence upon oxygen pressure should manifest
itself in a slowing of the reaction rate as alcohol oxidation
proceeds. As Fig 5-1 demonstrates, this is indeed the case, but
other factors may also contribute to lowering of catalyst activity
after several hundred turnovers. Adverse effects from water are
not likely (vide supra) nor is product aldehyde/ketone expected to
hinder the alcohol oxidations as alcohol solutions usually
contained 8.1 mmol 2-hexanone standard.
In their study of ruthenium tetroxide catalyzed oxidations of
organic compounds,^7 Sharpless, et. al, found the cationic oxo-
triruthenium carboxylate complex an effective catalyst for
oxidative cleavage of 1-octene by periodate only when acetonitrile
was included in the solvent system. Without acetonitrile, the
system was inactive, and this was attributed to coordinative
saturation by product pentanoic acid. Similarily, addition of
acetontrile to reaction mixtures in which catalysis had ceased
restored full catalyst activity. This strategy was applied to
the ruthenium catalyst employed in the present alcohol oxidations.
Addition of 10Í acetonitrile to an isopropanol/Ru^ system at 95
hours registered no increase in rate of acetone production
indicating that deactivation via the above mechanism is not
operative here.
Catalyst deactivation is important in terms of catalyst
longevity. For this reason and towards discerning the actual
active species, the catalyst solutions were examined after having
accomplished several hundred turnovers. A standard run, employing
0.036 mmol Ru^OCprop^CHgO^prop in 50 ml isopropanol was allowed
to proceed for 143 hours at 65° to accomplish about 1000

172
turnovers. At this point, the rate of acetone production was only
about 5% of that in the first half hour of reaction. The alcohol
solution remained clear and homogeneous. The usual color change
of dark blue green to olive green over the course of the reaction
was manifested in the electronic spectrum: free catalyst, in
methanol solution at 25°, A = 680( £ = 1800); 610 (1970); 380 (sh,
3250); 240 (sh, 11900). Catalyst in isopropanol, after 143 hours
at 65°, A = 680 ( e= 2160); 620 (2220); 360 (sh, 5470); 270
(20500).
In another experiment an isopropanol oxidation with
Ru^OtpropJgO^CO^prop was allowed to proceed for 45 hours at 65°
(135 turnovers, no rate decrease) and the catalyst recovered by
stripping the volatiles and vacuum drying (70°, 5 hr). Proton NMR
of the isolated green-black solid, though dirty, indicated the
presence of the starting catalyst material. FTIR revealed
intensity differences in the carboxylate ligand vibrations (Fig 5-
17). While these results do not identify the active catalyst they
do serve to indicate the species of interest is no longer strictly
the starting ruthenium triraer.
Finally, though the rate of oxidation could be increased
simply by raising the temperature, catalyst precipitation became a
problem. Running an isopropanol oxidation with
Ru^OCpropJgO^CO^prop at 100° produced 1015 turnovers within the
first 24 hours but catalyst precipitation began to occur at 6.5
hours (470 turnovers). This was the only alcohol oxidation run
above 65° and the only one in which any catalyst precipitation was
evident. Proton NMR of the precipitated solid was nondefinitive
though FTIR indicated carboxylate vibrations and UV-VIS

Fig 5-17. Ru^CKprop^CHgCO^prop in KBr (A) and catalyst isolated
from isopropanol solution after 45 hours (135 turnovers), in KBr
(B).

XTRANSUITTANCE
21.211 24.612 26. 013 31.414 34.614 36.215

175
demonstrated a lowering of catalyst concentration as reaction
proceeded. At 100°, then, deactivation occurs via precipitation
of some form of ruthenium carboxylate under reaction conditions.
One of the questions which remains in these studies is why
both the cationic Rudll)^ and mixed valence RuCIIDgClI)
complexes are equally good catalysts. Further, what is the nature
of the active catalyst and through what oxidation states do the
metal atoms shuttle to accomplish two electron oxidation of
alcohols? The response of the two trimers to oxidizing (C^) and
reducing (ROH) conditions can shed some light on both of these
questions.
The mixed valence (111)2(11) trimer is readily oxidized to the
cationic (III)^ trimer by C>2 in solution. Shown in Fig 5-18 is
the electronic spectrum of the (111)2(11) trimer in acetone under
nitrogen and after bubbling of oxygen through the solution for 10
minutes. The resulting spectrum agrees with that for the (III)^
complex. The (111)2(11) complex has been shown to revert slowly
114
to the (III)^ complex in the solid state upon air exposure.
The reverse reaction is not accomplished by alcohol in the
absence of oxygen. Shown in Fig 5-19 is the electronic spectrum
of Ru^OipropJgO^O^prop in water under nitrogen, after addition
of 200-fold excess of deaerated ethanol and after heating this
solution at 65° for 30 min. While spectral shifts are evident
with a new absorbance at 570 nm, this does not correspond to the
(III)2(H) complex.
A rational mechanistic proposal must be consistent with the
experimental evidence presented here and such a scheme is shown in
Fig 5-20. Consistent with the schemes advanced by Gagne (Fig 5-4)

Fig 5-18. Electronic spectrum of Ru^CKprop)^f^O)^
under nitrogen (A) and after passage of C>2 through so
10 minutes at 25° (B). Artifacts at 350 and 878 nm
optics changes.
in acetone
lution for
are due to

(WNiy ooq

Fig 5-19. Electronic spectrum of RuoO(prop)g(HpO)oprop in water
(3.2 X 10” M) under nitrogen (A), after addition of 200 parts
deaerated ethanol (6.8 X 10~^M) (B), followed by heating at 65°
for 30 min.(C).


°2
rgh2oh
II
RCHO + H20 RCH2OH
Fig 5-20. Proposed mechanism for catalyzed alcohol oxidative
dehydrogenation. Carboxylate bridges have been omitted for
clarity, and the ruthenium centers are represented by their
oxidation state numbers.

181
and Tovrog (Fig 5-5), this proposal relies upon the facile
disproportionation of Ru(III) (Eqn 5-3) to effect two electron
redox. The important difference is that this process is
intramolecular here. Some measure of catalyst activity is thus
probably due to elimination of an unfavorable entropy contribution
to this process. Each step of the mechanism will be discussed in
turn.
First, the mixed valence (IlDoClI) complex is readily
oxidized by oxygen in the absence of alcohol to give the cationic
(III)^ complex. The reverse reaction—reduction by alcohol to
form the (III)2(II) complex—must be followed rapidly by further
reaction since alcohol reduction of the cationic complex in the
absence of oxygen does not generate the electronic sepctrum of the
(III)2(II) intermediate. The proposed equilibrium accounts for
the utility of either the (III)^ or (111)2(11) complexes as
catalysts. Further, the observation that the ruthenium species
isolated from the catalyst solution gives IR and NMR signals
associated with the (III)^ species suggests the cationic complex
is the resting form of the active catalyst.
The disproportionation which follows has been established
for mononuclear Ru(III) fragments provided a Bronsted base is
available to accept the generated proton (Fig 5-5). Incorporation
of Bronsted acid or base would be expected to deliver a
corresponding slowing or enhancement, respectively, of this step
which may be manifested in the overall reaction rate. The
deleterious effects of perchloric and trifluoromethylsulfonic
acids and the positive effect of sodium ethoxide on the reaction
rates are consistent with this, though the negative effect of

182
2,6-lutidine is not. Alternatively, the central oxygen atom may
act as a proton acceptor to drive the disproportionation. If the
disporportionation step was rate determining, only Bronsted bases
stronger than the central oxygen atom would enhance this
reaction, while weaker bases would have little or no effect.
This proposal too is consistent with the rate enhancement brought
on by sodium ethoxide but fails to answer why 2,6-lutidine slows
the reaction instead of having no effect at all. Arguably, the
proposed II2IV intermediate represents just one way of
distributing electrons about a metal cluster in which
intramolecular electron transfer is expected to be rapid.
Designation of the IIjIV oxidation state is favored for two
reasons. First, the proposed hydroxy-alkoxy coordination would
favor the Ru(IV) state, assuming the carboxylate bridges remain
bound as in the starting complex. Second, protonation of the
central oxygen atom to form a hydroxy ligand removes a
potentially important pathway for intramolecular electron
transfer which could serve to stabilize the disporportionate
oxidation states.
Unlike the Ru(BPI) complexes studied by Gagne and Marks (Fig
5-6) or the ruthenium araraine complex examined by Tovrog (Fig 5-
7), the addition of base is not necessary to accomplish alcohol
oxidations with the ruthenium caraboxylate trimers. The two
former systems require proton acceptors to efffect ruthenium
disproportionation while the latter does not, an observation that
is consistent with facile electron transfer within the trimer
complex

183
Given protonation of the central oxygen atom, the hydroxy-
alkoxy intermediate which results is shown in the lower right of
the cycle. Here, coordinated alkoxide may undergo 6-hydride
elimination to form metal hydride and aldehyde (ketone). Proton
transfer from the hydride to the hydroxide would accomplish a two
electron reduction with production of one mole of water to give
the (11)^ species in the lower left of the cycle. The reactions
of alcohols with platinum metal species in the presence of base
offer a convenient route to metal hydride complexesj0^
Decomposition of metal alkoxides and metal amides to give metal
hydrides and oxidized ligands is well founded. 163*164 Dobson and
Robinson have reported NMR characterization of a metal hydride
intermediate in the dehydrogenation of alcohols catalyzed by
ruthenium and osmium complexes [MCOCOR^CCOXPPh^^] and attribute
its formation to 0-hydride elimination of coordinated
alkoxide.
The decomposition of the alkoxyrutheniura intermediate may also
follow a pathway similar to that for the decomposition of chromate
esters which is the key step in carbonyl compound formation from
the reaction of alcohol with a Cr(VI) oxo reagent1^ (Eqn 5-10).
Rx 0 0
\ /<-> //
R^CHOH + CrOo > C r Cr > R5C0 + Cr(IV) (5-10)
3 > \
R H 0 OH
A concerted intermediate is unlikely, however, given the trans
positions of the -OH and -OR ligands coordinated to Ru(IV)
(assuming that the carboxylate ligands remain bridging). A two
step reaction in which loss of the hydrogen atom as a proton,

184
followed by protonation of the hydroxy ligand to form water, is
more likely. The (II)^ species which is generated has been
isolated and characterized by Spencer and Wilkinson.^1*
Oxidation with single oxygen atom insertion is rapid in solution
in the presence of oxygen or pyridine-N-oxide, generating the
(III)2(II) species. In the proposed catalytic cycle, hydrogen
peroxide may also accomplish the oxidation step a3 shown by Fig
5-14 for the experiment with The observed oxygen pressure
dependence of the reaction rate (0.25 order in O2) suggests the
oxidation step may be rate determining or occur at a rate
competitive with the disproportionation step. In the context of
the proposed mechanism, support for competitive rates comes from
1) the observed acid/base dependence of the rate, all other
things being equal (including O2 pressure) and 2) observed oxygen
pressure dependence, all other things being equal (including no
added acid or base). In the proposed cycle, hydrogen peroxide is
not suggested as an intermediate reduced species as has been
proposed for the oxidative dehydrogenation of alcohols catalyzed
by other noble metal coraplexesJ^-148,150 Instead, oxygen is
suggested to be directly inserted into the (II)^ complex.
Similarly, hydrogen peroxide may accomplish the complex oxidation
step as suggested by the experiment with this oxidant. The rapid
production of acetone when hydrogen peroxide was included (Fig 5-
14) again suggests the oxygen atom insertion step as rate
determining.
Though no evidence is offered which rules out intermediate
formation of hydrogen peroxide, given the possibile intermediacy
of hydrogen peroxide in these oxidations and its utility as a

185
primary oxidant in a host of organic reactions, some attempts were
made to intercept any generated hydrogen peroxide using Fenton’s
reagent. In Fenton chemistry the combination of ferrous salts and
hydrogen peroxide in acidic media effectively oxidizes a wide
range of organic substrates via hydroxyl radical
intermediates. u Hydrogen abstraction from alcohols using this
reagent gives carbonyl, diol and coupling products. In the
following experiments the effect of Fenton's reagent in
conjunction with ruthenium catalyst was investigated towards
productdistribution and reaction rate.
Incorporation of Fe(Cl0 n) 2^ 2^ 6 and HCIO^ gave erratic
results as seen for identical runs (A,B) in Table 5-4 and Fig 5-
21. Though acetone production was initially accelerated, activity
quickly dropped off as a portion of the reaction solute
precipitated from solution. All experiments with Fe(II) gave an
orange brown solid within the first hour, presumed to be Fe2C>2.
Only when both Fe(II) and H+ were omitted was the reaction rate
constant over 24 hours. In all reactions acetone was the only
major product though transient production of propene (confirmed by
GCMS) was observed but not quantified when either Fe(II) or H+
were included. Stoichiometric oxidation of isopropanol by HCIOjj
cannot by itself account for accelerated acetone production as
shown by (F). Instead, a synergistic effect is observed for added
Fenton's reagent since neither RuCllD^ nor Fe(II), H+ alone can
match the reaction rate when both are used. The Fe(II) salt,
however, does not appear stable in this environment, surviving
only a few turnovers before precipitating. The least effective
system in this series was Fenton's reagent itself, presumably

186
Table 5-4. Isopropanol oxidation catalyzed by Fenton’s reagent
and Ru30(prop)g(H20)3prop.a
Ru(III)3b
Fe(II)c
H+ d
(mmol)
(mmol)
(mmol)
product(s)
TO/241
0.036
0.98
4.8
acetone
propene
283
0.035
1.00
4.8
acetone
propene
262
0.035
—
—
acetone
204
0.037
1 .02
—
acetone
propene
147
0.035
—
4.8
acetone
propene
64
0.99
4.8
acetone
propene
(32)f
a) Oxidation conducted in 50 ml isopropanol under 3 atra 0o at
65°. 2
b) Ru^OCprop^O^O^prop.
c) Fe(C104)2(H20)6.
d) 60$ HC104.
e) Moles acetone produced per mole Ru^ per 24 hours.
f) Based on same concentration of Ru^ in above reactions.

187

188
because no peroxide source was available. Because ruthenium
trimer eventually outpaces the other systems due to irreversible
precipitation of the iron species, this strategy was dropped as a
means for utilizing generated hydrogen peroxide.
C. Conclusion
In a unique demonstration of intramolecular synergy, the
trinuclear ruthenium carboxylates [Ru^OCO-jCRJgCO^]11* (R = CH^,
CgH^; L = H20, PPh^; n = 0, 1) have been shown to be effective
catalysts for the oxidative dehydrogenation of primary and
secondary alcohols. The four-electron reduction of dioxygen forms
two moles of water and two moles of aldehyde (ketone). Selective
formation of carbonyl product is accomplished in high yield with
no evidence of radical intermediates. Acid/base effects implicate
redox disproportionation in the catalytic cycle and a mechanistic
proposal accounts for this. The reaction is suggested to proceed
by disproportionation followed by 8-hydride elimination of
coordinated alkoxide; protonation of the central oxygen atom
produces water and oxygen reinsertion is accomplished with oxygen
or hydrogen peroxide. No evidence for intermediate hydrogen
peroxide was observed. High activity and the ability to operate
under neutral and acidic pH reflects the ease of accessing a
variety of oxidation states and the rapid intramolecular electron
transfer observed for complexes of the basic iron acetate
structure.
D. Experimental
1. Reagents and Equipment
Unless otherwise stated, all reagents used were reagent grade

189
or better, used without further purification. Manipulations
involving air sensitive and hygroscopic compounds were performed
under house nitrogen using a VAC dry box or standard Schlenk
techniques. Alcohols were stored over molecular sieves and passed
through a column of neutral alumina prior to use to remove
peroxide impurities. t-Butanol was recrystallized twice to
remove n-butanol impurities. Ferrothiocyanate reagent was
prepared according to the "Chemists Companion", Method 1.1^1
Ruthenium butyrate chloride (RugCbut^Cl) was obtained as a gift
from Josh TelserJ-^
GC analysis was performed on a Varian 3700 instrument fitted
with packed 8’ columns and dual flame ionization (FID) and thermal
conductivity (TCD) detectors. Column packings used were 51 DEGA
on Chromasorb P (organics) and Porapak T (water). Though not
used, 0.2Í Carbowax 1500 on Carbopak C proved a slightly superior
support for separation of alcohols and aldehydes. Peak areas and
retention times were automatically calculated with the
chromatogram using an HP 3390A integrating recorder. GCMS was
performed on a service basis by R. King at the University of
Florida mass spectroscopy facility using a DEGS packed column and
electron impact. Proton NMR spectra were collected on a JEOL-
FX100 Fourier transform instrument at 100 MHz. Electronic spectra
were performed on a PE 331 spectrophotometer and were background
corrected in all cases. IR spectra were collected on a Nicolet
5DXB FT spectrometer. Elemental analyses were performed on a
service basis by the University of Florida microanalytical
laboratory

190
2. Syntheses
Trisaquohexaprionato-u^-oxotrirutheniumCIII)propionate,
Ru30(prop)8(H20)3prop (A). A slightly modified procedure from
that reported was usedJ30,131 ■¡•0 a mixture of 50 ml ethanol
(200°) and 50 ml propionic acid was added 1.2 g sodium hydroxide.
Upon dissolution, 2g "RuCl^C^O)^" was added and the solution
refluxed under nitrogen for 4 hours during which time the orange
brown solution changed to green/black. The solution was cooled to
-78° for 3 hr, filtered and vacuum dried at 50° for 12 hours to
give crude oxotriruthenium propionate. Purification can be
accomplished by repeated (ca. four to five) cooling of MeOH/EtgO
solutions until no more unwanted Nacl or NaOAc precipitated.
Best results were obtained by liquid chromatography of MeOH
solutions of the crude propionate on a 4* column of Sephadex LH-
20 gel. The middle (deep blue-green) fraction was
rechromatographed, stripped of MeOH and dried (vac., 50°) 12 hr.
The final product is quite hygroscopic. The UV-VIS, NMR and IR
of oxotriruthenium propionate prepared in this manner agreed with
literature reports. Elemental analysis consistently gave trace
nitrogen despite all attempts at purification. Calculated for
Ru3C21H41018: C, 28.15; H, 4.67: N, 0.0. Found: C, 27.34; H,
4.32; N, 0.44.
Triaquohexaacetate-Ug-oxotrirutheniumdll) acetate,
Ru30(0Ac)g(H20)30Ac (B). The crude acetate was prepared as
reported.^30,131 -p0 a mixture of 50 ml ethanol (200°) and 50 ml
glacial acid was added 2 g "RuCl3(H20)3" and 2.4 g anhydrous
sodium acetate. Following reflux under nitrogen for 4 hr, the
solution was cooled to -20° for several hours, filtered, stripped

191
of volatiles and dried (vac, 50°) 12 hr. The crude acetate was
purified as above, best results again being obtained by Sephadex
chromatography. Satisfactory UV-VIS, NMR and IR characterization
was obtained. Calculated for Ru^C1 4H27O18: 21.38; H, 3.46;
N, 0.0. Found: C, 20.02; H, 3.36; N, 0.41.
Triaguphexapropionato-yp-oxotriruthenium (IID^CII),
Ru^OCprop^Ci^O^CC). The mixed valence propionate was prepared
as reported.^ A pressure bottle was charged with 2 g crude
(A), 30 ml water, 5 mg PtÜ2 and a stir bar. After flushing with
hydrogen, the reactor was pressurized to 40 psi hydrogen and
stirred 4 hr at 25°, throwing down a light olive solid. In a
glove bag under nitrogen, the reaction product was centrifuged
and washed with 4 times or until washing gave a clear
solution. The air sensitive complex was dried under vacuum (25°,
24 hr) and stored in the glove box. The mixed valence propionate
prepared in this manner gave satisfactory IR and UV-VIS spectra.
Tris(triphenylphosphine)hexapropionato-U2-°xotriruthe-
nium(111)2(11)» Ru^CKpr op) ^ ( P Ph^ )^( D). The mixed valence
propionate was prepared as reported, using a modified
procedure.11^ To a mixture of 0.72 g crude A and 0.50 g PPh^
under nitrogen was added 5 ml Ng saturated MeOH and the solution
stirred 20 hr. After storing the solution at 20 several days
the precipitated solid was collected, washed with ether,
dissolved in 8 ml benzene, filtered and EtOH added to bring the
volume to 50 ml. The solution was stored at ”42 2 days. The
fine green needles which formed were rinsed with cold EtoH and
dried (25 , vac) over silica gel 36 hr. Satisfactory UV-VIS and

192
IR spectra were obtained. Calculated for 3P3! C,
55.99; H, 4.89; N, 0.0. Found: C, 56.83; H, 5.05; N, 0.0.
3^ Oxidation Procedure
All pressurized reactions were conducted in a Parr 250 ral
glass pressure bottle fitted with a brass pressure head and
enclosed in a metal cage. The pressure head was equipped with a
gauge, relief valve, sampling valve, and connected directly to an
oxygen tank. A complete discription of the reactor and its use
has been reported.^7 The reactor was suspended in an oil bath
above a stir plate. A motor driven stirrer agitated the oil bath
and temperature was maintained by using a heating coil and
feedback circuit temperature controller which kept the temperature
within _+ 2° of the set temperature. A 1 ml Hamilton gas-tight
syringe with Leur-tip valve and 8" needle was used for
withdrawing solution aliquots from the reaction under pressure.
Sampling was accomplished as follows: The pressure head sampling
valve was opened and the needle (Leur valve closed) was inserted
through the pressure seal septum and through the valve, into the
reaction solution. The Leur valve was opened, about 0.2 ml
solution withdrawn and the Leur valve closed. The needle was
withdrawn and the pressure head valve closed. The collected
sample was immediately analyzed by GC and then stored in the
freezer until no longer needed.
In a typical experiment, an oven-dried pressure bottle was
charged with 50 ml alcohol and 1 ml 2 hexanone standard, the
solution incubated in the oil bath for 30 min, and the solution
sampled just before adding catalyst. A weighed amount of catalyst

193
(30-40 rag) was added and the weigh paper washed down with the warn
alcohol solution. The pressure head was immediately secured and
five purges performed by pressurizing to 40 psi with dioxygen and
exhausting to ambient pressure. The reactor was again pressurized
to 40 psi, and the inlet valve closed. The stir bar was set in
motion and this was considered the start of the reaction. The
pressure was recorded at the start and at every sampling.
Attempts to float the catalyst in plastic caps and initiate the
reaction by turning on the stir plate were unsuccessful as the
oxidation products decomposed the plastic caps used. With the
procedure used here the time elapsed between catalyst addition
and "start" of reaction was about 5 min, which is negligible over
a 24 hr reaction period.
4. Calculations
Calibration plots were constructed for determining yields of
oxidation products. Standard solutions of 1 ml standard and
various concentrations of aldehyde (ketone) in 50 ml alcohol were
used to construct a plot of mole ratio vs peak area ratio. The
integrated peak ratios from sample injections could thus be used
to obtain the molar concentration of ketone (aldehyde). Sampling
of the alcohol solution prior to adding the catalyst sometimes—but
not consistently—showed trace quantities of aldehyde (ketone)
when analyzed by GC. When this was the case, product yield was
corrected for this starting concentration. Turnovers were
calculated as mole of product per mole of Ru^ catalyst per a
given time period

CHAPTER VI
SUMMARY AND CONCLUSIONS
Several examples of the monitoring of bonding effects across
metal-metal bonds and the unique reactivity introduced by metal
atoms in close proximity have been explored. Ligand exchange
reactions of C^ÍC^CCF^)^ [(CH2CH2)2°]2 followed by electronic
spectroscopy demonstrate that donor ligands decrease the d
orbital overlap between the metal centers. Calorimetric
titrations demonstrate a lower enthalpy for second base exchange
as a result of transmission of donor effects from one metal
center to the next. The effect is less pronounced than for the
previously studied Rl^but)^ and Mo2(pfb)ij systems. Correlation
analysis of the calorimetric data reveals the Cr(II) centers are
significant Lewis acids, interacting with donor molecules in
primarily an electrostatic sense. Magnetic susceptibility
measurements on a range of ligand exchange species Cr2(tfa)^L2
demonstrate the rise in magnetic moment with donor strength and
support a direct exchange pathway for spin pairing.
EPR spectra of a range of base coordinated complexes B
Rf^Cp fb) ^(ni tro xy 1) demonstrated a weakening of the Rh-
(nitroxyl) bond with increasing donor strength whereas FTIR
measurements of B Rl^Cpfb^CO species demonstrated a
strengthening of the Rh-CO bond with increasing donor strength.
Correlation of the spectral pararaeterrs (g-value, carbonyl
194

195
stretching frequency) revealed the method dependence of
monitoring base binding effects. H (calorimetry), g (EPR) and
co (IR) all show different electrostatic/covalent responses to
base coordination across the Rh-Rh bond.
Attempts to extend these studies of the coordination of
transmission effects to neutral mixed-valence trimeric
carboxylates of formula M^OCC^CR^L^ (M = Fe, Co, Ru) focused on
the ligand dependence of their electronic spectra. For M = Fe,
the spectra were dominated obscuring charge transfer bands. For
M = Co, the product carboxylates could not be distinguished from
similar proposed dimeric species. For M = Ru, an envelope of
long wavelength transitions is red shifted with stronger donors;
these transitions have previously been suggested to be
responsible for intramolecular electron transfer.
In a unique demonstration of intramolecular synergy, the
trinuclear ruthenium carboxylates [Ru^OCC^CR^L^]n+ (R = CH^,
C2H^; L = H20, PPh^; n = 0,1) were shown to be effective
catalysts for the oxidative dehydrogenation of primary and
secondary alcohols. The four-electron reduction of dioxygen
forms two moles of water and two moles of aldehyde (ketone).
Selective formation of carbonyl product is accomplished in high
yield with no evidence of radical intermediates. Acid/base
effects implicate redox disproportionation in the catalytic cycle
and a mechanistic proposal accounts for this. The reaction is
suggested to proceed by disproportionation followed by -hydride
elimination of coordinated alkoxide; protonation of the central
oxygen atom produces water and oxygen reinsertion is accomplished
with oxygen or hydrogen peroxide. No evidence for intermediate

196
hydrogen peroxide was observed. High activity and the ability to
operate under neutral and acidic pH reflects the ease of
accessing a variety of oxidation states and the rapid
intramolecular electron transfer observed for complexes of the
basic iron acetate structure.

APPENDIX I
SPECTRAL AND CALORIMETRIC DATA

Cr^ (t fa) (Et ^0) 2 Et^POjj
Spectra
(Crp)
(Base)
0.008226
0
0.008216
0.0007348
0.008205
0.001468
0.008191
0.002489
0.008171
0.003946
0.008151
0.005393
0.008124
0.007267
0.008114
0.01016
0.008104
0.01305
0.008090
0.01707
0.008071
0.02281
0.008051
0.02851
0.008025
0.03588
e
0
e
1
£ 2
A5 10
A545
0.8600
1.088
0.8413
1.070
0.8113
1.055
0.7850
1.040
0.7513
1.020
0.7125
1.000
0.6663
0.9713
0.5925
0.9325
0.5363
0.9050
0.4925
0.8925
0.4925
0.8925
0.4925
0.8925
0.4938
0.8913
104.5
132.3
75.6
114.8
60.4
110.0
A605
A625
0.7650
0.5825
0.7663
0.5913
0.7825
0.6063
0.7975
0.6250
0.7884
0.6525
0.8450
0.6800
0.8725
0.7163
0.9325
0.7850
0.9850
0.8438
1.078
0.9213
1.078
0.9425
1.078
0.9425
1.068
0.9375
93.0
70.8
108.7
90.0
133.5
117.1
198

199
Cr2(tfa)4(Et20) + DMCA
Spectra
<<"N
o
||¡0
I'-'
(Base)
A520
A570 Ai
600 A650
0.008679
—
1.014
1.091
0.8490
0.4050
0.008584
0.001311
0.9840
1.073
0.8490
0.4155
0.008490
0.002593
0.9375
1.050
0.8490
0.4275
0.008310
0.005077
0.8610
1.022
0.8490
0.4455
0.008274
0.006551
0.8250
1.008
0.8595
0.4620
0.008214
0.009100
0.7605
0.9870
0.8700
0.4920
0.008111
0.01337
0.6960
0.9690
0.8835
0.5205
0.007946
0.02029
0.6660
0.9690
0.8955
0.5385
0.007922
0.02631
0.6195
0.9210
0.8505
0.5175
0.007874
0.03821
0.5625
0.8415
0.8415
0.4845
0.007780
0.06161
0.5385
0.7785
0.7185
0.4515
0.007628
0.09939
0.5130
0.7305
0.6720
0.4290
0.007411
0.1534
e 0
£ 1
e 2
0.4980
116.8 125.7 97.
85.9 118.7 108
83.8 121.9 112
8 46.7
.8 63.2
.7 67.8

200
Cr2(tfa)4(Et20)2 + DMTF
Spectra
(Cr2)
(Base)
A520
A570
A600
0.007925
—
0.9480
0.9885
0.7710
0.007825
0.001284
0.7995
0.8730
0.7140
0.007727
0.002535
0.6945
0.7905
0.6705
0.007538
0.004946
0.6375
0.7845
0.6390
0.007502
0.006947
0.6255
0.8055
0.7290
0.007430
0.01089
0.6105
0.8415
0.7830
0.007324
0.01667
0.6255
0.8370
0.8190
0.007155
0.02593
0.6270
0.8820
0.8340
0.007097
0.03269
0.6405
0.9045
0.8535
0.007091
0.04605
0.6570
0.9270
0.8760
0.006977
0.07436
0.6780
0.9360
0.8760
0.006823
1.163
0.6780
0.9405
0.8805
£ 0
124.2
129.5
101.1
e 1
109.3
135.2
114.2
e 2
88.4
138.2
126.2

201
Cr2(tfa)4(Et20)2 + DMSO
Spectra
(Cr2)
(Base)
A510
A550
A580
A620
0.00486
—
0.572
0.644
0.544
0.356
0.00480
0.000519
0.556
0.636
0.558
0.364
0.00474
0.00103
0.530
0.632
0.564
0.380
0.00464
0.00201
0.496
0.614
0.564
0.392
0.00462
0.00381
0.426
0.590
0.572
0.420
0.00459
0.00631
0.362
0.564
0.590
0.474
0.00453
0.0108
0.316
0.548
0.612
0.524
0.00447
0.0572
0.322
0.540
0.600
0.510
0.00447
0.0632
0.324
0.524
0.580
0.494
0.00444
0.0746
0.326
0.490
0.544
0.466
0.00438
0.0991
0.326
0.426
0.474
0.416
0.00429
0.135
0.304
0.330
0.368
0.336
e0
117.9
132.7
114.2
73.4
e1
79.0
125.0
128.6
99.7
e2
69.7
121.0
135.2
115.9

202
Cr2(tfa)2J(Et20)2 + Me^PO
Spectra
(Cr2)
(Base)
A505
A545
A605
A640
0.01164
0
1.350
1.818
1.284
0.774
0.01160
0.000714
1.316
1.800
1.304
0.798
0.01156
0.001397
1.280
1.782
1.322
0.824
0.01150
0.002317
1.240
1.756
1.336
0.844
0.01147
0.003251
1.200
1.734
1.358
0.874
0.01144
0.004179
1.154
1.706
1.382
0.904
0.01140
0.005562
1.094
1.664
1.410
0.946
0.01136
0.006937
1.038
1.624
1.438
0.988
0.01129
0.009202
0.942
1.548
1.438
0.988
0.01128
0.01064
0.852
1.478
1.524
1.128
0.01126
0.01207
0.762
1.404
1.568
1.204
0.01125
0.01350
0.674
1.334
1.604
1.274
0.01125
0.01635
0.514
1.184
1.676
1.420
0.01119
0.01918
0.474
1.118
1.646
1.412
0.01117
0.02200
0.454
1.022
1.512
1.314
e0
116.0
156.2
110.3
66.5
e 1
74.1
137.6
141.0
101.9
e 2
24.3
81.4
154.0
142.1

203
Ct2(tfa)4(Et20)2 + DMF
Spectra
(Cr0)
(Base)
A505
A545
A625
A600
0.008278
—
0.8232
1.097
0.5844
0.8184
0.008267
0.001038
0.7908
1.085
0.6024
0.834
0.008247
0.003106
0.7343
1.069
0.6396
0.864
0.008227
0.005165
0.7092
1.054
0.6672
0.888
0.008196
0.008233
0.6540
1.030
0.7104
0.9228
0.008166
0.01128
0.6060
0.1006
0.7488
0.954
0.008136
0.01430
0.5748
0.9900
0.7932
0.9828
0.008096
0.01830
0.5844
1.008
0.8088
1.0068
0.008056
0.02226
0.5904
1.013
0.8088
1.0092
0.008017
0.02617
0.5904
1.013
0.8008
1.008
0.007969
0.03102
0.5904
1.006
0.8016
1.0008
0.007874
0.04055
0.5832
0.9888
0.7944
0.9876
£
0
99.4
132.5
70.6
98.9
£
1
78.5
125.2
87.7
113.5
£
2
73.3
125.7
100.4
125.3

204
Cr2(tfa)4(Et20)2 + HMPA
Spectra
Cr2
(Base)
A520
A588
A620
A645
0.008830
—
1.034
0.9763
0.6613
0.4138
0.008818
0.0008612
0.9975
0.9763
0.6888
9.4438
0.008808
0.001720
0.9663
0.9763
0.7050
0.4650
0.008793
0.002918
0.9038
0.9763
0.7350
0.5413
0.008771
0.004624
0.8113
0.9663
0.7775
0.5963
0.008749
0.006319
0.7225
0.9531
0.8100
0.6475
0.008721
0.008514
0.5975
0.9325
0.8613
0.7175
0.008710
0.01269
0.4338
0.9125
0.9350
0.8225
0.008700
0.06185
0.3475
0.9563
1.049
0.9538
0.008685
0.02276
0.3413
0.9075
0.9925
0.9038
0.008657
0.03350
0.3225
0.7938
0.8613
0.7850
0.008615
0.04988
0.2963
0.6388
0.6763
0.6163
0.008572
0.06609
0.2838
0.5400
0.5638
0.5150
0.008511
0.09012
0.2750
0.4538
0.4638
0.4238
0.008409
0.1294
0.2750
0.4000
0.4038
0.3725
e 0
118.1
110.6
74.9
46.9
E1
68.6
112.0
103.3
94.3
e2
39.9
110.0
120.6
109.6

205
Cr2(tfa)4(Et20)2 + (MeO)3P
Calorimetry
(Cr2)
(Base)
Vol (ml)
Heat <
(cal)
0.003958
0.0007527
50.0463
0.035
0.003955
0.001374
50.0846
0.025
0.003952
0.002043
50.1259
0.025
0.0039^9
0.002705
50.1668
0.015
0.003942
0.004066
50.2511
0.057
0.003936
0.005388
50.3333
0.049
0.003929
0.006873
50.4260
0.025
0.003916
0.009480
50.5894
0.040
0.003896
0.01346
50.8407
0.014
0.003885
0.01590
50.9963
0
K1 1 X
108
K2 2 X
103
-AH,;,
0.792 MSD
0.08 CSD
0.06
kcal mol-1
“AH2:1
0.725 MSD
0.14 CSD
0.10
kcal mol-^

Cr2(tfa)4(Et20)2 + Me3PO
Calorimetry
206
(Cr2)
0.002938
0.002935
0.002933
0.002931
0.002927
0.002922
0.002917
0.002914
0.002909
0.002905
0.002900
0.002986
0.002888
1 X 107
K2 2 X 104
-AH1;1 5.95
-AH2:1 6.55
(Base)
0.0003646
0.0006325
0.0009406
0.001245
0.001871
0.002481
0.003165
0.003699
0.004366
0.004963
0.005582
0.006200
0.007327
MSD 0.35
MSD 0.57
Vol (ml)
55.0463
55.0846
55.1259
55.1668
55.2511
55.3333
55.4260
55.4985
55.5894
55.6709
55.7556
55.8407
55.9963
CSD 0.26
CSD 0.43
Heat (cal)
0.1118
0.0907
0.1082
0.0995
0.2213
0.2120
0.2272
0.1917
0.2113
0.2084
0.2149
0.0951
0.0356
kcal mol-1
kcal mol~1

207
Cr2(tfa)Jj(Et20)2 + DMF
Calorimetry
(Base) Vol (ml) Heat (cal)
(Cr2)
0.002224
0.002222
0.002219
0.002215
0.002212
0.002208
0.002205
0.002202
0.002199
0.002195
0.002192
K1 1 X 108
K2 1 X 104
-AH1;1 4.52
** AH2.i 3.01
0.0002620
0.0004784
0.0009419
0.001416
0.001876
0.002394
0.002798
0.003303
0.003754
0.004222
0.004690
MSD 0.39
MSD 0.69
55.0463
55.0846
55.1668
55.2511
55.3333
55.4260
55.4985
55.5894
55.6709
55.7556
55.8407
CSD 0.31
CSD 0.54
0.072
0.0566
0.0999
0.1056
0.1234
0.1363
0.1072
0.0648
0.0505
0.0281
0.0136
kcal mol-^
kcal mol-1

208
o
“1
ro
(tfa)
DMA
Calorimetry
(Cr2)
(Base)
Vol (ml)
Heat (cal)
0.005087
0.0007677
50.0463
0.230
0.005083
0.001402
50.0846
0.218
0.005079
0.002084
50.1259
0.208
0.005075
0.002759
50.1668
0.175
0.005067
0.004246
50.2511
0.333
0.005058
0.005495
50.3333
0.298
0.005049
0.007010
50.4260
0.291
0.005042
0.008191
50.4985
0.272
0.005033
0.009668
50.5984
0.213
0.005025
0.01099
50.6709
0.037
0.005016
0.01235
50.7556
0
0.005008
0.01372
50.8407
0
K1 1 X
109
k2 1 X
105
- AH1;1
5.26
MSD
0.36
CSD
0.24
kcal mol-^
- ah2;1
3.59
MSD
0.58
CSD
0.39
kcal raol“1

209
Cr2(tfa)4(Et20)2 + HMPA
Calorimetry
(Cr2)
(Base)
Vol (ml)
Heat (cal)
0.004478
0.0005933
50.0463
0.1823
0.004474
0.001083
50.0846
0.1295
0.004471
0.001611
50.1259
0.1673
0.004467
0.002133
50.1668
0.1444
0.004460
0.003204
50.2511
0.3611
0.004452
0.004246
50.3333
0.2792
0.004444
0.005418
50.4260
0.3382
0.004438
0.006329
50.4985
0.2642
0.004430
0.007470
50.5894
0.3338
0.004423
0.008490
50.6709
0.2801
0.004415
0.009546
50.7556
0.0229
0.004408
0.001060
50.8407
0
It, 1 X
1011
k2 1 X
106
" AH1:1
5.97
MSD
0.33 CSD
0.23 kcal mol-1
“ aH2:1
5.48
MSD
0.54 CSD
0.37 kcal raol-^

210
Cr2(tfa)ij(Et20)2 + Et^POjj
Calorimetry
(Cr2)
(Base)
Vol (ml)
Heat (cal)
0.004576
0.0006024
50.0463
0.1648
0.004572
0.001100
50.0846
0.1680
0.004569
0.001635
50.1259
0.1515
0.004565
0.002165
50.1668
0.0934
0.004557
0.003253
50.2511
0.2237
0.004550
0.004311
50.3333
0.1994
0.004541
0.005501
50.4260
0.1994
0.004535
0.006416
50.4985
0.1797
0.004527
0.007586
50.5894
0.1633
0.004519
0.008621
50.6709
0.1656
0.004512
0.009693
50.7556
0.0612
0.004504
0.01077
50.8407
0.0243
K1 1 X
1010
k2 1 X
106
“ AH1:1
4.37
MSD
0.27 CSD
0.18 kcal mol“^
- ah2;1
3.05
MSD
0.43 CSD
0.30 kcal mol-1

211
Cr2(tfa)4(Et20)2 + DMCA
Calorimetry
(Cr2)
0.003129
0.003127
0.003124
0.003122
0.003117
0.003113
0.003107
0.003103
0.003098
0.003094
0.003089
0.003084
K1 1 X 108
K2 3 X 103
-AH1;1 3.75
-AH2;1 2.66
(Base)
0.0004075
0.0007439
0.001106
0.001464
0.002201
0.002917
0.003722
0.004350
0.005098
0.005838
0.006564
0.007292
MSD 0.17
MSD 0.31
Vol (ml)
55.0463
55.0846
55.1259
55.1668
55.2511
55.3333
55.4260
55.4985
55.5894
55.6709
55.7556
55.8407
CSD 0.14
CSD 0.24
Heat (cal)
0.08353
0.07097
0.06824
0.07465
0.1518
0.1507
0.1284
0.09950
0.07695
0.03856
0.03793
0.03875
kcal mol“^
kcal raol“1

212
Cr2(tfa)4(Et20)2 + DMTF
Calorimetry
(Cr2)
(Base)
0.003477
0.0004049
0.003475
0.0007394
0.003472
0.001100
0.003469
0.001456
0.003464
0.002188
0.003459
0.002901
0.003453
0.003700
0.003449
0.004323
0.003443
0.005105
0.003438
0.005802
0.003433
0.006526
0.003428
0.007249
0.003418
0.008567
Vol (ml)
Heat (cal)
55.0463
0.0437
55.0486
0.0347
55.1259
0.0352
55.1668
0.0386
55.2511
0.0758
55.3333
0.0756
55.4260
0.0616
55.4985
0.0491
55.5894
0.0429
55.6709
0.0281
55.7556
0.0271
55.8407
0.0181
55.9963
0.0392
K1 1 X 105
K2 500
-AH1;1 1.89 MSD 0.066
-AH2;1 1.80
CSD 0.060
kcal mol”^
kcal mol“1
MSD 0.13
CSD 0.12

213
Cr2(tfa)4(Et20)2 + DMSO
Calorimetry
(Cr2)
0.004578
0.004574
0.004571
0.004567
0.004560
0.004553
0.004544
0.004538
0.004530
0.004523
0.004515
0.004508
K1 1 X 109
K2 6 X 104
-AH1;1 5.58
-AH2;1 3.95
(Base)
0.0006135
0.001120
0.001654
0.002205
0.003314
0.004392
0.005603
0.006547
0.007728
0.008782
0.009875
0.01097
MSD 0.40
MSD 0.66
Vol (ml)
51.0463
51.0846
51.1259
51.1668
51.2511
51.3333
51.4260
51.4985
51.5894
51.6709
51.7556
51.8407
CSD 0.27
CSD 0.45
Heat (cal)
0.1630
0.1443
0.1421
0.1436
0.3335
0.3133
0.2550
0.2669
0.2535
0.1308
0.0013
0
kcal mol-1
kcal raol“”'

214
Cr2(tfa)4(Et20)2 + DMA
Spectra
(Cr,)
(Base)
A510
A515.5
A600
A630
0.007988
. . _
0.8496
0.9048
0.7704
0.5064
0.007978
0.0008459
0.8340
0.8868
0.7788
0.5208
0.007970
0.001521
0.8028
0.8580
0.7944
0.5425
0.007960
0.002363
0.7788
0.8328
0.8088
0.5628
0.007948
0.003371
0.7428
0.7968
0.8280
0.5904
0.007932
0.004710
0.6972
0.7524
0.8496
0.6240
0.007913
0.006377
0.6552
0.7104
0.8712
0.6564
0.007889
0.008365
0.5964
0.6504
0.9048
0.7068
0.007880
0.01255
0.4884
0.5412
0.9708
0.8052
0.007870
0.01672
0.4620
0.5136
1.001
0.8460
0.007856
0.02254
0.4620
0.5124
1.003
0.8508
0.007837
0.03083
0.4620
0.5124
0.9972
0.8460
0.007818
0.03905
0.4632
0.5124
0.9816
0.8340
0.007793
0.04971
0.4668
0.5124
0.9576
0.8184
0.007755
0.06597
0.4704
0.5124
0.9336
0.8016
0.007699
0.09007
0.4740
0.5124
0.8964
0.7740
0.007608
0.1295
0.4764
0.5124
0.8568
0.7464
0.007431
0.2055
0.4740
0.4968
0.7848
0.6936
0.007100
0.3473
0.4608
0.4692
0.6516
0.5952
e
0
106.4
113.3
96.4
63.4
e
1
89.0
94.0
110.1
101.4
e
2
57.2
62.7
128.0
108.7

APPENDIX II
OPERATION OF CALORIMETER
Initial Status
Main Power (28): Should be turned on several hours before a run;
normally left on continually.
Slow/Fast (11): Down.
Hold /Capture (12): Up, to maintain the drift compensator at a
fixed value. Down when the baseline compensator is automatically
searching for the voltage to offset the signal drift.
One /Two (7): For single cell operation, only one heat output is
needed.
One /Two (26): For single cell operation, only one signal input
is needed.
Initial Connections
Cell Thermister < > Input One (23)
Cell Heater < > Heat One (7)
Recorder Input < > Recorder Output (15)
215

216
After cell is loaded and secured above the stir plate,
initial connections are made and stirrer turned on, the
apparatus is left for several (2-4) hours while the cel 1 and
contents come to thermal equilibrium with the environment.
During this time, it is best to minimize drafts and other
temperature changes. Thermal equilibrium has been achieved when
the recorder baseline is free of curvature.
Zero Bridge and Compensator
1. Press drift compensator reset (14).
2. Turn meter level switch (11) to low. While continuing to
press integrator reset (13)» adjust coarse (20), medium (21) and
fine (22) pontentiometers to zero meter. Repeat at meter
settings medium and high.
3. Disengage integrator reset (13).
4. At this point, meter will begin to drift. Flipping
capture/hold (12) to capture will allow the instrument to search
for the proper voltage to compensate for drift. In practice, it
was found easier to activate recorder and adjust offset to bring
the pen to the middle of the page. The zero drift potentiometer
(18) was then used to produce a non sloping baseline. Curvature
can be corrected by adjusting zero rate potentiometer (17) which
acts as a first order correction to the drift. When a linear
base line has been achieved, the calorimetric experiment can be
started

217
Calibration of the Heater
The calibration circuit is a series circuit consisting of a
power supply, a standard resistor, the cell heater and a
transistor switch controlled by a quartz clock. The heat
released into the cell may be ccalculated once the heater
resistance and IR drop are known. With the monitor switch (5)
set to H, the potential across the heater (VH) is sampled at the
connectors (6) with an accurate digital multimeter. The timer
start button (1) is depressed with the timer (3) set at half the
time used in the experimental calibration heats, and the digital
reading recorded at the end of the time period. This way, the
potential is an average, sampled halfway through an experimental
calibration heat. The heater resistance (R^) is measured across
the heater when it has been isolated from the circut.
p . vh2/rh
Heater Constant (cal / sec) = £ (Joules /_ sec)
4.186 (Joules / cal)

ROTARY SWITCH
® PUSH BUTTON
POTENTIOMETER
BANANA PLUG
O) BNC
Fig A-l.
Calorimeter front panel.
218

APPENDIX III
DERIVATION OF EQUATION 2-17
The four enthalpy terms, AH^, AH^, A and A Hg are
expressed as Af^.-j terras since they define adduct formation
between a 1:1 adduct and a second donor molecule.
A H1
Et20 - Cr2 - Et20
-AH1 = -AH2;1 = -AH1;1 - kEEt20 - kfCEt20
= EAEEt20 + CACEt20 " kEEt20 " k'CEt20
A H3
Et2Ü - Cr2 - B
-AH3 = -AH2;1 = -AH1;1
= eaeb + cAcB
" kEEt20 â–  k *CEt20
“ kEBEEt20 “ k'CBCEt20
i
B - Cr2 - EtgO
-AH5 = -AH2;1 = -AH1;1 - k E2 - k'CB
= EAEEt20 + CACEt20 “ kEBEEt20 ” k'CBCEt20
219

220
B - Cr2 - B
-AH6 = -AH.,., = -ah,., - kEg
= EaEb + CACB - kEg
k-c‘
k' Cg
The solvent corrected experimental enthalpies are given by
-AH,:t -
AH1
-ah.
-ah2:, .
ah5
-AH,
AH,.
1:1
-AHi;i = AH1 - AHj = Ea (Eg - EEt20) + ca^cB “ CEt20^
2 2
+ k(EE“t2Q - EEt20Eg) + k'(CEt20 ” CEt20CB^
define Eg = EEt20 - Eg
CB = CEt20 “ CB
-AHi;i = EAEB + CACB “ kEEt20EB “k,EEt20CB
= (Ea -kEEt20)Eg + (CA - k'CEt20)Cg
!2s1
-AH2;1 = AH5 -AH6 = ^A(EB“EEt20^ + CA(CB”CEt2(P
+ k(EgEEt2Q-Eg) + k'(CgCEt2Q - Cg)
-AH2;1 = EAEg + CACg - k Eg Eg - k'CgCg
= (EA “ kEB)El + (CA - k'CB> CB

221
The relationship between the first and second heats is given by
AH2;1 - (-AH1:1) = (Ea - kE0)EB + (CA-k’CB)CB
-(EA-kEEt20)EB - (CA-k'CEt20)C
= ^EEt20-EB^kEB + ^CEt20 “ CB^k'CB
= -k (Eg)2 - k1(CB)2
-A H2;1 = -A H1;1 - k(Eg)2 - k' (Cg)2 (2-17)
where Eg = Eg - EEt20
CB = CB " CEt20
â– k m

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BIOGRAPHICAL SKETCH
Carl Joseph Bilgrien was born November 20, 1959, in Pontiac,
Michigan, where he lived nine days. He spent the next seventeen
years in Milwaukee, and then, Sheboygan Falls, Wisconsin. He
returned to Michigan for four years, graduating from Michigan
State University in 1981 with a B.S. in chemistry. After one
year of graduate studies at the University of Illinois, he moved
to the University of Florida with the Drago research group. Carl
received his Ph.D. in chemistry in 1986 and like a salmon
returning upstream, returned to Michigan where he is currently
part of the research staff of Dow Corning Corp., Midland
231

I certify that I have read this study and that in ray 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.
y
Earl E. Mu^dhlitz, Jr. O
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.
DavTcT E. Richardson
Assistant Professor of Chemistry
I certify that I have read this study and that in my opinion
it conforms to acceptable standards of scholarly presentation and
is fully adequate, in scope and quality, as a dissertation for the
degree of Doctor of Philosophy.
Harry H. filaler
Distinguished Service Professor
Emeritus 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.
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 requirments for the degree of Doctor of
Philosophy.
August, 1986
Dean, Graduate School

UNIVERSITY OF

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