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Ionization, bonding, and solvation energetics of organometallic complexes

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Title:
Ionization, bonding, and solvation energetics of organometallic complexes
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Ryan, Matthew Francis, 1965-
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English
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xiv, 203 leaves : ill. ; 29 cm.

Subjects

Subjects / Keywords:
Electronics ( jstor )
Energy value ( jstor )
Enthalpy ( jstor )
Entropy ( jstor )
Ferrocenes ( jstor )
Free energy ( jstor )
Ionization ( jstor )
Ions ( jstor )
Ligands ( jstor )
Solvation ( jstor )
Chemistry thesis Ph. D
Dissertations, Academic -- Chemistry -- UF
Metal complexes ( lcsh )
Metallocenes ( lcsh )
Organometallic chemistry ( lcsh )
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bibliography ( marcgt )
non-fiction ( marcgt )

Notes

Thesis:
Thesis (Ph. D.)--University of Florida, 1993.
Bibliography:
Includes bibliographical references (leaves 194-202).
General Note:
Typescript.
General Note:
Vita.
Statement of Responsibility:
by Matthew F. Ryan.

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University of Florida
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University of Florida
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Copyright [name of dissertation author]. Permission granted to the University of Florida to digitize, archive and distribute this item for non-profit research and educational purposes. Any reuse of this item in excess of fair use or other copyright exemptions requires permission of the copyright holder.
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IONIZATION, BONDING, AND SOLVATION
ENERGETIC OF ORGANOMETALLIC COMPLEXES


By

MATTHEW F. RYAN


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
1993




























To my Grandparents












ACKNOWLEDGEMENTS


The sacrifices I have made over the past several years cannot be counted;

perhaps they should not even be considered. But sacrifice was a major factor towards

completion of this final work--a lot of time and a lot of energy and a lot of sacrifice

all concentrated for a sole purpose. Unusual things happen to a person who focuses so

long and hard on a single goal. The memories in between the battles become clouded

and unfocused. So many memories have amassed, most are happy, but some are sad

and lonely. However, the indecision and the doubt and the pure rage are paled by the

handful of triumphs. Perchance it is these triumphs that make the journey fulfilling.

I recall the time when I felt that I needed to understand more about science and

more importantly, the world around me. Although graduate school was not the sole

source of my salvation it was a beginning: the fresh start that I needed. When I

spoke with Professor Russell Drago in the spring of 1987 (I remember the

conversation well), I told him that if I was accepted to the University of Florida I

would succeed. With the support of Professor Drago and Professor Jack Kotz of

S.U.N.Y. Oneonta, I was admitted to U.F. in the fall of 1987. I am very grateful for

their support and the trust they had in me. I believe I was given an important break.

I would like to thank Professors David E. Richardson and John R. Eyler for all

their guidance. Dr. Eyler has helped me in many capacities and I am grateful to him






for all the discussions we have had. His confidence in me and my work has been very

important to me. David Richardson, my research advisor, has helped me develop

many skills necessary to be a scientist. Although it was sometimes a struggle, for

both him and me, I learned a great deal from him. His door was always open and he

always had time for discussions with me. I will always admire his work as an

educator and a mentor.

Several people assisted in various capacities with this work. I would like to

thank Dr. Allen Siedle, Dr. Mark Burk, Professor Charles Winter, Professor Russell

Hughes, Professor Dennis Lichtenberger, Dr. San Li, Professor William Weltner, and

my friends, Dr. Md. Nazrul I. Khan and Dr. Paul Sharpe for their help.

Throughout my graduate work, I have drawn strength from several vital

sources: my family, my friends, and Debbie Simpson. My family shared with me the

most exciting and frustrating moments. My parents understood and appreciated my

efforts to "catchup" and adjust to my new surroundings. When I spoke with them, any

discouragement I had was made insignificant; I spoke with them often because their

confidence and love are inspirational. My brothers and sisters, Mark, Joe, Cassandra,

and Andrea, were there to fill in the gaps. Their love for me was manifested in many

ways. I am grateful for all the things they have done for me for they have given me

so much and asked for nothing in return. I am blessed to have such a close family.

My friends Steve and Mike Messick helped me to persevere. Through Steve

and Mike, I learned how to focus my thoughts and emotions. The jaunts to Cape






Vincent (and points north) and California were extremely therapeutic. Although the

briar patch at the Tesla concert was not efficacious, Vegas was completely necessary.

When I came to U.F., I left behind several close friends, and I have remained

close with only a few people during my time in Florida. Jack Santos (soon to be a

father) and I have known each other forever. Although I may have neglected our

friendship occasionally, Jack always understood. He is a great friend.

I am very grateful to my friend John Moore for his help in the beginning and

his encouragement until the final moment. Before I moved to Florida, John found me

an apartment and, more importantly, introduced me to the U.F. Rugby Team soon after

my arrival. Rugby, in addition to supplying me with dozens of wonderful friendships,

was a vent for frustrations like no other I can imagine.

Steve Glatter and Steve Rubin have supplied me with everlasting laughter. In

addition to their advice and refreshing perspective on things in general, I would like to

thank them for their companionship throughout the years.

Mike Naugton and I were roommates for nearly five years. I enjoyed his

company, and I am glad we have remained friends.

Before I met Debbie Simpson, I often felt that I was isolated and alone. My

family and friends are all far away and the rugby team has a specific purpose. Debbie

is a wonderful friend, her smile is contagious, and her warmth and love never fade.

Debbie has always listened to me and has helped me to find the answers. We have

been through a great deal together, and our friendship, love, and respect for each other

is constantly growing with each passing challenge.












TABLE OF CONTENTS


ACKNOWLEDGEMENTS ........................................ iii

LIST OF TABLES .............................................. ix

LIST OF FIGURES............................................... xi

ABSTRACT ................................................ xiii

CHAPTERS

1 INTRODUCTION .......................................... 1

2 ADIABATIC IONIZATION ENERGIES, BOND DISRUPTION ENTHALPIES,
AND DIFFERENTIAL SOLVATION ENERGIES OF GAS-PHASE
METALLOCENES AND METALLOCENIUM IONS ................. 13

Introduction ............................................ 13
Reevaluation of Metallocene Free Energies of Ionization Based on
Electron-Transfer Equilibrium Studies ..................... 16
Insights into the Free Energy of Ionization of Ferrocene ............. 24
Intramolecular Entropy Changes of the Ferrocene/Ferrocenium Couple ... 34
Ionization Free Energies of Ruthenocene and Osmocene ............. 43
Free Energies of Ionization of Vanadocene, Manganocene and Nickelocene 44
Substituent Effects in Ferrocene Derivative Oxidations................ 45
Heterolytic and Homolytic Metal Ligand Bond Disruption Enthalpies of
Metallocenes and Metallocenium Ions ................... .. 49
Differential Solvation Free Energies for Metallocene Redox Couples ..... 54
Conclusions ... ............................ ............ 63
Experimental Methods ..................................... 65

3 SUBSTITUENT EFFECTS IN THE GAS-PHASE AND SOLUTION
IONIZATION AND ELECTRON-ATTACHMENT ENERGIES OF
ALKYLNICKELOCENES .................................... 70






Introduction ......................................... 70
Electron-Transfer Equilibrium Studies Involving Negative and
Positive Alkylnickelocene Ions .......................... 72
Alkyl Substituent Analyses for Positive and Negative Ions and Rationalization
of the Gas-Phase Trends for the Ionization and Electron Attachment
Free Energies ...................................... 81
Solvation Energetics of Nickelocene Cations and Anions ............. 91
Bond Disruption Enthalpies for Nickelocene Anion ................. 94
Conclusions ............................................ 94
Experimental Methods ......................... ............ 96

4 GAS-PHASE AND SOLUTION OXIDATION POTENTIALS OF
RUTHENOCENE DERIVATIVES .............................. 98

Introduction ... ......................................... 98
Results of the Electron-Transfer Equilibrium Reactions ............. 101
Evaluation of the Gas-Phase Free Energies of Ionization for a Series of
Ruthenocene Derivatives ............... .... ........... 108
Attempted Correlation of Ruthenocene Ionization Free Energies
with Taft oI Parameters .............................. 116
A New Parameter Scale for Cyclopentadienyl Substituents Based on Gas-Phase
Electron-Transfer Equilibrium Studies of Ruthenocenes ........ 123
Rates of Hydrogenolysis for Methylzirconocene Cations ............ 129
Comparisons of Gas-Phase Ionization Free Energies to Solution
Electrode Oxidation Potentials ....... ...... ....... 130
Determinations of Free Energies of Ionization in Solution from
Electrochemical Oxidation Potentials ..................... 137
Differential Solvation Free Energies for Several Ruthenocene/Ruthenocenium
Couples ........................................ 138
Application of the Born Model for Estimating Solvation Energetics
for Ruthenocene Oxidation Couples ...................... 140
Conclusions ........................................... 143
Experimental Methods .................................... 145

5 GAS-PHASE IONIZATION ENERGETIC, THERMOCHEMISTRY, AND
ELECTRON-TRANSFER KINETICS OF DECAMETHYLMETALLOCENES,
CHROMOCENE, AND COBALTOCENE ....................... 148

Introduction...................................... ...... 148
Gas-Phase Electron-Transfer Equilibrium Studies .................. 151
Electrochemical Studies for Some Decamethylmetallocenes. .......... 152
Bis(benzene)Chromium as a Reference Compound for Electron-Transfer
Equilibrium Investigations ............................ 156







Free Energies of Ionization for Some Decamethylmetallocenes and
Comparison to Photoelectron Spectroscopy Results ........... 159
Free Energies of Ionization for Chromocene and Cobaltocene ......... 163
Bond Disruption Enthalpies for Chromocene and Cobaltocene ......... 164
Evaluation of the Solvation Energetics for Decamethylmetallocenes
Chromocene and Cobaltocene .......................... 166
Electron-Transfer Kinetics................................... 171
Conclusions ........................................... 174
Experimental Methods .................................... 175

6 OVERVIEW OF EXPERIMENTAL METHODS AND PROCEDURES .... 178

Fourier Transform Ion Cyclotron Resonance Mass Spectrometry ...... 178
Measurement of Equilibrium Constants ......................... 183
Temperature Dependence Studies.............................. 187
Application of FTMS for the Study of Metal Complexes ............ 188


7 SUM M ARY .............................................. 191

REFERENCES .............................................. 194

BIOGRAPHICAL SKETCH ................. ................... 203












LIST OF TABLES


Table Page

2-1 Ionization Energetics Data for Some Metallocenes ................. 23

2-2 Calculated Entropies and Integrated Heat Capacities for Ferrocene and
Ferrocenium Ion at 298, 450, and 600 K ....................... 35

2-3 Vibrational Frequencies for Various Ferrocenium Salts ............. 37

2-4 Vibrational Frequency Data for Ferrocene and Ferrocenium Cation ..... 40

2-5 Mean Bond Disruption Enthalpies for Some Metallocenes ........... 55

2-6 Auxiliary Thermochemical Data Used in Thermochemical Cycles ...... 56

2-7 Electrochemical E Data and Differential Solvation Energies for Some
Metallocene Cp2M / Couples ............................... 58

3-1 Free Energies of Ionization and Electron Attachment ............... 79

3-2 Alkyl Substituent Parameters for Some Alkylnickelocene Complexes and
Free Energies for Reactions 3-4, 3-5, and 3-6 ................... 84

4-1 Values of AGi for Ruthenocene Derivatives and Other Data ........ 104

4-2 Ligand y* and y Parameters ................................ 105

4-3 Substituent Parameters for Selected Cyclopentadienyl Derivatives ..... 122

4-4 Electrode Potentials and Differential Solvation Free Energies for Some
Ruthenocene Derivatives ................................ 135

5-1 Ionization Energetics Data for Some Metallocenes and
Decamethylmetallocenes .................................. 154






Table Page

5-2 Electrochemical E, Data and Differential Solvation Energies of Some
Cp*2M+0 and Cp2M+O Couples ............................ 155

5-3 Average Bond Disruption Enthalpies for Chromocene and Cobaltocene 165

5-4 Electron-Transfer Kinetics for Some Metallocenes,
L2Ma + L2Mb L2Ma + L2Mb+ .......................... 172












LIST OF FIGURES


Figure Page

1-1 Potential well diagrams demonstrating vertical and adiabatic ionization
process. ................................................ 7

2-1 Log plot for electron-transfer reaction of Cp2Fe +DET' = DET + Cp2Fe+ 17

2-2 Electron-transfer equilibrium ladder for some metallocenes ............ 22

2-3 Van't Hoff plots for selected Metallocene electron-transfer
equilibrium couples ...................................... 25

2-4 Plot of AGi values (kcal mol"1) versus alkyl Taft parameters ......... 48

2-5 Thermochemical cycles used to determine bond disruption enthalpies and
differential free solvation energies for metallocenes ................ 52

2-6 Plot demonstrating periodic trends of ionization energies for the first transition
row metallocenes ........................................ 60

2-7 Structure of 18 e-1 vanadocenium complex ...................... 61

3-1 Molecular orbital diagrams for nickelocene anion, nickelocene, and nickelocene
cation ............................................... 73

3-2 Electron-transfer equilibrium ladder for ionizations for several alkylnickelocene
com plexes ......................................... ... 75

3-3 Electron-transfer equilibrium ladder for electron attachments for
alkylnickelocene complexes ................................ 76

3-4 Plots of AGi0 and AGa0 data versus Taft (o0I) parameters ........... 83

3-5 Plot of AGi and AGa0 data for some alkylnickelocene complexes versus 1(oy)
parameters ......... ..... ............................ 86






Figure Page

3-6 Plot of AG360 values (kcal mol-1) derived from equations 3-4, 3-5, and 3-6
versus (oij) values ...................................... 90

4-1 Electron-transfer equilibrium ladder for several ruthenocenes derivatives for the
process M M++e" .................................. 103

4-2 Plot of Ru 3d binding energies from reference 91 versus ETE AGi values for
several Cp*Ru-L complexes and ruthenocene ................... 115

4-3 Correlation of AGi values for several ruthenocene derivatives with Taft oI
param eters ............................................ 118

4-4 Plot of AGi values versus 7 parameters ................... ... 125

4-5 Plot of AGi0 values versus I(y) parameters .................... 128

4-6 Gas-phase rates of hydrogenolysis for several methylzirconocene cation
com plexes ............................................ 131

4-7 Structures of Cp*2Ru and ferrocene .................. ....... 142

5-1 Electron-transfer equilibrium ladder for several decamethylmetallocenes 153

5-2 High resolution He (I) photoelectron spectrum of bis(benzene)chromium in the
valence ionization region ................................. 158

5-3 Plot of AGi values for alkylferrocene and nickelocene complexes and versus
alkyl Taft oI parameters ................................. 161

5-4 Plot demonstrating periodic trends of ionization energies for the first transition
row decamethylmetallocenes ............................... 168

6-1 Orthorhombic ion trap used in a Fourier transform ion cyclotron resonance
mass spectrometer ...................................... 180

6-2 Schematic representation of a Fourier transform ion cyclotron resonance mass
spectrometer .......................................... 185

6-3 Van't Hoff plot of the CO/Kr electron-transfer equilibrium reaction ... 189











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

IONIZATION, BONDING, AND SOLVATION ENERGETIC
OF ORGANOMETALLIC COMPLEXES

By

Matthew F. Ryan

May, 1993

Chairperson: David E. Richardson
Major Department: Chemistry

Free energies of ionization (AGio) for Cp2M complexes (Cp = 15-

cyclopentadienyl; M = V, Cr, Mn, Fe, Co, Ni, Ru, Os) and Cp*2M complexes (Cp* =

r15-pentamethylcyclopentadienyl; M = Mn, Fe, Ni, Ru, Os) have been determined from

electron-transfer equilibrium methods by using Fourier transform ion cyclotron

resonance mass spectrometry. The AGi values for ferrocene derivatives,

alkylnickelocene derivatives and various ruthenocene complexes have also been

investigated through ETE studies.

Temperature dependence studies involving ETE of ferrocene with N,N'-

diethyltoluidine lead to values for the ionization enthalpy and entropy for ferrocene.

Experimental results of ASio(Cp2Fe) are compared to results from statistical

mechanical analyses. Thermochemical cycles are used to derive estimates of

heterolytic and homolytic bond M-Cp enthalpies for the first transition row

xiii





metallocenes and to derive estimates for the differential solvation free energies,

AAGsolv,, for the Cp2M+O couples. With the exception of Cp2V+0, first transition

row Cp2M+/0 couples are estimated to have AAG0solv values of -38 3 kcal mol1.

The average AAGsolv for the Cp*2M+/O couples is -26 3 kcal mol-1.

The AGi values and electron attachment free energies (AGa) for a series of

alkylnickelocene complexes (RCp)(R'Cp)Ni have been determined. The ionization

energies follow the expected trends (increased alkylation yields a decrease in AGi) but

the AGa values become more negative for R = Et and t-Bu and more positive for R =

Me. Several parameter schemes are used to interpret the AGio and AGaO data. Values

of AAGsolv, for the Cp2Ni+"/0 and (t-butylCp)2Ni+"l couples are discussed. From the

Cp2Ni+/0/ data, an average AAGsolvo value is estimated.

The AGi values and E estimates for series of ruthenocene derivatives have

been determined. A parameter scale that correlates electronic effects for Cp ligands,

rather than individual substituents, has been developed based on AGio values of the

ruthenocene complexes. The utility of the parameters for predicting reactivity of metal

complexes bearing Cp ligands is considered.











CHAPTER 1
INTRODUCTION


Thermochemical information is of primary importance to the understanding of

chemical reactivity. Reaction pathways and mechanisms can be better understood if

salient thermochemical data pertaining to reactant species or products exist.

Thermodynamics can inform us if a chemical reaction will proceed and can tell us the

efficiency of the reaction. Although information concerning the velocity of a reaction

can not be readily assessed from a thermodynamic analysis, theories of chemical

kinetics are based on the laws of thermodynamics.' Thus, the utility of

thermochemistry as a tool to both predict and comprehend chemical transformations is

essential to chemistry.

Organic chemistry has benefitted enormously from the depth and breadth of

available thermochemical data for simple and complex organic systems.2'3

Furthermore, the thermochemical foundation of organic chemistry allows complex

reaction mechanisms to be understood and reaction efficiencies to be maximized. The

utility of thermochemistry to predict reactivity based on known quantities is well

documented in organic chemistry.2'3 For instance, the derivation of parameterization

schemes by Hammett4 and Taft5 based on various equilibrium dissociation and

ionization constants elegantly demonstrates how thermochemistry can be used to

predict additional thermochemical information and interpret existing systems. The

1








quantification of field and resonance effects, wrought from extensive

experimentation,4'5 enables the direct application of thermochemistry to interpret the

energetic of chemical systems.

The application of thermochemistry to estimate additional thermochemical

information is undoubtedly dependent on accurate data. Although compilations of

thermochemical data are available for organic and simple inorganic molecules,6'7 the

situation for organometallic complexes is different. Because of the distinct differences

of organometallic complexes relative to organic systems, application of existing

organic thermochemistry to organometallics is limited. Inorganic thermochemical data

are primarily available for simple binary and tertiary molecules7 or main group

molecules,8 and parallels to large metal-centered complexes are inappropriate.

The thermochemistry of organometallic systems has gained increased attention

recently, driven by their increased importance in specialized areas of chemistry.9-11

The increased application of organometallic systems in the areas of catalysis and

material sciences has led to increased investigations into the thermodynamic

contributions to the reactivity of these systems. For example, knowledge of metal-

carbon and metal-hydrogen bond enthalpies is essential for the understanding of

catalytic reaction mechanisms that often involve metal-ligand bond cleavage.9

Bonding energetic of metal complex ions are important for heterogenous catalysis if a

complex ion is the active catalytic species.10 Additionally, knowledge concerning the

oxidation/reduction energetic of organometallic systems can be used to determine

solid-state lattice or binding energies for organometallic materials and ceramics.






3

Most thermochemical studies of organometallic complexes have focused on the

use of combustion or reaction calorimetry to determine bond energetic and heats of

formation in the condensed phase.12 Combustion calorimetry has been used to

estimate condensed-phase metal-ligand bond enthalpies for model organometallic

complexes (i.e. metal carbonyl and metal-arene complexes),13 and reaction calorimetry

has been used to derive metal-ligand bond enthalpies for metal complexes which are

useful as homogeneous catalysts.9 In spite of the accuracy of these techniques, there

are some disadvantages. If a reaction involves more than one metal-ligand bond

cleavage, the calorimetric study can yield ambiguous results. Further, the large

quantities of sample needed for calorimetry, along with the thermal instability of metal

complexes, renders traditional calorimeters ineffective for many organometallic
9 12
systems.9,12

Gas-phase investigations of metal complexes can lead to intrinsic properties of

a molecular species in the absence of solvent effects.9'14-17 For example, gas-phase

studies of isolated metal complexes or metal-complex ions can yield ionization

energetic data15'18-20 and bond dissociation enthalpies.15,16 Further, direct

comparison of the gas-phase energetic of metal species to their solution analogues

reveals solvent effects for the metal complex system.15-20

Gas-phase bond enthalpy studies have frequently involved coordinatively

unsaturated ionic metal species such as M+-R, where R = H, CH2, CH3 and so

forth.21-23 Although these studies are useful for estimating bare metal-substituent

bond enthalpies, direct solution analogues for the bare-metal ion species do not exist.








Bond enthalpies of the coordinatively unsaturated M+-R species may deviate

significantly from a related metal complex.9 Additionally, because the bare-metal

species lack supporting ligation, their stability and reactivity will not parallel that of a

comparable metal complex ion. The electronic properties of supporting ligands will

alter the bonding and solvation energetic for CP2MR+ complexes (Cp =

Tr5-cyclopentadienyl) relative to M+-R, for example.9'1

Numerous investigations on the stepwise bond energetic for organometallic

ions such as M(CO)n have been determined from mass spectrometric and

spectroscopic methods. For example, Norwood and coworkers have estimated the

sequential Fe+-CO bond strengths for ironpentacarbonyl by using a combination of

spectroscopic and mass spectrometric techniques.24 Recently, Schultz et al. have

measured the sequential Fe+-CO bond strengths for Fe(CO)5 by mass spectrometry.25

Further, Sunderlin et al. have recently reported the metal-ligand bond enthalpies for

various transition metal carbonyl anions in the gas-phase by using mass spectrometric

flow tube techniques.26

Another approach used to determine organometallic bond energetic has been

the application of thermochemical cycles which incorporate ionization and/or reduction

potentials.15'16 A limiting factor of this technique is the accuracy of the

organometallic oxidation/reduction data; auxiliary data needed for the cycles, such as

ligand heats of formation and metal atom ionization energies are typically accurately

known.8'27 Thermochemical cycles can also be used to determine solvation energetic

associated with metal complexes from direct comparisons of gas-phase and solution






5

oxidation/reduction energetics.15-20 Understanding solvation energetic is essential for

maximizing reaction efficiencies and understanding reaction pathways.28 Moreover,

analysis of the redox properties of metal-centered systems in the gas-phase and

solution can be utilized to interpret condensed-phase electrostatic interactions such as

ionic solvation effects,29 acid/base reactions,30 and electron-transfer phenomena.31

Ionization and reduction energetic are also useful for understanding and

developing periodic and group trends for an ensemble of metal complexes.16 For

example, the electronic effects of alkylation or ligand substitution can be readily

assessed by comparing relative redox potentials of a series of related metal

complexes.32 Then from an understanding of substituent electronic effects, complexes

can be modified to alter reactivity.5 Steric effects of substituents are easily

rationalized as they are primarily conceptional33 whereas electronic effects are less

obvious.2'3'5 Comparisons of ionization/reduction energies in the gas-phase enable

accurate assessment of substituent electronic effects. In the condensed phase, solvent

effects are predominant and often electronic effects in solution do not parallel trends

observed in the gas phase.5

The majority of data concerning ionization energetic for organometallic

complexes are either electrochemical potentials, E%,34 or vertical ionization energies,

vIP, measured by photoelectron spectroscopy.35 These data are not always appropriate

for deriving thermodynamic properties of ions near or at room temperature. For

example, a vertical ionization will not be a true assessment of a thermal ionization

process (free energy of ionization) if the equilibrium geometries of the ion and the








neutral are significantly different.27 Vertical ionization energies refer to an ionization

process in which the geometry of the parent ion is essentially the same as the ground

state neutral. Specifically, the most probable ionization transition will be that in which

the positions and moment of the nuclei remain unchanged.36 The ion is formed with

excess internal rotational and vibrational energy and is therefore not in its ground

state.27 A thermal ionization process, AGiT, which is similar to an adiabatic

ionization energy, or AH/,o, is a process in which both the neutral and the ion are in

their equilibrium geometries at temperature T.16 At temperatures greater than T = 0

K, there are entropy and heat capacity contributions to the ionization process.16'27 If

the structures of a neutral and an ion are significantly different, a vertical ionization

will not be an accurate assessment of the thermal ionization. Potential well diagrams

which represent both vertical and thermal ionization processes are shown in Figure

1-1. Although the vertical ionization process is the most probable transition,27'36 ion

relaxation (from ion-molecule collisions,14 for example), which removes excess

rotational and vibrational internal energy, results in a thermal ionization process. The

difference in vertical ionization energy and AGi is the ion relaxation energy given by

Er = vIP AGio.16

Electrochemistry is a useful method for evaluating redox potentials. However,

electrochemical potentials for many organometallics are irreversible, making

determination of a true E value difficult.34 Organometallic complexes that form

unstable solvated ions will produce electrochemical responses inconsistent with gas-

phase data. Although the vertical ionization energy for manganocene has been

















































Potential well diagrams demonstrating vertical and adiabatic ionization
processes. Arrows represent vertical ionization transitions. Figure A is
an example in which the equilibrium geometries of the ion and the
neutral are significantly different. Figure B represents a process in
which the geometries of the ion and neutral are similar. For this
example shown, vlP alP.


Figure 1-1








measured by photoelectron spectroscopy,35,37'38 the E value for Cp2Mn has not been

reported due to the instability of Cp2Mn+ in solution.34

The study of gas-phase electron-transfer equilibrium is a well-established

means for determining thermal ionization energies for organic39-42 and inorganic14-20

complexes at or near room temperature. The ability to trap ions and monitor an

electron-transfer equilibrium reaction is a powerful application of mass spectrometry.

Previous electron-transfer equilibrium studies have been reported by several groups

using various mass spectral techniques.15'39,42 An earlier study on the electron-

transfer equilibrium of the gas-phase ionization thermochemistry of ferrocene by using

pulsed high pressure mass spectrometry has been reported by Mautner.43

In this work thermal free energies of ionization for metallocenes and

metallocene derivatives have been determined from electron-transfer equilibrium

studies by using Fourier transform ion cyclotron mass spectrometry, FTMS.4447 From

the measured equilibrium constants for the electron transfer reaction, AGet values are

determined, and the free energies of ionization are estimated for the metallocenes for

the process shown below, where L represents a Cp ligand or a Cp derivative,



L2M -- L2M+ + e- (1-1)



where Cp = ir5-cyclopentadienyl. Ionization free energies for the first transition row

metallocenes, ruthenocene, and osmocene are presented.16 Metallocenes were chosen

because they represent the foundation of organometallic chemistry and have numerous







applications in the areas of homogeneous catalysis,9 material science,48 and nonlinear

optics.49

Following established methods,15'16 the ionization energetic data are

incorporated into thermochemical cycles to provide estimates of bond disruption

enthalpies for the first transition row metallocenes and metallocenium ions.16

Although bond enthalpies for the neutrals have been reported,13 accurate bond

enthalpy values for the metallocenium ions have only recently been established.16

Cyclopentadienyl ring substituent effects for metallocene complexes have been

studied by electrochemistry34 and occasionally by photoelectron spectroscopy.35

Relatively few gas-phase studies on the substituent effects for metal complexes have

been reported for comparison to the electrochemical potentials.50'51 In this work the

electron-transfer equilibrium method has been applied to a series of alkyl ferrocene16

and nickelocene19 derivatives to investigate the effects that substitution has on the gas-

phase redox energetic of metallocene complexes. Nickelocene is a useful complex

because it forms stable cations and anions in the gas-phase;19 therefore, effects of

alkylation on the positive and negative nickelocene complexes were studied. The AGi

and AGa0 values for the metallocene derivatives are correlated with alkyl substituent

parameters. Various parameterization schemes5,52'53 are considered to interpret and

understand the gas-phase investigations and electron-transfer energies. These data are

potentially useful in understanding and predicting ionization energies51 and optical

transition energies for metallocene derivatives.49 Optical transition energies are

important in the selection of chromophores for potentially useful nonlinear optical






10

devices. Since ligand substituents affect molecular orbital energies,5 metal complexes

can be designed with specific optical transitions based on knowledge derived from

gas-phase redox potentials.

The AGi values for a series of ruthenocene derivatives have been determined

from electron-transfer equilibrium reactions.20 The AGi values span over 2 eV within

the series. Because of the widespread variations of the Cp ligands, bulk ligand effects

were considered rather than individual ring substituent effects. A parameterization

scale for correlating electronic effects for Cp type ligands was developed.20 The

application of ligand substituent effects to understanding and predicting reactivity for

organometallics complexes is discussed.

Relatively little is known concerning the solvation energetic for organometallic

complexes.16 In order to better understand solvation effects of metal complex redox

couples, an understanding of the intrinsic (solvent free) electron-transfer chemistry is

important. Electrochemical potentials for several metallocene derivatives have been

measured in order to fully characterize the effects ligation and solvation has on the

redox chemistry of metal centered molecules. Thermochemical cycles, which

incorporate gas-phase and solution metallocene redox potentials,16'19'20 have been used

to derive estimates of solvation energetic. Electrostatic models for predicting

solvation energetic for spherical ionic species have been applied for the metallocene

complexes.15,54,55 Shortcomings and criticisms of the electrostatic model are

discussed in light of the experimentally derived estimates.









Throughout this work, several important objectives are examined. The

evaluation of accurate ionization and reduction potentials and detailed analyses of the

data for prototypical organometallic complexes are important objectives of this work.

For example, temperature-dependence studies have been determined to evaluate

entropy and enthalpies of ionization for several metallocene complexes.16

Comparisons of the experimental work to spectroscopically derived thermodynamic

parameters revealed specific contributions of the ionization process. Thus, the

thermodynamic origins of the ionization processes are critically analyzed.

The application of thermodynamics to predict reaction mechanisms and

reactivity is another major objective of this work. The thermodynamic parameters

presented in this work may be applied to other areas of research that utilize the myriad

of characteristics of organometallic complexes. Existing parameterization schemes52'53

for predicting reactivity have been successfully used to correlate thermodynamic data

for the metallocenes.16'19 Where established models fail to correlate, schemes have

been developed for predicting thermodynamic values for metallocene-type

complexes.20 The application and development of the parameterization schemes for

understanding and predicting chemical reactivity are discussed in detail.

A significant portion of this work is devoted to deriving and understanding

metallocene/metallocenium solvation effects. Relatively little is known concerning the

solvent effects for organometallic complexes. Solvation data reported here may help

to further develop a foundation for understanding solvation effects for other

organometallic complexes.








The development of methods for fully characterizing the thermodynamics of

organometallic complexes was explored. Fourier transform mass spectrometry has

proven to be a powerful tool towards fulfilling this goal. Electron-transfer equilibrium

techniques are an effective means of determining thermodynamic parameters. When

complimented by electrochemical studies34 and photoelectron spectroscopy,35 FTMS

studies allow for the full characterization of redox properties for organometallic

complexes. Further, the development of temperature-dependence techniques

establishes FTMS as a technique for deriving reaction entropies and enthalpies for

metal systems.












CHAPTER 2
ADIABATIC IONIZATION ENERGIES, BOND DISRUPTION ENTHALPIES,
AND SOLVATION FREE ENERGIES OF GAS-PHASE METALLOCENES
AND METALLOCENIUM IONS


Introduction


The majority of data concerning metallocene oxidation-reduction potentials is

in the form of vertical ionization energies measured by photoelectron spectroscopy34

and electrochemical potentials.35 Several experimental limitations of these two

techniques can cause uncertainty in the determination of thermodynamic reduction-

oxidation potentials. Many organometallic complexes have irreversible

electrochemical oxidation/reduction potentials35 which can lead to uncertain

assignments for Eh values. Vertical ionization energies may differ from adiabatic

potentials if the equilibrium geometries of the ion and the neutral are dissimilar.27'36

If the equilibrium geometry of the ion and the respective neutral are similar, then the

vertical potential will closely approximate the adiabatic potential which is referenced

at 0 K.27 However, even if photoelectron studies can accurately determine an

adiabatic potential, values for the ionization free energy or reduction and respective

enthalpy changes referenced at other temperatures must be estimated from statistical

thermodynamic analyses. Spectroscopic data, which include vibrational and structural

characterization, for metal complexes frequently do not exist for metal complexes.16








Electron-transfer equilibrium, ETE, is a powerful method for determining

thermal oxidation and reduction potentials for organic and inorganic species at ambient

temperatures.15,16'39'41 Kebarle and coworkers have used pulsed high pressure mass

spectrometry, PHPMS, to determine the free energies of electron attachment, AGa, for

organic compounds in the 0-3 eV range.42 In the Kebarle studies, the electron

attachment free energy of SO2 was used as the reference compound to anchor the

AGa0 values derived from electron-transfer equilibrium studies. An accurate value for

the electron affinity of SO2 has been determined from photodetachment studies by

Celotta et al.56 Additionally, sufficient spectroscopic data for SO2 exist for complete

statistical mechanical analyses of the enthalpy and entropy for electron attachment.57

Thus, all AGa0 values reported from Kebarle's laboratory are referenced to AGao(SO2)

at 423 K.

In this work, Fourier transform ion cyclotron resonance mass spectrometry4447

has been used to determine the free energies of ionization for some first transition row

metallocenes, ruthenocene, and osmocene. Derived AGio values are represented by

equation 2-1, where L denotes a Cp ligand or a substituted Cp ligand.



L2M(g) L2M+(g) + e" (2-1)



Ionization free energies for the compounds studied in this chapter have been

previously reported. 1318 However, the derived AGi values for the metallocenes

reported here differ from values of the previous study due to refinements in the free








energies of ionization for the organic reference compounds and more extensive

electron-transfer studies.16

Entropies and enthalpies of ionization for the metallocenes have been estimated

from investigating the temperature-dependence of the equilibrium constants for

selected metallocene reaction couples. Statistical mechanics has been used to estimate

the intramolecular entropy change for ferrocene. Available spectroscopic data for

Cp2Fe and Cp2Fe+, in addition to new vibrational frequency data for the ferrocenium

cation, were used to estimate values for the total ASio at several different temperatures.

Mautner has used PHPMS to determine the ionization free energy of ferrocene by

electron-transfer equilibrium in which alkylaniline compounds were used as reference

compounds.43 Additionally, Mautner was able to assess values for the entropy and

enthalpy of ionization by studying the temperature dependence studies of the measured

equilibrium constants. Thermodynamic parameters for ferrocene determined in this

work are compared to values reported by Mautner.

Thermochemical cycles were used to estimate differential solvation free

energies and bond disruption enthalpies.15'16,18 Values for heterolytic and homolytic

M-Cp bond cleavage have been previously presented.13 However, since bond

enthalpies derived from thermochemical cycles are dependent on AGio values, new

estimates are included which reflect refined AG1i data in addition to more accurate

thermochemical analyses. Differential solvation free energies are derived from direct

comparison of AGi values in the gas phase and solution.15'16'54 Estimates for








differential solvation free energies for several metallocene couples are compared to

values predicted by a simple electrostatic model.58

The effects of attached substituents to cyclopentadienyl and arene ring systems

have been studied previously by electrochemistry and photoelectron spectroscopy. An

earlier study demonstrated that free energies of ionization for alkylferrocene

derivatives correlate well with alkylbenzene analogues.13 In the present work, AGio

values derived from electron-transfer equilibrium studies for the alkylferrocene

complexes are correlated with alkyl substituent parameters. Alkyl substituent

parameters have been shown to correlate well with AGi data for organic compounds

and chromium coordination complexes,51 but have only recently been applied to

metallocenes.16'32 These data may prove beneficial for the interpretation and

prediction of physical properties for organometallic complexes such as ionization

energies or optical transition energies useful to photochemistry.49


Reevaluation of Metallocene Free Energies of Ionization
Based on Electron-Transfer Equilibrium Studies


Electron-transfer techniques have been described elsewhere.15'16'39'42 The

general electron-transfer equilibria shown in equation 2-2 were studied where L2M

represents a metallocene and R denotes a reference compound with known AGio value.

Ionization free energies for the reference compounds used in this work are typically

1 to 2 kcal mol1.27,40,59


L2M + R+ R + L2M+


(2-2)







































1.0


2.0 3.0
Reaction Time


4.0
(s)


5.0


6.0


Figure 2-1 Log plot for the electron-transfer reaction of Cp2Fe + DET+ = DET +
Cp2Fe+. Ion intensity units are arbitrary. DET = N,N-diethyltoluidine
and Fc = ferrocene.


100

i


1n


.,. 4*

+ NFc







m +
m ---- DET
m W


i- -


'


0.0







Figure 2-1 is an example of typical electron-transfer equilibrium reaction, the

ferrocene/N,N-diethyltoluidine couple in this example. The decay of ion signal over

time is due to diffusive loss of ions from the reaction cell of the mass spectrometer.

Equilibrium constants and subsequently electron-transfer reaction free energies, AGeto,

can be determined (equation 2-3) if the difference in the ionization

free energies of the two compounds is < 4 kcal mol'1.10 In equation 2-3, P denotes

the partial pressure and I represent the parent ion intensity for species in reaction 2-2.



AGeto = -RT In Kq= -RT In [(P(R)/P(L2M)) (I(L2M+)/I(R+))] (2-3)



Equilibrium between parent ions can actually be monitored for reactions with AGet0

values approaching 5 kcal mol'-;6 however, the partial pressure ratios must exceed

100, leading to large experimental uncertainties. The pressure of the minor component

would only be slightly above the background pressure (ca. 1-2 x 10-8 Torr). Since the

AGi value for R is known, the free energy of ionization of the metallocene at

temperature T can be determined from equation 2-4. Determination of AGet values

over a range of temperatures yields estimates of AHeto and ASet.



AGet,T = AGi,To(L2M) AGi,T(R) (2-4)



If the enthalpy and entropy of ionization of the reference compound are known,

corresponding values for the metallocenes can be derived for the appropriate








temperature range. All reactions temperatures in the present work used to derive

AGeto values are 350 K unless otherwise stated.


Reference Compounds Used in Electron Transfer Equilibrium Studies.


Appropriate reference compounds for ETE studies in the 6-7 eV range have

been studied by PHPMS and ion cyclotron resonance mass spectrometry.39'59 Lias et

al. have used ETE methods to determine the ionization energies of several aniline

derivatives which are used as reference compounds in the present work.59 Reference

compounds are anchored to the ionization potential value of N,N-dimethylaniline,

DMA, at 350 K.59 The ionization energy of DMA is anchored to the ionization

potentials of benzene and NO2,39 which have been determined spectroscopically.

Values of ASi for benzene and NO2 have been estimated from statistical mechanic

analyses;39 subsequently, accurate values of AGi,3500 have been determined.

Reevaluation of the ionization thermodynamics for the reference compounds

was necessary in order to determine AGi,3500 data for the metallocenes. Ionization

thermodynamics used in this work for the reference compounds are based on several

important assumptions:



(i) AHi(DMA) = 7.12 0.02 eV (164.2 0.5 kcal mol'1).27

(ii) ASi' for azulene results from ASelec. Therefore ASi = R In g = 1.38 cal

mol^K-1.40 The ionization of azulene is a singlet to doublet transition in

which g = 2.







(iii) Based on the reported ASet' value from ETE studies by Nelsen and

Mautner involving DMA and azulene, the ionization entropy for DMA is 2.3

cal moltIK1.

(iv) Based on the structural similarity, the ASi0 values for the aniline

derivatives used as reference compounds are equal to ASiO(DMA).

(v) Changes in the integrated heat capacities of the reference compounds and

their neutrals are negligible and will cancel. Therefore, derived AHio values

are assumed to be independent of temperature in the 350-500 K range.39



Assumption i has been used in other studies in which N,N-dimethylaniline has

been used as a reference compound.4043 The AGiO value used here is within 0.01

eV of the literature value27 and in agreement within experimental uncertainty of the

AGi value used by Nelsen and Mautner. Assumption ii is quite reasonable

considering the small vibrational changes expected for the ionization of azulene.

Azulene and its cation are expected to be structurally similar,16'40 therefore vibrational

and rotational contributions to the ionization entropy will be negligible. Translational

entropy changes for ions and their parent neutrals will be negligible since the two

species differ only by the mass of an electron. Aniline derivatives are also expected

to have small vibrational and rotational contributions to the total entropy change, and

ASiO = 2.3 cal mollK"1 (assumptions iii and iv) represents the electronic entropy

change (singlet to doublet process, g = 2) with some contribution for ASvib and ASror

Based on previous studies the ASet for the azulene/DMA couple40 was reported to be






21

0.9 cal molfK-1 which results in ASi(DMA) = 2.3 cal mol'lK1. The minor effects of

ACp on AHio are not unreasonable especially for larger molecules. For example,

AHi,350 Hi,O for benzene, determined from statistical mechanics, is only 0.2 kcal

mol"1 which is negligible when compared to the much larger AHio value.39 As will be

shown later, assumption v holds true for the metallocenes.


Ionization Energetics Data for Some Metallocenes


Figure 2-2 is an electron-transfer equilibrium ladder displaying all reaction

investigated in this chapter. This ladder is similar to an equilibrium ladder previously

reported.13 The AGi values in Figure 2-2 differ from those of an earlier study due to

the assumptions outlined above and more extensive ETE studies. Details involving

individual ETE have been discussed elsewhere.16 Several ETE reactions performed

with respect to an earlier reportl3 include the following reactions couples:

Cp2V/Cp2Fe, (EtCp)CpFe/Cp2Fe, (EtCp)CpFe/(n-butylCp)CpFe, and temperature

dependent studies of the equilibrium constant for Cp2Fe/DET and (EtCp)CpFe/Cp2Fe

couples. Numbers adjacent to arrows in Figure 2-2 denote AGet,3500 values for

individual ETE reactions. Reactions involving only reference compounds are included

in the equilibrium ladder to serve as checks on the literature values.59 Values in

parentheses are unchecked literature values. Table 2-1 lists the revised AGi values

for the Cp2M compounds in addition to the vertical ionization energies measured by

PES.35 An experimental error of 1.5 kcal mol-1 is assigned to each AGio value

based on experimental uncertainties of the reference compounds and errors in the













/A J


< -NMe2 163.4

NEt2
NEt2 160.1
0 NMe2159.4

)NEt2 156.3


164.6- Cp2Ru



160.6 Cp20s


154.5-


1


Cp2V


153.1 Cp2Fe
(CpCH2NMe2)CpFe
150.2 (EthylCp)CpFe
.149.1 (n-butylCp)CpFe
S148.9 (t-butylCp)CpFe
47.7 (MeCp)2Fe


N-N 145.6 4.2

1 IF143.8 Cp2Ni
1't 142.5 Cp2Mn







Figure 2-2 Free energy of ionization ladder for some metallocenes. Values of
AG,, lie adjacent to arrows and AGi0 values are next to the compound.
All values are in kcal mol- .


1








Ionization Energetics Data for Some Metallocenes.


Cp2M AGioab Ajioa ASi vIP (PES)a


Cp2V 154.5 -154 -0 0.3d 155.7e

Cp2Fe 153.1 157.2 11.6' 158.2g

Cp2Mn 142.5 159.38

Cp2Ni 143.8 -149.98

Cp2Ru 164.6 165.4 5.0h 171.88

Cp2Os 160.6 161.1 5.0h 164.9'


a. Units are kcal mol-1.
b. Determined from the derived free energy of ionization and the estimated entropy of
ionization at 350 K.
c. Units are cal mol-Kl1.
d. The ASi value is assumed to be equal ASelec only estimated from uncertain ion
ground state.
e. See reference 37.
f. Determined from a van't Hoff plot for the Cp2Fe/DET couple assuming
ASio(DET) = 2.3 cal mol'K-1.
g. See reference 35.
h. Value assumed to be equal to statistical mechanics value for ASi,3500(Cp2Fe).
i. See reference 36.


Table 2-1.






24

measured partial pressures of the neutral gases. Estimates of AHio and ASi for several

metallocenes are also included in Table 2-1 when sufficient data were available.

The AGi0 for Cp2Ni and Cp2Mn were referenced to the AGio for 1,1'-

bipyrrolidine which is anchored to the AGi value of DMA40 Nelsen has suggested a

value for AHi = 146.9 kcal molr1 for 1,1'-bipyrrolidine based upon PHPMS ETE

studies. The AHi value chosen was used as a reference for nickelocene and

manganocene in an earlier paper.16 From temperature dependent studies performed by

Nelsen and Mautner, a value for ASi = 3.2 cal mol-K-1 was reported.40 Based on the

assumptions stated earlier concerning the reference compounds, a value for AGi,3500 =

145.8 kcal mol"1 has been estimated. The ferrocene derivatives help extend the

equilibrium ladder from free energies of ionization of the alkylaniline compounds

down to manganocene and the internal agreement is very good ( 0.3 kcal mol'l).

The temperature dependence of the ET equilibrium constants are shown as

van't Hoff plots in Figure 2-3 and the derived thermodynamic parameters are also

presented in Figure 2-3. The ionization enthalpy and entropy values for ferrocene are

based on the derived thermodynamic parameters values from the van't Hoff plots.


Insights into the Free Energy of Ionization of Ferrocene


The electron-transfer equilibrium reaction of ferrocene and N,N-

diethyltoluidine, DET, has been studied by Mautner by using PHPMS; a value of

AGet,350 = -0.9 kcal mol"1 has been estimated for the couple.43 Mautner's study

covered a temperature range of 450 650 K, and AHet and ASet0 values have been
















6.0-


5.0-


4.0-


, 3.0-


2.0-


1.0-


0.0-
1.8








Figure 2-3


Cp2Fe/DET





A. EtFc/Fc








Cp2Ni/Cp2Mn

2.2 2.4 2.6 2.8 3
1000/T (K)







Van't Hoff plots for selected metallocene electron-transfer equilibrium
couples. Values of AHet and ASet0 for the reaction couples are as
follows. For the Cp2Fe/DET couple, AHet = +0.05 0.47 kcal mol1
and ASet = 9.3 1.1 cal mol'lK1. For the EtFc/Fc couple, Het and
ASet are 3.1 1.4 kcal mol"1 and 0.96 3.6 cal mol- K". For the
Cp2Ni/Cp2Mn couple, AHet = -3.2 1.5 kcal mol-1 and ASet = -5.7
3.0 cal mol'K"1. DET = N,N-diethyltoluidine, Fc = ferrocene, EtFc =
ethylferrocene.







extrapolated to 350 K, which is the temperature at which the majority of FTMS

studies were performed. From the present study, the estimated value of AGet,3500 for

reaction 2-5, is -3.2 0.5 kcal mor1 yielding a value of AGi,350(Cp2Fe) = 153.1 1

kcal molr. This value differs from a values from a previous study which reported

AGet = -2.8 kcal mol-1.18



Cp2Fe + DETI" DET Cp2Fe+ (2-5)



The larger equilibrium constant for equation 2-5 for the FTMS study is consistent with

the observation that ferrocene did not come to equilibrium with N,N-dimethyltoluidine,

DMT, where an estimated AGeto value > 5 kcal morl is expected. The Cp2Fe/DMT

reaction couple was studied by Mautner with an estimated AGetg = -3.9 kcal mol-1 at

429 K. Several electron-transfer equilibrium reactions that have been initially studied

by high pressure mass spectrometry have been repeated in our laboratory with good

accuracy. The electron-transfer reaction for p-cyanonitrobenzene/benzoquinone couple,

(p-CNNB/BQ), equation 2-6, was studied by Grimsrud et al.61 and AGet = -4.0 0.8

kcal mol1 was derived for the PHPMS study.



p-CNNB- + BQ BQ- + p-CNNB (2-6)



The same reaction was studied by using FTMS and the derived AGet0 = -3.8 kcal

mol-1 was determined, in good agreement with the PHPMS results.








The difference in the PHPMS and FTMS experiments for the Cp2Fe/DET

couple is much larger than expected for such comparisons and this led us to examine

the reaction couple in more detail. Mautner has determined the temperature

dependence of the equilibrium constants for reaction 2-5 in the 450 to 650 K range,43

therefore the same experiment was performed by using FTMS. The thermodynamic

parameters for the reaction couple derived from the PHPMS are AHet = -0.1 kcal

mol"1 and ASeto = 2.2 cal mol^K'1. An FTMS temperature dependence study would

serve to illustrate the origins of the differences for AGet between two studies and

would provide thermodynamic data that would overlap in temperature with the

PHPMS work.

The van't Hoff plot for the Cp2Fe/DET couple for reaction 2-5 is shown in

Figure 2-3. The temperature range in the present study is from 350 to 520 K.

Consistent with the PHPMS study, the reaction displays a minor temperature

dependence.43 The derived thermodynamic values for the reaction are AHet = 0.05

0.47 kcal mol-1 and ASeto = 9.3 1.1 cal mol^K"1 for the FTMS study. Values were

determined from linear regression and are reported at the 95% confidence limit. The

AHet0 value for both studies are within experimental error limits; but the ASet values

are significantly different, the FTMS values being much more positive, and this is

obviously the source of error in the AGet values.

Based on the assumptions outlined for the reference compounds, a value for

AHio = 157.2 1.5 kcal mol"1 (6.82 eV) is obtained for ferrocene. This value is 0.2

kcal mol"1 higher than the value obtained by Mautner because of the difference in the






28
chosen AHi(DET). The adiabatic ionization energy obtained for both studies is well

within experimental uncertainty and a value for AHio can be give as 6.82 0.08 eV.

The vertical ionization measured by photoelectron spectroscopy by various groups is

6.88 0.1 eV.35'37'38 An adiabatic ionization potential has not been reported for

ferrocene previously due to lack of vibrational fine structure in the PES valence

ionization manifold. Rabalias has reported vibrational fine structure for Cp2Fe with 35

meV separation,62 however this has not been consistently resolved by other groups.

The source for the difference in the ASio values of the PHPMS and the FTMS

studies is unclear; however, several experimental observations favor a higher value for

the entropy of ionization than that reported by Mautner. The temperature dependence

of other electron-transfer equilibria involving molecules and ions that have well

established thermodynamic constants have been examined by using FTMS, and the

entropy changes for the reaction couples are within 5 cal mol'K-1 of the expected

value.10 The difference in ASet value for the p-CNNB/BQ reaction couple (reaction

2-6) studied by PHPMS and FTMS is -4 cal mol^K-1.

The temperature dependence of the electron-transfer equilibrium reaction of

(EtCp)CpFe/Cp2Fe was examined to further assess our ASet0 values for the DET/Cp2Fe

reaction. Based on the assumptions made for the reference compounds, ASi'(DET) =

2.3 cal molK"1K, therefore the ASi(Cp2Fe) = 11.6 cal mol'1K1. Assuming that the

ASio for ethylferrocene is equivalent to that for ferrocene, the derived ASet for the

(EtCp)CpFe/Cp2Fe electron-transfer reaction should be -0 cal mol'lK1. The derived

entropy change for the reaction couple was found to be approximately zero within








experimental error. Parameters for the reaction are AHet = 3.06 1.43 kcal mol1l

and ASet = 0.96 3.6 cal molilK1 at the 95% confidence limit.

The organic reference compounds used in this work have AGi0 values that have

been determined from ETE methods from PHPMS40 and FTMS59 studies. Mautner

has determined AGi for 1,1-bipyrrolidine40a and the hydrazine has been linked to the

alkylaniline derivatives by ferrocene and its derivatives. The internal consistency of

the ladder is within 0.3 kcal mol"1. Since both the PHPMS and the FTMS studies

have used the same reference compounds in ETE measurements, and the internal

consistency of the present equilibrium ladder is well within experimental uncertainty,

the equilibrium constants obtained by FTMS for the metallocene are not expected to

have large random errors.


Origins of Possible Experimental Uncertainties in the Thermodynamic Constants
for Ferrocene


A possible explanation for the difference between the equilibrium constants for

the ETE reaction of DET/Cp2Fe would be inaccurate measurement of the parent ion

intensity ratios due to discrepancies in the ion detection. In PHPMS studies performed

by Mautner, ion sensitivities are frequently calibrated by comparing fragmentation

spectra with reference spectra to insure that the mass detector is not giving biased

sensitivities.63 Reported errors due to mass detection drift in the PHPMS are 0.3

kcal mol" for typical ETE reactions. In FTMS experiments, mass differences between

two ions can lead to differences in detection sensitivities especially if the mass

difference is extremely large.64 In this work, ion detection parameters were adjusted






30

to give the maximum signal for both ions. The total ion count (I(Cp2Fe+) + I(DE+))

was monitored with time and the change was determined to be small, ca. 10%. The

variation of the ion intensities during the approach to equilibrium is not expected to

significantly affect to the experimental uncertainties for our systems.

The most likely source of error is the measurement of the partial pressures of

the neutrals. Both techniques use different methods to determine pressures of the

reactant gases. It is therefore possible that systematic errors for one or both of the

methods exist, leading to a difference in the equilibrium constant. The difference in

the Keq values for the two experiments is equivalent to a factor of 25 in the pressure

ratios. Derived enthalpy changes are not dependent on pressure ratios; therefore, it is

of no surprise that the AHet values for the two studies are in agreement Pressures

are measured directly with an ion gauge in the FTMS studies. In the PHPMS work, a

solution containing known concentrations of reactants is introduced into a heated bulb

and partial pressures of the reactants are calculated.43'63 Measured partial pressures in

the FTMS studies are not expected to vary by more than 30%, resulting in a 0.2

kcal mol"1 error in the AGet0 value.63 In the PHPMS studies, pressures are monitored

by measuring Keq at various partial pressure ratios. Mautner reports that experimental

error due tp pressure fluctuations are 0.5 kcal mol'1.


Evaluation of the Electron-Transfer Reaction Rates for the CpgFe/DET Couple


The rates of the ion-molecule reactions for the Cp2Fe/DET couple were

examined to further support that the source of error between the PHPMS work and the








FTMS studies stems from inaccurate pressure determinations. Electron-transfer

reaction kinetics for ferrocene with DMT and DET have been investigated by using

PHPMS. Mautner reports values for kf for reaction 2-5 of 1.2 x 10-9 cm3 molec'1s-1

at 461 K and 1.7 x 10-9 cm3 molec's-1 at 429 K for the Cp2Fe/DMT couple (reaction

2-7). By using FTMS, kf = 2.5 0.5 x 10-10 cm3 molec-s"1 for reaction 2-5 was



Cp2Fe + DMT' -- DMT + Cp2Fe+ (2-7)



determined. The FTMS value of kf for reaction 2-7 was determined to be 1.3 0.3 x

10"10 cm3 molec-1s-1. Mautner's rate constant for reaction 2-5 is slightly faster than

the Langevin collision limit while the present values for kf is ca. 25% of the Langevin

collision limit. The two rate constants for reaction 2-5 differ by a factor of 5 and do

not resolve the disagreement in the derived AGi values for ferrocene.

For example, if the PHPMS value for Kq is assumed to be correct, then the

pressure ratio determined in the FTMS work, P(DET)/P(Cp2Fe), is too large by a

factor of -25. This would lead to an overestimation of kf by a factor of -25,

depending on the errors in the absolute partial pressures of DET and ferrocene. Thus,

the kf values for the two experiments would diverge if the pressure of ferrocene, for

example, was underestimated.

By examination of equation 2-8, if the pressure for ferrocene was incorrectly

underestimated, the equilibrium constant, and thus AGet0 would be too large since

there exists an inverse relationship between P(Cp2Fe) and Ke. Moreover, an






32

Kq= In [(P(DET)/P(Cp2Fe)) x (I(Cp2Fe)/I(DET)] (2-8)



underestimation of the reaction pressure of ferrocene would lead to a value for kf that

is slower than determined. Reaction pressure and kf are inversely related, therefore if

the total pressure of a reaction is actually greater than measured, the rate constant for

the reaction will be slower than the experimentally determined kf value. Thus, it

obvious that comparison of the rate constants for the Cp2Fe/DET couple do not

resolve the discrepancies for the two experiments.

Based on previous FTMS studies involving electron-transfer reactions of

metallocenes,14 it seems unlikely that such large errors in the measured partial

pressures could occur. The same electron-transfer methods have been applied to the

study of metallocene self-exchange and cross reactions.14 Estimated reaction

efficiencies for exothermic cross reactions were in the 0.5-1.5 range which suggests

that partial pressure errors may be incorrect by as much as 50%. Reaction efficiencies

are given by equation 2-9, where an estimate of kL for the metallocenes, the Langevin

collision rate, is 1.0 x 10-9 cm3 molec'ls-1.14



Efficiency = kf/kL (2-9)



An observed efficiency of ~0.2 is not unexpected for a reaction with low

exothermicity. The reported FTMS self-exchange rate constant for Cp2Fe+/0 is 2.7 x

10-10 cm3 molec^s"1 (0.27 efficient)14 which is consistent with the present kf for the






33
DET/Cp2Fe couple. Alternatively, large pressure errors in the PHPMS work also seem

improbable since ferrocene in not susceptible to thermal decomposition in the

temperature range used.65 Although there are distinct differences in the observed

forward rates for reaction 2-5, the kinetic data does not expose the origin for

difference in the two experiments.


A Brief Survey of Proton-Transfer Reactions of Ferrocene


The kinetics and thermodynamics of the protonation of ferrocene have been

studied by several groups by using PHPMS43'66 and ion cyclotron resonance mass

spectrometry.67 Recently, PHPMS proton-transfer equilibria studies performed by

Ikonomou and Kebarle assessed a value for AGBg for ferrocene of 195.2 1.0 kcal

mol-1 at 500 K.66 A value for AGB(Cp2Fe) = 195.0 1.5 kcal mol"- at 600 K has

been estimated from PHPMS data obtained by Mautner.43 The term AGBg is the gas-

phase bacisity (the free energy of proton attachment), and differs from the proton

affinity, which is an enthalpy change, AHBg, referenced at 298 K.

The proton-transfer reaction 2-10 was studied by using FTMS in order to gain

additional information concerning discrepancies in the AGi(Cp2Fe) values reported

here and by Mautner. The proton-transfer equilibrium reaction 2-10 has also been

studied by both Mautner43 and Ikonomou66 and the reported free energy changes are

-1.2 and -1.9 kcal mol-1 respectively. A value of AG3500 = -1.5 kcal mol"1 was


PyrroleH+ + Cp2Fe = Cp2FeH+ + Pyrrole


(2-10)






34

derived from FTMS studies of reaction 2-10 and this value is consistent with the two

PHPMS AG values. The difference in the AG values for the FTMS study and the

PHPMS work results in a factor of -1.5 in the measured equilibrium constants. A

factor of 25 in Kq for the proton-transfer reactions, which is the difference in the Keq

values for reaction 2-5, would yield a value of AG3500 ~3.6 kcal mol"1 for reaction 2-

10. As all three studies yield equivalent AGO values for reaction 2-10, evaluation of

proton-transfer reactions does yield information concerning the discrepancies in the

AGi values for ferrocene. Furthermore, the AGBg estimated by Beauchamp and

Stevens67 from ICRMS experiments is consistent with all three proton-transfer studies.


Intramolecular Entropy Changes for the Ferrocene/Ferrocenium Couple


In order to further understand the uncertain entropy change for the Cp2Fe+/O

couple, detailed statistical mechanical analyses have been performed to provide an

accurate estimate for ASi(Cp2Fe). Table 2-2 lists values for translational, rotational,

vibrational, and electronic entropies of ferrocene and ferrocenium ion at several

temperatures. Complete vibrational analyses for all 57 vibrational modes of ferrocene

have been reported by Bodenheimer and Low68 and the measured frequencies were

used in the vibrational entropy analysis. Although a complete vibrational analysis for

the ferrocenium cation has not been performed, sufficient vibrational frequency

assignments for Cp2Fe+ have been reported by several groups to allow for an estimate

of Sibo(CP2Fe+).69"71 Additionally, crystal structures for both ferrocene and















i^7


ON
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'-r












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C tr)

(Sl '/n













in kn
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C.,.
S






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a




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c


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a





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60

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ferrocenium cation have been reported providing information needed for determining

of the rotational entropy change for the Cp2Fe+/ couple.72'73.

Low-frequency metal-ligand vibrations can contribute significantly to entropy

changes of metal complex redox processes if M-L frequency shifts drastically change

upon oxidation or reduction. For example, changes in the skeletal Co-N vibrations for

the Co(NH3)62+/3+ redox couple result in an estimated entropy change of 12.6 cal

mol'K'1.74 The change in the asymmetric T1u vibration alone accounts for -3 cal

mol-K1K or -25% of the total skeletal vibrational entropy change for the

Co(NH3)62+/3+ couple.74

The doubly-degenerate Cp-Fe-Cp bend has not been reported previously, but

the corresponding mode for ferrocene has a relatively low frequency, u22 = 179

cm'1.68 Because of the importance of this low frequency skeletal mode, rather than

estimate a value for v22, the Cp-Fe-Cp frequency for ferrocenium ion was measured.

Our far-infrared studies have conformed )22 for ferrocene and tentatively assign a

value of 135 cm"1 for u22(Cp2Fe+). The Fe-Cp distance increases upon oxidation, and

this distance increase is expected to result in a lowering of the frequency for the v22

metal-ligand bending mode.72 This single shift in the bending mode contributes 1.0

cal mol'K"1 to the ASi(Cp2Fe+/o) at 298 K. The results of the far-IR studies for

several ferrocenium salts are given in Table 2-3.

An estimate of ASetO for the electron-transfer equilibrium reaction at 450 K can

be determined by combining the calculated entropy of ionization for ferrocene (4.9 cal

mol1'K'1) with the assumed value of ASi(DET) = 2.3 cal mol1'K1 giving a value for







Table 2-3. Vibrational Frequencies For Various Ferrocenium Salts


Compound Frequency (u22) S0(22)a
Cp-Fe-Cp Bend cal mol-K-1


(Cp2Fe+)(C1) 135 cm-1 5.7

(Cp2Fe+)(BF4-) 140 cm-1 5.6

(Cp2Fe+)(PF6) 130 cm-1 5.8

Cp2Fe 179 cm-1 4.7


a. Contribution of )22 doubly degenerate mode to the vibrational entropy at 298 K.








ASeto = 2.8 cal mol'K-1 for reaction 2-5. Assuming that the statistical mechanics

calculations provide a good estimate for ASi"(Cp2Fe), this analysis supports the

PHPMS temperature dependence study of reaction 2-5. However, use of solid-state

vibrational frequencies for gas-phase ferrocenium ion may not be valid, therefore the

estimated value for ASi0 may be different from the true value. Gas-phase vibrational

frequencies for metal-ligand modes may be significantly shifted to lower frequencies

relative to the solid-state since crystal lattice stabilization is lost in the gas-phase.


Vibrational Entropy Calculations for the Ferrocene/Ferrocenium Couple


Since only a limited number of published vibrational frequencies for Cp2Fe+

exist, several frequencies for the cation were estimated from the known vibrational

frequency changes of ferrocene and ferrocenium. From known vibrational frequencies

of the ion and neutral that were measured (Table 2-4), the percent change in the

vibrational force constant was calculated. It was assumed that for a specific

vibrational mode, the percent change for an anologous mode which has not yet been

measured would be equivalent. For example, the symmetric ring-distortion mode, 1g28,

for both the cation and the neutral have been measured,68 where U28 is 597 cm- for

ferrocene68 and 471 cm-1 for Cp2Fe+.69 The corresponding force constants are 3.26

mdyne A-' and 2.04 mdyne A-' respectively. Equation 2-11 was used to determine

vibrational force constants from measured frequencies where a is converted to energy,

E, p is the reduced mass for the vibrational mode, k is the vibrational force constant,

and h is Plank's constant75






39
E = (h/2n)*(k/p)" (2-11)



Thus, as demonstrated by the scheme below, an estimate for the vibrational force

constant for the unsymmetrical 034 mode, which is the asymmetric analogue to v28,

was calculated. Note that the v34 and the v28 modes correspond to degenerate low-

energy ligand distortions perpendicular to the principle molecular axis.



k(Cp2Fe) for 128 = 3.26 mdyne A-1

k(Cp2Fe+) for U28 = 2.04 mdyne A1

Ak28 = 0.63 and assume Ak28 = Ak34 (degenerate modes)

k(Cp2Fe) for U34 = 2.97 mdyne A-'

k(Cp2Fe+) for J28 = (0.63 2.97) mdyne A-' = 1.86 mdyne A1

From equation 2-11, 134 = 450 cm-1.



Vibrational entropy contributions for individual frequencies were calculated

from equation 2-12. The value of u is dependent on the frequency of the vibration

and the absolute temperature, shown in equation 2-13. From equation 2-12, entropies

for all vibrational frequencies were calculated for ferrocene and ferrocenium ion.

Table 2-4 lists the frequencies used in the vibrational entropy calculations; literature



Svib = R[(u/eu 1) In (1 e-u)] (2-12)

u = 1.4387(u/T) (2-13)








Table 2-4 Vibrational Frequency Data for Ferrocene and Ferrocenium Cation


No.a Cp2Fe(u)b CP2Fe+(u) Cp2Fe+ Reference

1 3110 3110 Estimated
2 814 805 Estimated
3 1102 1102 Estimated
4 309 304 Ref. 70
5 1255 1258 Estimated
6 c c c
7 1250 1250 Estimated
8 3103 3100 Ref. 69
9 820 782 Ref. 69
10 1110 1110 Ref. 69
11 478 418 Ref. 69
12 3068 3068 Estimated
13 998 1005 Estimated
14 844 833 Estimated
15 1410 1412 Estimated
16 389 398 Ref. 70
17 3077 3100 Ref. 69
18 1005 1017 Ref. 69
19 855 846 Ref. 69
20 1410 1420 Ref. 69
21 492 490 Ref. 69
22 179 135 d
23 3100 3100 Estimated
24 1191 1192 Ref. 69
25 1058 1058 Ref. 69
26 1356 1356 Estimated
27 897 874 Ref. 69
28 597 471 Ref. 70
29 3085 3085 Estimated
30 1189 1192 Ref. 69
31 1055 1055 Ref. 69
32 1351 1351 Estimated
33 885 874 Ref. 69
34 569 450 Estimated

a. Vibrational frequency assignments from reference 68.
b. All Cp2Fe vibrational frequencies from reference 68.
c. Torsional vibration determined as internal rotation.
d. Measured frequency.








sources are included when available. All frequencies for ferrocene used in this work

were measured by Bodenheimer and Low.68

Vibrational entropy calculations based on spectroscopic data assume that

molecular vibrations behave as harmonic oscillators, therefore the displacement force

is equal to the restoring force.55 Atomic systems typically deviate from harmonic

conditions and corrections for anharmonicity must be considered for more complete

calculations. However, the harmonic formula is an accurate first approximation for

Svib quantities.


Translational and Rotational Contributions to the Entropies


The translation entropy of an ideal gas is dependent on the molecular weight

and the temperature of the species. For ionization processes, the molecular weight of

the ion and the neutral are essentially equal, therefore the translational contribution at

temperature T to ASi is 0 cal mol'lK1. In order to report accurate ASi estimates for

Cp2Fe and Cp2Fe+, translational entropies were determined for completeness. The

Sackur-Tetrode equation76 (equation 2-14) was used to estimate Stra values, where



tans0 = R(3/2 In M + 5/2 In T) 2.315 cal mol'K-1 (2-14)



M is the molecular weight, R is the ideal gas constant, and T is the temperature.

Rotational entropies were calculated by assuming that ferrocene behaves as a

linear rigid rotor and that the Cp rings are point masses of 65 a.m.u. Rotational








entropies are dependent on the moments of inertia I of the molecule, therefore by

choosing the Fe atom to lie at the center of gravity of the molecule, the moments of

inertia for Cp2Fe were calculated. Since the Fe atom is defined to lie at the origin of

the cartesian axes, I, and Iy are equal (corresponding to rotations perpendicular to the

principle axis). The rotational moment of inertia along the principle axis was defined

as Iz. Here, the Cp rings where considered as mass with five-fold rotational symmetry

rather than point masses. The product of the three principle moments of inertia yields

a determinant D which is used in estimating Srot (equation 2-15).55 The symmetry

number a for ferrocene is 10 based on D5d molecular symmetry.72 Note that D5h

symmetry, for an eclipsed structure, will also yield a = 10.



Srot = R(1/2 In D x 10117 + In T In o) 0.033 (2-15)



Because of the low energy barrier, the Cp rings of ferrocene freely rotate at

ambient temperatures.72 Therefore, rather than consider rotation of the Cp ligands as a

torsional vibration, the internal rotational entropy was estimated as a restricted internal

rotation. The potential barrier of rotation V = 0.9 kcal mor.72 The moment of

inertia for the system is given as Ir = (z)2/2z, where Iz is the rotational moment of

inertia along the principle axis. Therefore, using the method outlined by Brewer and

Pitzer,55 Sinto for was estimated. The internal rotational entropy for a restricted

rotation is calculated from the entropy of free rotation along the principle axis minus






43

the restriction energy associated with rational barrier. Since Iz for both ferrocene and

ferrocenium are equal, ASint = 0 cal mol'K-1.


Electronic Entropy Considerations for the Ferrocene/Ferrocenium Couple.


The electronic energy level separations of a molecule are usually large, and at

ambient temperatures only the ground state is thermally populated.55 The electronic

partition function q is used to determine Seec values where qe = g, the spin degeneracy

of the ground state. Thus, in the absence of thermally accessible energy levels, the

electronic entropy is given by S = R In(g). In the case of ferrocene, the spin

multiplicity is 1 and Sele = 0 cal mol^K'1. Ferrocenium is a 2E complex, however

the 2E state is spit due to spin-orbit coupling with ~700 cm- separation.77 Since the

split states are thermally accessible in the 298 600 K range, the value of qelec was

evaluated which includes the theraml energy of the split electronic states. The

electronic entropy of Cp2Fe+ is given as Selec,T = EthermT + R In(g), where EthermT is

the energy of the thermal population of the split states at temperature T.


Ionization Free Energies of Ruthenocene and Osmocene


Discussion of the ionization free energies of ruthenocene and osmocene has

been reported previously.16 The observed trend in AGi values for the iron triad is

consistent with vertical ionizations35'78 (IP Cp2Fe < Cp2Os < Cp2Ru).

Lichtenberger and Copenhaven were able to obtain vibrational fine structure for

the first ionization manifold of osmocene.78 An average spacing of 42.1 meV for the






44
spacing of vibrational energy levels in the first ionization manifold for the cation was

used to determine the adiabatic ionization energy for osmocene of 161.1 kcal mo'1.

Assuming the ASi0 for osmocene is equal to ASio(Cp2Fe) estimated from statistical

mechanics, the estimated osmocene alP for Cp2Os is 162 kcal molr1. Thus the two

techniques are within experimental error.


Free Energies of Ionization for Vanadocene, Manganocene, and Nickelocene


The sharpness of the first ionization band in the PES of vanadocene indicates

that the difference in the equilibrium geometries of the ion and neutral are small.37'38

A recent PES of Cp2V assigns the vIP = 155.7 0.1 kcal mol'.16 As the first

ionization manifold is a sharp band (width = 0.19 eV), the aIP is closely approximated

by the vIP. The AGi value from ETE studies is 154.5 1.5 kcal mor1. Assuming

ASi0 is predominately the electronic entropy change for vanadocene (ASvib ~0 based

on the PES) the AHi is estimated at 154 2 kcal mol"1. The estimated AHLi value is

slightly less than the estimated alP but lies within the band envelope.

The AGi for nickelocene is in agreement with PES data; however,35 the AGi

of manganocene is -18 kcal morl lower than the reported vIP values.35,37'38 The

relaxation energy for Cp2Mn is large compared to other metallocenes. Manganocene

exists as a high-spin complex with only a small percentage of the complex in a low-

spin configuration.79 From PES studies the vIP of the low-spin complex has been

assigned (-144 kcal mol1) and is in agreement with the AGiO(Cp2Mn).








The temperature dependence of the Cp2Mn/Cp2Ni ETE reaction couple was

investigated to assess the expected negative intramolecular entropy change for the

Cp2Mn+/0 couple. The van't Hoff plot in Figure 2-3 indicates that the Cp2Mn/Cp2Ni

ETE reaction couple is strongly temperature dependent with a negative ASet0. Details

of the Cp2Mn/Cp2Ni ETE couple have been presented elsewhere.13 The origin of the

large negative entropy change is primarily attributed to ASio(Cp2Mn). The estimated

change in the metal to ring-centroid distance for Cp2Mn/CP2Mn+ is -0.25 A80

resulting in a large change in vibrational and rotational entropy for the Cp2Mn"+/

couple. The Mn-C distance will decrease upon oxidation leading to a negative ASvib.

As spin degeneracy is conserved, ASelec(Cp2Mn+/0) is essentially zero. However, due

to spin orbit coupling of the 3E state of the ion, ASeec for manganocene will be

greater than that of manganocenium ion. This is consistent with a net negative

entropy change expected for Cp2Mn+/0. The rotational entropy change is estimated to

contribute 0.6 cal mol'K1 to ASi based on the estimated change in the Mn-Cp

distance accompanying ionization. Thus estimated vibrational, rotational, and

electronic entropy changes for manganocene are all negative and are consistent with

the negative ASet0 observed for the Cp2Mn/Cp2Ni ETE couple.16


Substituent Effects in Ferrocene Derivative Oxidations


Although extensive photoelectron spectroscopy studies of the ionization

energetic of metallocenes have been reported,35 little is known concerning the effects

substituents have on metallocene thermochemistry. Therefore, various ferrocene






46

derivatives have been studied by ETE methods in order to assess the effect alkylation

of the cyclopentadienyl rings has on the ionization potentials of metallocenes. Free

energy of ionization data for several alkylferrocene complexes have been reported

earlier16 and a AGi value for ethylferrocene is reported here. Values of AGG differ

from other values previously reported13 due to modifications in the AGio values of the

reference compounds.

It is known from PES data that dimethylation or permethylation of

metallocenes lower the ionization energies.35 The vertical ionization energy of

ferrocene (6.88 eV) is -0.2 eV greater than the vIP of 1,1'-dimethlyferrocene35 and

1 eV more endoergic than the vIP of decamethylferrocene.37 Attachment of alkyl

groups on the metallocene rings stabilizes the molecular cation relative to the neutral

compound, thus lowering the free energy of ionization of the derivative relative to the

parent metallocene.32,35 Electron-transfer equilibrium results for alkylferrocene

derivatives studied in the present work are shown in Figure 2-2.

Molecular ionization potentials of organic81 and organometallic compounds32'51

have been correlated previously with Taft substituent parameters. The aliphatic oI

parameters were derived originally for substituted acetic acids, and they have been

used successfully to correlate IP data.51'52'81 For example, a plot of IP for benzene-R

chromiumtricarbonyl derivatives versus oI(R), where R is an attached substituent,

shows a strong correlation line.51 The equation used by Levitt and coworkers for the

Taft analyses is given in eq 2-16, where IP(M-R) and IP(M) are the ionization

potentials for the derivative and the parent compound respectively, oI is the Taft






47

IP(M-R) = aal(R) + IP(M) (2-16)



parameter for R, and aI is the slope of the line. The Taft parameter for H is zero thus

the substituent effects are referenced to hydrogen.5 The slope aI indicates the

sensitivity of the ionization process to the change in substituents.81

A plot employing equation 2-16 for the ferrocene derivatives is shown in

Figure 2-4. The slope of the line is 57 6 kcal mol-1 which can be compared to that

for the (RBz)Cr(CO)3 compounds, a, = 34.6 kcal mol-1, and alkylbenzenes, a, = 109.3

kcal mol-1.51 Therefore, alkylferrocene ionization energies are nearly twice as

sensitive to changes in alkyl substituents as the chromium arene complexes but are

less affected than the alkylbenzenes.51 The AGio data here is consistent with the

results of Matsumura-Inoue and coworkers who correlated PES data of Cp2Fe

derivatives with Taft parameters.32 The differences in substituent effect sensitivity for

various parent compounds can be rationalized by several factors, including proximity

of substituents to the site of ionization.81 Additionally, electronic coupling between

the alkyl a orbitals and the ionized molecular orbital may direct changes in the

ionization potentials with respect to the parent compound.16 It should be noted that

although the observed substituent effects in alkylferrocenes follow the trend of

expected "electron-releasing" ability of the alkyl substituents, the observed sensitivity

of ionization energies to substitution will not necessarily hold for other

metallocenes.16'19 Electron loss originates from different valence molecular orbitals

for the metal complexes.35

























.Q 152

LU




0-
u-
c)



u 145




14E












Figure 2-4


-0.06 -0.04
Alkyl Taft Parameter


Plot of AG, values (kcal mol"1) versus alkyl Taft parameters (or) for
several ferrocene derivatives. Asterisk indicates new Taft parameter for
CH2N(CH3)2.






49

Further, the electron-releasing character of the alkyl groups observed for the ionization

of neutral metallocenes does not necessarily apply in other circumstances. For

example, the substituent effects for electron attachment to neutral metal complexes has

been shown to be different than for ionization process.19'82

With respect to the construction of Taft parameter correlations such as that in

Figure 2-4, it is notable that parameters for substituents with low ionization potentials

can be derived from the data for ferrocene derivative ionizations. For R = CH2NMe2,

the first ionization of the benzene derivative removes an electron from the nitrogen

lone-pair orbital and not from the benzene ring, which would be a more endoergic

process. The ionization of benzene is 9.25 eV,27 which is significantly more

endoergic than the ionization potential of N,N-dimethylaminobenzene (7.69 eV).27

However, in the case of (CpCH2NMe2)(Cp)Fe, the ionization occurs at the same site

as in the other alkylferrocenes, thus allowing derivation of a oI parameter for the

substituent (-0.049 0.013).


Heterolytic and Homolytic Metal-Ligand Bond Disruption Enthalpies for
Metallocenes and Metallocenium Ions


Bond disruption enthalpies for several metallocenes have been reported

previously16 therefore only a brief discussion concerning M-Cp bond enthalpies will

be give here. Thermochemical cycles have been used to derive estimates of M-Cp

bond disruption enthalpies.10'16 Since AH values depend on ionization energetic

data, M-Cp bond enthalpies reported here differ from previously reported13 values due

to more accurate free energy of ionization data for the metallocenes and more accurate







assessment of auxiliary thermochemical data.16 Additionally, more detailed error

analysis was performed for the heterolytic and homolytic bond disruption enthalpies

and in most cases error limits were found to be lower than previously reported.13


Application of Thermochemical Cycles to Derive Estimates for Metallocene and
Metallocenium Bond Disruption Enthalpies


Thermochemical cycles for ionization processes of the metallocenes and their

corresponding ions used to derive bond enthalpies data are shown in Figure 2-5. The

bottom portion of the figure is a cycle used to derive solvation energetic for the

metallocenes which will be discussed later. Similar cycles have been used by

Buckingham and Sargeson to derive crude estimates for metal complex thermodynamic

quantities.83 Related thermodynamic quantities have recently been derived for

coordination complexes and complex ions by use of analogous thermochemical

cycles.84 In Figure 2-5, the term AX represents any thermodynamic function

therefore entropy, enthalpy or free energy data can be incorporated in the cycles to

derive thermodynamic values.

In this work, average heterolytic bond disruption enthalpies, half the AH

expressed in reactions 2-17 and 2-18, will be denoted as AHheto. Homolytic metal-

ligand bond cleavage will be denoted as AHhom0 and represents AAH for reactions

2-19 and 2-20. From the thermochemical cycles used here only average bond



Cp2M(g) M2+(g) + 2Cp'(g) (2-17)

Cp2M(g) M3+(g) + 2Cp-(g) (2-18)






51

Cp2M(g) -> M(g) + 2Cp(g) (2-19)

Cp2M+(g) M+(g) + 2Cp(g) (2-20)



disruption enthalpies can be obtained. Further, enthalpy values for M-Cp bond

cleavage for consecutive ring cleavage will not be equivalent. For heterolytic

cleavage, removal of the first Cp- ligand will be less endothermic than removal of the

second ligand due to an increase in the electrostatics between the positively charged

metal center and the anionic Cp ligand. Similarly, AHhom for the first M-Cp

homolytic cleavage will not be equal to A/homo for the second M-Cp cleavage.

Homolytic bond energies for ferrocene have been studied by pyrolysis

techniques.85 The activation energy for the decomposition of ferrocene, reaction 2-21,

was monitored and based on the measured Ea, the first bond dissociation bond



Cp2Fe -- CpFe + Cp (2-21)



enthalpy was estimated to be 95 kcal mol .85 Consequently, removal of the second Cp

ring is less endothermic by approximately 30 kcal mol'1 than removal of the first

ligand. Faulk and Dunbar have used photodissociation methods to arrive at a value of

85 7 kcal mol"1 for the first homolytic cleavage of Cp2Fe+. Therefore, the AIhomo

for the second cleavage is more endothermic than the first dissociation. Increased

electrostatic attraction for the second Fe+-Cp cleavage accounts for an increase in the

second A-hom0 for Cp2Fe+.86





















M2+ (g) + 2Cp- (g) A


AX*het(Cp2M,g)


M (g)





AX*sub(Cp2M,g)
Cp2M (s)

AXsoI(CP2Ms)
C


XXi(M2+,g)


M3+ (g) + 2Cp- (g) + e-


AXohet(CP2M+,g)


+ 2Cp- (g)


AXi(Cp2M,g)


AXsoiv(Cp2M,g


p2M (soln)


AX*i(Cp2M,soln)


M+ (g) + 2Cp- (g) + e-

AXohom(Cp2M+,g)

Cp2M+ (g) + e-

AXsov,(Cp2M+,g)


Cp2M+ (soln) + e-


Thermochemical cycles used to determine bond disruption enthalpies
and differential solvation free energies for metallocenes. The upper
portion of the cycle yields estimates for the average homolytic and
heterolytic bond disruption enthalpies for M-Cp cleavage. Comparison
of AGij(g) and AGi(soln) in lower portion of the cycle ields estimates
of differential solvation energies, AAGsolv, for Cp2M+ couples.


Figure 2-5








Combining Free Energy and Enthalpy Data in Thermodynamic Cycles


Free energy of ionization data has been used to derive heterolytic and

homolytic bond disruption enthalpies for metallocenium ions. The validity of

combining AGi data for the metallocenes with enthalpy data referenced at other

temperature, i.e. 0 or 298 K, is dependent on the accuracy of substituting AGi,350

values for AHi,2980. This approximation relies on the premise that the entropy of

ionization is small with respect to the larger values of AH at temperature T.

Manganocene is expected to have the largest | ASi j of the metallocenes studied in

this work.16 Assuming |ASi = 12 cal mol'K-1, the entropy contribution to the free

energy of ionization at 350 K is 4 kcal mol"1. Furthermore, even with such a large

entropy change, the error estimated for manganocene is not expected to exceed 2 kcal

mol1 per bond. Generally, the assumption that AGi,3500 = AHi,2980 for the other

metallocenes is acceptable since entropy contributions will be small with respect to the

larger values for the free energy of ionization. When compared to the absolute values

for heterolytic and homolytic dissociation, the error introduced by assuming AGi,350

AHi,298 will small, ca. 2-3% for the homolytic bond enthalpies and less than 1% for

the heterolytic enthalpies.

Heat capacity corrections are also expected to be small on going from 350 to

298 K. The heat capacity terms for ferrocene has been determined by statistical

mechanics and the difference in the AHi values is only 0.1 kcal mol-1 from 350 to

298 K.16








Heterolytic and Homolytic Bond Disruption Enthalpies for Metallocenes and
Metallocenium Ions


Bond disruption enthalpies for the Cp2V, Cp2Mn, Cp2Fe, and Cp2Ni have been

reported previously.13 However, as mentioned earlier, values reported here have been

refined due to more accurate AGi values and auxiliary data. Derived bond enthalpies

for the metallocenes are presented in Table 2-5. Auxiliary data16 used in the

thermochemical cycles are presented in Table 2-6. Error limits for homolytic and

heterolytic dissociations take into account errors in the heats of formation of the

neutral27 and ionic species16 and errors in the free energies of ionization. Bond

enthalpies for ruthenocene, osmocene, and the substituted metallocenes are not

reported due to lack of reliable thermochemical data for AHf of the neutral

organometallics and the alkylated cyclopentadienyl compounds.


Differential Solvation Free Energies for Metallocene Redox Couples


Differential solvation free energies, AAGsolv", for several metallocene redox

couples have been determined through the application of thermochemical cycles

(Figure 2-5) by combining Eh data at 298 K to AGi,3500 values. As mentioned earlier,

because of the relatively small entropy effects for the metallocene Cp2M+/0 oxidation

couples, the errors introduced by substituting AGi,3500 for AGi,2980 are expected to be

less than 1 kcal mol"1. Values for E have been used to derive AGio(Cp2M)(soln);

from Figure 2-5, values of AAGsolv have been derived from the lower thermochemical







Table 2-5 Mean Bond Disruption Enthalpies for Some Metallocenes.


CP2M Hhet[MC2+P-Cp-I


FHhet[M3+_Cpia


L'Khom[M-CP.I'


AHhom[M+-Cp.]a


V 303 3 563 4 95 2 95 3

Mn 286 3 604 5 59 2 74 4

Fe 318 3 593 4 79 1 91 3

Ni 326 3 659 4 67 2 83 3


a. Units are kcal mol-1







Table 2-6. Auxiliary Thermochemical Data Used in Thermochemical Cycles.


Process V Mn Fe Ni

AHf[Cp2M]b 49 2a 66 2 58 1 85 1

AHIfM]c 123 2 67 2 99 2 103 2

AIf[M+c 278 2 238 2 281 2 279 2

AHIy[M2+4 616 2 599 2 654 2 698 2

AHf[M3+]c 1292 2 1376 2 1361 2 1509 2

AHo[Cp]b = 58 1

AHfo[Cp-]b= 19.6 4

a. Units are kcal mol-.
b. See reference 27.
c. See reference 8.






57

cycle. A negative value of AAGsolv0 represents decreased exoergicity for the reduction

of a metallocenium ion in solution compared to the gas phase. An analysis of the

estimation of absolute electrode potentials for redox couples in solution has been given

earlier,54 and a similar approach has been used for the derivation of AAGsolvo

quantities in this work. Specifically, a value of 4.44 V has been used for the absolute

potential of the standard hydrogen electrode, ENHE, and no corrections for liquid

junction potentials have been applied to the E data for the metallocenes.54

Thermochemical cycles were used in the derivation of the absolute potential of the

normal hydrogen electrode. The largest source of error introduced in the derivation of

ENHE was the free energy change for the solvation of the proton, AGsolv0, which can

not be measured directly.87 The value used for AGsolv(H+) is based on

electrochemical experiments and the reliability of this value has been discussed

elsewhere.54'84 In addition, the stationary electron convention is used for both the gas-

phase and solution thermochemistry, although near 298 K the thermal electron

convention yields similar results for AGio values.88

The electrochemical E values used in the estimates of AAGsolvG values are

shown in Table 2-7. Equation 2-22 was used to derive estimate for AGi(soln) for the

metallocenes. The value Er is the potential of the reference electrode

relative to the normal hydrogen electrode, n is the number of electrons transferred in


AGi(Cp2M)(soln) = -nF[(E(CP2M) + Ere + ENHE0]


(2-22)










Electrochemical E Data and Differential Solvation Energies for Some
Metallocenes Cp2M+/o couples.


Cp2M+/o

V

Cr

Fe

Co

Ni

Ru

Ru

Os

Os

Mn


E a (solvent)

-0.55c,d (THF)

-0.67c (CH3CN)

0.31c (CH3CN)

-0.94c (CH3CN)

-0.09c (CH3CN)

1.03' (CH2C12)

0.78d'g (CH3CN)

0.86' (CH2C12)

0.75d,g (CH3CN)

(0.13)h


a. Values reported in volts using 0.1 M Bu4NPF6 as supporting electrolyte against
SCE, except ruthenocene in 0.1 M Bu4TFPB against Ag/AgCl and osmocene in
0.1 M Bu4NBF4 against SCE.
b. Units are kcal mol1. Estimated error limits 4 kcal mol1.
c. See reference 89.
d. Irreversible oxidation potential.
e. Estimated from AGji values. See Chapter 5.
f. See reference 90.
g. See reference 92.
h. Estimated from data in Figure 2-6 and reported against SHE.


Table 2-7.


AGio(soln)b

95

92

115

86

106

131

126

127

125

(105)h


-AAGsolvb

60

36e

38

38e

38

33

38

34

36

(38)h








the electrochemical process (le' oxidations for the metallocenes) and F is Faraday's

constant. Most of the E values were obtained from a single literature source89 and

were measured under common experimental conditions. The solvent is acetonitrile for

all quoted E1 values except for vanadocene (THF) and ruthenocene and osmocene

(CH2Cl2).90,91 Values of E for ruthenocene and osmocene are also reported in

CH3CN.92 Table 2-7 presents the derived differential solvation energies for the

metallocenes and the corresponding solution free energies of oxidation for comparison

to the gas-phase ionization energies. Differential solvation free energies were

calculated from equation 2-23, which originates from the thermochemical cycles in

Figure 2-5. A plot of free energy of ionization data versus the first transition row



-AAGsolvo (Cp2M+/0 = AGi(Cp2M)(g) AGi(Cp2M)(soln) (2-23)



metallocenes is presented in Figure 2-6 and serves to demonstrates the periodic trends

for AGio(Cp2M) in solution and the gas-phase. The average value for AAGGsolv for the

first transition row metallocenes is -38 2 kcal mol"1, excluding vanadocene. The

AAGsolvo values for cobaltocene and chromocene were based on AGi values

determined from FTMS ETE studies and are discussed in Chapter 5. The obvious

exception to the observed trend in differential solvation energies (Figure 2-6) is noted

for Cp2V oxidation,89 which has a AAGsolvo value more negative by -20 kcal mol-1

relative to the other first transition row metallocenes. This additional stabilization of

the cation can be attributed to inner sphere coordination


















IS-
180



160-


-=
E

.140-


W


0

C
o 120
o
c

8 100'
0
"a


80-



60 2-








Figure 2-6


r Mn e o I
Cr Mn Fe Co Ni


Plot demonstrating the periodic trend of ionization energies for the first
transition row metallocenes. Gas-phase data (filled squares) include
AGivalues from this work. Solution AGi values (crosses) were
determined through thermochemical cycles. The solvent is CH3CN
except for Cp2V in THF. An estimate of AGi(soln) for manganocene
(versus SCE) is included (open square).


Gas-Phase


-








of solvent following oxidation.89 From the trends in Figure 2-6, a prediction can be

made for the unknown Ei for Cp2Mn, -0.13 V vs. SCE.

Gieger has reported that electrochemical oxidation of CpV is quasi-reversible

in tetrahydrofuran, TIF, which is consistent with there being no significant structural

changes in going from a d3 to d2 metal complex.9 However, the correlation of

AGi(g) versus AGi(soln) clearly indicates that Cp2V lies outside the fit; thus, a

structural variation can not be ruled out. A structural change, similar to that shown in

Figure 2-7, could explain the lack of correlation for Cp2V. It can be rationalized that

in an attempt to increase the electron density around the electrophilic metal



7+


V 1





Figure 2-7 Structure of an 18e"1 vanadocenium-THF complex


center, polar solvent molecules coordinate to the metal complex. The addition of two

THF molecules, for example, would lead to the formation of an 18 e' cationic

Cp2V-2THF complex.

A comparison of AAGsolv0 results to values predicted by dielectric continuum

theory suggests that the solvation thermochemistry of Cp2M+/0 couples can be

adequately modelled by the Born charging model.54 The Born equation determines






62

the change in electrostatic free energy, AGei0 when a charge on an electrostatic sphere

of radius ref is transferred in a vacuum to a sphere of equivalent volume in a solution

of dielectric constant D. For equation 2-24, the Born equation, z denotes fundamental

charge of the ion (here 1+). The definition of the Born equation is directly

comparable to the concept of differential solvation energy defined by equation 2-23.54



AGeli = (-166z2/reff)(1 -1/D) kcal mol-' (2-24)



The Born model neglects the actual work required for an ion to pass from a vacuum

through the solvent barrier. However, this work is usually quite small compared to the

values for the differential solvation free energy.55

From crystallographic data, the radii for Cp2Fe and Cp2Ni are estimated to be

3.9 and 3.7 A respectively.72 The reff value obtained from the Born equation for a

AGel = -38 kcal mol-1 in acetonitrile is 4.3 A. From another point of view, the

structural model radius (3.9 A) predicts a AAGsolv, value of -41 kcal mol"1. This

close agreement between the structurally estimated radii and the thermochemical radii

is consistent with relatively small specific interactions between solvent and metal

complex as well as the compact structure of metallocenes. The same conclusions were

obtained by Krishtalik et al. who used a AGi value of ferrocene based on the

photoelectron spectrum.93 For comparison, in the tris(acetylacetonate) metal

complexes, where polar solvent molecules can interpenetrate between the chelating

bidentate ligands, the experimental solvation energy is approximately twice the value






63

predicted from the structural model. The Born model predicts a value for rff of 2.9 A

for Ru(acac)30/1 based on a value of 57.5 kcal mol"' for AAGsol .84 However, the

crystallographic radius (maximum radius taken from Ru metal center to furthest ligand

proton) for the ruthenium complex is -6 A. A shortcoming of the Born model is that

it assumes the charge is evenly distributed over the entire sphere which is not

necessarily true, especially for large metal complexes

Equation 2-24 predicts that AAGisolv will be increased by -3 kcal mol"1 for reff

= 3.9 A when acetonitrile, D = 36, is replaced by methylene chloride, D = 9. From

the reversible potentials for ferrocene, ruthenocene, and osmocene in CH2Cl2 given by

Hill et al.,90 lower solvation energetic for AAGsolv values (-34 1 kcal mol-1) are

estimated for the three couples. However, the difference in the AAGsolvo values

derived in the different solvents are in good agreement and are consistent with

experimental AAGsolv' values determined through thermochemical cycles.


Conclusions


Free energies of ionization have been determined for a number of gas-phase

metallocenes. These ETE data complement and extend information on the oxidation

energies of metallocenes obtained previously by electrochemistry34 and photoelectron

spectroscopy.35 Further, the AGi values for the metallocenes are in agreement with

the vertical ionization energies measured by PES by several groups.37,38,78 The free

energy of ionization for manganocene is in agreement with the vIP for the low-spin

complex. The temperature dependence of the Cp2Ni/Cp2Mn couple is consistent with







64

a large negative entropy change for the manganocene ionization process. Comparison

of ETE data with PES spectra for Cp2V and Cp2Os yield the same values for adiabatic

ionization potentials within experimental error.16'78 Overall, the good agreement of

AG1i data with PES data indicates that large experimental errors do not exist for the

ETE studies. With the exception of vanadocene, correlation AGi values in the gas

phase and solution for the first transition row metallocenes is linear. Inner-sphere

solvent coordination can be used to rationalize the disparity of the vanadocene couple

relative to the other metallocenes.89 An estimate for E for manganocene has been

made based on E values for the other first transition row metallocenes.

Detailed investigations of the ionization free energy for ferrocene support a

value of AGio = 153.1 1.5 kcal mol'1. Temperature dependent ETE studies on

selected equilibria have established experimental enthalpy and entropy of ionization

values for ferrocene. Although the present data establish a relatively large value of

ASio(Cp2Fe), statistical mechanical analyses and previous PHPMS results43 suggest

that a smaller ASi0 value (-4-5 cal mol-1K') is appropriate. The positive entropy

change for ferrocene ionization can be attributed to roughly equal contributions from

intramolecular vibrational entropy changes and changes in electronic entropy. The

vibrational contribution to the entropy change accounts for over 50% of the total ASi'

at 298 K for the Cp2Fe+0 couple. Temperature dependent investigation of the ETE

reaction couple EtFc/Cp2Fe indicate that the ASet0 is -0 cal moll^K1 which indicates

that the ASi for ethylferrocene and ferrocene are equivalent. The ETE reaction couple









Cp2V/Cp2Fe indicates that the AGio of ferrocene is less than vanadocene which

supports the lower AGio(Cp2Fe) value.

The AGi data have been incorporated into thermochemical cycles to allow

estimations of bond disruption enthalpies for selected gas-phase metallocenium ions.

Estimates of differential solvation energies for several Cp2M+/0 couples have also been

derived from thermochemical cycles. An average value of AAGsolvo for the

metallocenes is -38 2 kcal mol"1 with CH3CN as the solvent. Solvation of a gas-

phase metallocene decreases the ionization energy by a relatively constant amount

(-1.6 eV with acetonitrile as solvent). Solvation is therefore secondary to metal

ligation in determining the potential of Cp2M+/o couple relative to the ionization of the

corresponding M2 (g) ion (ligation of M2" by two Cp- ligands reduces the ionization

energy by ca. 13 1 eV).

The Born equation, used for deriving estimates of differential solvation

energies, has been applied to the metallocenes. Values of AGelO are in agreement with

AAGsolvO(Cp2M+/O) values which demonstrates that the metallocenes have compact

molecular structures. Specific solvent interactions such as inner-sphere coordination of

solvent are minimal for the metallocenes.


Experimental Methods


Electron-Transfer Equilibrium Studies


Electron-transfer equilibrium studies were performed by using a Nicolet FT/MS

1000 Fourier transform ion cyclotron resonance mass spectrometer as previously









described.15'20 Briefly, the pressure of each compound was adjusted to establish a

workable pressure ratio to allow equilibria to be monitored. The time dependence of

parent ions formed from neutral molecules of known partial pressure was monitored as

the molecular ions underwent electron transfer with neutrals. Reactions were typically

followed to ca. 5 or more seconds. Apparent equilibrium was generally attained in

less than two seconds.

Reference compounds were sublimed into the FTMS high vacuum chamber

through a precision leak valve. The vapor pressure of most metallocene samples was

sufficient to allow for direct introduced into the high vacuum chamber through a

second leak valve. Ruthenocene, Cp2Ru, and osmocene, Cp2Os, were introduced by

using a heated solids probe positioned adjacent to the reaction cell. The FTMS

reaction cell was typically 350 K as measured by an Omega RTD thin film detector.

Positive ions were produced by electron impact at 9-12 eV with beam times ranging

between 5 and 25 ms. Ionization of metallocenes and organic occurs with some

fragmentation of the molecular ion videe infra); other unwanted ions are formed by

ion-molecule reactions. Prior to study of electron-transfer reactions, several ion

ejections were required in order to select only the parent ions.

Since reactions were followed for at least 5 s, all ions formed as a result of El

were assumed to be effectively thermalized through ion-molecule collisions. At a

neutral pressure of 10-6 Torr at 350 K, a typical metallocenium ion will undergo ca.

30 collisions s-1, which is believed sufficient to remove much of the excess rotational

and vibrational energy present due to the ionization process.14'16 Approach to









equilibrium was followed from endoergic and exoergic directions. Prior to reaction,

one of the parent ions was ejected from the reaction cell and the population change of

both parent ions was monitored at set time intervals. Equilibrium was deemed to have

been achieved when the ratio of the two parent ion populations reached a constant

value within experimental error.

Partial pressures of the various neutrals were determined by using an ion gauge

calibrated with an MKS baratron capacitance manometer (in the 10-5 torr range)

extrapolated to experimental conditions. In order to approach dynamic pressure

equilibrium throughout the vacuum chamber, the 300 L s-1 pumping speed of the

diffusion pump connected to the high vacuum chamber was reduced to ca. 75 L s1.

Neutral gas pressures were calibrated for all reactants in open (75 L s"1) and closed

(no pumping) systems. It has been shown that partial pressure is independent of

neutral gas leak rate. A calibrated ion gauge connected to a Granville-Phillips

controller was positioned at the site of the reaction cell with the magnetic field off,

thus providing a field free vacuum system. The pressure measured at the middle of

the vacuum chamber where the reaction cell is located, was equivalent to metallocene

pressures determined at the remote ion gauge following pressure calibrations.94


Temperature Dependence Studies


The temperature dependencies of electron-transfer equilibria were investigated

by using a customized cell heater designed to heat a 1" x 1" x 1%" analyzer cell. The

heater consisted of two Macor plates (1" x 2" x 14") attached to the long sides of the







68
reaction cell. Macor sheets (6" x 6"x 14") were purchased from Astro Met Inc. Ni-Cu

wire (0.015" diameter, purchased from Omega Industries) was wrapped around the

external Macor plates and was resistively heated by using an Omega digital

temperature controller (maximum output current (~5 A) resulted in temperatures of

-520 K). Cell temperatures were measured using an Omega RTD thin film detector

fastened to the analyzer cell. Additionally, the entire high vacuum chamber was

heated by using the vacuum bake-out system in order to minimize radiative

temperature loss to the vacuum chamber walls.

Following the measurement of Keq at a set temperature, the cell heater and

bake-out were allowed to cool to a lower temperature and the entire system was

allowed to equilibrate at the new temperature for 30 min. Experimental

reproducibility was tested by following the temperature dependence of K as the

reaction cell temperature was increased from 350 to 500 K.

The cell temperature was measured before and after each reaction and usually

fluctuated 2 K during a single experiment. Typically, reactions were repeated three

times at a single temperature. Linear regression and statistical analyses of the all

measured equilibrium constants provided error limits at the 95% confidence level for

reported AHet and ASetO values.


Metallocenes and Reference Compounds


Metallocenes were purchased through Strem Chemicals except for ferrocene

and ruthenocene (Aldrich). No further purification was required except for Cp2Mn









which was resublimed prior to use. For temperature dependence studies, a sublimed

sample of ferrocene was used. Organic reference compounds were purchased from

Aldrich except N,N'-diethyl-p-toluidine (Alfa Chemicals). A sample of 1,1'-

bipyrrolidine40a was donated by Professor Stephen Nelsen from the University of

Wisconsin. Organic reference compounds were used without further purification.

Liquid samples were degassed through several freeze-pump-thaw cycles prior to use.


Far-infrared Spectroscopy.


Ferrocene salts in Table 2-3 were prepared according to literature procedures.95

The hexafluorophosphate and the tetrafluoroborate salts were prepared by dissolution

of pure ferrocene in concentrated sulfuric acid, followed by dilution of the dark blue

solution with water. The solution was filtered and then an aqueous solution of Bu4NX

was added, where X = PF6 or BF4. The precipitate was filtered and washed with

water until the washings were clear. The chloride salt was obtained by distilling

ferrocene in concentrated HCI for several days. The blue precipitate was filtered and

washed with water. Infrared spectroscopy in the 400 3200 cm-1 region confirmed

the compounds to be ferrocenium salts.69-71 The samples were prepared as dilute 13

mm polyethylene pellets. Far-infrared spectra were recorded using a Bruker IFS 113V

spectrophotometer in the 50 750 cmn' spectral region.












CHAPTER 3
SUBSTITUENT EFFECTS IN THE GAS-PHASE AND SOLUTION IONIZATION
AND ELECTRON ATTACHMENT ENERGIES OF ALKYLNICKELOCENES.


Introduction


Alkyl substituent effects have been studied extensively in organic chemistry

with emphasis towards understanding acidities,30a basicities,30b and reactivity of

organic systems.5'53 The correlation between structure and reactivity of alkyl

substituents has led to detailed explanations and an interpretation of chemical

reactivity and chemical equilibria.2-5 Methods that correlate structure-reactivity

relationships allow for quantitative interpretation of various electronic perturbations

of substituents relative to a parent molecular-frame. Subsequently,5 chemical

transformations for yet unknown species can be determined based on known

substituent effects.

Gas-phase investigations have been very effective in evaluating intrinsic

substituent effects.16'19'20'32'37'91 Specific electronic effects may be masked in

condensed-phase; therefore, intrinsic effects observed in solution studies may be

significantly modified or even reversed relative to gas-phase studies.5 Additionally,

the results may vary from solvent to solvent. Nonpolar solvents may yield

approximations of intrinsic substituent electronic effects, while strongly coordinating








solvent may inhibit substituent electronic effects altogether.5'53 To truly resolve

substituent effects in various solvents, gas-phase data must be used as a reference.

Relatively few gas-phase studies of alkyl substituents effects of metal

complexes have appeared,1618'35'91 and little is known concerning the relative alkyl

effects in the thermochemistry of metal complex redox couples.32,34 Comparisons of

the gas-phase and solution results address solvent effects of metal redox couples, and

allow for estimates of differential solvation energetic of metal complexes to be

made.15'16 Additionally, such studies help extend parameterization schemes derived

from organic systems to the area of inorganic chemistry.

From photoelectron spectra35 and electrochemical studies,32'34 it is commonly

assumed that alkyl groups on coordinated ligands are electron-donating in metal

complex redox processes. Specifically, alkyl groups stabilize the oxidized form of a

complex relative to the reduced species. In order to further explore and understand

alkyl substituent effects for transition metal compounds, gas-phase electron-transfer

equilibrium studies have been performed, by using FTMS,4447 to determined free

energies of electron-attachment (AGa0) and ionization (AGi0) for a series of alkylated

nickelocene complexes.19 The results indicate, similar to organic systems,2,3 that alkyl

groups are not always electron-donating in organometallic redox processes.19 Models

that include polarizability effects, in addition to more traditional inductive effects,

must be used to interpret the data.5'52'53

Nickelocene is a useful parent compound for these studies because it forms

stable anions and cations in the gas-phase18'19 and solution.34,89 Furthermore,








oxidation and reduction of nickelocene involves the same elg* set of molecular

orbitals.96 A two-electron oxidation process for negative nickelocene ion is shown in

Figure 3-1 with accompanying molecular orbital diagrams of the 3d valence orbitals

for the Cp2Ni+/" complexes. A one-electron model suggests, because the same

molecular orbital of nickelocene is both oxidized and reduced, perturbations in the

orbital energies due to alkyl substitution on the Cp rings are expected to be similar for

the ionization and electron attachment processes. That is to say, the difference in the

ionization energies and electron affinities of an alkylated complex relative to

nickelocene should be comparable.36 However, this is not the case and an a more

flexible model must be used to explain the trends in the experimental data.

Solution phase redox studies were performed for comparison to the gas-phase

data. Differential solvation energies for some alkylated nickelocene complexes have

been derived from thermochemical cycles. Values for AMG 0so for the Cp2Ni+/0 and

the Cp2NiO/- couples are discussed.


Electron-Transfer Equilibrium Studies Involving Negative
and Positive Alkylnickelocene Ions


The electron-transfer equilibrium method has been discussed previously.16,18

Procedures for determining free energies of ionization and free energies of electron

attachments are similar. The equilibrium constants were determined for the reactions

shown in equations 3-1 and 3-2 where RCp and R'Cp represent alkylated

cyclopentadienyl ligands and X, for these examples, denotes a reference compounds










Cc2 c2


-H-
44-
Il--H--


A- I


-1-4


Ni
zcN


9-i-


Figure 3-1 Molecular orbital diagrams for nickelocene anion, nickelocene, and
nickelocene cation.


'Cc;D
Ni
4






74

with known a AGiO or AGa0 value. Ion intensities and partial pressures were measured

directly during the ETE experiment. Thus, the equilibrium constants and reaction free

energies for the reactions 3-1 and 3-2 can be determined (see equation 2-3).



(RCp)(R'Cp)Ni + X+ = X + (RCp)(R'Cp)Ni+ (3-1)

(RCp)(R'Cp)Ni + X- = X + (RCp)(R'Cp)Ni' (3-2)



Free energy ladders for gas-phase electron-transfer equilibria studied in this

work are shown in Figures 3-2 and 3-3. Derived AGi and AGa (electron attachment

free energy) values are referenced at -350 K as measured by an RTD thin film

detector. The AGi, values, for the process (RCp)(R'Cp)Ni -- (RCp)(R'Cp)Ni+ + e-, in

Figure 3-2 are anchored to the AGi values of nickelocene, manganocene, and

bis(benzene)chromium. The AGi values for Cp2Mn and Cp2Ni have been reported

previously.16 Electron-attachment free energies (for the process M + e- -- M-) in

Figure 3-3 are anchored to the AGa0 value of azulene and nickelocene.18

The AGi value for Bz2Cr was based on the photoelectron spectrum which has

an extremely sharp first ionization band with a peak maximum at 5.47 eV.19 Because

of the sharpness of the band, the vertical ionization energy closely approximates the

adiabatic IP (see Figure 1-1B). The AGio(Bz2Cr) = 125.7 kcal mol1 was estimated by

assuming alP = vlP = 126.1 kcal mol"1. An estimate for ASi0 was determined by

assuming only the electronic entropy contribution was important. Thus,

ASvib ASroto ~ 0 cal mol'1 K1. The electronic entropy change associated










CN-N 145.6 1



C2Mn 142 1.3
Cp2Mn 142.5 -t1---


143.8


Ni


141.5 EthylNc


139.6 (MethylCp)2Ni


138.2 (EthylCp)2Ni


136.4 (t-butylCp)2Ni


125.6


121.2 Cp*2Ni


Figure 3-2 Electron-transfer equilibrium ladder for ionizations for several
alkylnickelocene complexes for the process M -4 M + e'. Values of
AGio ( 1.5 kcal mol 1) for the nickelocene complexes are to the right
of the ladder and AGeto values for individual ETE reaction are adjacent
to the arrows. The AGet value for the (MeCp)2Ni/(t-BuCp)2Ni couple
is not within the expected 0.5 kcal mol-1 experimental error limit.


Cr

C]3


































CO 17.3


Ni

21.2 t-butyl

(EthylCp)2Ni
20.3
20.2 (EthylCp)CpNi

19.7 CpgNi


19.0 (MeCp)2Ni


Electron-transfer equilibrium ladder for electron-attachments for
alkylnickelocene complexes for the process M + e" -4 M'. Values of
AGa, ( 1.5 kcal mol1-) for the nickelocene complexes are to the right
of the ladder and AGetC values for individual ETE reaction are adjacent
to the arrows. The AGa for Cp*2Ni is an estimated value.


t-butyl


Figure 3-3








with an 1A to 2A transition is ASelec = R In 2 = 1.4 cal mol"' K-1. Further details

concerning ASi and AGio of Bz2Cr are discussed in Chapter 5.

Electron-transfer equilibrium reactions were repeated several times to insure

reproducibility. Cross checks were performed when possible to check the internal

consistency of the derived AGi and AGa0 values. The experimental uncertainty in the

individual electron-transfer equilibrium reactions is 0.5 kcal mol-1. Electron

attachment and ionization free energies are reported with 1.5 kcal mol"1 error due

largely to errors in the AG values of the reference compounds.

Nickelocene is expected to have the highest electron affinity of the first

transition row metallocenes since it has the lowest reduction potential.89 This is

further substantiated in the observation that no other metallocene forms negative ions

in the gas phase by low energy electron impact or chemical ionization.

Decamethylnickelocene could not be brought to electron-transfer equilibrium with any

reference compound. Further, Cp*2Ni- was not observed from electron impact or

chemical ionization. Attempts to ionize Cp*2Ni with electron-transfer reagents such as

azulene (AGa = 17.3 kcal mol-1)27 and C6F6 (AGa0 = 12 kcal mol-1)27 were

unsuccessful. The electron affinity of Cp*2Ni was estimated from the difference in

the AGao values of (MeCp)2Ni and nickelocene by assuming that the electronic effect

of the methyl groups is additive. The AGao value of Cp*2Ni puts it at the bottom of

the equilibrium ladder for compounds that have had AGa values determined from the

electron-transfer equilibrium method.42








Results of the ETE studies for the cations and anions are presented in Table

3-1. The AGi value for nickelocene is included as a reference for the alkylated

complexes. The AGao for nickelocene was determined from ETE studies with

azulene.18 Nickelocene was brought to equilibrium with azulene eight times from

both exoergic and endoergic directions and therefore serves as a second reference for

ETE studies of the negative ions. It is worth mentioning for historic reasons that the

author used the azulene/nickelocene negative ion couple as a training project to learn

how to operate a Nocolet FT/MS-1000. The not-so-serendipitous electron-transfer

equilibrium that ensued developed into a cavalcade of valuable experiments and the

present dissertation.

It is clear from the data in Table 3-1 that increasing the size and number of

alkyl substituents for nickelocene decreases the free energy of ionization. A similar

effect was observed for the alkylferrocene derivatives.16 In contrast to the ionization

energy data, trends in the AGa values do not consistently reflect an increase in an

"electron-donating" effect. Moreover, the ethyl and t-butyl groups lead to an increase

in the electron affinity relative to H on the Cp rings. However, two methyl groups

lower the electron affinity relative to nickelocene. The shifts in the AGa values

relative to nickelocene are in general small, but appear to be significantly larger than

the error estimated for the electron-transfer free energies (-0.5 kcal mol-) derived

from electron-transfer equilibrium experiments. The absolute free energies have larger

errors as mentioned earlier of 1.5 kcal mot1.








Table 3-1 Free Energies of Ionization and Electron


No. L L'

1 Cp Cp

2 MeCp MeCp

3 EtCp Cp

4 EtCp EtCp

5 t-BuCp t-BuCp

6 C5Me5 C5Me5

a. Units are kcal mol1.
b. Estimated error in absolute values
c. Estimated value.


AG.oab


AGioa,b

143.8

139.6

141.5

138.2

136.4

121.2


1.5 kcal mol1.


-AGaOab

19.7

19.0

20.2

20.3

21.2

(~16)c


Attachment







80
Note that by convention, electron affinities are expressed as positive values (the

affinity of an electron to be attracted to the nucleus of an atom) although they

represent an exoergic property. For application in thermochemical cycles, -AGa0 and

-AHa values are incorporated as negative values,15 however in the discussion the

electron affinities and free energies of electron attachment, the negative sign is

dropped.

An increase in AGa0 for larger alkyl groups relative to R = Me is a well

documented effect for organic systems.98'99 For example, the electron affinities of

alkoxy radicals, RO*, increase in the order R = Me < Et < n-Pr < t-Bu.98 However,

in solution the trends are reversed due to solvent effects.5 The solution acidity of

methanol is greater than that of ethanol, but in the gas phase, ethanol is more acidic

than methanol and even water.100 For p-benzoquinone (BQ) derivatives, the electron

affinity of 2,6-di-tert-butyl-BQ is -1 kcal morl greater than that of 2,6-dimethyl-BQ.

The AGa values for the series of methylated benzoquinones compounds decrease

monotonically by -2 kcal mol"1 per methyl group from the methyl to the tetramethyl

derivative.27 Although the alkyl effects for these systems are quite subtle relative to

alkyl effects for positive ions, the gas-phase results demonstrate that large alkyl groups

can stabilize anions in simple saturated and conjugated systems.100 0' However,

methyl groups tend to destabilize the electron affinities. The lower electron affinity

for (MeCp)2Ni relative to nickelocene is consistent with the usual destabilization of

anions by methyl substitution.42 Thus, from the above analysis, methyl groups can be






81

described as intrinsically electron-donating functions, stabilizing cationic systems and

destabilizing anionic complexes.


Alkyl Substituent Analyses for Positive and Negative Ions and
Rationalization of the Gas-Phase Trends for the Ionization
and Electron Attachments Free Energies


Numerous alkyl substituent parameter schemes have been developed to fit

chemical reactivity to an electrostatic models.4'5'52'53 The parameters are based on the

premise that any substituent R in place of a reference, hydrogen for example, may

alter the bonding, reactivity and overall chemical characteristics of the parent

molecule.5 Substituent schemes based on a single parameters or several parameters

have been used to correlate energy perturbations for chemical systems relative to a

parent reference.50-52 The single parameter model quantitatively predicted shifts in the

AGi values of the alkylferrocene complexes (Chapter 2) relative to ferrocene with 3

kcal molT1 accuracy.6

The Taft model used in the ferrocene analysis employed oI parameters which

were used to assess inductive effects of alkyl parameters. Generally, the parameters

incorporate several electronic effects (i.e, field, polarizability, resonances) thus

separation of the specific electronic effects are not accounted for in this model. Field

effects refer to a charge/dipole, or dipole/dipole electronic interaction transmitted

through space or a through polarizable bond.5 Polarization effects pertain to a

charge/induced-dipole or dipole/induced-dipole interaction. Polarization effects are







82
more strongly distance dependent than field effects.5 However, the relative magnitude

of the two effects are also important in understanding the overall substituent effects.

A single5'53 and a two-parameter102 model based on previously derived

schemes were used to fit the AGi and AGa0 data for the alkylnickelocene complexes.

Free energies of ionization and electron attachment in Table 3-1 were plotted against

ao parameters. Figure 3-4 is a plot of the free energy data versus the alkyl Taft

parameters. Equation 3-3a was used to fit the AGi data to the Taft ao parameters and

equation 3-3b was used to fit the AGa values to the oy parameters. The alkyl

substituents used in the correlations are shown in Table 3-2. Additivity of the

parameters is assumed for the fits. The values of pI is the sensitivity parameter for the



AGi(RCp)(R'Cp)Ni = pI( La) + AGi(Cp2Ni) (3-3a)

AGao(RCp)(R'Cp)Ni = pi(a0i) + AGao(Cp2Ni) (3-3b)



Taft analyses. All fits used in substituent parameter analyses correspond to shifts in

the free energy data relative to nickelocene, where R = H, for the processes shown

below. The free energies for equations 3-4 and 3-5 are the stabilization energy or

destabilization energy for the substituents. Thus AG340 is the stabilization for the



(RCp)(R'Cp)Ni + Cp2Ni+ = Cp2Ni + (RCp)(R'Cp)Ni+ (3-4)

(RCp)(R'Cp)Ni + Cp2Ni' = Cp2Ni + (RCp)(R'Cp)Ni- (3-5)




















O
m E


R=R'=ethyl ~
i R=R=methyl
\ CW
\ R-ethyl;R'=H ,
,/ -15 U_
\ R=R'=H

\ \ E
\ \ v

.-2 0 <-
b 2, 0




-0.12 -0.08 -0:04 O 5
Sum of Alkyl Taft Parameters







Plots of AGio and AGaG data versus Taft X(aI) parameters. (a)
Ionization data (left scale, squares) are plotted as -AGi values for L2Ni
-- L2Ni+ + e'. (b) Electron attachment data (right scale, triangles) are
plotted for the process L2Ni + e" -- L2Ni'. The best fit line for the AGa
fit is drawn for all data except (MeCp)2Ni.


0 F
E -140






L.
w -145-
CD

r-
0
N
73 -150
o














Figure 3-4









Table 3-2 Alkyl Substituent Parameters for Some Alkylnickelocene Complexes
and Free Energies for Reactions 3-4, 3-5 and 3-6.


No.a I CI 1a, AG34ob,c AG350b,c


1 0


AG ob,C
36


0 0


2 -0.092 -0.70

3 -0.055 -0.49

4 -0.110 -0.98

5 -0.148 -1.5

6 -0.46 -3.5


-4.2

-2.3

-5.6

-7.4

-22.5


0.7

-0.5

-0.6

-1.5

(-3.5)


-4.9

-1.8

-5.0

-5.9

(-26)


yjb,d ypb.d


0 0

2.5 1.7

1.4 0.9

2.5 3.1

3.0 4.4


a. Compound numbers taken from Table 3-1
b. Units are kcal mol1 .
c. Estimated error for free energy is 0.4 kcal mol1.
d. The derived I and P values in kcal mol" for the individual substituents are
H, I = P = 0 (defined); Me, I = 1.2, P = 0.9; Et, I = 1.3, P = 1.3;
t-Bu, I = 1.5, P = 2.2







85

cations and AG350 is the stabilization energy for anions relative to nickelocene, except

Me which is destabilizing. The direction of the slopes are opposite because, in

general, alkylation of nickelocene stabilizes formation of the cations and anions. For

the positive ions, electron-attachment becomes less exoergic as alkylation increases

since the cations are stabilized relative to nickelocene. Clearly, the effects of the alkyl

groups are different on the electron attachment energies compared to the ionization

energies. The ionization data fit with equation 3-3a yields a good fit, with a

correlation coefficient of r = 0.997. The parameter p, (= 49.9 kcal mol"1) is the

slope of the line and reflects the sensitivity of nickelocene ionization potentials to

alkylation. This value is comparable to pI for the alkylferrocenes of 57 kcal

mol1.16 Conversely, the same parameters provide an unacceptable fit of the electron

attachment data (r = 0.51). The lack of correlation of the methyl derivative is

primarily responsible for the poor fit of the AGaG data to equation 3-3b. The pI value

for the negative ion data is 6 6 kcal mol1.

A single parameter model was developed by Hehre et al. based on

polarizability effects,52 0a parameters, of R. The (a parameters have been used to

successfully fit the proton-transfer free energies of various cationic and neutral acids.

Fits to the alkylnickelocene AGi and AGa data (Figure 3-5) result in good correlation

for the ionization free energies (r = 0.994) but only poor correlation for the electron

attachment free energies (r = 0.63). The o( parameters are included in Table 3-2.

The slopes for the plots are 5.0 0.4 kcal mol1 for the ionization and 0.8 0.6 kcal

mol1 for the reductions. Based on the observed lack of correlation, a single


















-1
-135-



R=1

S-140-

0

C
w -145-

u-
c
.o

"E -150-
A




-155
-1 .60











Figure 3-5


t-Bu E

R=R'= Et --10

S\ R=R'=Me "
C)
R=H;R'=Et u,,

-.-- -15 u-
SR= R'=H -



-20c
0




j35
b _


-1.20 -0.80 -0.40 -0.05

Sum of ca.parameters









Plot of AGi and AGa data for some alkylnickelocene complexes versus
((og) parameters. (a) Ionization data (left scale, squares) are plotted as
-AGi values for L2Ni -- L2Ni+ + e-. (b) Electron attachment data (right
scale, triangles) are plotted for the process L2Ni + e- -> L2Ni'. Plot of
AGa versus a. is drawn for all points except (MeCp)2Ni.




Full Text

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FILES



IONIZATION, BONDING, AND SOLVATION
ENERGETICS OF ORGANOMETALLIC COMPLEXES
By
MATTHEW F. RYAN
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
1993

To my Grandparents

ACKNOWLEDGEMENTS
The sacrifices I have made over the past several years cannot be counted;
perhaps they should not even be considered. But sacrifice was a major factor towards
completion of this final work-a lot of time and a lot of energy and a lot of sacrifice
all concentrated for a sole purpose. Unusual things happen to a person who focuses so
long and hard on a single goal. The memories in between the battles become clouded
and unfocused. So many memories have amassed, most are happy, but some are sad
and lonely. However, the indecision and the doubt and the pure rage are paled by the
handful of triumphs. Perchance it is these triumphs that make the journey fulfilling.
I recall the time when I felt that I needed to understand more about science and
more importantly, the world around me. Although graduate school was not the sole
source of my salvation it was a beginning: the fresh start that I needed. When I
spoke with Professor Russell Drago in the spring of 1987 (I remember the
conversation well), I told him that if I was accepted to the University of Florida I
would succeed. With the support of Professor Drago and Professor Jack Kotz of
S.U.N.Y. Oneonta, I was admitted to U.F. in the fall of 1987. I am very grateful for
their support and the trust they had in me. I believe I was given an important break.
I would like to thank Professors David E. Richardson and John R. Eyler for all
their guidance. Dr. Eyler has helped me in many capacities and I am grateful to him

for all the discussions we have had. His confidence in me and my work has been very
important to me. David Richardson, my research advisor, has helped me develop
many skills necessary to be a scientist. Although it was sometimes a struggle, for
both him and me, I learned a great deal from him. His door was always open and he
always had time for discussions with me. I will always admire his work as an
educator and a mentor.
Several people assisted in various capacities with this work. I would like to
thank Dr. Allen Siedle, Dr. Mark Burk, Professor Charles Winter, Professor Russell
Hughes, Professor Dennis Lichtenberger, Dr. San Li, Professor William Weltner, and
my friends, Dr. Md. Nazrul I. Khan and Dr. Paul Sharpe for their help.
Throughout my graduate work, I have drawn strength from several vital
sources: my family, my friends, and Debbie Simpson. My family shared with me the
most exciting and frustrating moments. My parents understood and appreciated my
efforts to "catchup" and adjust to my new surroundings. When I spoke with them, any
discouragement I had was made insignificant; I spoke with them often because their
confidence and love are inspirational. My brothers and sisters, Mark, Joe, Cassandra,
and Andrea, were there to fill in the gaps. Their love for me was manifested in many
ways. I am grateful for all the things they have done for me for they have given me
so much and asked for nothing in return. I am blessed to have such a close family.
My friends Steve and Mike Messick helped me to persevere. Through Steve
and Mike, I learned how to focus my thoughts and emotions. The jaunts to Cape
tv

Vincent (and points north) and California were extremely therapeutic. Although the
briar patch at the Tesla concert was not efficacious, Vegas was completely necessary.
When I came to U.F., I left behind several close friends, and I have remained
close with only a few people during my time in Florida. Jack Santos (soon to be a
father) and I have known each other forever. Although I may have neglected our
friendship occasionally, Jack always understood. He is a great friend.
I am very grateful to my friend John Moore for his help in the beginning and
his encouragement until the final moment. Before I moved to Florida, John found me
an apartment and. more importantly, introduced me to the U.F. Rugby Team soon after
my arrival. Rugby, in addition to supplying me with dozens of wonderful friendships,
was a vent for frustrations like no other I can imagine.
Steve Glatter and Steve Rubin have supplied me with everlasting laughter. In
addition to their advice and refreshing perspective on things in general, I would like to
thank them for their companionship throughout the years.
Mike Naugton and I were roommates for nearly five years. I enjoyed his
company, and I am glad we have remained friends.
Before I met Debbie Simpson, I often felt that I was isolated and alone. My
family and friends are all far away and the rugby team has a specific purpose. Debbie
is a wonderful friend, her smile is contagious, and her warmth and love never fade.
Debbie has always listened to me and has helped me to find the answers. We have
been through a great deal together, and our friendship, love, and respect for each other
is constantly growing with each passing challenge.
v

TABLE OF CONTENTS
ACKNOWLEDGEMENTS iii
LIST OF TABLES ix
LIST OF FIGURES xi
ABSTRACT xiii
CHAPTERS
1 INTRODUCTION 1
2 ADIABATIC IONIZATION ENERGIES, BOND DISRUPTION ENTHALPIES,
AND DIFFERENTIAL SOLVATION ENERGIES OF GAS-PHASE
METALLOCENES AND METALLOCENIUM IONS 13
Introduction 13
Reevaluation of Metallocene Free Energies of Ionization Based on
Electron-Transfer Equilibrium Studies 16
Insights into the Free Energy of Ionization of Ferrocene 24
Intramolecular Entropy Changes of the Ferrocene/Ferrocenium Couple ... 34
Ionization Free Energies of Ruthenocene and Osmocene 43
Free Energies of Ionization of Vanadocene, Manganocene and Nickelocene 44
Substituent Effects in Ferrocene Derivative Oxidations 45
Heterolytic and Homolytic Metal Ligand Bond Disruption Enthalpies of
Metallocenes and Metallocenium Ions 49
Differential Solvation Free Energies for Metallocene Redox Couples 54
Conclusions 63
Experimental Methods 65
3 SUBSTITUENT EFFECTS IN THE GAS-PHASE AND SOLUTION
IONIZATION AND ELECTRON-ATTACHMENT ENERGIES OF
ALKYLNICKELOCENES 70
vi

Introduction 70
Electron-Transfer Equilibrium Studies Involving Negative and
Positive Alkylnickelocene Ions 72
Alkyl Substituent Analyses for Positive and Negative Ions and Rationalization
of the Gas-Phase Trends for the Ionization and Electron Attachment
Free Energies 81
Solvation Energetics of Nickelocene Cations and Anions 91
Bond Disruption Enthalpies for Nickelocene Anion 94
Conclusions 94
Experimental Methods 96
4 GAS-PHASE AND SOLUTION OXIDATION POTENTIALS OF
RUTHENOCENE DERIVATIVES 98
Introduction 98
Results of the Electron-Transfer Equilibrium Reactions 101
Evaluation of the Gas-Phase Free Energies of Ionization for a Series of
Ruthenocene Derivatives 108
Attempted Correlation of Ruthenocene Ionization Free Energies
with Taft Gj Parameters 116
A New Parameter Scale for Cyclopentadienyl Substituents Based on Gas-Phase
Electron-Transfer Equilibrium Studies of Ruthenocenes 123
Rates of Hydrogenolysis for Methylzirconocene Cations 129
Comparisons of Gas-Phase Ionization Free Energies to Solution
Electrode Oxidation Potentials 130
Determinations of Free Energies of Ionization in Solution from
Electrochemical Oxidation Potentials 137
Differential Solvation Free Energies for Several Ruthenocene/Ruthenocenium
Couples 138
Application of the Bom Model for Estimating Solvation Energetics
for Ruthenocene Oxidation Couples 140
Conclusions 143
Experimental Methods 145
5 GAS-PHASE IONIZATION ENERGETICS, THERMOCHEMISTRY, AND
ELECTRON-TRANSFER KINETICS OF DECAMETHYLMETALLOCENES,
CHROMOCENE, AND COBALTOCENE 148
Introduction 148
Gas-Phase Electron-Transfer Equilibrium Studies 151
Electrochemical Studies for Some Decamethylmetallocenes 152
Bis(benzene)Chromium as a Reference Compound for Electron-Transfer
Equilibrium Investigations 156
vii

Free Energies of Ionization for Some Decamethylmetallocenes and
Comparison to Photoelectron Spectroscopy Results 159
Free Energies of Ionization for Chromocene and Cobaltocene 163
Bond Disruption Enthalpies for Chromocene and Cobaltocene 164
Evaluation of the Solvation Energetics for Decamethylmetallocenes
Chromocene and Cobaltocene 166
Electron-Transfer Kinetics 171
Conclusions 174
Experimental Methods 175
6 OVERVIEW OF EXPERIMENTAL METHODS AND PROCEDURES .... 178
Fourier Transform Ion Cyclotron Resonance Mass Spectrometry 178
Measurement of Equilibrium Constants 183
Temperature Dependence Studies 187
Application of FTMS for the Study of Metal Complexes 188
7 SUMMARY 191
REFERENCES 194
BIOGRAPHICAL SKETCH 203
viii

LIST OF TABLES
Table
2-1
2-2
2-3
2-4
2-5
2-6
2-7
3-1
3-2
4-1
4-2
4-3
4-4
5-1
Page
Ionization Energetics Data for Some Metallocenes 23
Calculated Entropies and Integrated Heat Capacities for Ferrocene and
Ferrocenium Ion at 298, 450, and 600 K 35
Vibrational Frequencies for Various Ferrocenium Salts 37
Vibrational Frequency Data for Ferrocene and Ferrocenium Cation 40
Mean Bond Disruption Enthalpies for Some Metallocenes 55
Auxiliary Thermochemical Data Used in Thermochemical Cycles 56
Electrochemical £L Data and Differential Solvation Energies for Some
Metallocene CpjlVT^0 Couples 58
Free Energies of Ionization and Electron Attachment 79
Alkyl Substituent Parameters for Some Alkylnickelocene Complexes and
Free Energies for Reactions 3-4, 3-5, and 3-6 84
Values of AGj° for Ruthenocene Derivatives and Other Data 104
Ligand y* and y Parameters 105
Substituent Parameters for Selected Cyclopentadienyl Derivatives 122
Electrode Potentials and Differential Solvation Free Energies for Some
Ruthenocene Derivatives 135
Ionization Energetics Data for Some Metallocenes and
Decamethylmetallocenes 154
IX

Table Page
5-2 Electrochemical E^ Data and Differential Solvation Energies of Some
Cp*2M+/0 and Cp2M+/0 Couples 155
5-3 Average Bond Disruption Enthalpies for Chromocene and Cobaltocene . 165
5-4 Electron-Transfer Kinetics for Some Metallocenes,
L2Ma+ + L2Mb -> + L2Mb+ 172
x

LIST OF FIGURES
Figure
Page
1-1 Potential well diagrams demonstrating vertical and adiabatic ionization
process 7
2-1 Log plot for electron-transfer reaction of Cp2Fe +DET+ = DET + Cp2Fe+ 17
2-2 Electron-transfer equilibrium ladder for some metallocenes 22
2-3 Van’t Hoff plots for selected Metallocene electron-transfer
equilibrium couples 25
2-4 Plot of AG° values (kcal mol'1) versus alkyl Taft parameters 48
2-5 Thermochemical cycles used to determine bond disruption enthalpies and
differential free solvation energies for metallocenes 52
2-6 Plot demonstrating periodic trends of ionization energies for the first transition
row metallocenes 60
2-7 Structure of 18 e'1 vanadocenium complex 61
3-1 Molecular orbital diagrams for nickelocene anion, nickelocene, and nickelocene
cation 73
3-2 Electron-transfer equilibrium ladder for ionizations for several alkylnickelocene
complexes 75
3-3 Electron-transfer equilibrium ladder for electron attachments for
alkylnickelocene complexes 76
3-4 Plots of AG¡° and AGa° data versus Taft E(aj) parameters 83
3-5 Plot of AG¡° and AG° data for some alkylnickelocene complexes versus Z(ca)
parameters 86
xi

Figure
Page
3-6 Plot of AG36° values (kcal mol'1) derived from equations 3-4, 3-5, and 3-6
versus X(Oj) values 90
4-1 Electron-transfer equilibrium ladder for several ruthenocenes derivatives for the
process M —> M+ + e' 103
4-2 Plot of Ru 3d binding energies from reference 91 versus ETE AG¡° values for
several Cp*Ru-L complexes and ruthenocene 115
4-3 Correlation of AG¡° values for several ruthenocene derivatives with Taft Oj
parameters 118
4-4 Plot of AGj° values versus y* parameters 125
4-5 Plot of AGj° values versus Z(y) parameters 128
4-6 Gas-phase rates of hydrogenolysis for several methylzirconocene cation
complexes 131
4-7 Structures of Cp*2Ru and ferrocene 142
5-1 Electron-transfer equilibrium ladder for several decamethylmetallocenes 153
5-2 High resolution He (I) photoelectron spectrum of bis(benzene)chromium in the
valence ionization region 158
5-3 Plot of AG° values for alkylferrocene and nickelocene complexes and versus
alkyl Taft Gj parameters 161
5-4 Plot demonstrating periodic trends of ionization energies for the first transition
row decamethylmetallocenes 168
6-1 Orthorhombic ion trap used in a Fourier transform ion cyclotron resonance
mass spectrometer 180
6-2 Schematic representation of a Fourier transform ion cyclotron resonance mass
spectrometer 185
6-3 Van’t Hoff plot of the CO/Kr electron-transfer equilibrium reaction ... 189
xn

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
IONIZATION, BONDING, AND SOLVATION ENERGETICS
OF ORGANOMETALLIC COMPLEXES
By
Matthew F. Ryan
May, 1993
Chairperson: David E. Richardson
Major Department: Chemistry
Free energies of ionization (AG¡°) for Cp2M complexes (Cp = r|5-
cyclopentadienyl; M = V, Cr, Mn, Fe, Co, Ni, Ru, Os) and Cp*2M complexes (Cp* =
r|5-pentamethylcyclopentadienyl; M = Mn, Fe, Ni, Ru, Os) have been determined from
electron-transfer equilibrium methods by using Fourier transform ion cyclotron
resonance mass spectrometry. The AG¡° values for ferrocene derivatives,
alkylnickelocene derivatives and various ruthenocene complexes have also been
investigated through ETE studies.
Temperature dependence studies involving ETE of ferrocene with N,N’-
diethyltoluidine lead to values for the ionization enthalpy and entropy for ferrocene.
Experimental results of AS¡°(Cp2Fe) are compared to results from statistical
mechanical analyses. Thermochemical cycles are used to derive estimates of
heterolytic and homolytic bond M-Cp enthalpies for the first transition row
xiii

metallocenes and to derive estimates for the differential solvation free energies,
AAGSoiv°> for the Cp2M+/0 couples. With the exception of Cp2V+/0, first transition
row Cp2M+/0 couples are estimated to have AAGsolv° values of -38 ± 3 kcal mol'1.
The average AAGsolv° for the Cp*2M+^° couples is -26 ± 3 kcal mol'1.
The AGj° values and electron attachment free energies (AGa°) for a series of
alkylnickelocene complexes (RCp)(R’Cp)Ni have been determined. The ionization
energies follow the expected trends (increased alkylation yields a decrease in AG¡°) but
the AGa° values become more negative for R = Et and t-Bu and more positive for R =
Me. Several parameter schemes are used to interpret the AGj° and AGa° data. Values
of AAGsolv° for the Cp2Ni+/0/' and (t-butylCp)2Ni+/0/' couples are discussed. From the
Cp2Ni+/0/" data, an average AAGsolv° value is estimated.
The AG° values and estimates for series of ruthenocene derivatives have
been determined. A parameter scale that correlates electronic effects for Cp ligands,
rather than individual substituents, has been developed based on AG¡° values of the
ruthenocene complexes. The utility of the parameters for predicting reactivity of metal
complexes bearing Cp ligands is considered.
xiv

CHAPTER 1
INTRODUCTION
Thermochemical information is of primary importance to the understanding of
chemical reactivity. Reaction pathways and mechanisms can be better understood if
salient thermochemical data pertaining to reactant species or products exist.
Thermodynamics can inform us if a chemical reaction will proceed and can tell us the
efficiency of the reaction. Although information concerning the velocity of a reaction
can not be readily assessed from a thermodynamic analysis, theories of chemical
kinetics are based on the laws of thermodynamics.1 Thus, the utility of
thermochemistry as a tool to both predict and comprehend chemical transformations is
essential to chemistry.
Organic chemistry has benefitted enormously from the depth and breadth of
available thermochemical data for simple and complex organic systems.2,3
Furthermore, the thermochemical foundation of organic chemistry allows complex
reaction mechanisms to be understood and reaction efficiencies to be maximized. The
utility of thermochemistry to predict reactivity based on known quantities is well
documented in organic chemistry.2,3 For instance, the derivation of parameterization
schemes by Hammett4 and Taft5 based on various equilibrium dissociation and
ionization constants elegantly demonstrates how thermochemistry can be used to
predict additional thermochemical information and interpret existing systems. The
1

2
quantification of field and resonance effects, wrought from extensive
experimentation,4,5 enables the direct application of thermochemistry to interpret the
energetics of chemical systems.
The application of thermochemistry to estimate additional thermochemical
information is undoubtedly dependent on accurate data. Although compilations of
thermochemical data are available for organic and simple inorganic molecules,6,7 the
situation for organometallic complexes is different. Because of the distinct differences
of organometallic complexes relative to organic systems, application of existing
organic thermochemistry to organometallics is limited. Inorganic thermochemical data
n
are primarily available for simple binary and tertiary molecules or main group
molecules,8 and parallels to large metal-centered complexes are inappropriate.
The thermochemistry of organometallic systems has gained increased attention
recently, driven by their increased importance in specialized areas of chemistry.9'11
The increased application of organometallic systems in the areas of catalysis and
material sciences has led to increased investigations into the thermodynamic
contributions to the reactivity of these systems. For example, knowledge of metal-
carbon and metal-hydrogen bond enthalpies is essential for the understanding of
catalytic reaction mechanisms that often involve metal-ligand bond cleavage.9
Bonding energetics of metal complex ions are important for heterogenous catalysis if a
complex ion is the active catalytic species.10 Additionally, knowledge concerning the
oxidation/reduction energetics of organometallic systems can be used to determine
solid-state lattice or binding energies for organometallic materials and ceramics.

3
Most thermochemical studies of organometallic complexes have focused on the
use of combustion or reaction calorimetry to determine bond energetics and heats of
formation in the condensed phase. Combustion calorimetry has been used to
estimate condensed-phase metal-ligand bond enthalpies for model organometallic
complexes (i.e. metal carbonyl and metal-arene complexes), and reaction calorimetry
has been used to derive metal-ligand bond enthalpies for metal complexes which are
useful as homogeneous catalysts.9 In spite of the accuracy of these techniques, there
are some disadvantages. If a reaction involves more than one metal-ligand bond
cleavage, the calorimetric study can yield ambiguous results. Further, the large
quantities of sample needed for calorimetry, along with the thermal instability of metal
complexes, renders traditional calorimeters ineffective for many organometallic
9 12
systems. ’
Gas-phase investigations of metal complexes can lead to intrinsic properties of
a molecular species in the absence of solvent effects.9,14'17 For example, gas-phase
studies of isolated metal complexes or metal-complex ions can yield ionization
energetics data15,18'20 and bond dissociation enthalpies.15,16 Further, direct
comparison of the gas-phase energetics of metal species to their solution analogues
reveals solvent effects for the metal complex system.15'20
Gas-phase bond enthalpy studies have frequently involved coordinatively
unsaturated ionic metal species such as M+-R, where R = H, CH2, CH3 and so
forth. Although these studies are useful for estimating bare metal-substituent
bond enthalpies, direct solution analogues for the bare-metal ion species do not exist.

4
Bond enthalpies of the coordinatively unsaturated M+-R species may deviate
significandy from a related metal complex.9 Additionally, because the bare-metal
species lack supporting ligation, their stability and reactivity will not parallel that of a
comparable metal complex ion. The electronic properties of supporting ligands will
alter the bonding and solvation energetics for Cp2MR+ complexes (Cp =
ri5-cyclopentadienyl) relative to M+-R, for example.9'11
Numerous investigations on the stepwise bond energetics for organometallic
ions such as MCCO)^ have been determined from mass spectrometric and
spectroscopic methods. For example, Norwood and coworkers have estimated the
sequential Fe+-CO bond strengths for ironpentacarbonyl by using a combination of
spectroscopic and mass spectrometric techniques.24 Recently, Schultz et al. have
measured the sequential Fe+-CO bond strengths for Fe(CO)5 by mass spectrometry 25
Further, Sunderlin et al. have recently reported the metal-ligand bond enthalpies for
various transition metal carbonyl anions in the gas-phase by using mass spectrometric
flow tube techniques.26
Another approach used to determine organometallic bond energetics has been
the application of thermochemical cycles which incorporate ionization and/or reduction
potentials.15,16 A limiting factor of this technique is the accuracy of the
organometallic oxidation/reduction data; auxiliary data needed for the cycles, such as
ligand heats of formation and metal atom ionization energies are typically accurately
known.8,27 Thermochemical cycles can also be used to determine solvation energetics
associated with metal complexes from direct comparisons of gas-phase and solution

5
oxidation/reduction energetics.15'20 Understanding solvation energetics is essential for
maximizing reaction efficiencies and understanding reaction pathways. Moreover,
analysis of the redox properties of metal-centered systems in the gas-phase and
solution can be utilized to interpret condensed-phase electrostatic interactions such as
ionic solvation effects,29 acid/base reactions,30 and electron-transfer phenomena.31
Ionization and reduction energetics are also useful for understanding and
developing periodic and group trends for an ensemble of metal complexes.16 For
example, the electronic effects of alkylation or ligand substitution can be readily
assessed by comparing relative redox potentials of a series of related metal
'V)
complexes. Then from an understanding of substituent electronic effects, complexes
can be modified to alter reactivity.5 Steric effects of substituents are easily
rationalized as they are primarily conceptional whereas electronic effects are less
obvious.2,3,5 Comparisons of ionization/reduction energies in the gas-phase enable
accurate assessment of substituent electronic effects. In the condensed phase, solvent
effects are predominant and often electronic effects in solution do not parallel trends
observed in the gas phase.5
The majority of data concerning ionization energetics for organometallic
complexes are either electrochemical potentials, E^,34 or vertical ionization energies,
vIP, measured by photoelectron spectroscopy.35 These data are not always appropriate
for deriving thermodynamic properties of ions near or at room temperature. For
example, a vertical ionization will not be a true assessment of a thermal ionization
process (free energy of ionization) if the equilibrium geometries of the ion and the

6
97
neutral are significantly different. Vertical ionization energies refer to an ionization
process in which the geometry of the parent ion is essentially the same as the ground
state neutral. Specifically, the most probable ionization transition will be that in which
'Xfi
the positions and momenta of the nuclei remain unchanged. The ion is formed with
excess internal rotational and vibrational energy and is therefore not in its ground
state.27 A thermal ionization process, AT°, which is similar to an adiabatic
ionization energy, or AHx 0, is a process in which both the neutral and the ion are in
their equilibrium geometries at temperature 7\16 At temperatures greater than T = 0
K, there are entropy and heat capacity contributions to the ionization process.16,27 If
the structures of a neutral and an ion are significantly different, a vertical ionization
will not be an accurate assessment of the thermal ionization. Potential well diagrams
which represent both vertical and thermal ionization processes are shown in Figure
1-1. Although the vertical ionization process is the most probable transition,27,36 ion
relaxation (from ion-molecule collisions,14 for example), which removes excess
rotational and vibrational internal energy, results in a thermal ionization process. The
difference in vertical ionization energy and AG ° is the ion relaxation energy given by
Et = vIP - AG¡°.16
Electrochemistry is a useful method for evaluating redox potentials. However,
electrochemical potentials for many organometallics are irreversible, making
determination of a true E^ value difficult.34 Organometallic complexes that form
unstable solvated ions will produce electrochemical responses inconsistent with gas-
phase data. Although the vertical ionization energy for manganocene has been

7
Figure 1-1 Potential well diagrams demonstrating vertical and adiabatic ionization
processes. Arrows represent vertical ionization transitions. Figure A is
an example in which the equilibrium geometries of the ion and the
neutral are significantly different. Figure B represents a process in
which the geometries of the ion and neutral are similar. For this
example shown, vIP ~ alP.

8
measured by photoelectron spectroscopy,35,37,38 the E^ value for Cp2Mn has not been
reported due to the instability of Cp2Mn+ in solution.34
The study of gas-phase electron-transfer equilibrium is a well-established
means for determining thermal ionization energies for organic39"42 and inorganic14"20
complexes at or near room temperature. The ability to trap ions and monitor an
electron-transfer equilibrium reaction is a powerful application of mass spectrometry.
Previous electron-transfer equilibrium studies have been reported by several groups
using various mass spectral techniques.15,39,42 An earlier study on the electron-
transfer equilibrium of the gas-phase ionization thermochemistry of ferrocene by using
pulsed high pressure mass spectrometry has been reported by Mautner.43
In this work thermal free energies of ionization for metallocenes and
metallocene derivatives have been determined from electron-transfer equilibrium
studies by using Fourier transform ion cyclotron mass spectrometry, FTMS.44-47 From
the measured equilibrium constants for the electron transfer reaction, AGet° values are
determined, and the free energies of ionization are estimated for the metallocenes for
the process shown below, where L represents a Cp ligand or a Cp derivative,
L^M -4 L2M+ + e" (1-1)
where Cp = ri5-cyclopentadienyl. Ionization free energies for the first transition row
metallocenes, ruthenocene, and osmocene are presented.16 Metallocenes were chosen
because they represent the foundation of organometallic chemistry and have numerous

applications in the areas of homogeneous catalysis,9 material science,48 and nonlinear
49
optics.
Following established methods,15,16 the ionization energetics data are
incorporated into thermochemical cycles to provide estimates of bond disruption
enthalpies for the first transition row metallocenes and metallocenium ions.16
Although bond enthalpies for the neutrals have been reported,13 accurate bond
enthalpy values for the metallocenium ions have only recently been established.16
Cyclopentadienyl ring substituent effects for metallocene complexes have been
studied by electrochemistry34 and occasionally by photoelectron spectroscopy.35
Relatively few gas-phase studies on the substituent effects for metal complexes have
been reported for comparison to the electrochemical potentials.50'51 In this work the
electron-transfer equilibrium method has been applied to a series of alkyl ferrocene16
and nickelocene19 derivatives to investigate the effects that substitution has on the gas
phase redox energetics of metallocene complexes. Nickelocene is a useful complex
because it forms stable cations and anions in the gas-phase;19 therefore, effects of
alkylation on the positive and negative nickelocene complexes were studied. The A
and AG° values for the metallocene derivatives are correlated with alkyl substituent
c C9 c'j
parameters. Various parameterization schemes ’ ’ are considered to interpret and
understand the gas-phase investigations and electron-transfer energies. These data are
potentially useful in understanding and predicting ionization energies51 and optical
transition energies for metallocene derivatives.49 Optical transition energies are
important in the selection of chromophores for potentially useful nonlinear optical

10
devices. Since ligand substituents affect molecular orbital energies,5 metal complexes
can be designed with specific optical transitions based on knowledge derived from
gas-phase redox potentials.
The AG° values for a series of ruthenocene derivatives have been determined
from electron-transfer equilibrium reactions. The AG ° values span over 2 eV within
the series. Because of the widespread variations of the Cp ligands, bulk ligand effects
were considered rather than individual ring substituent effects. A parameterization
scale for correlating electronic effects for Cp type ligands was developed.20 The
application of ligand substituent effects to understanding and predicting reactivity for
organometallics complexes is discussed.
Relatively little is known concerning the solvation energetics for organometallic
complexes.16 In order to better understand solvation effects of metal complex redox
couples, an understanding of the intrinsic (solvent free) electron-transfer chemistry is
important. Electrochemical potentials for several metallocene derivatives have been
measured in order to fully characterize the effects ligation and solvation has on the
redox chemistry of metal centered molecules. Thermochemical cycles, which
incorporate gas-phase and solution metallocene redox potentials,16,19,20 have been used
to derive estimates of solvation energetics. Electrostatic models for predicting
solvation energetics for spherical ionic species have been applied for the metallocene
complexes.15,54,55 Shortcomings and criticisms of the electrostatic model are
discussed in light of the experimentally derived estimates.

11
Throughout this work, several important objectives are examined. The
evaluation of accurate ionization and reduction potentials and detailed analyses of the
data for prototypical organometallic complexes are important objectives of this work.
For example, temperature-dependence studies have been determined to evaluate
entropy and enthalpies of ionization for several metallocene complexes.16
Comparisons of the experimental work to spectroscopically derived thermodynamic
parameters revealed specific contributions of the ionization process. Thus, the
thermodynamic origins of the ionization processes are critically analyzed.
The application of thermodynamics to predict reaction mechanisms and
reactivity is another major objective of this work. The thermodynamic parameters
presented in this work may be applied to other areas of research that utilize the myriad
of characteristics of organometallic complexes. Existing parameterization schemes52,53
for predicting reactivity have been successfully used to correlate thermodynamic data
for the metallocenes.16,19 Where established models fail to correlate, schemes have
been developed for predicting thermodynamic values for metallocene-type
complexes. The application and development of the parameterization schemes for
understanding and predicting chemical reactivity are discussed in detail.
A significant portion of this work is devoted to deriving and understanding
metallocene/metallocenium solvation effects. Relatively little is known concerning the
solvent effects for organometallic complexes. Solvation data reported here may help
to further develop a foundation for understanding solvation effects for other
organometallic complexes.

12
The development of methods for fully characterizing the thermodynamics of
organometallic complexes was explored. Fourier transform mass spectrometry has
proven to be a powerful tool towards fulfilling this goal. Electron-transfer equilibrium
techniques are an effective means of determining thermodynamic parameters. When
complimented by electrochemical studies34 and photoelectron spectroscopy,35 FTMS
studies allow for the full characterization of redox properties for organometallic
complexes. Further, the development of temperature-dependence techniques
establishes FTMS as a technique for deriving reaction entropies and enthalpies for
metal systems.

CHAPTER 2
ADIABATIC IONIZATION ENERGIES, BOND DISRUPTION ENTHALPIES,
AND SOLVATION FREE ENERGIES OF GAS-PHASE METALLOCENES
AND METALLOCENIUM IONS
Introduction
The majority of data concerning metallocene oxidation-reduction potentials is
in the form of vertical ionization energies measured by photoelectron spectroscopy34
and electrochemical potentials.35 Several experimental limitations of these two
techniques can cause uncertainty in the determination of thermodynamic reduction-
oxidation potentials. Many organometallic complexes have irreversible
electrochemical oxidation/reduction potentials which can lead to uncertain
assignments for £j¿ values. Vertical ionization energies may differ from adiabatic
97 9A
potentials if the equilibrium geometries of the ion and the neutral are dissimilar. ’
If the equilibrium geometry of the ion and the respective neutral are similar, then the
vertical potential will closely approximate the adiabatic potential which is referenced
at 0 K. However, even if photoelectron studies can accurately determine an
adiabatic potential, values for the ionization free energy or reduction and respective
enthalpy changes referenced at other temperatures must be estimated from statistical
thermodynamic analyses. Spectroscopic data, which include vibrational and structural
characterization, for metal complexes frequently do not exist for metal complexes.16
13

14
Electron-transfer equilibrium, ETE, is a powerful method for determining
thermal oxidation and reduction potentials for organic and inorganic species at ambient
temperatures.15,16,39,41 Kebarle and coworkers have used pulsed high pressure mass
spectrometry, PHPMS, to determine the free energies of electron attachment, AGa°, for
organic compounds in the 0-3 eV range.42 In the Kebarle studies, the electron
attachment free energy of S02 was used as the reference compound to anchor the
AG° values derived from electron-transfer equilibrium studies. An accurate value for
the electron affinity of S02 has been determined from photodetachment studies by
Celotta et al.56 Additionally, sufficient spectroscopic data for S02 exist for complete
statistical mechanical analyses of the enthalpy and entropy for electron attachment.57
Thus, all AG° values reported from Kebarle’s laboratory are referenced to AGa°(S02)
at 423 K.
In this work, Fourier transform ion cyclotron resonance mass spectrometry44'47
has been used to determine the free energies of ionization for some first transition row
metallocenes, ruthenocene, and osmocene. Derived AG¡° values are represented by
equation 2-1, where L denotes a Cp ligand or a substituted Cp ligand.
L2M(g) -> L2M+(g) + e‘ (2-1)
Ionization free energies for the compounds studied in this chapter have been
previously reponed.13,18 However, the derived AG¡° values for the metallocenes
reponed here differ from values of the previous study due to refinements in the free

15
energies of ionization for the organic reference compounds and more extensive
electron-transfer studies.16
Entropies and enthalpies of ionization for the metallocenes have been estimated
from investigating the temperature-dependence of the equilibrium constants for
selected metallocene reaction couples. Statistical mechanics has been used to estimate
the intramolecular entropy change for ferrocene. Available spectroscopic data for
Cp2Fe and Cp2Fe+, in addition to new vibrational frequency data for the ferrocenium
cation, were used to estimate values for the total A5¡° at several different temperatures.
Mautner has used PHPMS to determine the ionization free energy of ferrocene by
electron-transfer equilibrium in which alkylaniline compounds were used as reference
compounds.43 Additionally, Mautner was able to assess values for the entropy and
enthalpy of ionization by studying the temperature dependence studies of the measured
equilibrium constants. Thermodynamic parameters for ferrocene determined in this
work are compared to values reported by Mautner.
Thermochemical cycles were used to estimate differential solvation free
energies and bond disruption enthalpies.15,16,18 Values for heterolytic and homolytic
M-Cp bond cleavage have been previously presented.13 However, since bond
enthalpies derived from thermochemical cycles are dependent on AG^ values, new
estimates are included which reflect refined AG ° data in addition to more accurate
thermochemical analyses. Differential solvation free energies are derived from direct
comparison of AG¡° values in the gas phase and solution.15,16,54
Estimates for

16
differential solvation free energies for several metallocene couples are compared to
CQ
values predicted by a simple electrostatic model.
The effects of attached substituents to cyclopentadienyl and arene ring systems
have been studied previously by electrochemistry and photoelectron spectroscopy. An
earlier study demonstrated that free energies of ionization for alkylferrocene
derivatives correlate well with alkylbenzene analogues. In the present work, AG°
values derived from electron-transfer equilibrium studies for the alkylferrocene
complexes are correlated with alkyl substituent parameters. Alkyl substituent
parameters have been shown to correlate well with AG¡° data for organic compounds
and chromium coordination complexes,51 but have only recently been applied to
metallocenes.16,32 These data may prove beneficial for the interpretation and
prediction of physical properties for organometallic complexes such as ionization
energies or optical transition energies useful to photochemistry.49
Reevaluation of Metallocene Free Energies of Ionization
Based on Electron-Transfer Equilibrium Studies
Electron-transfer techniques have been described elsewhere.15,16'39'42 The
general electron-transfer equilibria shown in equation 2-2 were studied where L2M
represents a metallocene and R denotes a reference compound with known AG¡° value.
Ionization free energies for the reference compounds used in this work are typically
± 1 to 2 kcal mol'1 27,40,59
LjM + R+ ** R + L2M+
(2-2)

Intensity
17
Figure 2-1 Log plot for the electron-transfer reaction of Cp2Fe + DET4- = DET +
Cp2Fe+. Ion intensity units are arbitrary. DET = N,N-diethyitoluidine
and Fc = ferrocene.

18
Figure 2-1 is an example of typical electron-transfer equilibrium reaction, the
ferrocene/N,N-diethyltoluidine couple in this example. The decay of ion signal over
time is due to diffusive loss of ions from the reaction cell of the mass spectrometer.
Equilibrium constants and subsequently electron-transfer reaction free energies, AG °,
can be determined (equation 2-3) if the difference in the ionization
free energies of the two compounds is < 4 kcal mol'1.10 In equation 2-3, P denotes
the partial pressure and / represent the parent ion intensity for species in reaction 2-2.
AGet° = -RT In ATeq= -RT In [(P(R)//)(L2M)) * (7(L2M+)//(R+))] (2-3)
Equilibrium between parent ions can actually be monitored for reactions with AGet°
values approaching 5 kcal mol'1;60 however, the partial pressure ratios must exceed
100, leading to large experimental uncertainties. The pressure of the minor component
would only be slightly above the background pressure (ca. 1-2 x 10‘8 Torr). Since the
AG° value for R is known, the free energy of ionization of the metallocene at
temperature T can be determined from equation 2-4. Determination of AGet° values
over a range of temperatures yields estimates of A//et° and A5C[°.
AGeCr° = AGiT°(L2M) - AGiT°(R) (2-4)
If the enthalpy and entropy of ionization of the reference compound are known,
corresponding values for the metallocenes can be derived for the appropriate

19
temperature range. All reactions temperatures in the present work used to derive
AGet° values are 350 K unless otherwise stated.
Reference Compounds Used in Electron Transfer Equilibrium Studies.
Appropriate reference compounds for ETE studies in the 6-7 eV range have
been studied by PHPMS and ion cyclotron resonance mass spectrometry.39,59 Lias et
al. have used ETE methods to determine the ionization energies of several aniline
derivatives which are used as reference compounds in the present work.59 Reference
compounds are anchored to the ionization potential value of N,N-dimethylaniline,
DMA, at 350 K.59 The ionization energy of DMA is anchored to the ionization
potentials of benzene and N02, which have been determined spectroscopically.
Values of AS¡° for benzene and N02 have been estimated from statistical mechanic
analyses;39 subsequently, accurate values of AG¡ 350° have been determined.
Reevaluation of the ionization thermodynamics for the reference compounds
was necessary in order to determine AG¡ 35q° data for the metallocenes. Ionization
thermodynamics used in this work for the reference compounds are based on several
important assumptions:
(i) A//j°(DMA) = 7.12 ± 0.02 eV (164.2 ± 0.5 kcal mol'1).27
(ii) A5j° for azulene results from ASelec. Therefore AS¡° = R In g = 1.38 cal
mol^K"1.40 The ionization of azulene is a singlet to doublet transition in
which g = 2.

20
(iii) Based on the reported ASet° value from ETE studies by Nelsen and
Mautner involving DMA and azulene, the ionization entropy for DMA is 2.3
cal mol^K'1.
(iv) Based on the structural similarity, the AS ° values for the aniline
derivatives used as reference compounds are equal to A5j°(DMA).
(v) Changes in the integrated heat capacities of the reference compounds and
their neutrals are negligible and will cancel. Therefore, derived AH ° values
are assumed to be independent of temperature in the 350-500 K range.39
Assumption i has been used in other studies in which N,N-dimethylaniline has
been used as a reference compound.40,43 The AGj° value used here is within ± 0.01
eV of the literature value and in agreement within experimental uncertainty of the
AG¡° value used by Nelsen and Mautner. Assumption ii is quite reasonable
considering the small vibrational changes expected for the ionization of azulene.
Azulene and its cation are expected to be structurally similar,16,40 therefore vibrational
and rotational contributions to the ionization entropy will be negligible. Translational
entropy changes for ions and their parent neutrals will be negligible since the two
species differ only by the mass of an electron. Aniline derivatives are also expected
to have small vibrational and rotational contributions to the total entropy change, and
A5j° = 2.3 cal mol^K'1 (assumptions iii and iv) represents the electronic entropy
change (singlet to doublet process, g = 2) with some contribution for ASvib and ASrot.
Based on previous studies the ASet° for the azulene/DMA couple40 was reported to be

21
0.9 cal mol^K'1 which results in A5j°(DMA) = 2.3 cal mol^K'1. The minor effects of
ACp on AH° are not unreasonable especially for larger molecules. For example,
A//¡ 35Q - //j q for benzene, determined from statistical mechanics, is only 0.2 kcal
mol'1 which is negligible when compared to the much larger AH° value.39 As will be
shown later, assumption v holds true for the metallocenes.
Ionization Energetics Data for Some Metallocenes
Figure 2-2 is an electron-transfer equilibrium ladder displaying all reaction
investigated in this chapter. This ladder is similar to an equilibrium ladder previously
reported.13 The AG¡° values in Figure 2-2 differ from those of an earlier study due to
the assumptions outlined above and more extensive ETE studies. Details involving
individual ETE have been discussed elsewhere.16 Several ETE reactions performed
with respect to an earlier report13 include the following reactions couples:
Cp2V/Cp2Fe, (EtCp)CpFe/Cp2Fe, (EtCp)CpFe/(n-butylCp)CpFe, and temperature
dependent studies of the equilibrium constant for Cp2Fe/DET and (EtCp)CpFe/Cp2Fe
couples. Numbers adjacent to arrows in Figure 2-2 denote AGeu35Q° values for
individual ETE reactions. Reactions involving only reference compounds are included
in the equilibrium ladder to serve as checks on the literature values.59 Values in
parentheses are unchecked literature values. Table 2-1 lists the revised AG ° values
for the Cp2M compounds in addition to the vertical ionization energies measured by
o c 1
PES. An experimental error of ± 1.5 kcal mol is assigned to each AG° value
based on experimental uncertainties of the reference compounds and errors in the

22

167.8
(4.4)
<0}-NMe2
163.4
160.1
^S>-NMe2
159.4 *
-^Q>-NEt2
156.3--
t 1 -2
164.6-
3.3
2.7
160.6
11.1
3.1
(3.8)
3.2 “
0.8
2.9 4.7
1.1
•KA
1.4
N—N
145.6
i
4.2
1-8Í i
’
153.1-
150.2-
^,149.1 -
148.9
147.7 -
—j 43.8-
13—142.5.
Cp2Ru
Cp2Os
Cp2V
Cp2Fe
.(CpCH2NMe2)CpFe
(EthylCp)CpFe
(n-butylCp)CpFe
(t-butylCp)CpFe
(MeCp)2Fe
Cp2Ni
Cp2Mn
Figure 2-2 Free energy of ionization ladder for some metallocenes. Values of
AGe[° lie adjacent to arrows and AG ° values are next to the compound.
All values are in kcal mol'1 .

23
Table 2-1. Ionization Energetics Data for Some Metallocenes.
Cp2M
AG^b
AH°*
A5ioC
vIP (PES)a
Cp2V
154.5
-154
-0 ± 0.3d
155.7e
Cp2Fe
153.1
157.2
11.6f
158.2g
Cp2Mn
142.5
-
-
159.3g
Cp2Ni
143.8
-
-
149.9s
Cp2Ru
164.6
165.4
5.01*
171.8g
Cp2Os
160.6
161.1
5.011
164.9*
a. Units are kcal mol'1.
b. Determined from the derived free energy of ionization and the estimated entropy of
ionization at 350 K.
c. Units are cal mol^K'1.
d. The AS ° value is assumed to be equal ASelec° only estimated from uncertain ion
ground state.
e. See reference 37.
f. Determined from a van’t Hoff plot for the Cp2Fe/DET couple assuming
A5i°(DET) = 2.3 cal mol^K'1.
g. See reference 35.
h. Value assumed to be equal to statistical mechanics value for A5¡ 35Q°(Cp2Fe).
i. See reference 36.

24
measured partial pressures of the neutral gases. Estimates of AH° and A5¡° for several
metallocenes are also included in Table 2-1 when sufficient data were available.
The AGj° for Cp2Ni and Cp2Mn were referenced to the AG¡° for 1,1’-
bipyrrolidine which is anchored to the AG¡° value of DMA.40 Nelsen has suggested a
value for &H ° = 146.9 kcal mol'1 for 1,1’-bipyrrolidine based upon PHPMS ETE
studies. The AH ° value chosen was used as a reference for nickelocene and
manganocene in an earlier paper.16 From temperature dependent studies performed by
Nelsen and Mautner, a value for A5¡° = 3.2 cal mol^K'1 was reported.40 Based on the
assumptions stated earlier concerning the reference compounds, a value for AG¡ 350o =
145.8 kcal mol'1 has been estimated. The ferrocene derivatives help extend the
equilibrium ladder from free energies of ionization of the alkylaniline compounds
down to manganocene and the internal agreement is very good (± 0.3 kcal mol'1).
The temperature dependence of the ET equilibrium constants are shown as
van’t Hoff plots in Figure 2-3 and the derived thermodynamic parameters are also
presented in Figure 2-3. The ionization enthalpy and entropy values for ferrocene are
based on the derived thermodynamic parameters values from the van’t Hoff plots.
Insights into the Free Energy of Ionization of Ferrocene
The electron-transfer equilibrium reaction of ferrocene and N,N-
diethyltoluidine, DET, has been studied by Mautner by using PHPMS; a value of
AGet,35o° = '0-9 kcal mol'1 has been estimated for the couple.43 Mautner’s study
covered a temperature range of 450 - 650 K, and A//et° and ASct° values have been

In(Keq)
25
Figure 2-3 Van’t Hoff plots for selected metallocene electron-transfer equilibrium
couples. Values of AHet° and ASel° for the reaction couples are as
follows. For the Cp2Fe/DET couple, AHe° = +0.05 ± 0.47 kcal mol'1
and ASet° = 9.3 ± 1.1 cal mol^K'1. For the EtFc/Fc couple, Ht° and
ASc° are 3.1 ± 1.4 kcal mol'1 and 0.96 ± 3.6 cal mol^K'1. For the
Cp2Ni/Cp2Mn couple, Atfet° = -3.2 ±1.5 kcal mol'1 and A5et° = -5.7 ±
3.0 cal mol^K'1. DET = N,N-diethyltoluidine, Fc = ferrocene, EtFc =
ethylferrocene.

26
extrapolated to 350 K, which is the temperature at which the majority of FTMS
studies were performed. From the present study, the estimated value of AGet 350° for
reaction 2-5, is -3.2 ± 0.5 kcal mol'1 yielding a value of AGi>35o0(Cp2Fe) = 153.1 ± 1
kcal mol'1. This value differs from a values from a previous study which reponed
AGet° = -2.8 kcal mol'1.18
Cp2Fe + DET* ** DET Cp2Fe+ (2-5)
The larger equilibrium constant for equation 2-5 for the FTMS study is consistent with
the observation that ferrocene did not come to equilibrium with N,N-dimethyltoluidine,
DMT, where an estimated AGet° value > 5 kcal mol'1 is expected. The Cp2Fe/DMT
reaction couple was studied by Mautner with an estimated AGet° = -3.9 kcal mol'1 at
429 K. Several electron-transfer equilibrium reactions that have been initially studied
by high pressure mass spectrometry have been repeated in our laboratory with good
accuracy. The electron-transfer reaction for p-cyanonitrobenzene/benzoquinone couple,
(p-CNNB/BQ), equation 2-6, was studied by Grimsrud et al.61 and AGet° = -4.0 ± 0.8
kcal mol'1 was derived for the PHPMS study.
p-CNNB' + BQ * BQ- + p-CNNB (2-6)
The same reaction was studied by using FTMS and the derived AGet° = -3.8 kcal
mol'1 was determined, in good agreement with the PHPMS results.

27
The difference in the PHPMS and FTMS experiments for the Cp2Fe/DET
couple is much larger than expected for such comparisons and this led us to examine
the reaction couple in more detail. Mautner has determined the temperature
dependence of the equilibrium constants for reaction 2-5 in the 450 to 650 K range,43
therefore the same experiment was performed by using FTMS. The thermodynamic
parameters for the reaction couple derived from the PHPMS are AHc° = -0.1 kcal
mol'1 and ASet° = 2.2 cal mol^K*1. An FTMS temperature dependence study would
serve to illustrate the origins of the differences for AGet° between two studies and
would provide thermodynamic data that would overlap in temperature with the
PHPMS work.
The van’t Hoff plot for the Cp2Fe/DET couple for reaction 2-5 is shown in
Figure 2-3. The temperature range in the present study is from 350 to 520 K.
Consistent with the PHPMS study, the reaction displays a minor temperature
dependence.43 The derived thermodynamic values for the reaction are AH ° = 0.05 ±
0.47 kcal mol'1 and ASet° = 9.3 ± 1.1 cal mol^K'1 for the FTMS study. Values were
determined from linear regression and are reponed at the 95% confidence limit. The
AHe° value for both studies are within experimental error limits; but the ASet° values
are significantly different, the FTMS values being much more positive, and this is
obviously the source of error in the AGet° values.
Based on the assumptions outlined for the reference compounds, a value for
AH° = 157.2 ± 1.5 kcal mol'1 (6.82 eV) is obtained for ferrocene. This value is 0.2
kcal mol'1 higher than the value obtained by Mautner because of the difference in the

28
chosen A/Z^CDET). The adiabatic ionization energy obtained for both studies is well
within experimental uncertainty and a value for AHx 0 can be give as 6.82 ± 0.08 eV.
The vertical ionization measured by photoelectron spectroscopy by various groups is
6.88 ±0.1 eV.35,37,38 An adiabatic ionization potential has not been reported for
ferrocene previously due to lack of vibrational fine structure in the PES valence
ionization manifold. Rabalias has reported vibrational fine structure for Cp2Fe with 35
meV separation,62 however this has not been consistently resolved by other groups.
The source for the difference in the AS° values of the PHPMS and the FTMS
studies is unclear; however, several experimental observations favor a higher value for
the entropy of ionization than that reported by Mautner. The temperature dependence
of other electron-transfer equilibria involving molecules and ions that have well
established thermodynamic constants have been examined by using FTMS, and the
entropy changes for the reaction couples are within ± 5 cal mol^K'1 of the expected
value.10 The difference in ASet° value for the p-CNNB/BQ reaction couple (reaction
2-6) studied by PFIPMS and FTMS is ~4 cal mol^K'1.
The temperature dependence of the electron-transfer equilibrium reaction of
(EtCp)CpFe/Cp2Fe was examined to further assess our A5et° values for the DET/Cp2Fe
reaction. Based on the assumptions made for the reference compounds, A5’i°(DET) =
2.3 cal mol^K'1, therefore the A5¡°(Cp2Fe) = 11.6 cal rnol^K'1. Assuming that the
AS ° for ethylferrocene is equivalent to that for ferrocene, the derived ASet° for the
(EtCp)CpFe/Cp2Fe electron-transfer reaction should be ~0 cal mol^K'1. The derived
entropy change for the reaction couple was found to be approximately zero within

29
experimental error. Parameters for the reaction are AHe° = 3.06 ± 1.43 kcal mol'1
and ASet° = 0.96 ± 3.6 cal mol^K'1 at the 95% confidence limit.
The organic reference compounds used in this work have AG¡° values that have
been determined from ETE methods from PHPMS40 and FTMS59 studies. Mautner
has determined AG¡° for l,l-bipyrrolidine40a and the hydrazine has been linked to the
alkylaniline derivatives by ferrocene and its derivatives. The internal consistency of
the ladder is within 0.3 kcal mol'1. Since both the PHPMS and the FTMS studies
have used the same reference compounds in ETE measurements, and the internal
consistency of the present equilibrium ladder is well within experimental uncertainty,
the equilibrium constants obtained by FTMS for the metallocene are not expected to
have large random errors.
Origins of Possible Experimental Uncertainties in the Thermodynamic Constants
for Ferrocene
A possible explanation for the difference between the equilibrium constants for
the ETE reaction of DET/Cp2Fe would be inaccurate measurement of the parent ion
intensity ratios due to discrepancies in the ion detection. In PHPMS studies performed
by Mautner, ion sensitivities are frequently calibrated by comparing fragmentation
spectra with reference spectra to insure that the mass detector is not giving biased
sensitivities.63 Reponed errors due to mass detection drift in the PHPMS are ± 0.3
kcal mol'1 for typical ETE reactions. In FTMS experiments, mass differences between
two ions can lead to differences in detection sensitivities especially if the mass
difference is extremely large.64 In this work, ion detection parameters were adjusted

30
to give the maximum signal for both ions. The total ion count (7(Cp2Fe+) + /(DET*))
was monitored with time and the change was determined to be small, ca. 10%. The
variation of the ion intensities during the approach to equilibrium is not expected to
significantly affect to the experimental uncertainties for our systems.
The most likely source of error is the measurement of the partial pressures of
the neutrals. Both techniques use different methods to determine pressures of the
reactant gases. It is therefore possible that systematic errors for one or both of the
methods exist, leading to a difference in the equilibrium constant. The difference in
the A”cq values for the two experiments is equivalent to a factor of 25 in the pressure
ratios. Derived enthalpy changes are not dependent on pressure ratios; therefore, it is
of no surprise that the A7/et° values for the two studies are in agreement Pressures
are measured directly with an ion gauge in the FTMS studies. In the PHPMS work, a
solution containing known concentrations of reactants is introduced into a heated bulb
and partial pressures of the reactants are calculated.43,63 Measured partial pressures in
the FTMS studies are not expected to vary by more than ± 30%, resulting in a ± 0.2
kcal mol'1 error in the AGet° value.63 In the PHPMS studies, pressures are monitored
by measuring K at various partial pressure ratios. Mautner reports that experimental
error due tp pressure fluctuations are ± 0.5 kcal mol'1.
Evaluation of the Electron-Transfer Reaction Rates for the Cn2Fe/DET Couple
The rates of the ion-molecule reactions for the Cp2Fe/DET couple were
examined to further support that the source of error between the PHPMS work and the

31
FTMS studies stems from inaccurate pressure determinations. Electron-transfer
reaction kinetics for ferrocene with DMT and DET have been investigated by using
PHPMS. Mautner reports values for kf for reaction 2-5 of 1.2 x 10'9 cm3 molec'V1
at 461 K and 1.7 x 10‘9 cm3 molec'V1 at 429 K for the Cp2Fe/DMT couple (reaction
2-7). By using FTMS, kf = 2.5 ± 0.5 x 10'10 cm3 molec'V1 for reaction 2-5 was
Cp2Fe + DMT* -> DMT + Cp2Fe+ (2-7)
determined. The FTMS value of kf for reaction 2-7 was determined to be 1.3 ± 0.3 x
10'10 cm3 molec'V1. Mautner’s rate constant for reaction 2-5 is slightly faster than
the Langevin collision limit while the present values for kf is ca. 25% of the Langevin
collision limit. The two rate constants for reaction 2-5 differ by a factor of 5 and do
not resolve the disagreement in the derived AG ° values for ferrocene.
For example, if the PHPMS value for Keq is assumed to be correct, then the
pressure ratio determined in the FTMS work, PlDETy/yCp^e), is too large by a
factor of ~25. This would lead to an overestimation of kf by a factor of ~25,
depending on the errors in the absolute partial pressures of DET and ferrocene. Thus,
the kf values for the two experiments would diverge if the pressure of ferrocene, for
example, was underestimated.
By examination of equation 2-8, if the pressure for ferrocene was incorrectly
underestimated, the equilibrium constant, and thus AGet° would be too large since
there exists an inverse relationship between P(Cp2Fe) and K^. Moreover, an

32
Kt = In [(P(DET)/P(Cp2Fe)) x (/(Cp2Fe+)//(DET+))]
(2-8)
underestimation of the reaction pressure of ferrocene would lead to a value for kf that
is slower than determined. Reaction pressure and k{ are inversely related, therefore if
the total pressure of a reaction is actually greater than measured, the rate constant for
the reaction will be slower than the experimentally determined kf value. Thus, it
obvious that comparison of the rate constants for the Cp2Fe/DET couple do not
resolve the discrepancies for the two experiments.
Based on previous FTMS studies involving electron-transfer reactions of
metallocenes,14 it seems unlikely that such large errors in the measured partial
pressures could occur. The same electron-transfer methods have been applied to the
study of metallocene self-exchange and cross reactions.14 Estimated reaction
efficiencies for exothermic cross reactions were in the 0.5-1.5 range which suggests
that partial pressure errors may be incorrect by as much as 50%. Reaction efficiencies
are given by equation 2-9, where an estimate of kL for the metallocenes, the Langevin
collision rate, is 1.0 x 10‘9 cm3 molec'V1.14
Efficiency = kf/kL
(2-9)
An observed efficiency of ~0.2 is not unexpected for a reaction with low
exothermicity. The reported FTMS self-exchange rate constant for Cp2Fe+^° is 2.7 x
10'10 cm3 molec'V1 (0.27 efficient)14 which is consistent with the present kf for the

33
DET/Cp2Fe couple. Alternatively, large pressure errors in the PHPMS work also seem
improbable since ferrocene in not susceptible to thermal decomposition in the
temperature range used.65 Although there are distinct differences in the observed
forward rates for reaction 2-5, the kinetic data does not expose the origin for
difference in the two experiments.
A Brief Survey of Proton-Transfer Reactions of Ferrocene
The kinetics and thermodynamics of the protonation of ferrocene have been
studied by several groups by using PHPMS43,66 and ion cyclotron resonance mass
spectrometry.67 Recently, PHPMS proton-transfer equilibria studies performed by
Ikonomou and Kebarle assessed a value for AGB° for ferrocene of 195.2 ± 1.0 kcal
mol'1 at 500 K.66 A value for AGB°(Cp2Fe) = 195.0 ± 1.5 kcal mol'1 at 600 K has
been estimated from PHPMS data obtained by Mautner.43 The term AGB° is the gas-
phase bacisity (the free energy of proton attachment), and differs from the proton
affinity, which is an enthalpy change, AHB°, referenced at 298 K.
The proton-transfer reaction 2-10 was studied by using FTMS in order to gain
additional information concerning discrepancies in the AG¡°(Cp2Fe) values reported
here and by Mautner. The proton-transfer equilibrium reaction 2-10 has also been
studied by both Mautner43 and Ikonomou66 and the reported free energy changes are
-1.2 and -1.9 kcal mol'1 respectively. A value of AG350o = -1.5 kcal mol'1 was
PyrroleH4 + Cp2Fe = Cp2FeH+ + Pyrrole
(2-10)

34
derived from FTMS studies of reaction 2-10 and this value is consistent with the two
PHPMS AG° values. The difference in the AG° values for the FTMS study and the
PHPMS work results in a factor of ~1.5 in the measured equilibrium constants. A
factor of 25 in ATeq for the proton-transfer reactions, which is the difference in the Kcq
values for reaction 2-5, would yield a value of AG3500 ~3.6 kcal mol'1 for reaction 2-
10. As all three studies yield equivalent AG° values for reaction 2-10, evaluation of
proton-transfer reactions does yield inforamtion concerning the discrepancies in the
AGj° values for ferrocene. Furthermore, the AGB° estimated by Beauchamp and
Stevens67 from ICRMS experiments is consistent with all three proton-transfer studies.
Intramolecular Entropy Changes for the Ferrocene/Ferrocenium Couple
In order to further understand the uncertain entropy change for the Cp2Fe+/0
couple, detailed statistical mechanical analyses have been performed to provide an
accurate estimate for AS¿°(Cp2Fe). Table 2-2 lists values for translational, rotational,
vibrational, and electronic entropies of ferrocene and ferrocenium ion at several
temperatures. Complete vibrational analyses for all 57 vibrational modes of ferrocene
have been reported by Bodenheimer and Low68 and the measured frequencies were
used in the vibrational entropy analysis. Although a complete vibrational analysis for
the ferrocenium cation has not been performed, sufficient vibrational frequency
assignments for Cp2Fe+ have been reported by several groups to allow for an estimate
of 5vib°(Cp2Fe+).69'71 Additionally, crystal structures for both ferrocene and

Table 2-2. Calculated Entropies and Integrated Heat Capacities for Ferrocene and Fcrroccnium Ion at 298, 450, and 600 K.
Cp2Fe (298K)
Cp2Fe+ (298K)
Cp2Fe (450K)
Cp2Fe+(450K)
Cp2Fc (600K)
Cp2Fc+ (600K)
co a
° irans
41.6
41.6
43.6
43.6
45.0
45.0
co a,c
^ rot
24.3
24.4
25.5
25.6
26.4
26.6
co a
° int rot
5.24
5.24
5.75
5.75
6.12
6.12
co a,d
^ vib
14.88
17.19
30.17
32.97
45.84
48.89
co a,e,f
° elec
Alg
1A
Aig
N.
:E2S
q = 1
q = 2.07
q = 1
q = 2.22
q = i
q = 2.55
S° = 0
S° = 1.7
S° = 0
S° = 2.02
s° = 0
S° = 2.25
co a
° total
86.0
90.1
105.0
109.9
123.4
128.9
AH°X - AH°0Kb
5.3
5.7
12.3
13.0
21.7
22.6
a. Units are cal mol 'K'1.
b. Units are kcal mol.
c. Entropy of rotations calculated by assuming Cp-Fe behaves as a linear rigid rotor.
d. Vibrational entropy was calculated for all vibrational modes.
e- S°ckT = -RIn(q) + £lherm,T-
f. The zE2g orbitals are split due to spin-orbit coupling (-700 cm ). See reference 77.

36
ferrocenium cation have been reponed providing information needed for determining
of the rotational entropy change for the Cp2Fe+//° couple.72,73.
Low-frequency metal-ligand vibrations can contribute significantly to entropy
changes of metal complex redox processes if M-L frequency shifts drastically change
upon oxidation or reduction. For example, changes in the skeletal Co-N vibrations for
the Co(NH3)62+/3+ redox couple result in an estimated entropy change of 12.6 cal
mol^K'1.74 The change in the asymmetric Tlu vibration alone accounts for ~3 cal
mol^K'1 or ~25% of the total skeletal vibrational entropy change for the
Co(NH3)62+/3+ couple.74
The doubly-degenerate Cp-Fe-Cp bend has not been reported previously, but
the corresponding mode for ferrocene has a relatively low frequency, d22 = 179
1 ¿"Q
cm . Because of the importance of this low frequency skeletal mode, rather than
estimate a value for v22, the Cp-Fe-Cp frequency for ferrocenium ion was measured.
Our far-infrared studies have conformed \)22 for ferrocene and tentatively assign a
value of 135 cm'1 for u22(Cp2Fe+). The Fe-Cp distance increases upon oxidation, and
this distance increase is expected to result in a lowering of the frequency for the
metal-ligand bending mode.72 This single shift in the bending mode contributes 1.0
cal mol^K'1 to the ASi°(Cp2Fe't70) at 298 K. The results of the far-IR studies for
several ferrocenium salts are given in Table 2-3.
An estimate of ASet° for the electron-transfer equilibrium reaction at 450 K can
be determined by combining the calculated entropy of ionization for ferrocene (4.9 cal
mol^K'1) with the assumed value of AS^DET) = 2.3 cal mol^K"1 giving a value for

37
Table 2-3. Vibrational Frequencies For Various Ferrocenium Salts
Compound
Frequency (t)22)
Cp-Fe-Cp Bend
S°(u22)a
cal mol^K'1
(Cp2Fe+)(Cl')
135 cm'1
5.7
(Cp2Fe+)(BF4-)
140 cm'1
5.6
(Cp2Fe+)(PF6-)
130 cm'1
5.8
Cp2Fe
179 cm'1
4.7
a. Contribution of u22 doubly degenerate mode to the vibrational entropy at 298 K.

38
A5et° = 2.8 cal mol^K'1 for reaction 2-5. Assuming that the statistical mechanics
calculations provide a good estimate for A5i°(Cp2Fe), this analysis supports the
PHPMS temperature dependence study of reaction 2-5. However, use of solid-state
vibrational frequencies for gas-phase ferrocenium ion may not be valid, therefore the
estimated value for ASj° may be different from the true value. Gas-phase vibrational
frequencies for metal-ligand modes may be significantly shifted to lower frequencies
relative to the solid-state since crystal lattice stabilization is lost in the gas-phase.
Vibrational Entropy Calculations for the Ferrocene/Ferrocenium Couple
Since only a limited number of published vibrational frequencies for Cp2Fe+
exist, several frequencies for the cation were estimated from the known vibrational
frequency changes of ferrocene and ferrocenium. From known vibrational frequencies
of the ion and neutral that were measured (Table 2-4), the percent change in the
vibrational force constant was calculated. It was assumed that for a specific
vibrational mode, the percent change for an anologous mode which has not yet been
measured would be equivalent. For example, the symmetric ring-distortion mode, t)2g,
for both the cation and the neutral have been measured,68 where "u28 is 597 cm-1 for
ferrocene68 and 471 cm'1 for Cp2Fe+.69 The corresponding force constants are 3.26
mdyne Á'1 and 2.04 mdyne Á'1 respectively. Equation 2-11 was used to determine
vibrational force constants from measured frequencies where u is convened to energy,
E, p is the reduced mass for the vibrational mode, k is the vibrational force constant,
75
and h is Plank’s constant

39
E = (hlln^iklvf1 (2-11)
Thus, as demonstrated by the scheme below, an estimate for the vibrational force
constant for the unsymmetrical u34 mode, which is the asymmetric analogue to t^g,
was calculated. Note that the u34 and the \)28 m°des correspond to degenerate low-
energy ligand distortions perpendicular to the principle molecular axis.
&(Cp2Fe) for u28 = 3.26 mdyne Á"1
£(Cp2Fe+) for u28 = 2.04 mdyne Á'1
A^g = 0.63 and assume A/^g = M34 (degenerate modes)
&(Cp2Fe) for u34 = 2.97 mdyne Á'1
£(Cp2Fe+) for t)28 = (0.63 * 2.97) mdyne Á'1 = 1.86 mdyne Á"1
From equation 2-11, d34 = 450 cm"1.
Vibrational entropy contributions for individual frequencies were calculated
from equation 2-12. The value of u is dependent on the frequency of the vibration
and the absolute temperature, shown in equation 2-13. From equation 2-12, entropies
for all vibrational frequencies were calculated for ferrocene and ferrocenium ion.
Table 2-4 lists the frequencies used in the vibrational entropy calculations; literature
5vib° = R[(u/tu - 1) - In (1 - e"u)] (2-12)
u = 1.4387(u/7’)
(2-13)

40
Table 2-4 Vibrational Frequency Data for Ferrocene and Ferrocenium Cation
No.a
Cp2FeCu)b
Cp2Fe+(u)
Cp2Fe+ Reference
1
3110
3110
Estimated
2
814
805
Estimated
3
1102
1102
Estimated
4
309
304
Ref. 70
5
1255
1258
Estimated
6
c
c
c
7
1250
1250
Estimated
8
3103
3100
Ref. 69
9
820
782
Ref. 69
10
1110
1110
Ref. 69
11
478
418
Ref. 69
12
3068
3068
Estimated
13
998
1005
Estimated
14
844
833
Estimated
15
1410
1412
Estimated
16
389
398
Ref. 70
17
3077
3100
Ref. 69
18
1005
1017
Ref. 69
19
855
846
Ref. 69
20
1410
1420
Ref. 69
21
492
490
Ref. 69
22
179
135
d
23
3100
3100
Estimated
24
1191
1192
Ref. 69
25
1058
1058
Ref. 69
26
1356
1356
Estimated
27
897
874
Ref. 69
28
597
471
Ref. 70
29
3085
3085
Estimated
30
1189
1192
Ref. 69
31
1055
1055
Ref. 69
32
1351
1351
Estimated
33
885
874
Ref. 69
34
569
450
Estimated
a. Vibrational frequency assignments from reference 68.
b. All Cp9Fe vibrational frequencies from reference 68.
c. Torsional vibration determined as internal rotation.
d. Measured frequency.

41
sources are included when available. All frequencies for ferrocene used in this work
¿o
were measured by Bodenheimer and Low.
Vibrational entropy calculations based on spectroscopic data assume that
molecular vibrations behave as harmonic oscillators, therefore the displacement force
is equal to the restoring force.55 Atomic systems typically deviate from harmonic
conditions and corrections for anharmonicity must be considered for more complete
calculations. However, the harmonic formula is an accurate first approximation for
Svib° quantities.
Translational and Rotational Contributions to the Entropies
The translation entropy of an ideal gas is dependent on the molecular weight
and the temperature of the species. For ionization processes, the molecular weight of
the ion and the neutral are essentially equal, therefore the translational contribution at
temperature T to AS° is 0 cal mol^K'1. In order to report accurate AS° estimates for
Cp2Fe and Cp2Fe+, translational entropies were determined for completeness. The
Sackur-Tetrode equation76 (equation 2-14) was used to estimate 5^^° values, where
5trans° = RW2 ^ M + 5/2 In 7^ - 2.315 cal mol^K'1 (2-14)
M is the molecular weight, R is the ideal gas constant, and T is the temperature.
Rotational entropies were calculated by assuming that ferrocene behaves as a
linear rigid rotor and that the Cp rings are point masses of 65 a.m.u. Rotational

42
entropies are dependent on the moments of inertia / of the molecule, therefore by
choosing the Fe atom to he at the center of gravity of the molecule, the moments of
inertia for Cp2Fe were calculated. Since the Fe atom is defined to lie at the origin of
the cartesian axes, Ix and /y are equal (corresponding to rotations perpendicular to the
principle axis). The rotational moment of inertia along the principle axis was defined
as /z. Flere, the Cp rings where considered as mass with five-fold rotational symmetry
rather than point masses. The product of the three principle moments of inertia yields
a determinant D which is used in estimating Srot° (equation 2-15).55 The symmetry
number o for ferrocene is 10 based on D5d molecular symmetry.72 Note that D5h
symmetry, for an eclipsed structure, will also yield a = 10.
5rot° = R( 1/2 In D x 10117 + In T - In a) - 0.033 (2-15)
Because of the low energy barrier, the Cp rings of ferrocene freely rotate at
ambient temperatures.72 Therefore, rather than consider rotation of the Cp ligands as a
torsional vibration, the internal rotational entropy was estimated as a restricted internal
rotation. The potential barrier of rotation V = 0.9 kcal mol'1.72 The moment of
inertia for the system is given as IT = (7z)2/2/z, where /z is the rotational moment of
inertia along the principle axis. Therefore, using the method outlined by Brewer and
Pitzer,55 5int° for was estimated. The internal rotational entropy for a restricted
rotation is calculated from the entropy of free rotation along the principle axis minus

43
the restriction energy associated with rational barrier. Since Iz for both ferrocene and
ferrocenium are equal, = 0 cal mol^K'1.
Electronic Entropy Considerations for the Ferrocene/Ferrocenium Couple.
The electronic energy level separations of a molecule are usually large, and at
ambient temperatures only the ground state is thermally populated.55 The electronic
partition function q is used to determine Selec values where qe = g, the spin degeneracy
of the ground state. Thus, in the absence of thermally accessible energy levels, the
electronic entropy is given by S = R ln(g). In the case of ferrocene, the spin
multiplicity is 1 and Selec = 0 cal mol^K'1. Ferrocenium is a 2E complex, however
the “E state is spit due to spin-orbit coupling with ~700 cm' separation. Since the
split states are thermally accessible in the 298 - 600 K range, the value of <7elec was
evaluated which includes the theraml energy of the split electronic states. The
electronic entropy of Cp2Fe+ is given as Selec T = ¿theim.T + ^ where ¿^enn,T *s
the energy of the thermal population of the split states at temperature T.
Ionization Free Energies of Ruthenocene and Osmocene
Discussion of the ionization free energies of ruthenocene and osmocene has
been reported previously.16 The observed trend in AG° values for the iron triad is
OC *70
consistent with vertical ionizations ’ (IP Cp2Fe < Cp20s < Cp2Ru).
Lichtenberger and Copenhaven were able to obtain vibrational fine structure for
the first ionization manifold of osmocene.78 An average spacing of 42.1 meV for the

44
spacing of vibrational energy levels in the first ionization manifold for the cation was
used to determine the adiabatic ionization energy for osmocene of 161.1 kcal mol'1.
Assuming the ASj° for osmocene is equal to ASj0(Cp2Fe) estimated from statistical
mechanics, the estimated osmocene aEP for Cp2Os is 162 kcal mol'1. Thus the two
techniques are within experimental error.
Free Energies of Ionization for Vanadocene, Manganocene, and Nickelocene
The sharpness of the first ionization band in the PES of vanadocene indicates
that the difference in the equilibrium geometries of the ion and neutral are small.37'38
A recent PES of Cp2V assigns the vIP = 155.7 ±0.1 kcal mol'1.16 As the first
ionization manifold is a sharp band (width = 0.19 eV), the alP is closely approximated
by the vIP. The AG° value from ETE studies is 154.5 ±1.5 kcal mol"1. Assuming
AS° is predominately the electronic entropy change for vanadocene (A5vib ~0 based
on the PES) the AH° is estimated at 154 ± 2 kcal mol"1. The estimated AH° value is
slightly less than the estimated alP but lies within the band envelope.
The AG ° for nickelocene is in agreement with PES data; however,35 the AG °
of manganocene is -18 kcal mol'1 lower than the reported vIP values.35’37,38 The
relaxation energy for Cp2Mn is large compared to other metallocenes. Manganocene
exists as a high-spin complex with only a small percentage of the complex in a low-
spin configuration. From PES studies the vIP of the low-spin complex has been
assigned (-144 kcal mol'1) and is in agreement with the AG¡°(Cp2Mn).

45
The temperature dependence of the Cp2Mn/Cp2Ni ETE reaction couple was
investigated to assess the expected negative intramolecular entropy change for the
Cp2Mn+/0 couple. The van’t Hoff plot in Figure 2-3 indicates that the Cp2Mn/Cp2Ni
ETE reaction couple is strongly temperature dependent with a negative ASet°. Details
of the Cp2Mn/Cp2Ni ETE couple have been presented elsewhere.13 The origin of the
large negative entropy change is primarily attributed to A5j°(Cp2Mn). The estimated
change in the metal to ring-centroid distance for Cp2Mn/Cp2Mn+ is -0.25 Á80
resulting in a large change in vibrational and rotational entropy for the Cp2Mn+/0
couple. The Mn-C distance will decrease upon oxidation leading to a negative AAvib°.
As spin degeneracy is conserved, ASelec(Cp2Mn+,/0) is essentially zero. However, due
to spin orbit coupling of the 3E state of the ion, A5elec for manganocene will be
greater than that of manganocenium ion. This is consistent with a net negative
entropy change expected for Cp2Mn+/0. The rotational entropy change is estimated to
contribute 0.6 cal mol^K'1 to AS° based on the estimated change in the Mn-Cp
distance accompanying ionization. Thus estimated vibrational, rotational, and
electronic entropy changes for manganocene are all negative and are consistent with
the negative AAet° observed for the Cp2Mn/Cp2Ni ETE couple.16
Substituent Effects in Ferrocene Derivative Oxidations
Although extensive photoelectron spectroscopy studies of the ionization
ic
energetics of metallocenes have been reported, little is known concerning the effects
substituents have on metallocene thermochemistry. Therefore, various ferrocene

46
derivatives have been studied by ETE methods in order to assess the effect alkylation
of the cyclopentadienyl rings has on the ionization potentials of metallocenes. Free
energy of ionization data for several alkylferrocene complexes have been reported
earlier16 and a AG° value for ethylferrocene is reported here. Values of AG ° differ
from other values previously reported13 due to modifications in the AG ° values of the
reference compounds.
It is known from PES data that dimethylation or permethylation of
oc
metallocenes lower the ionization energies. The vertical ionization energy of
ferrocene (6.88 eV) is ~0.2 eV greater than the vIP of Ll’-dimethlyferrocene35 and
1 eV more endoergic than the vIP of decamethylferrocene.37 Attachment of alkyl
groups on the metallocene rings stabilizes the molecular cation relative to the neutral
compound, thus lowering the free energy of ionization of the derivative relative to the
parent metallocene.32,35 Electron-transfer equilibrium results for alkylferrocene
derivatives studied in the present work are shown in Figure 2-2.
Molecular ionization potentials of organic81 and organometallic compounds32,51
have been correlated previously with Taft substituent parameters. The aliphatic Oj
parameters were derived originally for substituted acetic acids, and they have been
used successfully to correlate IP data.51,52,81 For example, a plot of IP for benzene-R
chromiumtricarbonyl derivatives versus Gt(R), where R is an attached substituent,
shows a strong correlation line.51 The equation used by Levitt and coworkers for the
Taft analyses is given in eq 2-16, where IP(M-R) and IP(M) are the ionization
potentials for the derivative and the parent compound respectively, G¡ is the Taft

47
IP(M-R) = ajOjíR) + EP(M) (2-16)
parameter for R, and aj is the slope of the line. The Taft parameter for H is zero thus
the substituent effects are referenced to hydrogen.5 The slope indicates the
sensitivity of the ionization process to the change in substituents.81
A plot employing equation 2-16 for the ferrocene derivatives is shown in
Figure 2-4. The slope of the line is 57 ± 6 kcal mol'1 which can be compared to that
for the (RBz)Cr(CO)3 compounds, aj = 34.6 kcal mol'1, and alkylbenzenes, aj = 109.3
kcal mol'1.51 Therefore, alkylferrocene ionization energies are nearly twice as
sensitive to changes in alkyl substituents as the chromium arene complexes but are
less affected than the alkylbenzenes.51 The AG¡° data here is consistent with the
results of Matsumura-Inoue and coworkers who correlated PES data of Cp2Fe
derivatives with Taft parameters.32 The differences in substituent effect sensitivity for
various parent compounds can be rationalized by several factors, including proximity
of substituents to the site of ionization. Additionally, electronic coupling between
the alkyl a orbitals and the ionized molecular orbital may direct changes in the
ionization potentials with respect to the parent compound.16 It should be noted that
although the observed substituent effects in alkylferrocenes follow the trend of
expected "electron-releasing" ability of the alkyl substituents, the observed sensitivity
of ionization energies to substitution will not necessarily hold for other
metallocenes.16,19 Electron loss originates from different valence molecular orbitals
for the metal complexes.

Ferrocene Ionization Free Energies
48
Figure 2-4 Plot of AG ° values (kcal mol'1) versus alkyl Taft parameters (ar) for
several ferrocene derivatives. Asterisk indicates new Taft parameter for
CH2N(CH3)2.

49
Further, the electron-releasing character of the alkyl groups observed for the ionization
of neutral metallocenes does not necessarily apply in other circumstances. For
example, the substituent effects for electron attachment to neutral metal complexes has
been shown to be different than for ionization process. ’
With respect to the construction of Taft parameter correlations such as that in
Figure 2-4, it is notable that parameters for substituents with low ionization potentials
can be derived from the data for ferrocene derivative ionizations. For R = CH2NMe9,
the first ionization of the benzene derivative removes an electron from the nitrogen
lone-pair orbital and not from the benzene ring, which would be a more endoergic
process. The ionization of benzene is 9.25 eV, which is significantly more
endoergic than the ionization potential of N,N-dimethylaminobenzene (7.69 eV).27
However, in the case of (CpCH2NMe2)(Cp)Fe, the ionization occurs at the same site
as in the other alkylferrocenes, thus allowing derivation of a substituent (-0.049 ± 0.013).
Heterolytic and Homolytic Metal-Ligand Bond Disruption Enthalpies for
Metallocenes and Metallocenium Ions
Bond disruption enthalpies for several metallocenes have been reported
previously16 therefore only a brief discussion concerning M-Cp bond enthalpies will
be give here. Thermochemical cycles have been used to derive estimates of M-Cp
bond disruption enthalpies.10,16 Since AH0 values depend on ionization energetics
data, M-Cp bond enthalpies reported here differ from previously reported13 values due
to more accurate free energy of ionization data for the metallocenes and more accurate

50
assessment of auxiliary thermochemical data.16 Additionally, more detailed error
analysis was performed for the heterolytic and homolytic bond disruption enthalpies
and in most cases error limits were found to be lower than previously reported.13
Application of Thermochemical Cycles to Derive Estimates for Metallocene and
Metallocenium Bond Disruption Enthalpies
Thermochemical cycles for ionization processes of the metallocenes and their
corresponding ions used to derive bond enthalpies data are shown in Figure 2-5. The
bottom portion of the figure is a cycle used to derive solvation energetics for the
metallocenes which will be discussed later. Similar cycles have been used by
Buckingham and Sargeson to derive crude estimates for metal complex thermodynamic
quantities.83 Related thermodynamic quantities have recently been derived for
coordination complexes and complex ions by use of analogous thermochemical
cycles. In Figure 2-5, the term AX° represents any thermodynamic function
therefore entropy, enthalpy or free energy data can be incorporated in the cycles to
derive thermodynamic values.
In this work, average heterolytic bond disruption enthalpies, half the AH°
expressed in reactions 2-17 and 2-18, will be denoted as A//het°. Homolytic metal-
ligand bond cleavage will be denoted as A//hom° and represents XAAH° for reactions
2-19 and 2-20. From the thermochemical cycles used here only average bond
(2-17)
Cp2M(g) -4 M2+(g) + 2Cp(g)
Cp2M+(g) M3+(g) + 2Cp(g)
(2-18)

Cp2M(g) -> M(g) + 2Cp(g)
Cp2M+(g) -» M+(g) + 2Cp(g)
51
(2-19)
(2-20)
disruption enthalpies can be obtained. Further, enthalpy values for M-Cp bond
cleavage for consecutive ring cleavage will not be equivalent. For heterolytic
cleavage, removal of the first Cp' ligand will be less endothermic than removal of the
second ligand due to an increase in the electrostatics between the positively charged
metal center and the anionic Cp ligand. Similarly, AHhom° for the first M-Cp
homolytic cleavage will not be equal to AHhom° for the second M-Cp cleavage.
Homolytic bond energies for ferrocene have been studied by pyrolysis
techniques.85 The activation energy for the decomposition of ferrocene, reaction 2-21,
was monitored and based on the measured £a, the first bond dissociation bond
Cp2Fe —» CpFe + Cp (2-21)
enthalpy was estimated to be 95 kcal mol . Consequently, removal of the second Cp
ring is less endothermic by approximately 30 kcal mol'1 than removal of the first
ligand. Faulk and Dunbar have used photodissociation methods to arrive at a value of
85 ± 7 kcal mol'1 for the first homolytic cleavage of Cp2Fe+. Therefore, the AHhom°
for the second cleavage is more endothermic than the first dissociation. Increased
electrostatic attraction for the second Fe+-Cp cleavage accounts for an increase in the
second AHhom° for Cp2Fe+.86

52
M2+ (g) + 2Cp- (g) AX°¡(M2+,g) m3+ (g) + 2Cp- (g) + e-
Figure 2-5 Thermochemical cycles used to determine bond disruption enthalpies
and differential solvation free energies for metallocenes. The upper
portion of the cycle yields estimates for the average homolytic and
heterolytic bond disruption enthalpies for M-Cp cleavage. Comparison
of AG¡°(g) and AG¡°(soln) in lower portion of the cycle yields estimates
of differential solvation energies, AAGS0|V°, for Cp2M+/0 couples.

53
Combining Free Energy and Enthalpy Data in Thermodynamic Cycles
Free energy of ionization data has been used to derive heterolytic and
homolytic bond disruption enthalpies for metallocenium ions. The validity of
combining AGj° data for the metallocenes with enthalpy data referenced at other
temperature, i.e. 0 or 298 K, is dependent on the accuracy of substituting AG¡ 350°
values for AHx -,98°. This approximation relies on the premise that the entropy of
ionization is small with respect to the larger values of AH° at temperature T.
Manganocene is expected to have the largest |A5j°| of the metallocenes studied in
this work.16 Assuming |A5j°| = 12 cal mol^K'1, the entropy contribution to the free
energy of ionization at 350 K is 4 kcal mol'1. Furthermore, even with such a large
entropy change, the error estimated for manganocene is not expected to exceed 2 kcal
mol'1 per bond. Generally, the assumption that AG¡ 350° = AHx 29g0 for the other
metallocenes is acceptable since entropy contributions will be small with respect to the
larger values for the free energy of ionization. When compared to the absolute values
for heterolytic and homolytic dissociation, the error introduced by assuming AG; 350° =
A//j 298° will small, ca. 2-3% for the homolytic bond enthalpies and less than 1% for
the heterolytic enthalpies.
Heat capacity corrections are also expected to be small on going from 350 to
298 K. The heat capacity terms for ferrocene has been determined by statistical
mechanics and the difference in the AH ° values is only 0.1 kcal mol'1 from 350 to
16
298 K.

54
Heterolytic and Homolytic Bond Disruption Enthalpies for Metallocenes and
Metallocenium Ions
Bond disruption enthalpies for the Cp2V, Cp2Mn, Cp2Fe, and Cp2Ni have been
reported previously. However, as mentioned earlier, values reported here have been
refined due to more accurate AG ° values and auxiliary data. Derived bond enthalpies
for the metallocenes are presented in Table 2-5. Auxiliary data16 used in the
thermochemical cycles are presented in Table 2-6. Error limits for homolytic and
heterolytic dissociations take into account errors in the heats of formation of the
neutral27 and ionic species16 and errors in the free energies of ionization. Bond
enthalpies for ruthenocene, osmocene, and the substituted metallocenes are not
reported due to lack of reliable thermochemical data for AHf° of the neutral
organometallics and the alkylated cyclopentadienyl compounds.
Differential Solvation Free Energies for Metallocene Redox Couples
Differential solvation free energies, AAGsolv°, for several metallocene redox
couples have been determined through the application of thermochemical cycles
(Figure 2-5) by combining E^ data at 298 K to AG¡ 350° values. As mentioned earlier,
because of the relatively small entropy effects for the metallocene Cp2M+/0 oxidation
couples, the errors introduced by substituting AG¡ 350o for AG¡ 298° are expected to be
less than 1 kcal mol'1. Values for E^ have been used to derive AGi°(Cp2M)(soln);
from Figure 2-5, values of AAGsolv
° have been derived from the lower thermochemical

55
Table 2-5 Mean Bond Disruption Enthalpies for Some Metallocenes.
Cp2M
A//het[M2+-Cp']a
A//het[M3+-Cp']a
^h0m[M-Cp]a
A^homt^-Cp]
V
303 ± 3
563 ± 4
95 ± 2
95 ± 3
Mn
286 ± 3
604 ± 5
59 ± 2
74 ± 4
Fe
318 ± 3
593 ± 4
79 ± 1
91 ± 3
Ni
326 ± 3
659 ± 4
67 ± 2
83 ± 3
a. Units are kcal mol'1

56
Table 2-6. Auxiliary Thermochemical Data Used in Thermochemical Cycles.
Process
V
Mn
Fe
Ni
AHf°[Cp2M]b
49 ± 2a
66 ± 2
58 ± 1
85 ± 1
AHf°[M]c
123 ± 2
67 ± 2
99 ± 2
103 ± 2
AHf°[M+]c
278 ± 2
238 ± 2
281 ± 2
279 ± 2
AHf°[M2+]c
616 ± 2
599 ± 2
654 ± 2
698 ± 2
AHf°[M3+]c
1292 ± 2
1376 ± 2
1361 ± 2
1509 ± 2
AHf°[Cp]b = 58 ± 1
AHf°[Cp‘]b = 19.6 ± 4
a. Units are kcal mol'1.
b. See reference 27.
c. See reference 8.

57
cycle. A negative value of AAGsolv° represents decreased exoergicity for the reduction
of a metallocenium ion in solution compared to the gas phase. An analysis of the
estimation of absolute electrode potentials for redox couples in solution has been given
earlier,54 and a similar approach has been used for the derivation of AAGsolv°
quantities in this work. Specifically, a value of 4.44 V has been used for the absolute
potential of the standard hydrogen electrode, £nhe°, an<^ n0 corrections for liquid
junction potentials have been applied to the E^ data for the metallocenes.54
Thermochemical cycles were used in the derivation of the absolute potential of the
normal hydrogen electrode. The largest source of error introduced in the derivation of
Z?NHE° was the free energy change for the solvation of the proton, AGsolv°, which can
not be measured directly.87 The value used for AGsolv°(H+) is based on
electrochemical experiments and the reliability of this value has been discussed
elsewhere.54,84 In addition, the stationary electron convention is used for both the gas-
phase and solution thermochemistry, although near 298 K the thermal electron
oo
convention yields similar results for AG ° values.
The electrochemical E^ values used in the estimates of AAGsolv° values are
shown in Table 2-7. Equation 2-22 was used to derive estimate for AGj°(soln) for the
metallocenes. The value £ref° is the potential of the reference electrode
relative to the normal hydrogen electrode, n is the number of electrons transferred in
AGi°(Cp2M)(soln) = -rc£[(£^(Cp2M) + £ref° + E^]
(2-22)

58
Table 2-7. Electrochemical E^ Data and Differential Solvation Energies for Some
Metallocenes Cp2M+/0 couples.
Cp2M+/0
£i¿a (solvent)
AG¡°(soln)b
-AAG°solvb
V
-0.55c,d (THF)
95
60
Cr
-0.67c (CH3CN)
92
36e
Fe
0.3 lc (CH3CN)
115
38
Co
-0.94c (CH3CN)
86
38e
Ni
-0.09° (CH3CN)
106
38
Ru
1.03f (CH2C12)
131
33
Ru
0.78d,g (CH3CN)
126
38
Os
0.86f (CH2C12)
127
34
Os
0.75d’g (CH3CN)
125
36
Mn
(0.13)h
(105)h
(38)h
a. Values reported in volts using 0.1 M Bu4NPF6 as supporting electrolyte against
SCE, except ruthenocene in 0.1 M Bu4TFPB against Ag/AgCl and osmocene in
0.1 M Bu4NBF4 against SCE.
b. Units are kcal mol'1. Estimated error limits ± 4 kcal mol'1.
c. See reference 89.
d. Irreversible oxidation potential.
e. Estimated from AG¡° values. See Chapter 5.
f. See reference 90.
g. See reference 92.
h. Estimated from data in Figure 2-6 and reported against SHE.

59
the electrochemical process (le' oxidations for the metallocenes) and F is Faraday’s
constant. Most of the Ex/l values were obtained from a single literature source89 and
were measured under common experimental conditions. The solvent is acetonitrile for
all quoted Ex/l values except for vanadocene (THF) and ruthenocene and osmocene
(CH2CI2).90’91 Values of E^ for ruthenocene and osmocene are also reponed in
CH3CN.92 Table 2-7 presents the derived differential solvation energies for the
metallocenes and the corresponding solution free energies of oxidation for comparison
to the gas-phase ionization energies. Differential solvation free energies were
calculated from equation 2-23, which originates from the thermochemical cycles in
Figure 2-5. A plot of free energy of ionization data versus the first transition row
-AAGsolv° (Cp2M+/0) = AGi°(Cp2M)(g) - AGi°(Cp2M)(soln) (2-23)
metallocenes is presented in Figure 2-6 and serves to demonstrates the periodic trends
for AGj0(Cp2M) in solution and the gas-phase. The average value for AAGsolv° for the
first transition row metallocenes is -38 ± 2 kcal mol'1, excluding vanadocene. The
AAGsolv° values for cobaltocene and chromocene were based on AG ° values
determined from FTMS ETE studies and are discussed in Chapter 5. The obvious
exception to the observed trend in differential solvation energies (Figure 2-6) is noted
for Cp2V oxidation,89 which has a AAGsolv° value more negative by ~20 kcal mol'1
relative to the other first transition row metallocenes. This additional stabilization of
the cation can be attributed to inner sphere coordination

Metallocene Oxidation Energy (kcal/mol)
60
Figure 2-6 Plot demonstrating the periodic trend of ionization energies for the first
transition row metallocenes. Gas-phase data (filled squares) include
AGj°values from this work. Solution AG ° values (crosses) were
determined through thermochemical cycles. The solvent is CH3CN
except for Cp2V in THF. An estimate of AG¡0(soln) for manganocene
(versus SCE) is included (open square).

61
QQ #
of solvent following oxidation. From the trends in Figure 2-6, a prediction can be
made for the unknown for Cp2Mn, -0.13 V vs. SCE.
Gieger has reported that electrochemical oxidation of Cp-V is quasi-reversible
in tetrahydrofuran, TIFF, which is consistent with there being no significant structural
changes in going from a d to d~ metal complex. However, the correlation of
AG¡°(g) versus AG^soln) clearly indicates that Cp2V lies outside the fit; thus, a
structural variation can not be ruled out. A structural change, similar to that shown in
Figure 2-7, could explain the lack of correlation for Cp2V. It can be rationalized that
in an attempt to increase the electron density around the electrophilic metal
Figure 2-7 Structure of an 18c'1 vanadocenium-THF complex
center, polar solvent molecules coordinate to the metal complex. The addition of two
THF molecules, for example, would lead to the formation of an 18 e' cationic
Cp2V-2THF complex.
A comparison of AAGsolv° results to values predicted by dielectric continuum
theory suggests that the solvation thermochemistry of Cp2M+/0 couples can be
adequately modelled by the Bom charging model.54 The Bom equation determines

62
the change in electrostatic free energy, AGel° when a charge on an electrostatic sphere
of radius reff is transferred in a vacuum to a sphere of equivalent volume in a solution
of dielectric constant D. For equation 2-24, the Bom equation, z denotes fundamental
charge of the ion (here 1+). The definition of the Bom equation is directly
comparable to the concept of differential solvation energy defined by equation 2-23.54
AGel° = (-166z2/reff)(l -1/D) kcal mol'1 (2-24)
The Bom model neglects the actual work required for an ion to pass from a vacuum
through the solvent barrier. However, this work is usually quite small compared to the
values for the differential solvation free energy.55
From crystallographic data, the radii for Cp2Fe and Cp2Ni are estimated to be
3.9 and 3.7 Á respectively.72 The reff value obtained from the Bom equation for a
AGel° = -38 kcal mol'1 in acetonitrile is 4.3 Á. From another point of view, the
structural model radius (3.9 Á) predicts a AAGsolv° value of -41 kcal mol'1. This
close agreement between the structurally estimated radii and the thermochemical radii
is consistent with relatively small specific interactions between solvent and metal
complex as well as the compact structure of metallocenes. The same conclusions were
obtained by Krishtalik et al. who used a AG° value of ferrocene based on the
photoelectron spectrum.93 For comparison, in the tris(acetylacetonate) metal
complexes, where polar solvent molecules can interpenetrate between the chelating
bidentate ligands, the experimental solvation energy is approximately twice the value

63
predicted from the structural model. The Bom model predicts a value for reff of 2.9 Á
for Ru(acac)30/' based on a value of 57.5 kcal mol'1 for AAGsolv0.84 However, the
crystallographic radius (maximum radius taken from Ru metal center to furthest ligand
proton) for the ruthenium complex is ~6 Á. A shortcoming of the Bom model is that
it assumes the charge is evenly distributed over the entire sphere which is not
necessarily true, especially for large metal complexes
Equation 2-24 predicts that AAGsolv° will be increased by ~3 kcal mol'1 for reff
= 3.9 Á when acetonitrile, D = 36, is replaced by methylene chloride, D =9. From
the reversible potentials for ferrocene, ruthenocene, and osmocene in CH2C12 given by
Hill et al.,90 lower solvation energetics for AAGsolv° values (-34 ± 1 kcal mol'1) are
estimated for the three couples. However, the difference in the AAGsolv° values
derived in the different solvents are in good agreement and are consistent with
experimental AAGsolv° values determined through thermochemical cycles.
Conclusions
Free energies of ionization have been determined for a number of gas-phase
metallocenes. These ETE data complement and extend information on the oxidation
energies of metallocenes obtained previously by electrochemistry34 and photoelectron
o c
spectroscopy. Further, the AG¡° values for the metallocenes are in agreement with
the vertical ionization energies measured by PES by several groups.37,38,78 The free
energy of ionization for manganocene is in agreement with the vIP for the low-spin
complex. The temperature dependence of the Cp2Ni/Cp2Mn couple is consistent with

64
a large negative entropy change for the manganocene ionization process. Comparison
of ETE data with PES spectra for Cp2V and Cp2Os yield the same values for adiabatic
ionization potentials within experimental error. ’ Overall, the good agreement of
AG ° data with PES data indicates that large experimental errors do not exist for the
ETE studies. With the exception of vanadocene, correlation AG¡° values in the gas
phase and solution for the first transition row metallocenes is linear. Inner-sphere
solvent coordination can be used to rationalize the disparity of the vanadocene couple
relative to the other metallocenes.89 An estimate for E^ for manganocene has been
made based on Ex/l values for the other first transition row metallocenes.
Detailed investigations of the ionization free energy for ferrocene support a
value of AG ° = 153.1 ± 1.5 kcal mol"1. Temperature dependent ETE studies on
selected equilibria have established experimental enthalpy and entropy of ionization
values for ferrocene. Although the present data establish a relatively large value of
ASj°(Cp2Fe), statistical mechanical analyses and previous PHPMS results43 suggest
that a smaller AS¡° value (~4-5 cal mol^K'1) is appropriate. The positive entropy
change for ferrocene ionization can be attributed to roughly equal contributions from
intramolecular vibrational entropy changes and changes in electronic entropy. The
vibrational contribution to the entropy change accounts for over 50% of the total AS °
at 298 K for the Cp2Fe+/0 couple. Temperature dependent investigation of the ETE
reaction couple EtFc/Cp2Fe indicate that the ASet° is ~0 cal mol^K'1 which indicates
that the AS° for ethylferrocene and ferrocene are equivalent. The ETE reaction couple

65
Cp2V/Cp2Fe indicates that the AG ° of ferrocene is less than vanadocene which
supports the lower AG¡0(Cp2Fe) value.
The AGj° data have been incorporated into thermochemical cycles to allow
estimations of bond disruption enthalpies for selected gas-phase metallocenium ions.
Estimates of differential solvation energies for several Cp2M+/0 couples have also been
derived from thermochemical cycles. An average value of AAGsolv° for the
metallocenes is -38 ± 2 kcal mol"1 with CH3CN as the solvent. Solvation of a gas-
phase metallocene decreases the ionization energy by a relatively constant amount
(~1.6 eV with acetonitrile as solvent). Solvation is therefore secondary to metal
ligation in determining the potential of Cp2M+/0 couple relative to the ionization of the
corresponding M2+ (g) ion (ligation of M2+ by two Cp* ligands reduces the ionization
energy by ca. 13 ± 1 eV).
The Bom equation, used for deriving estimates of differential solvation
energies, has been applied to the metallocenes. Values of AGel° are in agreement with
AAGsoivO(Cp2M+/0) values which demonstrates that the metallocenes have compact
molecular structures. Specific solvent interactions such as inner-sphere coordination of
solvent are minimal for the metallocenes.
Experimental Methods
Electron-Transfer Equilibrium Studies
Electron-transfer equilibrium studies were performed by using a Nicolet FT/MS
1000 Fourier transform ion cyclotron resonance mass spectrometer as previously

66
described.15'20 Briefly, the pressure of each compound was adjusted to establish a
workable pressure ratio to allow equilibria to be monitored. The time dependence of
parent ions formed from neutral molecules of known partial pressure was monitored as
the molecular ions underwent electron transfer with neutrals. Reactions were typically
followed to ca. 5 or more seconds. Apparent equilibrium was generally attained in
less than two seconds.
Reference compounds were sublimed into the FTMS high vacuum chamber
through a precision leak valve. The vapor pressure of most metallocene samples was
sufficient to allow for direct introduced into the high vacuum chamber through a
second leak valve. Ruthenocene, Cp2Ru, and osmocene, Cp2Os, were introduced by
using a heated solids probe positioned adjacent to the reaction cell. The FTMS
reaction cell was typically 350 K as measured by an Omega RTD thin film detector.
Positive ions were produced by electron impact at 9-12 eV with beam times ranging
between 5 and 25 ms. Ionization of metallocenes and organics occurs with some
fragmentation of the molecular ion (vide infra); other unwanted ions are formed by
ion-molecule reactions. Prior to study of electron-transfer reactions, several ion
ejections were required in order to select only the parent ions.
Since reactions were followed for at least 5 s, all ions formed as a result of El
were assumed to be effectively thermalized through ion-molecule collisions. At a
neutral pressure of 10‘6 Torr at 350 K, a typical metallocenium ion will undergo ca.
30 collisions s'1, which is believed sufficient to remove much of the excess rotational
and vibrational energy present due to the ionization process.14,16 Approach to

67
equilibrium was followed from endoergic and exoergic directions. Prior to reaction,
one of the parent ions was ejected from the reaction cell and the population change of
both parent ions was monitored at set time intervals. Equilibrium was deemed to have
been achieved when the ratio of the two parent ion populations reached a constant
value within experimental error.
Partial pressures of the various neutrals were determined by using an ion gauge
calibrated with an MKS baratron capacitance manometer (in the 10'5 torr range)
extrapolated to experimental conditions. In order to approach dynamic pressure
equilibrium throughout the vacuum chamber, the 300 L s'1 pumping speed of the
diffusion pump connected to the high vacuum chamber was reduced to ca. 75 L s'1.
Neutral gas pressures were calibrated for all reactants in open (75 L s'1) and closed
(no pumping) systems. It has been shown that partial pressure is independent of
neutral gas leak rate. A calibrated ion gauge connected to a Granville-Phillips
controller was positioned at the site of the reaction cell with the magnetic field off,
thus providing a field free vacuum system. The pressure measured at the middle of
the vacuum chamber where the reaction cell is located, was equivalent to metallocene
pressures determined at the remote ion gauge following pressure calibrations.94
Temperature Dependence Studies
The temperature dependencies of electron-transfer equilibria were investigated
by using a customized cell heater designed to heat a 1" x 1" x 13A" analyzer cell. The
heater consisted of two Macor plates (1 Vi" x 2" x 14") attached to the long sides of the

68
reaction cell. Macor sheets (6" x 6"x 14") were purchased from Astro Met Inc. Ni-Cu
wire (0.015" diameter, purchased from Omega Industries) was wrapped around the
external Macor plates and was resistively heated by using an Omega digital
temperature controller (maximum output current (~5 A) resulted in temperatures of
-520 K). Cell temperatures were measured using an Omega RTD thin film detector
fastened to the analyzer cell. Additionally, the entire high vacuum chamber was
heated by using the vacuum bake-out system in order to minimize radiative
temperature loss to the vacuum chamber walls.
Following the measurement of K^ at a set temperature, the cell heater and
bake-out were allowed to cool to a lower temperature and the entire system was
allowed to equilibrate at the new temperature for 30 min. Experimental
reproducibility was tested by following the temperature dependence of /if as the
reaction cell temperature was increased from 350 to 500 K.
The cell temperature was measured before and after each reaction and usually
fluctuated ±2 K during a single experiment. Typically, reactions were repeated three
times at a single temperature. Linear regression and statistical analyses of the all
measured equilibrium constants provided error limits at the 95% confidence level for
reported A//et° and A5e[° values.
Metallocenes and Reference Compounds
Metallocenes were purchased through Strem Chemicals except for ferrocene
and ruthenocene (Aldrich). No further purification was required except for Cp2Mn

69
which was resublimed prior to use. For temperature dependence studies, a sublimed
sample of ferrocene was used. Organic reference compounds were purchased from
Aldrich except N,N’-diethyl-p-toluidine (Alfa Chemicals). A sample of 1,1’-
bipyrrolidine40a was donated by Professor Stephen Nelsen from the University of
Wisconsin. Organic reference compounds were used without further purification.
Liquid samples were degassed through several freeze-pump-thaw cycles prior to use.
Far-infrared Spectroscopy.
Ferrocene salts in Table 2-3 were prepared according to literature procedures.95
The hexafluorophosphate and the tetrafluoroborate salts were prepared by dissolution
of pure ferrocene in concentrated sulfuric acid, followed by dilution of the dark blue
solution with water. The solution was filtered and then an aqueous solution of Bu4NX
was added, where X = PF6 or BF4. The precipitate was filtered and washed with
water until the washings were clear. The chloride salt was obtained by distilling
ferrocene in concentrated HC1 for several days. The blue precipitate was filtered and
washed with water. Infrared spectroscopy in the 400 - 3200 cm'1 region confirmed
the compounds to be ferrocenium salts.69'71 The samples were prepared as dilute 13
mm polyethylene pellets. Far-infrared spectra were recorded using a Bruker EFS 113V
spectrophotometer in the 50 - 750 cm'1 spectral region.

CHAPTER 3
SUBSTITUENT EFFECTS IN THE GAS-PHASE AND SOLUTION IONIZATION
AND ELECTRON ATTACHMENT ENERGIES OF ALKYLNICKELOCENES.
Introduction
Alkyl substituent effects have been studied extensively in organic chemistry
with emphasis towards understanding acidities,303 basicities,30b and reactivity of
organic systems. ’ The correlation between structure and reactivity of alkyl
substituents has led to detailed explanations and an interpretation of chemical
reactivity and chemical equilibria. Methods that correlate structure-reactivity
relationships allow for quantitative interpretation of various electronic perturbations
of substituents relative to a parent molecular-frame. Subsequently,5 chemical
transformations for yet unknown species can be determined based on known
substituent effects.
Gas-phase investigations have been very effective in evaluating intrinsic
substituent effects. 16d9,20,32,37,91 Specific electronic effects may be masked in
condensed-phase; therefore, intrinsic effects observed in solution studies may be
significantly modified or even reversed relative to gas-phase studies.5 Additionally,
the results may vary from solvent to solvent. Nonpolar solvents may yield
approximations of intrinsic substituent electronic effects, while strongly coordinating
70

71
r rq
solvent may inhibit substituent electronic effects altogether. To truly resolve
substituent effects in various solvents, gas-phase data must be used as a reference.
Relatively few gas-phase studies of alkyl substituents effects of metal
complexes have appeared,16,18’35’91 and little is known concerning the relative alkyl
effects in the thermochemistry of metal complex redox couples.32,34 Comparisons of
the gas-phase and solution results address solvent effects of metal redox couples, and
allow for estimates of differential solvation energetics of metal complexes to be
made.15,16 Additionally, such studies help extend parameterization schemes derived
from organic systems to the area of inorganic chemistry.
From photoelectron spectra35 and electrochemical studies,32,34 it is commonly
assumed that alkyl groups on coordinated ligands are electron-donating in metal
complex redox processes. Specifically, alkyl groups stabilize the oxidized form of a
complex relative to the reduced species. In order to further explore and understand
alkyl substituent effects for transition metal compounds, gas-phase electron-transfer
equilibrium studies have been performed, by using FTMS,44'47 to determined free
energies of electron-attachment (AGa°) and ionization (AG¡°) for a series of alkylated
nickelocene complexes.19 The results indicate, similar to organic systems,2,3 that alkyl
groups are not always electron-donating in organometallic redox processes.19 Models
that include polarizability effects, in addition to more traditional inductive effects,
must be used to interpret the data.5,52,53
Nickelocene is a useful parent compound for these studies because it forms
stable anions and cations in the gas-phase18,19 and solution.34,89
Furthermore,

72
oxidation and reduction of nickelocene involves the same elg* set of molecular
orbitals.96 A two-electron oxidation process for negative nickelocene ion is shown in
Figure 3-1 with accompanying molecular orbital diagrams of the 3d valence orbitals
for the Cp2Ni+/0/‘ complexes. A one-electron model suggests, because the same
molecular orbital of nickelocene is both oxidized and reduced, perturbations in the
orbital energies due to alkyl substitution on the Cp rings are expected to be similar for
the ionization and electron attachment processes. That is to say, the difference in the
ionization energies and electron affinities of an alkylated complex relative to
nickelocene should be comparable.36 However, this is not the case and an a more
flexible model must be used to explain the trends in the experimental data.
Solution phase redox studies were performed for comparison to the gas-phase
data. Differential solvation energies for some alkylated nickelocene complexes have
been derived from thermochemical cycles. Values for AAGsolv° for the Cp2Ni+/0 and
the Cp2Ni0/' couples are discussed.
Electron-Transfer Equilibrium Studies Involving Negative
and Positive Alkylnickelocene Ions
The electron-transfer equilibrium method has been discussed previously.16,18
Procedures for determining free energies of ionization and free energies of electron
attachments are similar. The equilibrium constants were determined for the reactions
shown in equations 3-1 and 3-2 where RCp and R’Cp represent alkylated
cyclopentadienyl ligands and X, for these examples, denotes a reference compounds

73
e2g
a ,
lg
^g
Figure 3-1 Molecular orbital diagrams for nickelocene anion, nickelocene, and
nickelocene cation.

74
with known a AG ° or AG° value. Ion intensities and partial pressures were measured
directly during the ETE experiment. Thus, the equilibrium constants and reaction free
energies for the reactions 3-1 and 3-2 can be determined (see equation 2-3).
(RCp)(R’Cp)Ni + X+ = X + (RCp)(R’Cp)Ni+
(3-1)
(RCp)(R’Cp)Ni + X' = X + (RCpXR’Cp)Ni'
(3-2)
Free energy ladders for gas-phase electron-transfer equilibria studied in this
work are shown in Figures 3-2 and 3-3. Derived AG° and AG° (electron attachment
free energy) values are referenced at ~350 K as measured by an RTD thin film
detector. The AG¡° values, for the process (RCp)(R’Cp)Ni —» (RCp)(R’Cp)Ni+ + e", in
Figure 3-2 are anchored to the AG ° values of nickelocene, manganocene, and
bis(benzene)chromium. The AG¡° values for Cp2Mn and Cp2Ni have been reported
previously.16 Electron-attachment free energies (for the process M + e' —» M') in
1 R
Figure 3-3 are anchored to the AG° value of azulene and nickelocene.
The AGj° value for Bz^r was based on the photoelectron spectrum which has
an extremely sharp first ionization band with a peak maximum at 5.47 eV.19 Because
of the sharpness of the band, the vertical ionization energy closely approximates the
adiabatic IP (see Figure 1-1B). The AGi°(Bz2Cr) = 125.7 kcal mol'1 was estimated by
assuming alP = vIP = 126.1 kcal mol'1. An estimate for AS¡° was determined by
assuming only the electronic entropy contribution was important. Thus,
A5vib° ~ A5rot° ~ 0 cal mol'1 K'1. The electronic entropy change associated

75
Figure 3-2
M - N
145.6
1.8
A 1.3A
Cp2Mn 142.5 —^
143.8
Ni
4.2
141.5 EthyINc
139.6 (MethylCp)2Ni
1.4
1.8
3.3
2.5
1 138.2 (EthylCp)2N¡
t t
136.4 (t-butylCp)2N¡
Cr 125.6-
CÉD
4.4
121.2 Cp*2Ni
Electron-transfer equilibrium ladder for ionizations for several
alkylnickelocene complexes for the process M —> + e'. Values of
AG ° (± 1.5 kcal mol’1) for the nickelocene complexes are to the right
of the ladder and AGet° values for individual ETE reaction are adjacent
to the arrows. The AGet° value for the (MeCp)2Ni/(t-BuCp)2Ni couple
is not within the expected ± 0.5 kcal mol'1 experimental error limit.

76
t-butyl
1.1
1 .
L i
-
i
1.3
0.0-
L
i
i
t
1
' '
0.5
-
'
i
0.7
1
L i
f
L
'
1.3
r
-
2.4
'
Ni
21-2^r^/^
(EthylCp)2N¡
20.3'
20.2 (EthylCp)CpNi
19.7 Cp2Ni
19.0 (MeCp)2N¡
Figure 3-3 Electron-transfer equilibrium ladder for electron-attachments for
alkylnickelocene complexes for the process M + e' —» M'. Values of
AGa° (± 1.5 kcal mol'1) for the nickelocene complexes are to the right
of the ladder and AGet° values for individual ETE reaction are adjacent
to the arrows. The AG° for Cp*2Ni is an estimated value.

77
with an 1A to 2A transition is ASelec° = R In 2 = 1.4 cal mol'1 K'1. Further details
concerning A5¡0 and AG¡° of Bz^Cr are discussed in Chapter 5.
Electron-transfer equilibrium reactions were repeated several times to insure
reproducibility. Cross checks were performed when possible to check the internal
consistency of the derived AG° and AGa° values. The experimental uncertainty in the
individual electron-transfer equilibrium reactions is ± 0.5 kcal mol'1. Electron
attachment and ionization free energies are reported with ± 1.5 kcal mol'1 error due
largely to errors in the AG° values of the reference compounds.
Nickelocene is expected to have the highest electron affinity of the first
transition row metallocenes since it has the lowest reduction potential.89 This is
further substantiated in the observation that no other metallocene forms negative ions
in the gas phase by low energy electron impact or chemical ionization.
Decamethylnickelocene could not be brought to electron-transfer equilibrium with any
reference compound. Further, Cp*2Ni' was not observed from electron impact or
chemical ionization. Attempts to ionize Cp*2Ni with electron-transfer reagents such as
azulene (AGa° = 17.3 kcal mol'1)27 and C6F6 (AGa° = 12 kcal mol'1)27 were
unsuccessful. The electron affinity of Cp*2Ni was estimated from the difference in
the AGa° values of (MeCp)2Ni and nickelocene by assuming that the electronic effect
of the methyl groups is additive. The AGa° value of Cp*2Ni puts it at the bottom of
the equilibrium ladder for compounds that have had AGa° values determined from the
electron-transfer equilibrium method.42

78
Results of the ETE studies for the cations and anions are presented in Table
3-1. The AG ° value for nickelocene is included as a reference for the alkylated
complexes. The AG° for nickelocene was determined from ETE studies with
azulene.18 Nickelocene was brought to equilibrium with azulene eight times from
both exoergic and endoergic directions and therefore serves as a second reference for
ETE studies of the negative ions. It is worth mentioning for historic reasons that the
author used the azulene/nickelocene negative ion couple as a training project to learn
how to operate a Nocolet FT/MS-1000. The not-so-serendipitous electron-transfer
equilibrium that ensued developed into a cavalcade of valuable experiments and the
present dissertation.
It is clear from the data in Table 3-1 that increasing the size and number of
alkyl substituents for nickelocene decreases the free energy of ionization. A similar
effect was observed for the alkylferrocene derivatives.16 In contrast to the ionization
energy data, trends in the AGa° values do not consistently reflect an increase in an
"electron-donating" effect. Moreover, the ethyl and t-butyl groups lead to an increase
in the electron affinity relative to H on the Cp rings. However, two methyl groups
lower the electron affinity relative to nickelocene. The shifts in the AG° values
relative to nickelocene are in general small, but appear to be significantly larger than
the error estimated for the electron-transfer free energies (~0.5 kcal mol'1) derived
from electron-transfer equilibrium experiments. The absolute free energies have larger
errors as mentioned earlier of ± 1.5 kcal mol'1.

79
Table 3-1 Free Energies of Ionization and Electron Attachment
No.
L
L’
AGioa’b
-A GaoSUb
1
Cp
Cp
143.8
19.7
2
MeCp
MeCp
139.6
19.0
3
EtCp
Cp
141.5
20.2
4
EtCp
EtCp
138.2
20.3
5
r-BuCp
r-BuCp
136.4
21.2
6
C5Me5
C5Me5
121.2
(~16)c
a. Units are kcal mol'1.
b. Estimated error in absolute values ±1.5 kcal mol'1.
c. Estimated value.

80
Note that by convention, electron affinities are expressed as positive values (the
affinity of an electron to be attracted to the nucleus of an atom) although they
represent an exoergic property. For application in thermochemical cycles, -AG° and
-AH° values are incorporated as negative values,15 however in the discussion the
electron affinities and free energies of electron attachment, the negative sign is
dropped.
An increase in AGa° for larger alkyl groups relative to R = Me is a well
documented effect for organic systems.98,99 For example, the electron affinities of
alkoxy radicals, RO, increase in the order R = Me < Et < n-Pr < t-Bu.98 However,
in solution the trends are reversed due to solvent effects.5 The solution acidity of
methanol is greater than that of ethanol, but in the gas phase, ethanol is more acidic
than methanol and even water.100 For p-benzoquinone (BQ) derivatives, the electron
affinity of 2,6-di-reri-butyl-BQ is ~1 kcal mol'1 greater than that of 2,6-dime thy 1-BQ.
The AG° values for the series of methylated benzoquinones compounds decrease
monotonically by ~2 kcal mol'1 per methyl group from the methyl to the tetramethyl
derivative.27 Although the alkyl effects for these systems are quite subtle relative to
alkyl effects for positive ions, the gas-phase results demonstrate that large alkyl groups
can stabilize anions in simple saturated and conjugated systems.100,101 However,
methyl groups tend to destabilize the electron affinities. The lower electron affinity
for (MeCp)2Ni relative to nickelocene is consistent with the usual destabilization of
anions by methyl substitution.42 Thus, from the above analysis, methyl groups can be

81
described as intrinsically electron-donating functions, stabilizing cationic systems and
destabilizing anionic complexes.
Alkyl Substituent Analyses for Positive and Negative Ions and
Rationalization of the Gas-Phase Trends for the Ionization
and Electron Attachments Free Energies
Numerous alkyl substituent parameter schemes have been developed to fit
chemical reactivity to an electrostatic models.4,5,52,53 The parameters are based on the
premise that any substituent R in place of a reference, hydrogen for example, may
alter the bonding, reactivity and overall chemical characteristics of the parent
molecule 5 Substituent schemes based on a single parameters or several parameters
have been used to correlate energy perturbations for chemical systems relative to a
parent reference.50'52 The single parameter model quantitatively predicted shifts in the
AGj° values of the alkylferrocene complexes (Chapter 2) relative to ferrocene with ± 3
kcal mol'1 accuracy.16
The Taft model used in the ferrocene analysis employed at parameters which
were used to assess inductive effects of alkyl parameters. Generally, the parameters
incorporate several electronic effects (i.e, field, polarizability, resonances) thus
separation of the specific electronic effects are not accounted for in this model. Field
effects refer to a charge/dipole, or dipole/dipole electronic interaction transmitted
through space or a through polarizable bond.5 Polarization effects pertain to a
charge/induced-dipole or dipole/induced-dipole interaction. Polarization effects are

82
more strongly distance dependent than field effects.5 However, the relative magnitude
of the two effects are also important in understanding the overall substituent effects.
A single5,53 and a two-parameter102 model based on previously derived
schemes were used to fit the AG ° and AG° data for the alkylnickelocene complexes.
Free energies of ionization and electron attachment in Table 3-1 were plotted against
Gj parameters. Figure 3-4 is a plot of the free energy data versus the alkyl Taft
parameters. Equation 3-3a was used to fit the AG ° data to the Taft o¡ parameters and
equation 3-3b was used to fit the AG° values to the G¡ parameters. The alkyl
substituents used in the correlations are shown in Table 3-2. Additivity of the
parameters is assumed for the fits. The values of pl is the sensitivity parameter for the
AGi°(RCp)(R’Cp)Ni = pjCEoj) + AGi°(Cp2Ni) (3-3a)
AGa°(RCp)(R’Cp)Ni = pjfZGj) + AGa°(Cp2Ni) (3-3b)
Taft analyses. All fits used in substituent parameter analyses correspond to shifts in
the free energy data relative to nickelocene, where R = H, for the processes shown
below. The free energies for equations 3-4 and 3-5 are the stabilization energy or
destabilization energy for the substituents. Thus AG34° is the stabilization for the
(RCp)(R’Cp)Ni + Cp2Ni+ = Cp2Ni + (RCp)(R’Cp)Ni+ (3-4)
(RCp)(R’Cp)Ni + Cp2Ni‘ = Cp2Ni + (RCpXR’Cp)Ni'
(3-5)

-Ionization Free Energy (kcal/mol)
83
Figure 3-4 Plots of AG¡° and AGa° data versus Taft E(aj) parameters, (a)
Ionization data (left scale, squares) are plotted as -AG ° values for L2Ni
—> L2Ni+ + e". (b) Electron attachment data (right scale, triangles) are
plotted for the process L^Ni + e' —» L^Ni'. The best fit line for the AGa
fit is drawn for all data except (MeCp)2Ni.
Electron Attachment Free Energy (kcal/mol)

84
Table 3-2 Alkyl Substituent Parameters for Some Alkylnickelocene Complexes
and Free Energies for Reactions 3-4, 3-5 and 3-6.
No.a
lOj
AG34ob’c
A G35“b-C
A036°bX
^/b.d
XPb-d
1
0
0
0
0
0
0
0
2
-0.092
-0.70
-4.2
0.7
-4.9
2.5
1.7
3
-0.055
-0.49
-2.3
-0.5
-1.8
1.4
0.9
4
-0.110
-0.98
-5.6
-0.6
-5.0
2.5
3.1
5
-0.148
-1.5
-7.4
-1.5
-5.9
3.0
4.4
6
-0.46
-3.5
-22.5
(~3.5)
(-26)
a. Compound numbers taken from Table 3-1
b. Units are kcal mol'1 .
c. Estimated error for free energy is ± 0.4 kcal mol'1.
d. The derived / and P values in kcal mol'1 for the individual substituents are
H, / = P = 0 (defined); Me, / = 1.2, P = 0.9; Et, / = 1.3, P = 1.3;
t-Bu, / = 1.5, P = 2.2

85
cations and AG35° is the stabilization energy for anions relative to nickelocene, except
Me which is destabilizing. The direction of the slopes are opposite because, in
general, alkylation of nickelocene stabilizes formation of the cations and anions. For
the positive ions, electron-attachment becomes less exoergic as alkylation increases
since the cations are stabilized relative to nickelocene. Clearly, the effects of the alkyl
groups are different on the electron attachment energies compared to the ionization
energies. The ionization data fit with equation 3-3a yields a good fit, with a
correlation coefficient of r = 0.997. The parameter pj (= 49.9 ± kcal mol'1) is the
slope of the line and reflects the sensitivity of nickelocene ionization potentials to
alkylation. This value is comparable to for the alkylferrocenes of 57 ± kcal
mol'1.16 Conversely, the same parameters provide an unacceptable fit of the electron
attachment data (r = 0.51). The lack of correlation of the methyl derivative is
primarily responsible for the poor fit of the AG° data to equation 3-3b. The pl value
for the negative ion data is 6 ± 6 kcal mol'1.
A single parameter model was developed by Hehre et al. based on
polarizability effects, ca parameters, of R. The oa parameters have been used to
successfully fit the proton-transfer free energies of various cationic and neutral acids.
Fits to the alkylnickelocene AG° and AGa° data (Figure 3-5) result in good correlation
for the ionization free energies (r = 0.994) but only poor correlation for the electron
attachment free energies (r = 0.63). The oa parameters are included in Table 3-2.
The slopes for the plots are 5.0 ± 0.4 kcal mol'1 for the ionization and 0.8 ± 0.6 kcal
mol"1 for the reductions. Based on the observed lack of correlation, a single

Ionization Free Energy (kcal/mol)
86
Figure 3-5 Plot of AG° and AG° data for some alkylnickelocene complexes versus
Z(aa) parameters, (a) Ionization data (left scale, squares) are plotted as
-AG¡° values for L^Ni —» L2Ni+ + e". (b) Electron attachment data (right
scale, triangles) are plotted for the process L^Ni + e' —» L2Ni'. Plot of
AGa versus Ga is drawn for all points except (MeCp)2Ni.
Electron Attachment Free Energy (kcal/mol)

87
paramter model is not adequate to correlate the alkyl effects both of the AG ° and
AG° values. The failure for the electron attachment data to fit with either of the
a
single parameter models can be largely attributed to the experimentally observed
effects of methyl in contrast to t-butyl and ethyl. Both Oj and ca values predict the
same direction of effect for all three groups.50'53
The trends in the free energy data require a more complex model for
interpreting alkyl substituent effects for the nickelocenes than one based only on a
single parameter. Apparently, more than one electronic effect is responsible for the
disparity in the correlations for the AG ° and AG° data with the Oj and oa values. A
simple alternative model combines the effects of the polarizability of the alkyl groups
with the inductive effects. A quantitative model has been developed by Taft and
workers to separate inductive and polarizability effects for gas-phase proton transfer
equilibria involving a number of alcohols (ROH) and their corresponding bases (RO')
and acids (ROH2+).102 However, the successful application of oa parameters to the
same data lead Hehre et al. to question the necessity of a two-parameter model.52 The
application of the two-parameter model102 was used for a comparison to the other fits
and to uncover the origin of the disparate effects for methyl relative to other alkyl
groups.
The polarizability and the inductive effects relative to R = H are separated by
examining the hypothetical gas-phase equilibrium reaction described by equation 3-6
where Cp2Ni (R = H) is used as the reference compound. Furthermore, the free

88
energy for the hypothetical equilibria is given by the difference in the free energies of
equations 3-4 and 3-5. The polarization effect (P) for R relative the H arises from
(RCp)(R’Cp)Ni‘ + Cp2Ni+ ^ Cp2Ni' + (RCp)(R’Cp)Ni+ (3-6)
AG
36
ag34° - ag35°
(3-7)
greater charge-induced dipole stabilization of both the cation and the anion.
Polarization effects are not destabilizing since the induced-dipole interaction is brought
about by a charge localization on a species remote to the substituent.5 It is assumed
that the polarizable alkyl groups will stabilize the nickelocene cations and anions
equivalently. Since the same molecular orbital is both electron donor and acceptor
orbital, the assumption that polarization stabilization for the positive and negative
1 A 9
metal complex ions is the same is easily rationalized. An inductive effect (/) is an
electron-releasing effect and will therefore destabilize the negative ions relative to a
parent (i.e. (MeCp)2Ni* relative to Cp2Ni') but stabilize cation complexes (i.e.
(MeCp)2Ni+ relative to Cp2Ni+).102 It is also assumed that / effects, similar to P, will
be approximately equivalent for both the positive and negative ions.102 Consequently,
for the equations 3-4 and 3-5, P is considered positive, while I has opposite effects in
both processes. The free energies for equations 3-4 and 3-5 can be written as -AG34°
- / + P and -AG35° - -/ + P. By subtracting these two equations, AG36° = -21 for
equation 3-6 is derived. The derived P and 7 parameters for the alkylnickelocene

89
compounds are shown in Table 3-2. As expected, correlation of -AG36° values (21)
with Oj parameters (Figure 3-6) yields a good fit (r = 0.995).
The resulting analysis leads to the conclusion that the inductive and
polarizability contributions to shifts in the AG° and AG° values for alkylnickelocene
are similar in magnitude. The relatively small changes in the AG° values can be
attributed as the difference in the polarizability and inductive effects {P - /), both of
which are of comparable magnitude. The polarizability and inductive effects for the
positive ions are additive (P + Í) resulting in a larger overall sensitivity to alkyl
substitution.
The separation of polarizability and inductive effects for the acid and base
equilibria of the gas-phase alkyl alcohols indicated that P effects are significantly more
important than / effects in evaluating the relative gas-phase acidities and basicities of
the ROH compounds. For example, the PH ratio derived for ethyl and t-butyl ROH
derivatives are 5.4 and 3.7 respectively.102 The P!I ratio for the alkylnickelocene
complexes for this study are Et =0.9 and t-Bu = 1.5. The / parameters for the two
gas-phase studies are equivalent, however the P values derived for the gas-phase
alcohols are several times larger than P values derived for the nickelocenes.
According to Hehre et al., the separation of effects may not be valid for the ROH
compounds since the oa parameters provide and acceptable fit to the data.53
However, it may be necessary to use a two parameter model when the magnitudes of
the polarization and inductive effects are comparable.

90
Figure 3-6
Plot of AG36° values (kcal mol'1) derived from equations 3-4, 3-5, and
3-6 versus X(Gj) values. Compound numbers are given in Table 3-1.

91
The two parameter scheme suggests that the alkyl group polarizability is much
smaller for the nickelocenes than the effects of alkyl groups for the ROH
compounds.102 This can be easily rationalized for the nickelocenes because of the
increased distance of the charged center to the R group (~3.3 Á)72 in comparison to
RO\ for example. The polarizability has a 1/r4 effect in this case, which can be
compared to a 1/r2 dependence for inductive effects.5 The rapid decrease of P and I
effects with increasing distance is consistent with the present results.
Solvation Energetics of Nickelocene Cations and Anions
As demonstrated in Chapter 2, values for oxidation and reduction of a
compound can be related to the gas-phase free energy of ionization or electron
attachment through thermochemical cycles. The gas-phase and solution values are
related through the differential solvation energetics of the ion and the neutral.16'20
The AAGsolv° for the Cp2Ni+/0 has been discussed in the previous chapter. The
AAGsolv° value for the Cp2Ni0/" couple has been determined and relationship of the
two values is considered.
Nickelocene has an electrochemical reversible one-electron oxidation potential
in various solvents. The reduction of nickelocene is only quasi-reversible in DMF
of THF at low temperatures.34,89 The earlier electrochemical studies by Holloway and
Geiger of nickelocene have been repeated.89 The cathodic peak potential for the
reduction of Cp2Ni was found to be relatively insensitive to temperature over a wide
temperature range (-210 to 296 K). Similarly, the E^ value for (t-butylCp)2Ni0/" was

92
also found to be independent of temperature over the same temperature range.
Electrochemical potentials were measured by using ferrocene as an internal
standard.103 The A£,¿ value relative to £^(Fc+/0) for the oxidation of nickelocene is
0.42 ±0.1 V compared to a value of 0.40 V measured by Holloway.89 An estimate of
the Afor quasi-reversible reduction potential for nickelocene determined from
£i¿(Fc+/0) - Ep. is 2.4 V ± 0.2 V. Decomposition of the Cp2Ni' species at
temperatures approaching 273 K was a complicating factor that was circumvented at
lower temperatures; therefore, the low temperature reduction potential (-210 K) for
nickelocene were used to estimate the Evalue.
F'*'
From the gas-phase and solution phase data, the differential solvation energies
for the 0/- couple was derived by using equation 3-8. The AAGsolv° value for the
Cp2Ni°/' couple is -40 ± 5 kcal mol'1. The AAGsolv° for the Cp2Ni+/0 couple from
Chapter 2 is -38 ± 5 kcal mol'1.16
AACsolv° = ^[£pc + %HE + ¿ref] ' AGf(g) (3-8)
Both values are exoergic since the solvation energy of the ion will be greater than the
solvation energy of the neutral. The estimated A£^ from E^ ox - ~2-0 * 0-2
V in THF. From the gas-phase and solution data for the Cp2Ni+/,° and Cp2Ni0//*
couples, it is possible to extract the average differential solvation energy,104
AAGsoiv av° by using equation 3-9. Note that the average differential solvation free
energy determined

93
AAGsolv>av° = -V4[-F(V> - £red°) + (AGj° + AGa°)] (3-9)
by equation 3-9 does not require knowledge of the absolute potential of the reference
electrode.16 The absolute potential of the normal hydrogen electrode has been a
source of confusion because some of the thermodynamic quantities associated with the
process lAH2(g) —> H+(aq) + e'(aq) are not accurately known.54 Therefore, the
application of equation 3-9 allows for estimates of AAGsolv° values without detailed
analyses of solvent effects or reference electrode potentials. The AAGsolv av° value
of -39 ± 3 kcal mol’1 obtained from the electrochemical potentials measured in this
work is in agreement with the AAGsolv° (Cp2Ni+/0) estimated in Chapter 2 based on an
estimate of the absolute potential for the Cp2Ni+/0 couple in acetonitrile.16 Thus, it
appears that the AAGsolv° values for the cationic and anionic couples of nickelocene
are the equivalent which is consistent with the Bom model used in Chapter 2.
A similar analysis for (t-butylCp)2Ni was not as easily accomplished.
Although the oxidation of (t-butylCp)2Ni is reversible in THF at ambient temperatures,
the reduction wave was irreversible even at -200 K. Moreover, because of the
observed irreversible reduction wave, the trends in the redox responses relative to
nickelocene were inconsistent with the gas-phase data. The oxidation (A£,¿ = 0.52 ±
0.1 versus £^[Fc])was cathodically shifted by 100 mV, however, the reduction
potential was slightly anodically shifted by -50 mV (A= 2.35 ± 0.2 versus
£i¿[Fc])both relative nickelocene. Increased alkylation of nickelocene is expected to
shift the reduction potentials cathodically due primarily to solvent effects associated

94
with the t-butyl groups. The differential solvation energy for the (t-butylCp)2Ni+/0
couple is -33 ± 5 kcal mol'1. The estimated Afor (t-butylCp)2Ni is 1.8 ± 0.1 V
resulting in a AAGsolvav° of -37 ± 3 kcal mol’1, which is in agreement with the
average differential solvation energy for nickelocene.
Bond Disruption Enthalpies for Nickelocene Anion
Estimates of the average heterolytic and homolytic Ni-Cp bond disruption
enthalpies for Cp2Ni' have been determined from thermochemical cycles as previously
described.16,18,84 Values for the heats of formation of Ni(g) and Ni+(g) and Ni'(g)
were determined from known thermodynamic quantities. ’ The reported
experimental errors include uncertainties in the AG° value for nickelocene in addition
to errors in the literature values for the A//f° quantities. The average bond disruption
enthalpies are A//hom°(Cp2Ni') = 64 ± 3 kcal mol'1 and A//het°(Cp2Ni') = 126 ± 5
kcal mol'1.
Conclusions
The alkylnickelocene studies have allowed the first survey of alkyl substituent
effects for electron attachment and ionization energetics in which the same molecular
orbital can be considered both the acceptor and donor orbital. The data for the
alkylnickelocene complexes demonstrate that gas-phase electron affinities of metal
complexes can be increased by increased alkylation. However, due to larger inductive
effects versus polarization effects, methyl groups destabilize formation of nickelocene

95
negative ions. The destabilization of methylated nickelocene complexes was further
demonstrated by lack of an observed ion signal for Cp*2Ni" and the absence of an
electrochemical response for reduction of Cp*2Ni.39,105 For comparison, the gas-phase
AGa° values for some ruthenium [5-diketonate complexes42 are consistent with the
present conclusions that increased alkylation yield an increase in electron affinity.
Thus, alkyl groups are not intrinsically electron-donating with respect to gas-phase
redox thermochemistry of metallocenes16,19 or coordination complex ions.84
Specifically, electron attachment and ionization free energies are not always
lowered by alkyl groups.
An electrostatic model was successfully used to rationalize the trends in the
AGj0 and AGa° values for the nickelocene complexes. Single parameter models can be
used to interpret free energy of ionization data for the alkyl metallocene
compounds, ’ ’ however, electron attachment data fail to correlate using either (Jj or
Ga parameters. The separation of inductive and polarizability effects102 revealed that
the two electronic effects are comparable in magnitude for the nickelocene complexes.
Furthermore, I and P effects are additive when considering shifts in the AG ° values
relative to nickelocene but the effects approximately cancel for the negative ions.
Overall, the shifts in the AG¡° values relative to nickelocene are greater than
comparable shifts in the AGa° values.
Analysis of the differential solvation free energies for one-electron oxidation
and reduction of nickelocene shows that both couples have similar differential
solvation energies. The AAGsolv av° values estimated for (t-butylCp)2Ni is -37 ± 5

96
kcal mol"1, and is consistent with the similar value for nickelocene (-39 ± 5 kcal
mol"1). The difference in solvation relative to nickelocene can be rationalized due to
the lipophilicity of the t-butyl groups. The similar AAGsolv° values for the nickelocene
+/0 and 0/- redox couples are consistent with the Bom model. According to the Bom
model, AAGsolv° values are dependent on ionic charge {q1) and ionic radius (r) in
addition to the dielectric constant of the solvent.54 Since both the size and charge
parameters are equal for the Cp2Ni+/0 and Cp2Ni0/~ couples, the differential solvation
energies, according to the Bom model should be equal in the same solvent.
Experimental Procedures
Positive and negative ion electron-transfer equilibrium studies were performed
on a Fourier transform ion cyclotron resonance mass spectrometer as previously
described.14'20 The methods for positive and negative ion studies are similar with a
few important exceptions. Most obvious is that positive trapping potentials were used
for positive ion studies and negative trapping potentials were employed for the study
of negative ions. Electron impact energies ranged from 8-12 volts for the positive
ions studies however negative ions were generated from low energy electron impact.
Cyclic voltammetry studies were performed with a PAR system (Models
173,. 75). A platinum button working electrode and a KC1 saturated Ag/AgCl
reference electrode were used. Electrochemical experiments were performed in THF
containing 0.1 M Bu4NPF6 and solutions were freshly prepared prior to each
experiment. Tetrahydrofuran was distilled from Na/benzophenone. The supporting

97
electrolyte Bu4NPF6 was recrystallized several times from an acetone/ethanol mixture
and dried in vacuo overnight. Electrolytic solutions containing the nickelocenes were
either prepared in an inert atmosphere glove box or by injecting millimolar solutions
of a nickelocene/THF solution directly into the cell through a septum. Sublimed
samples of ferrocene and nickelocene were used.
Low temperature electrochemical studies were performed by submerging the
entire electrochemical cell into solvent/C02 baths: acetone (-72 °C), isopropanol
(-55 °C), CC14 (-23 °C), and ice water. By varying the solvent/C02 ratio, intermediate
temperatures were obtained. The electrochemical cell was allowed to equilibrate at a
temperature for several minutes before scans were recorded. Typically, temperature
variation of ± 5 K during each individual scan were measured by using a low
temperature alcohol thermometer.
Azulene was purchased from Aldrich chemicals and used without further
purification. Nickelocene and Bz^Cr were purchased from Strem Chemicals and were
purified by vacuum sublimation. The alkylated nickelocene complexes were prepared
according to literature methods106 and purified by preparative HPLC.

CHAPTER 4
GAS-PHASE AND SOLUTION OXIDATION POTENTIALS OF RUTHENOCENE
DERIVATIVES
Introduction
Variation of the ligand environment around a metal center is an effective
means of evaluating the electronic effects of ligands with respect to a parent
compound.107 Modifications of Cp ligands of a metal complex can lead to
modifications in reactivity. For example, substitution of Cp* for Cp is well-known for
increasing the steric congestion around a metal center and therefore influence the
reactivity of the complex; however, Cp* substitution for Cp will also alter the
electronic character metal complex.105,108 Since steric effects have little influence on
the oxidation-reduction potential of molecular species when variations are small,
determined AG¡° or AGa° changes can be attributed to the effect of the supporting
ligands. Electrochemistry and photoelectron spectroscopy have been used to
investigate the electronic effects of varying Cp ligation for metallocenes and
metallocene derivatives. However, as mentioned earlier, the effectiveness of these two
techniques for evaluation of metal complex oxidation-reduction energies can be limited
by several influencing factors. Electrochemical potentials for metal complexes may be
irreversible therefore the measured E^ values may not be accurate assessments of the
thermodynamic potential. Vertical ionization energies measured by photoelectron
98

99
spectroscopy are only accurate assessments of the adiabatic ionization if the
97
equilibrium geometry of the ion and the neutral are small.
In Chapter 2, the thermal free energy of ionization for ruthenocene was
reported (AG ° = 164.6 ± 1.5 kcal mol'1) at 350 K.16 This value of AG¡° differs from
a previously reported value due to corrections made for A5j° values for the reference
compounds. The electrochemical oxidation of ruthenocene is irreversible under most
experimental conditions.91 For example, an irreversible 2e' process is observed for
Cp2Ru in CH3CN containing 0.1 M TEAP and in CH2C12 containing 0.1 M
Bu4NPF691 Recently, Mann and coworkers have reported a reversible oxidation
potential for ruthenocene in CH2C12 containing tetrabutylammonium tetrakis(3,5-
bis(trifluoromethyl)phenyl)borate, Bu4N+TFPB‘. The ruthenocene cation is extremely
susceptible to secondary reaction in common solvent/electrolyte systems thus the
extremely weakly noncoordinating anion TFPB' reduces the susceptible to secondary
nucleophilic reactions of Cp2Ru+.90 The majority of ruthenocene derivatives
demonstrate irreversible electrochemical oxidation potentials in common
electrochemical solvent systems.91 A reversible oxidation for decamethylruthenocene
has been reported in CH2C12 containing 0.1 M Bu4NPF6 by several workers.90,91,92
The lack of reversible electrochemical oxidations for most ruthenocene
derivatives dictates the need for accurate thermochemical ionization potentials.
Electron-transfer equilibrium techniques15'20 have been used to determine the free
energies of ionization for a broad series of ruthenocene derivatives by using Fourier
transform mass spectrometry.44'47 Electron-transfer equilibrium methods have been

100
described in the previous chapters. Free energies of ionization for the ruthenocene
derivatives are represented by equation 4-1 and are reported at 350 K.
Cp2Ru —> Cp2Ru+ + e‘ (4-1)
Ruthenocene derivatives were of the type LL’Ru where L and L’ are
cyclopentadienyl ligands. Ligand variations for the ruthenocene derivatives include
perhalogenation of a Cp ring, attachment of electronic withdrawing groups (i.e. N02
and CF3) and fused-ring systems.
Cyclic voltammetry was used to reevaluate the oxidation potentials for many of
the ruthenocene derivatives.91 Oxidation values reported to be irreversible for
several of the ruthenocenes were observed to quasi-reversible under our experimental
conditions. Additionally, E^ values for several ruthenocene complexes are reported
for the first time. Several solvent/electrolyte systems were used for selected
complexes in order to carefully characterize the electrochemical oxidation potentials.
Comparison of the free energies of ionization and electrochemical oxidations
potentials leads to insight concerning solvent effects of the ruthenocene oxidation
potentials. Thermochemical cycles which incorporate AG¡° values in both the gas
phase and solution have been used to determine differential solvation free
energies,16,20,54,84 AAGsolv°, for the ruthenocene complexes. Values for AAGsolv° are
compared to an electrostatic model for predicting solvation energetics.20’54

101
Direct comparison of AG¡° values for the series of ruthenocene LL’Ru
complexes with AG¡0(Cp2Ru) has lead to information concerning the electronic effects
of the various Cp derivatives. Ligand effects have been quantified through the
development of a new set of Cp ligand parameters which use Cp and Cp* as
references.20 Implications for use of such a parameter scale for predicting metal
complex reactivity are considered.
Results of the Electron-Transfer Equilibrium Reactions
Details concerning electron-transfer equilibrium techniques have been given in
Chapter 2. Reactions were followed for 5-10 s to confirm ETE and were monitored
from both endoergic and exoergic directions to insure that equilibrium was
independent to the approach to ETE. The general electron-transfer equilibrium
reaction is given in 4-2 was studied, where LL’Ru denotes a ruthenocene complex
LL’Ru + R+ = R + LL’Ru+ (4-2)
and R denotes a reference compound with a known AG¡° value. Figure 4-1 is an
electron-transfer equilibrium ladder displaying all reactions involving the ruthenocene
complexes. Similar to other ladders in this work, values adjacent to arrows represent
AGet,35o° values f°r individual ETE reactions. All AG° values he adjacent to the
compound in Figure 4-1 and are also summarized in Table 4-1. Compound

102
abbreviations used throughout this chapter are also given in Table 4-1 and Cp ligand
abbreviations are given in Table 4-2.
The ruthenocene derivatives exhibit a wide range of AG¡° values spanning over
2 eV. Various organic compounds and metallocenes were used as reference
compounds for ETE studies.16,27,39,40,59 Details concerning the AGj° values of the
reference compounds used for ruthenocene ETE studies are given in Chapter 2.
Aniline and benzene derivatives used as reference compounds have ionization
potentials that are anchored to the AGi350° values of N,N-dimethylaniline (DMA)40,59
and benzene27,39 which are 163.4 ± 0.5 and 212.4 ± 0.5 kcal mol'1 respectively. The
free energies of ionization of the metallocene reference compounds were determined
from ETE studies and are anchored to DMA. The AG¡° for ruthenocene (1) has been
reponed previously16,20 and is used here for comparison to other ruthenocene AG¡°
values.
In several cases, compounds were brought to equilibrium with more than one
reference compound or another metallocene as a check on the internal consistency of
the ladder. In general, the consistency throughout the ladder is good. For example, 1
and 5 were both brought to equilibrium with DMA and each other; it can be observed
from the ladder that the expected difference in the AG° values of the ruthenocene
complexes is obtained from the ETE study. In another example,
bis(indenyl)ruthenocene 8 was brought to ETE with ferrocene and ethylferrocene. The
difference in the AG¡° values for the ferrocene compounds was previously
determined16 to be 2.9 ± 0.5 kcal mol'1 and the observed difference from the two ETE

103
CF-
Cp'
Ru H
CF,
CF
p-Xylene 152.8-
m-Cresol 190.4
m-Toluidine
\
172.2
192.
2.6'
1°5
CF
CF,
Cp'
I
Ru CF,
CF
CF,
OSi(Et)3
171.7-
14 Cp*(C5Br5)Ru
167.8-
\
2.7
o-
/ V N(CH3)2 2^
165.1:
163.4
. \ 170.8 Cp’(C5F5)Ru 13
-165.4 Cp’(C5Cl5)Ru 5
8
164.6.
1.2 Cp2Ru 1
V>N(Et)2-
ch3/ N(CH3)2
3.3
160.1
1.8
161.9
L
.156.3
CH3Vj^N(Et)2"
9 Cp'(TMSCp)Ru —151.3:
0.7
3-’
3.2
j 3.9
4
i
1.6-
4 2.4
0.8
i
-i-
2.9
1.0
4
4.7
4
0.8
'
t
158.4 (TMSCp)2Ru 7
Cp2Fe
153.1 ^CpCp'Ru 2
1 2 $
151.0 ^ Cp(ethylCp)Fe
150.2^
149.4
148.2—Cp(n-butylCp)Fe x '\^J
Ru 10
4.2
t
i.'M
(X6
1
â–  4
4.6
r
147.0—(MeCp)2Fe
Cp2Ni
, 143.8/
143.1 Cp'(n5-Fluorenyl)Ru 11
U.O * a
\
• Cp2Mn
137.9 Cp'2Ru 12
Figure 4-1 Electron-transfer equilibrium ladder for several ruthenocene derivatives
for the process M —» M+ + e\ Values adjacent to arrows denote AGet°
values at T - 350 K for individual ETE couples. Ruthenocene AG¡°
values (kcal mol'1) are anchored to AG¡°(DMA) and AG¡° of benzene.
The AG¡° value for 3 is a bracketed value.

104
Table 4-1 Values AG ° for Ruthenocene Derivatives and Other Data
LL’Rua
No.
AGi°(exp.)b,c
XPSd
AGj°(eq 4-9)bl'
(Cp)(Cp)Ru
1
164.6
280.7
164.6
(Cp)(Cp*)Ru
2
152.3
280.2
151.6
(TTFMH)(Cp*)Ru
3
192
192
(TTFMOSi)(Cp*)Ru
4
171.7
171.7
(C5Cl5)(Cp*)Ru
5
165.4
280.8
165.4
(N02Cp)(Cp*)Ru
6
161.9
161.9
(TMSCp)(TMSCp)Ru 7
158.4
160.2
(Ind)(Ind)Ru
8
151.0
153.9
(TMSCp)(Cp*)Ru
9
151.3
149.4
(Ind)(Cp*)Ru
10
149.4
280.1
146.4
(Flu)(Cp*)Ru
11
143.1
279.6
143.1
(Cp*)(Cp*)Ru
12
137.9
279.9
138.6
(C5F5)(Cp)Ru
13
170.8
170.8
(C5Br5)(Cp)Ru
14
165.1
165.1
a. Ligand abbreviations given in Table 4-2.
b. Units are kcal mol'1.
c. Error limits are ± 2 kcal mol'1 except 3 (± 5 kcal mol'1).
d. Core binding energies measured by XPS are taken from reference 91.
e. Only equation 4-9 data shown. Values of AG¡° predicted from equation 4-8 will be
exactly equal to the experimental AG¡° values.

105
Table 4-2 Ligand y* and y Parameters
Ligand Abbreviation
H
S‘(MC)3
CH, CH3
TTFMH
TTFMOSi
C5F5
C5CI5
C5Br5
N02Cp
Cp
TMSCp
Ind
Flu
Cp*
Y*(L)a
2.76
1.31
1.28
0.91
0.89
0.67
0
-0.07
-0.20
-0.64
-1
7(L)a
3.10
1.55
1.47
1.06
1.03
0.79
0
-0.17
-0.41
-0.65
-1
a. Ligand parameters for Cp and Cp* are 0 and -1 respectively by definition.

106
studies involving 8 is 3.4 kcal mol'1, within the ± 0.5 kcal mol'1 experimental error.
It was not always possible to do cross checks since many of the organometallics had
to be introduced into the mass spectrometer by use of a heated solids probe.
Therefore, only reference compounds that were sufficiently volatile to be sublimed via
a heated leak valve were used for checks of internal consistency.
For the ruthenocene derivatives, error limits of ± 2 kcal mol'1 in the AG°
values are reported. Errors in the free energies of ionization originate from two
sources in the experiment. Error limits in the reference compounds, which are
typically ± 1 kcal mol'1 for the organic reference compounds,27,39,40,59 although some
compounds have errors as high as ± 2 kcal mol'1, and ± 1.5 kcal mol'1 for the
metallocene reference compounds.16 Since the majority of AG° values used in the
equilibrium studies are referenced to other compounds, errors are propagated
throughout the system as the accuracy of the reference compounds decreases. Errors
in the measured K values for individual ETE experiments are the second source of
experimental uncertainty. Neutral partial pressure measurements of the reactant gases
and ion detection measurements introduce errors in the quantities of approximately
± 0.5 kcal mol'1 (this represents an error of 50% in the measured equilibrium
constant). The partial pressure measurements account for the largest experimental
uncertainty in evaluation of Meq values. However, an error as large as 30% in the
partial pressure of the neutral gases is only a ± 0.2 kcal mol'1 error in Kpn values.
Errors in the measured ion intensities are not expected to exceed 10%. Considering all

107
possible sources of error in the evaluation of AG ° values for the ruthenocene
derivatives, an overall ± 2 kcal mol uncertainty is reported
All free energy values in Figure 4-1 were obtained from direct ETE reactions
involving at least one reference compounds. However, definitive equilibrium for
Cp*(TTFMH)Ru (3) could not be established with either m-cresol or p-xylene.
Therefore, a bracketed value for AG ° of 3 is reponed with ± 5 kcal mol'1 error. The
bracketed value for 3 was obtained from sequential exothermic electron-transfer
reactions as demonstrated by equations 4-3 and 4-4. Both reaction were observed
p-xylene+ + 3 —» 3+ + p-xylene (4-3)
3+ + m-cresol —> m-cresol+ + 3 (4-4)
to completion in that no significant back electron transfer in the endothermic direction
was conclusively observed. From the electron-transfer bracketing experiments the
AG;0 value of 3 is estimated from the average AG¡° values of the two organic
compounds. The reaction rate of 3 with m-cresol was slow (/q- for reaction 4-4 was
4.7 ± 1.4 x 10'11 cm3 molec'V1) and this may explain why electron-transfer
equilibrium was not observed.61’80 For comparison, compound 4 was brought to
equilibrium with m-toluidine and the rate constant reaction 4-5 of is 4.8 x ± 1.4 10’10
'l i i
cm molec s .
4 + m-toluidine+ —» m-toluidine + 4+
(4-5)

108
Evaluation of the Gas-Phase Free Energies of Ionization for a Series
of Ruthenocene Derivatives
As all of the neutral ruthenocene complexes are d6 metal compounds with *Alg
ground states, variation in the ionization potentials of the derivatives with respect to
1 is due solely to shifts in the molecular orbital energies resulting from electronic
effects of the ancillary Cp ligands.20 Therefore, comparison of the AG° values for the
series of compounds need only include electronic effects of the substituents or the
ligands since all the ruthenocene complexes are electronically degenerate.35 As the
substituents become more electron withdrawing, the energy of the valence orbitals will
c rn
decrease, i.e. harder to remove an electron from a molecular orbital. Conversely,
as substituents become more electron donating, the energy of the valence molecular
C C'J
orbitals are raised, thus the complexes will be more easily oxidized.
Overall, the experimentally derived free energies of ionization for the
ruthenocene complexes follow the expected trends in that compounds bearing electron-
withdrawing substituents such as -Cl, -F, or -CF3 have AG¡° values greater than
ruthenocene and compounds with electron-donating substituents have lower AG¡°
values.5,20,50'53 The observation for electron-withdrawing groups to increase the
ionization potential with respect to the parent ruthenocene has been observed by Burk
and coworkers109 and by Gassman and Winter.91 Burk measured the oxidation
potentials of 3 and 4 by using cyclic voltammetry and observed the oxidation
potentials to be anodically shifted relative to ruthenocene;109 however, the compounds
demonstrated irreversible electrochemistry. Gassman and Winter measured the Ru 3d

1Ü9
core binding energies by using X-ray photoelectron spectroscopy (XPS) of a series of
ruthenocenes and noted the Cp*(C5Cl5)Ru 3d core binding energy is 0.1 eV greater
than that for ruthenocene.91 Alternatively, Gassman also observed that ruthenocenes
with electron-donating functions on the Cp rings had lower Ru 3d core binding
energies than ruthenocene. For example, the binding energy of ruthenocene is 280.7
eV while the binding energy of decamethvlruthenocene is 279.9 eV. The core binding
energies are reported to be accurate within 0.01 eV.91
Ionization Free Energies of Ruthenocenes with Electron-Withdrawing Ligands
The compounds studied that bear electron withdrawing groups are
Cp*(TTFMH)Ru (3), Cp*(TTFMOSi)Ru (4), and Cp*(N02Cp)Ru (6). The
halogenated complexes Cp*(C5X5)Ru, where X = F, Cl, Br will be discussed in a
separate section. The fluorinated complexes 3 and 4 have AG ° values greater than
ruthenocene.20 Although the AG¡° value for 6 is lower than AG¡°(Cp2Ru),
Cp*(N02Cp)Ru possesses an electron-withdrawing N02Cp ligand, which is countered
by the polarizable pentamethylcyclopentadienyl ligand.
Complexes 3 and 4 both possess extremely electron-withdrawing
trifluoromethyl groups.109,110,111 A survey of reported ionization potentials for several
benzene derivatives demonstrates that the CF3 group is generally more electron-
withdrawing than -CN or fluorine. The ionization potential of CF^CgHg = 223.3 kcal
mol'1 which is 1.5 kcal mol'1 greater than cyanobenzene and 11 kcal mol'1 greater
27
than fluorobenzene.

110
Burk and workers first prepared the TTFMH ligand in the interest of
1 HQ 119
developing highly electrophilic oxidative-resistent metal complexes.
Additionally, the solubility and reactivity of fluorinated complexes is quite different
with respect to the nonfluorinated analogues. Gassman and Winter prepared the
first metallocene bearing the trifluoromethyl function as (CF3Cp)2Fe and demonstrated
the powerful electron-withdrawing nature of the complex by using X-ray PES.110 The
Fe 2p core binding energy of (CF3Cp)2Fe is 708.6 eV compared to the binding energy
of Cp2Fe = 708.0 eV.111 Recently, Gassman and workers prepared the novel ligand
tetramethyl(trifluoromethyl)Cp and demonstrated that electronically, Me4CF3Cp is
equivalent to Cp. Specifically, the Ru 3d core binding energies of (Me4CF3Cp)CpRu
and ruthenocene are equal. Based on Gassman’s data, one CF3 group is as electron
withdrawing as four methyl groups are electron donating, when the functions are
bonded to a coordinated ligand.111
Gassman’s results are consistent with the AGj° value of 3. The TTFMH ligand
is approximately three times more electron withdrawing as the Cp* ligand is electron
donating. Therefore, a methyl group lowers the ionization potential of ruthenocene by
ca. 2.5 kcal mol'1 per group, based on AGj° data for CpCp*Ru and Cp*2Ru. Thus, if
four methyl groups have the opposite effect to one CF3 group, a CF3 function on a Cp
ring of ruthenocene should increase the AGj° relative to Cp2Ru by ~10 kcal mol'1.
Therefore, the TTFMH ligand of 3 should increase the AG¡° relative to ruthenocene by
~40 kcal mol'1 and the Cp* ligand is expected to decrease AG¡° ~13 kcal mol'1; an

Ill
estimated AG ° value for 3 of 192 kcal mol'1 which is the equal to the AG¡° value
estimated from electron-transfer bracketing experiments of 3.
The free energy of ionization of 4 is lower than AG¡°(3) by ~20 kcal mol"1.
The addition of the triethylsiloxyl function in place of the a hydrogen for TTFMOSi
accounts for the lower ionization potential. The siloxyl group is electron donating and
i rvQ
counters the electron-withdrawing nature of the CF3 groups.
The AGj° value of 6 is less endoergic than ruthenocene due to the strong
destablizing inductive effect of the nitro group. The AG° value of CpCp*Ru is 10
kcal mol'1 less than AG¡°(6) indicating that the N02 groups is as electron withdrawing
as a CF3 group. Both functions stabilize ruthenocene towards oxidation by -10
kcal mol'1.
Ruthenocenes Bearing a Perhalogenated Cyclopentadienyl Ligand
Three ruthenocene derivatives with a supporting perhalogenated Cp ligand have
been studied. The pentafluorocyclopentadienyl complex, which represents the first
metal complex possessing the r)5-C5F5 ligand, was recently synthesized by Hughes and
Cumow.114 The bromo complex was prepared by Winter and coworkers114 by
successive permercuration/perbromination of Cp. Compound 5 was prepared by the
method of Gassman and Winter.91 The interest in perhalogenated cyclopentadienyl
ligands is in their electron-withdrawing properties, in addition to novel solubility and
reactivity relative to Cp.109,113 Little is known concerning the electronic properties of
perhalogenated Cp ligands 91,109,115 thus electron-transfer equilibrium methods can be

112
used to quantify the effects these ligands have on the physical and chemical properties
of metal complexes.
In all three cases, the perhalogenated ligand stabilizes the ruthenocene
complexes towards oxidation. The electronic effects of the bromo and chloro ligands
are equivalent,3,5 as AG¡°(5) = 165.4 ± 2 kcal mol'1 is only slightly greater than that
of AGj°(14) = 165.1 kcal mol'1. For comparison, the ionization potential
chlorobenzene is 2 kcal mol'1 more endoergic than the IP of bromobenzene.27,39 The
perhalogenated bromo and chloro ligands are essentially opposite in electronic effects
to Cp*. Since the AG ° values of complexes 5 and 14 are similar to AGj0(Cp2Ru), it
can be concluded that the electron-donating nature of the Cp* ligand is countered by
the electron-withdrawing character of the halogenated ligands.
The measured AGet° values for electron-transfer equilibrium of Cp*(C5F5)Ru
with reference compounds azulene (AG¡° = 167.8 ±1.0 kcal mol'1) and m-toluidine
(AG¡°= 172.2 ± 1.5 kcal mol’1) are +2.6 and -1.0 kcal mol'1. The AG¡° value for the
for complex 13 was determined to be 170.8 ± 1.5 kcal mol'1.
The perfluoro ligand was found to have electronic effects comparable to the
perchlorocyclopentadienyl ligand. Replacing Cl for F results in an increase in AG¡° by
~1 kcal mol'1 per F. Compared to Cl, F is expected to increase the ionization
potential due it lower polarizabilty, however 7i-resonance effects of F compensate53 for
the polarizabilty in the case of ruthenocene complexes. Resonance donation of the F
lone-pair electrons into the Cp ring diminishes the expected large positive inductive

113
effect. The overall result is that the AG ° values for the chloro and fluoro compounds
are equivalent.
Structurally, the metal complexes bearing perhalogenated ligands are ostensibly
comparable to ruthenocene and decamethylruthenocene. In both Cp2Ru and Cp*2Ru,
the Cp rings are eclipsed and crystal structure data for 5 demonstrates that the methyl
groups and the chlorides are also eclipsed; however, an indication of the electron-
withdrawing ability of the C5C15 ligand can be inferred from the ruthenium-to-ring-
centroid distance. The metal center is unsymmetrically displaced closer to the
electron-withdrawing ligand. The Ru-Cp* distance is 1.82 Á while the Ru-C5Cl5
distance is 1.78 Á.91 This indicates that Ru is being held closer to the more electron-
deficient ligand, possibly as a result of the electron-donating ability of Cp*.
Free Energies of Ionization of Ruthenocenes Derivatives with Fused-Ring Systems and
Electron-Donating Ligands
Gassman and Winter noted that Cp*FluRu was more easily oxidized than
Cp*2Ru based on electrochemical measurements and Ru 3d core binding energies
measured by XPS; thus, Ru was determined to be more electron donating than Cp*.91
Because the reported electrochemistry of 11 is irreversible,91 and the core binding
energies may not reflect valence ionization potentials, their conclusions are in doubt.
Comparison of the free energies of ionization of the two compounds indicates that 12
has a lower AG¡° value than 11, which is inconsistent with Gassman’s findings.
Moreover, decamethylruthenocene has the lowest AG° value of all the ruthenocene
compounds studied. Based on the adiabatic ionization potentials, Ru does not

114
stabilize formation of Cp*RuL+ to the extent that Cp* does. The reported
electrochemical potentials91 may follow a different trend due to different solvation
energies of 11 versus 12 or because of the irreversible electrochemistry of 11 does not
yield an accurate assessment of the true thermodynamic potential. Definitive
interpretation of the XPS results with the electron-transfer equilibrium data is not
possible to due the limited amount of data that can be compared. Figure 4-2 is a plot
of XPS binding energies91 and AG° data for several ruthenocene complexes. It is
obvious from the plot that the correlation fails for Cp* and Flu.
Electron-transfer equilibrium of Cp*FluRu and Cp*2Ru (reaction 4-6) was
attempted but the reaction was exothermic favoring Cp*FluRu+ even at pressure
Cp*2Ru + Cp*FluRu+ —» Cp*2Ru+ + Cp*FluRu (4-6)
ratios approaching 100. The entropies of ionization for these two compounds have not
been determined. Significant differences in the AS° values could realistically reverse
the trend in the A//i350° values. However, in order to account for a 5 kcal mol'1
difference in the AH° values, a difference in the AS¡° values of ~15 cal mol'1 K'1 is
required. There is no obvious reason for an entropy difference this large. The thermal
ionization free energies reported here probably reflect the electronic effects that would
be expected in chemical reactions of metal complexes with these supporting ligands.
The AG ° values of CpCp*Ru and Cp*2Ru are consistent in that the electronic
effects of the Cp* ligand are additive. Thus the AG° value for 2 is -12 kcal mol'1

Ru 3d5/2 Binding Energy (eV)
115
6.2 6.4 6.6 6.8
Ionization Free Energy (eV)
Figure 4-2 Plot of Ru 3d binding energies from reference 91 versus ETE AG®
values for several Cp*Ru-L complexes and ruthenocene.

116
less than ruthenocene and the AG¡° value of 12 is ~26 kcal mol'1 less than
ruthenocene. From these data an electron-donating effect of ca. 2.5 kcal mol'1 per
methyl group can be estimated.
The AG ° of Cp*IndRu is only 3 kcal mol'1 less than that of 2 indicating that
the indenyl ligand stabilizes a positive charge only slightly more relative to Cp. The
AG¡0 for (Ind)2Ru confirms that the indenyl ligand is only moderately electron
donating. Gassman and Winter concluded that the electronic properties of the indenyl
species were between cyclopentadienyl and pentamethylcyclopentadienyl91 and the
AG° data for compounds 8 and 10 is consistent with their results.
Electron-transfer equilibrium for 7 was established with N,N-diethylaniline and
N,N-dimethyltoluidine with good internal consistency (0.6 kcal mol'1 difference in the
measured AGet° value and the expected AGet° value, based on published AG¡° data for
the reference compounds). Electron-transfer equilibrium for 9 was established with
ferrocene and ethylferrocene. Similar to the indenyl and permethyl ruthenocene
derivatives, compounds with TMSCp ligation demonstrate an additive electronic effect.
From the AGj° data for 7 and 9, a trimethylsilyl group is estimated to be roughly as
electron donating as a methyl group. The electronic effects of other (TMS)nCp
ligands, where n= 1, 2, or 3, are consistent with the present observations.116
Attempted Correlation of Ruthenocene Ionization Free Energies with
Taft Gj Parameters
Levitt has demonstrated that correlation of ionization potentials with Taft
inductive substituent constants exhibits a linear frce-energy relation.50,51 Correlation

117
of PES ionization energy data for chromium arene complexes (C6H5R)Cr(CO)3 with
Taft inductive parameters (r = 0.996)51 and chromium acetylacetonato Cr(acac)3
complex with Hammett substituent constants (r = 0.99 8)50 generated excellent fits.
This allows the ionization energies of compounds that do not form stable gas-phase
species or are difficult to isolate or purify to be calculated fairly accurately from a
simple linear function.
Correlation of AG ° values for alkyl ferrocene16 and nickelocene19 derivatives
were demonstrated in earlier chapters. The fits of the AG¡° data to the established o¡
parameters showed strong correlation (r = 0.997 for ferrocene and r = 0.994 for
nickelocene) however a similar fit for the ruthenocene complexes was not as
successful. An attempt to correlate several of the ruthenocene complexes with o¡
parameters with is shown in Figure 4-3. A correlation of the entire series was not
possible because g¡ parameters do not exist for all the substituents such as OSiEt3 and
the fused-ring systems. The poor correlation of AG¡° values in Figure 4-3 indicates
that these parameters are not suitable for predicting Cp ligand effects in the
ruthenocene derivative ionizations. As expected from pervious correlations, the
alkylated derivatives Cp2Ru, CpCp*Ru and Cp*2Ru yield a linear fit; however,
compounds with electron-withdrawing groups (Cp*(N02Cp)Ru, Cp*(C5Cl5)Ru, and
Cp*((CF3)4C5H)Ru, for example) do not increase the ionization energy relative to
ruthenocene to the extent that is predicted by the or parameters. The line drawn in
Figure 4-3 is for the best fit of compounds 1, 2, and 12.

118
Figure 4-3 Correlation of AG ° values for several ruthenocene derivatives with Taft
Oj parameters. The best-Fit line is drawn for AG¡° values versus Z(Gj)
for the methylated complexes only (1, 2, 12). Compounds numbers
defined in Table 4-1 and Figure 4-1

119
The slope of the line (the p value referred in equation 4-7) in Figure 4-3 is 58 ± 1
kcal mol'1 which is indicative of the sensitivity of ruthenocene to alkyl
AG¡°(LL’Ru)= p(IOj) + AG¡°(Cp2Ru) (4-7)
substitution. Ruthenocene is equally as sensitive as ferrocene (p = 57 ± 1 kcal mol’1)
to alkyl substitution.
Experimental data suggest that the published value of for TMS (-0.10 - 0.11)
n c
may be too large. ’ The correlation of Gj values with AGj° data for Cp*(TMSCp)Ru
and (TMSCp)2Ru indicates that the predicted AG¡° values for the TMSCp complexes
are consistently lower than the experimental electron-transfer equilibrium by ~5 kcal
mol'1. From the estimated AG¡° values, a TMS function is relatively as electron-
donating as a methyl group. Gassman and workers recently estimated that the TMS
function is 1.25 times more electron donating than a methyl group.116 The Oj value
for Me is -0.046, hence a Gj value for TMS of -0.06 would predict AG¡° values of 7
and 9 in better agreement with the experimental AG¡° value and would also be more
consistent with the data of Gassman and workers.116
The predicted AG ° value for Cp^CjCl^Ru based on the slope for the
methylated compounds is 120 kcal mol'1 more endoergic than the AG¡° derived from
ETE. Similarly, the value for the Cp*(C5Br5)Ru is also 120 kcal mol'1 lower than
predicted from a Taft analysis. The absolute values of the Taft parameter for Cl (Gj =
0.46) is an order of magnitude larger than the methyl parameter (Gj = -0.046) yet the

120
experimental AG¡° values for ruthenocene and Cp*(C5Cl5) are equivalent. This infers
that a Cl group is approximately as electron-withdrawing as a methyl group is
electron-donating for the ruthenocenes. A similar comparison is also true for the
bromo and the fluoro derivative. The predicted AG¡° value for 13 is -125 greater than
the experimental value. Note that the AG ° data for compounds 5 and 14 are
consistent with the relative ordering of the Taft parameters in that the Gj for Br and Cl
are +0.44 and +0.46 respectively. The AGi°(Cp*(C5F5)Ru) is also in agreement with
the ordering of the halogen Gj parameters. Although the Taft parameter analysis is
inappropriate for these ruthenocene derivatives, the trend of the predicted order is
consistent with trend of the experimental AG¡° data.
The small electron-withdrawing nature of the halogens is consistent with the
observations of Levitt and Levitt who found that substituting Cl or Br for an arene
ring proton in (C6H6)Cr(CO)3 results in only a 0.1 eV increase in the ionization
energy and F substitution results in a 0.2 eV increase. The moderate increase is
attributed to lone-pair 7i-resonance effects which back-donate electron density to the
aromatic ring. This effects has been well documented for simple benzene
"5 *5 S'?
systems. ’ ’
The parameters used for fitting the gas-phase data have been described by Taft
and workers and include in addition to Gj parameters, Ga parameters for polarizability
effects and gr+ used to predict 7t-resonance donation effects.5 The published
parameters for several substituents in Figure 4-3 are given in Table 4-3. A negative
value for a parameter indicates that the electronic effect will stabilize a decrease in

121
electron density at a reaction center and lead to a lowering in the ionization potential.
The field effects of Cl and F are essentially equal so only the polarizabilty and
resonance effects will contribute to the difference of the substituents. Fluorine,
because of its compact size, is less polarizable than Cl, but the n resonance effect of
F, which is larger than that of chlorine (Table 4-3), compensate substantially in the
case of the ruthenocene derivatives. The overall result is that the AG¡° values for the
complexes are similar. A parallel argument can be used for the bromo complex also.
Note that the van der Waals radius of F is 0.5 Á smaller than of chlorine and the
electronic effects are similar.6
It is also notable that the observed substituent effects for F (13) and CF3 (5)
are reversed from what would be predicted by the Oj parameters. The observed shift
in the ionization potentials of 3 and 6 (replacing H for CF3 or N02) are also much
smaller than predicted by Taft parameters but the AGj° values do not deviate from
experiment to the extent of the halogenated complexes. The estimated AG¡° values for
3 and 6 are 50 and 25 kcal mol"1 greater than the ETE values respectively. Neither
the CF3 group nor the N02 group are capable of n resonance back donation to the Cp
ring (cR+ = 0 for CF3 and N02).53 The larger electronic effect of CF3 compared to F
is consistent with a model of 7t lone-pair resonance donation for the halogens but not
CF3 or N02.115 Such effects are well-known from studies of organic substituent
effects. For example, the ionization potential of CF3-C6H5 is 11 kcal mol"1 greater
than for F-C6H5,27 yet the parameter for F is similar to Gj(CF3). The ordering of
the benzene analogue ionizations is consistent with the n resonance model.

122
Table 4-3 Substituent Parameters for Selected Cyclopentadienyl Derivatives
Ligand
Substituent
°ia
aFb
aR+b
C5(CF3)4H
cf3
0.42
-0.25
0.44
0.0
c5h4no2
no2
0.65
-0.26
0.65
0.0
c5f5
F
0.5
0.13
0.44
-0.25
c5ci5
Cl
0.46
-0.43
0.45
-0.17
C5Br5
Br
0.44
-0.25
c5h4tms
TMS
-0.1
-0.72
-0.02
0.0
Cp
H
0C
OF
0C
0C
Cp*
Me
-0.046
-0.35
0
-0.08
a. Values for parameters obtained from reference 2.
b. Values for parameters obtained from reference 5a.
b. By definition.

123
A New Parameter Scale for Cyclopentadienyl Substituents Based on
Gas-Phase Electron-Transfer Equilibrium Studies of Ruthenocenes
Given the absence of parameters for many of the ruthenocene derivatives 1-14,
particularly the fused-ring ligands fluorenyl and indenyl, and the lack of general
correlation with Oj parameters as noted in the previous section, a new parameter scale
for cyclopentadienyl derivatives has been developed.20 The overall parameters for the
Cp derivatives are based on the electronic effects of the ligand rather than the
individual ring substituents. The parameters relate specifically to the stabilization or
the destabilization in the oxidation of ruthenocene derivatives. It is expected that the
parameters will correlate the tendency of the ligands to promote processes that involve
a positive charge buildup at a metal center in metal complexes bearing these ligands.
To establish the new Cp ligand parameters, which will be denoted as y values,
the scale must be anchored to some reference values. Since the Taft parameters are
anchored to Me and h,3,5,50'55 it seems logical that the new y parameters be anchored
to the cyclopentadienyl and pentamethylcyclopentadienyl. The parameter values y(Cp)
= 0 and y(Cp*) = -1.00 are arbitrarily assigned.20 The negative sign indicates that the
y parameter will stabilize an increase in positive charge at a metal center or, in this
case, result in a lowering of the ionization potential of the ruthenocene complex. A
positive sign indicates that the parameter will destabilize an increase in positive charge
at a metal center. Alternatively, the ligands may also be used for both positive and
negative electronic effects however the application of y parameters for negative ions is
presently untested.

124
Two approaches have been used to assign ligand parameters to the various
ligands. The first approach uses only data for the compounds bearing a
pentamethylcyclopentadienyl ligand, Cp*RuL derivatives were L denotes the ligand for
which these parameters are to be derived. The ligand parameters for this scheme are
designated as y* value in which the asterisk corresponds to the Cp*-Ru base.
Equation 4-8 was used to obtain the y* parameters. The value of the sensitivity
AG¡°(Cp*Ru-L) =a*[7*(L)] + AGi°(CpCp*Ru) (4-8)
= 14.4[y*(L)] + 152.3 (kcal mof1)
parameter a* = 14.4 kcal mof1 is fixed by the anchored values of Cp and Cp*
(0 and -1.00 respectively). The a* value can be compared to the p value in a Taft
analysis, as both parameters are determined from the slope of the correlation. A plot
of the y* parameters and AG ° data for the ruthenocene complexes based on equation
4-8 is given in Figure 4-4. The derived ligand parameters were presented in Table
4-2. Because of the manner in which the y* values were derived, all the points
intrinsically yield a perfectly linear relationship. The plot is intended as graphical
representation of the relative electronic effects of the Cp ligands and is not a result of
two independent sets of data. The free energies of ionization for the ruthenocene were
used to calculate y* values from equation 4-8, and not the reverse as is the case with
the Taft Gj analysis. The ligands to the right of the dashed line in Figure 4-4
destabilize the oxidation of the ruthenocene complex Cp*Ru-L relative to CpCp*Ru.

125
Figure 4-4
Plot of AG ° values versus y ligand parameters derived from the
equation AG¡°(Cp*Ru-L) = 14.4y + AG¡°(CpCp*Ru) (equation 4-8).

126
A drawback to the above parameter scale is that the 7* values are based on
AG° values for only one system. Errors in the derived 7* values are not obvious
since the values are based solely on a single AG¡° value. Moreover, the additivity of
the parameters has not been substantiated.
A second approach for deducing ligand parameters makes use of the
homoleptic compounds (L^Ru) in this study such as Ind2Ru and (TMSCp)2Ru.
Incorporating all the gas-phase AG ° values requires that the electronic effects be an
additive property of the ligands. The second set of parameters derived from equation
4-9 are simply referred to as 7 parameters. The utility of the parameter scheme based
on equation 4-9 is that it tests the validity of ligand additivity.
AGj°(LL’Ru) = a(y(L) + y(L’)) + AGi°(Cp2Ru) (4-9)
The best-fit value (a = 13.0 kcal mol'1) is based on the AG ° values for Cp2Ru,
CpCp*Ru and Cp*2Ru. The experimental AG¡° values from Table 4-1 were then used
to derive best-fit 7 values for the other Cp ligands from equation 4-9. The 7 values
are presented Table 4-2 with the 7* parameters. In many cases a 7 value was derived
from a single AG¡° value since the ligand appears on a single compound only. The 7
parameters the indenyl and TMSCp derivatives were determined from an iterative fit
of the AG ° values of the metal complexes bearing these ligands. For these two
ligands, the 7 value was varied until the sum of the difference in the predicted AG¡°
values and the ETE derived AG¡° data was minimized. The indenyl and TMSCp

127
ruthenocene complexes test the assumption that the electronic effects of the ligands is
additive.
The correlation of AG ° data with the (y(L) +y(L’) sum for compounds 1-14 is
shown in Figure 4-5. From the correlation shown if Figure 4-5 and the predicted and
experimental AG ° in Table 4-1, it is apparent that the ligand additivity assumptions is
approximately substantiated. Predicted values of AG ° are typically within 2-3 kcal
mol'1 of experimental values. As mentioned earlier, the additivity assumption has not
been tested for all of the ligands.
Included in Table 4-1 are the XPS Ru 3d core binding energies of several
ruthenocene complexes measured by Gassman.91 Comparisons of Gassman’s data
with electron-transfer equilibrium derived AG¡° values have been discussed in previous
section. The general correlation, except for the reversal in the trends for 11 and 12, of
the XPS data with AGj° values is good (Figure 4-2). Subsequendy, because of the
strong correlation in Figure 4-2, the y parameters are also expected to correlate well
with the 3d binding energies. A plot of 3d XPS data versus y parameters for several
complexes L^ZrC^ complexes116 also demonstrates good correlation (r = 0.986);
therefore, the core binding energies of the Ru and Zr systems demonstrate reliable fits
with the ligand parameters which indicates that ionization energetics data determined
independent of this work can be correlated with the new Cp ligand y parameters.
The y and y* parameters are anticipated to be useful in predicting electronic
effects for different organometallic systems other than the simple metallocenes.

128
Figure 4-5
Plot of AG ° values versus Z(y) ligand parameters derived from the
equation AG¡°(LrRu-L2) = 13.0(7(1.,) + -y(L2)] + AG¡°(Cp2Ru)
(equation 4-9).

129
The application of these parameters in predicting reactivity of organometallic systems
that have catalytic properties is of great interest.
Rates of Hvdrogenolvsis of Methylzirconocene Cations
The rates of hydrogenolysis, ^H’ for methylzirconocene cation and several
methylzirconocenium derivatives have been examined. The purpose of the study was
to demonstrate that the Cp ligand parameters derived from free energies of ionization
for ruthenocene complexes can be applied to interpret and predict chemical reactivity.
The hydrogenolysis reaction for the methylzirconocene cations is convenient because
steric effects are not expected to be important.14 The insertion of H2 into the Zr-CH3
bond is sterically undemanding compared to other reaction involving larger
substrates.14,117 For example, reactions of Cp*2ScCH3 with alkenes are significantly
influenced by the steric effects of the Cp* ligands.118 The general reaction shown
below (reaction 4-10) was studied in the gas-phase by using FTMS. Ancillary Cp
ligands are denoted by L. Methylzirconocene cations were generated from low energy
L2Zr-CH3+ + H2 -> L2ZrH+ + CH4 (4-10)
electron-impact of the L^ZrtCH^ neutral which eliminates CH3.14 The application of
FTMS to study homogeneous organometallic catalysts in the gas-phase has been
described by Christ et al.14 Similar experimental procedures were used for the present
hydrogenolysis studies reported here.

130
From the data in Figure 4-6, it is apparent that the electronic effects of
supporting ligation influence the rates of hydrogenolysis of the zirconocene cationic
complexes. A similar observation on the reactivity of zirconocene complexes was
observed by Spaleck and coworkers.119 In their studies, the orientation of ancillary
indenyl ligands in several bis(indenyl)zirconium derivatives permitted for open
coordination sites to the metal center. Subsequently, the observed rates of propylene
polymerization by the Inc^ZrCl-, complexes were assumed to be primarily influenced
by the electronic characteristics of the ligands.119
The rates of hydrogenolysis for only a few complexes were examined.
Therefore, a meaningful correlation of data with the Cp ligand parameters was not
possible. However, the study serves to illustrate that their exists a measurable
electronic effect for the zirconocene cationic complex reactivities.14,119 Developing
structure/reactivity relationships is fundamentally important in organometallic
chemistry and further study of the gas-phase rates of hydrogenolysis for the L^Zr-
CH3+ may serve towards achieving that goal.
Comparisons of Gas-Phase Ionization Free Energies to Solution
Electrode Oxidation Potentials
As the vast majority of organometallic chemistry is studied in the condensed
phase, understanding the effects solvent has on the energetics of organometallic
reaction is fundamentally important. Direct comparison of gas-phase and solution
oxidations quantifies the effects solvent has upon the ionization process.16'18'20'54'84

131
+
3.7E-11
1.3E-11
1.4E-12
Slow/Decomposed
Figure 4-6 Gas-phase rates of hydrogenolysis for several methylzirconocene cation
complexes. Numbers adjacent to structures are kH values (cm3molec'1s'1).
Measurement of the kH value for Flu2ZrCH3+ was complicated by an
inefficient reaction rate and thermal decomposition of the Zr complex.

132
In Chapter 2, thermodynamic cycles (Figure 2-5) that include AG ° data in the gas
phase and solution were used to determine differential solvation free energies,
AAGsolv° values. Similar cycles have been used here to determine the differential
solvation energies for the ruthenocene derivatives. Electrochemical E^ values
measured by cyclic voltammetry are compared to gas-phase AG ° values determined
from electron-transfer equilibrium reactions. Derived AAGsolv° values from
thermochemical cycles will be discussed following analysis of the electrochemical
oxidation potentials.
Oxidation potentials for several of the ruthenocene derivatives have been
previously reported.90,91,109 However, as many of the reported values were based
on irreversible cyclic voltammetry measurements,90'91 reevaluation of the oxidations
was performed in order to evaluate ambiguous E^ assignments therefore
electrochemical oxidation potentials for a many of the ruthenocene derivatives were
measured here by using cyclic voltammetry. For complexes that are irreversibly
oxidized, rather than estimate E^ values, anodic peak potentials £pa are reported.
Accurate assessment of £pa potentials is not possible based solely on literature E^
values. Furthermore, in order to evaluate the E^ data for the entire series of
complexes under common reaction conditions, reevaluation of the electrode potentials
was necessary in order to establish more accurate comparisons of the oxidation
potentials for the series of compounds. Literature potentials were determined under
various reaction conditions such as different supporting electrolytes and reference
electrodes.90,91,109 In many cases, oxidation potentials that were reported irreversible

133
were found to be either reversible or quasi-reversible. A quasi-reversible potential is
used to refer to the situation in which the ratio of the intensities of the anodic and
cathodic peaks is not unity and both cathodic and anodic waves are undoubtedly
present. An irreversible potential describes an electrochemical process in which either
the cathodic potential of an oxidation or the anodic potential of a reduction is not
definitively observed due to decomposition of the ionized species and/or solvent
complications.120,121
All potentials were measured against the oxidation potential of ferrocene, which
was used as a le" internal standard.103 For aqueous electrochemistry, the normal
hydrogen electrode (NHE) or saturated calomel electrode (SCE) are universally
accepted reference standards.121 Not all compounds, including the metallocenes,89 are
soluble or stable in aqueous media, therefore E^ potentials must be measured in
alternative solvent systems. However, for nonaqueous systems there is no universally
accepted reference standard. The advantage of using ferrocene as an internal
standard for nonaqueous electrochemistry is that the Cp2Fe+y,° couple is reversible in a
number of different solvent systems. Gagne and workers have demonstrated that
although the formal potential for a metal redox couple may be shifted when different
reference electrodes or solvent systems are used, the potential difference relative to the
Cp2Fe+/0 couple was constant103 Different solvents have different junction potentials,
which are usually unknown, and diverse resistivities which can result in complicated
wave shapes, i.e. seemingly irreversible behavior.103 By comparing the A£p or E^

134
value of a reaction couple to £i¿(Cp2Fe+y,°), these various experimental limitations to
nonaqueous electrochemistry can be circumvented.
Table 4-4 lists electrochemical £pa and E^ values for compounds 1-14.
Reversible potentials were not observed for all compounds; therefore, only £pa values
relative to ferrocene are reported. Additionally, literature E^ values are included when
available for comparison. All electrochemical potentials reponed here were measured
in methylene chloride containing 0.1 M tetrabutylammonium hexafluorophosphate at a
platinum disk electrode (working electrode) with a KCL saturated Ag/AgCl reference
electrode. All values are reported versus the Cp2Fe+/0 couple which was observed to
be reversible for all measurements. The ruthenocene oxidations are expected to
proceed by a le' process to the corresponding cation.90 However, for cases were the
oxidation was irreversible, a rapid 2e' process is probable. ’ For example,
ruthenocene is irreversible most solvent/electrolyte systems.90 An observed 2e*
process is thought to result from the initial oxidation of Cp2Ru+ followed by rapid
disproportionation.90,92 It is likely that the larger metal-ring separation relative to
Cp2Fe+ enhances the nucleophilic attack of Cp2Ru+ by the supporting electrolyte or
solvent.
Recently, Mann and workers reported the reversible oxidation of ruthenocene in
CH2C12 containing 0.1 M tetrabutylammonium tetrakis[3,5-bis(trifluoromethyl)phenyl]
borate, TBA'TFPB', as a noncoordinating electrolyte.90 The reversible
electrochemistry for Cp2Ru in the presence of the strongly noncoordinating
TBA'TFPB' is consistent with the concept that ruthenocene is susceptible to

135
Table 4-4 Electrode Potentials and Differential Solvation Free Energies for Some
Ruthenocene Derivatives.
No.a
t? b,c
£pa
F b
AGi°(soln)d,e
-AAGsolv°d,f
1
0.58
0.56h,i
130
34
0
2
0.35
0.23*
126
27
+7
3
1.3*
148
44
-10
4
1.12*
144
28
+6
5
1.11
1.03h
141
24
+10
6
0.78
122
40k
-6
9
0.35
113
38k
-4
10
0.23
0.12h
123
26
+7
11
0.21
112
21k
+13
12
0.15
0.0911
120
18
+16
13
1.07
140
31k
+3
14
1.03
0.95h
139
26
+ 12
a. Compound numbers defined in Table 4-1.
b. Values in volts relative to Eft ox(Cp2Fe).
c. Measured in CH2C12/0.1 M Bu4NPF6. See Experimental Methods.
d. Values are kcal mol'1.
e. From equation 4-11.
f. From equation 4-12a.
g. From equation 4-12b. AAGsolv°(Fc+/0) = -34 kcal mol'1 in CH2C12.
h. Eft value reported reported for quasi-reversible potential.
i. See reference 91.
j. See reference 109.
k. £pp for ferrocene used to estimate Eft from £pa value.

136
nucleophilic attack by coordinating anions. The oxidation of decamethylruthenocene is
reversible in CH2C12 with most common electrolytes, conceivably due to steric
hinderance of the metal center by the Cp* ligands. Reversible E^ values for
unsubstituted ruthenocenes are obtained with only the most noncoordinating anions.
Anions such as PF6', C104', and BF4‘ contribute to irreversible electrochemical
90 91
responses.
The oxidation of Cp*(C5F5)Ru (13) was measured in CH2C12 containing 0.1 M
Bu4NPF6 by Richardson et al. An irreversible 2e' potential was observed for the 13
with £pa = 1.07 V versus £)¿(Cp2Fe+/0). Error limits for the oxidation potentials are
given in Table 4-4. At a platinum disk electrode, the peak potential for the 13 was
approximately twice that for ferrocene, consistent with a 2e' process. The oxidation of
13 is also irreversible when the weakly nucleophilic tetrakis[3,5-bis(trifluoromethyl)
phenyl]borate (TFPB) is the anion of the electrolyte (£pa = 0.98 V). The Epa value for
13 can be compared to the £pa = 1.11 V for Cp*(C5Cl5) (5) in CH2C12/Bu4NPF6 and
1.08 V with TBA+TFPB' as the supporting electrolyte. Interestingly, the oxidation of
5 is quasi-reversible in methylene chloride with E^ ~1.0 V vs. Fc+/0 and the AE^
value was independent of supporting electrolytes used in this work. The bromo
derivative also has a quasi-reversible oxidation potential in CH2Cl2/Bu4NPF6 with
E^ = 0.80 V versus Fc+/0. The observed reversible responses for 5 and 14 support
that ruthenocenes with bulky ligation prevent nucleophilic attack while unhindered
ruthenocenes are more susceptible to follow-up reactions subsequent to oxidation.

137
The van der Waals radii of H and F are equivalent6 and a 2e' oxidation both
ruthenocene and 13 is observed. The van der Waals radii of Cl is larger than that of F
but slightly smaller than CH3. Thus, the steric restrictions of a methyl group are
expected to be similar to a chlorine or bromine, which may account for the observed
reversible oxidation potentials for 5 and 14. In this manner, the iodo complex,
Cp*(C5I5)Ru,114 is predicted to give a quasi-reversible electrochemical response.
The order of the £’pa values for the three halogenated complexes is dissimilar to
the AG ° values in that the chloro derivative has a higher oxidation potential than the
fluoro derivative. The irreversible oxidation of 13 however may account for the
observed trend and the order of the true thermodynamic E° values may be parallel to
the gas-phase AG ° values.
Determination of Free Energies of Ionization in Solution
from Electrochemical Oxidation Potentials
Free energies of ionization in solution have been derived previously for
metallocenes16,19,20 and coordination complex ions.84 Half-cell potentials for the
metal complex redox couples referenced to a standard aqueous reference electrode can
be referenced to the absolute potential of the normal hydrogen electrode by assuming a
potential for NHE.54 From energy cycles, it can be shown that FT* + e' = !AH2(g) has
an absolute thermodynamic potential of 4.44 V. Through the Faraday relationship,
AG° = -nFE°, estimates of AG¡°(soln) can be made.54 Equation 4-11,
AGj°(soln) = nF(E^ + £ref +
(4-11)

138
was used to estimate AG¡°(soln) values for the ruthenocene derivatives studied.
Oxidation potentials reported in Table 4-4 were corrected to include Avalues
relative to £°(Cp2Fe+/0). Therefore, the ruthenocene oxidation potentials were
corrected by measuring the Ato ferrocene (¿s^fFc-^0) = 0.46 V in CH2C12,89,103 and
then adding the potentials of the hydrogen and reference electrodes. Irreversible
electrochemical potentials were used to estimate AG¡°(soln) values for several
ruthenocene complexes, accordingly £pa values were used when data unavailable.
Values of AG^Csoln) reported in Table 4-4 are considered absolute thermodynamic
potentials for the process M(soln) = M+(soln) + e\ where e' is a treated as a gas-phase
electron.54
Differential Solvation Free Energies for Several Ruthenocene/
Ruthenocenium Couples
Differential solvation free energies, AAGsolv°, for the ruthenocene couples are
reported in Table 4-4. Values of AAGsolv° were determined from the AG¡°(soln)
estimates. In an effort to establish a reference standard for estimates of solvation free
energies from metal redox couples, AAGsolv° values referenced to
¿¿Gsolv°(Cp2Fe*' couple was chosen as reference standard for AAGsolv° values for several reasons. The
gas-phase16 and solution oxidation potentials34,89,103 are well established values,
therefore AAGsolv°(Cp2Fe+/0) is accurately known in CH2C12 and CH3CN. Further, as
the Cp2Fe+/0 couple is reversible in many solvents,103 shifts in the redox potentials for
metal complexes (AE^) in different systems are can conveniently be compared relative

139
to ferrocene. Additionally, the experimental differential solvation free energy for the
ferrocene couple is modeled accurately by the dielectric continuum theory indicating
that ferrocene is nearly spherical in solution.116 Thus, ion-solvent interaction for
Cp2Fe+/0 are expected to be dependent on outer-sphere electrostatic interactions and
not specific solvent interactions such as ion/solvent-dipole interactions isolated at the
metal center. From this, inference to the ion-solvent interactions for metal redox
couples can be made with respect to a pseudo-spherical well-modeled le' redox
couple.
Equation 4-12a was used to derive AAGsolv° values for the ruthenocene
derivatives. From equation 4-12b, AAGsolv° values relative to AAGsolv°(Cp2Fe+/0)
were derived. By convention, AAGsolv° values are reported as negative values to
indicate the amount of stabilization energy a solvated ion gains with respect
-AAGsolv° = AG¡0(g) - AGj°(soln) (4-12a)
5pc = AACso1v° - AAGsolv°(Cp2Fe+/0) (4-12b)
to the parent neutral. Negative 5pc values indicate that the solvation energy of the
ruthenocene couple is greater (more exoergic) than Cp2Fe+/0. For oxidation couples
where the 5pc value is positive, a decrease in solvation stabilization is observed
relative to Cp2Fe+^°. For example, the AAGsolv° value for Cp*2Ru+/0 is -18 kcal/mol
and 5pc = +16 kcal mol'1 in methylene chloride. In the gas phase, the methyl
functions of the Cp* ligand can be thought of as a solvent layer around ruthenocene;

therefore, in a dielectric continuum which would be an additional solvent layer,
Cp*2Ru and Cp*2Ru+ are essentially equally solvated.
140
Application of the Bom Model for Estimating Solvation Energetics for
Ruthenocene Oxidation Couples
The Bom model for predicting solvation free energies has been applied to
metallocene Cp2M+/0 couples in Chapter 2. Values of AGel° determined from the
Bom equation are equivalent to AAGsolv° estimates from thermochemical
cycles.16,54,84 For all cases save the Cp2V+/0 couple, an experimental value for
AAGsolv° = -38 ± 2 kcal mol'1 was obtained for the first transition row metallocenes.16
The Bom model is successful in estimating the change in electrostatic free energies,
AGel°, for the metallocene +/0 couples within 2 kcal mol'1 when acetonitrile (D = 36)
is the solvent. The Bom equation (AGel° = (166z2/reff)(l-l/Z))) predicts AGel°(L2M+/,°)
values for ruthenocene and ferrocene couples of -34 kcal mol'1 when methylene
chloride (D = 9) is the solvent; however, AGel° values for several ruthenocene
derivatives are inconsistent with experimental AAGsolv° data. The AAGsolv° value for
Cp*2Ru = -20 kcal mol'1 but AGel°(Cp*2Ru+/*°) = -28 kcal mol'1 in methylene
chloride. The overestimation of AGel°(Cp*2Ru+/0) from equation 4-13 may indicate
that a simple electrostatic model may not apply for nonspherical or asymmetrical
complexes. The crystallographic radius for Cp*2Ru,124 estimated from the Ru center
to the van der Waals radii of the methyl protons, used to estimate AGel°(Cp*2Ru+/0) is
5.2 Á. Because the Bom model assumes that ions are charged spheres, a radius of 5.2
Á may inappropriate if decamethylruthenocene is not a spherical complex.

141
The AAGsolv° value for the asymmetric complex Cp*IndRu is -26 kcal mol'1
and the AGcl° value is estimated to be -26 kcal mol'1 with methylene chloride as the
solvent. The value reff = 5.6 Á was estimated based on the average distance from Ru
to the protons of Cp* (5.2 Á) and indenyl (6.0 Á to the furthest benzoid proton)
ligands.91 However, the AGel° value for Cp*FluRu based on the same crystallographic
radius for Cp*IndRu is 6 kcal mol'1 more exoergic than the experimental AAGsolv°
value. The oxidation of 11 is irreversible and E^ versus £,^(Cp2Fe+/,°) was used to
estimate a value for AAGsolv° which may explain the disparity of the values. Using
E^(ll) = 0.34 V measured by Gassman and Winter,91 a value for AAGsolv°(ll) in
agreement with the AGcl° is obtained. For the halogenated complexes, the Bom model
gives approximate estimates for solvation energetics. The ligands C5C15 and C5Br5
are expected to be structurally similar to a pentamethylcyclopentadienyl in terms of
steric bulk, reff = 5.1 Á for Cp*(C5Cl5)Ru (5.2 Á to the methyl proton and 5.0 Á to
the chloride)91 and AAGsolv° of Cp*(C5Cl5)Ru is -24 ± 3 kcal mol'1 and AGel° = -28
kcal mol'1.
An alternative to using the maximum radius of a molecule to estimate AAGel°
values, may be to use the molecular volume if the complex deviates significantly from
a sphere. By comparing the volumes of two equally charged species an approximate
ratio of relative AAGel° values may be assessed. For example, Cp*2Ru better qualifies
as a cylindrical molecule rather than a sphere. The Ru-Cp* distance is 1.8 Á and the
van der Waals radii of carbon is 1.85 Á resulting in a total molecular length of 7.3 Á,
compared to a molecular radius of 5.2 Á. Rough schematic representations

142
Figure 4-7 Structures of Cp*2Ru and ferrocene. The molecular representations help
to illustrate the relation of size to solvation energy. The differential
solvation energy of the Cp*2Ru+/0 couple is 16 kcal mol'1 less than
AAGsoiv(Cp2Fe)+'<’.

143
of the structures of Cp*2Ru and ferrocene are shown in Figure 4-7. The structures,
although not exact scale representations, demonstrate distinct differences in the
geometry of the two metal complexes and helps to illustrate the origin for the large
positive 5pc value. Estimates of the molecular volumes for Cp*2Ru and ferrocene are
approximately 530 Á3 and 260 Á3 respectively.72,124 The ratio of the molecular
volumes is consistent with the relative AAGsolv° values for the two complexes.
Although this is a crude approximation, it serves to demonstrate a possible source of
the disparate values in estimated AGcl° values.
Conclusions
Gas-phase electron-transfer equilibria have been used to determine free energies
of ionization of various ruthenocene derivatives at 350 K. The results provide an
accurate measure of the intrinsic electronic effects of Cp substituent effects on
ruthenocene oxidations. The trends in the ETE AG¡° values are consistent with
electrochemical potentials and Ru 3d core binding energies91 reported earlier;
however, the correlation fails in some cases. Most notably is the reversal of the AG¡°
values for Cp*2Ru and Cp*RuRu relative to the electrochemistry on the Ru 3d
binding energies.91 Therefore, the fluorenyl ligand does not stabilize the oxidation of
ruthenocene to extant that Cp* does. The stabilizing ability of Ru is only -60% of
the Cp* electronic effect.
Perhalogenated Cp* ligands do not stabilize ruthenocene towards oxidation to
the extent that is predicted by organic substituent parameters. The electronic effect of

144
the C5C15 is balanced by the electron-donating effects Cp* ligand, subsequently
AGi°(Cp*(C5Cl5)) ~ AGi°(Cp2Ru). The C5F5 ligand is only 1.5 more time electron-
withdrawing than Cp* is electron-donating. The moderate electron-withdrawing nature
of the C5F5 ligand is consistent with the findings of Paprott and Seppelt.115 From
equilibrium reactions, they determined that C5F5H was only slightly more acidic than
CpH but less acidic than CF3CH2OH.115 The small perfluorination effects can be
attributed to n resonance effects of the F lone pair electrons.5,115 For comparison, the
C(CF3)4H ligand is strongly electron-withdrawing which can be rationalized because
CF3 groups have electron-withdrawing qualities only.5
A Cp ancillary ligand parameter scale has been developed for the ligands
investigated here. The parameters, termed y and y* parameters depending on the
derivation used, should provide the basis for correlating electronic effects of Cp
ancillary ligands on the physical and chemical properties of other organometallic
complexes. Such parameters may provide insight concerning the electronic effects of
various Cp ligands on the activity and selectivity of homogeneous organometallic
catalysts bearing these ligands. The initial gas-phase studies of the rates of the
hydrogenolysis for several L2Zr-CH3+ cations were aimed at achieving this goal. The
application of the Cp ligand parameters towards predicting reactivity for such catalytic
systems, in addition to alkene polymerization and C-H activation, is of great interest.
The Cp ligand parameters have been derived from gas-phase experiments
absent of solvent effects. Application of these parameters in the condensed phase may
not parallel the gas-phase trends. Modifications of the condensed phase trends due to

145
solvent effects associated with the metal complexes and structural differences (ring
slippage, dimerization, or solvent coordination) must be considered.
Electrochemical studies for several of the ruthenocenes indicated that potentials
reported to be irreversible91 were found to be reversible or quasi-reversible under our
conditions outlined here. The £pa for Cp*2Ru was cathodically shifted relative to
£pa(Cp*FluRu) supporting the gas-phase trends in the AG¡°values. Differential
solvation energies for the complexes were consistent with AAGsolv° values derived for
other metallocenes. However, the AAGsolv° values for several of the ruthenocene are
not accurately predicted by the Bom model. This demonstrates that structural
deviations and charged localization may account for the disparity in the solvation
energetics for non-spherical metal complexes.
Experimental Methods
Details of the electron-transfer equilibrium technique were provided
earlier. ' ’ J Ruthenocene derivatives were sublimed from a solids insertion probe
to provide pressures in the 1 x 10’7 - 1 x 10"6 Torr range. For several of the
ruthenocene derivatives additional heating of the solids probe was necessary to
generate pressures of the neutrals in the 10 Torr range. Reference compounds were
admitted into the high vacuum chamber through a precision leak valve. Molecular
ions were generated through electron impact with beam voltages adjusted between 10 -
14 eV in order to generate ions in good yield and to minimize production of fragment

146
ions. Electron-transfer equilibrium was usually establish in ~ 1 s and was followed for
an additional 5-10 s.
Pressures were measured directly by using a nude ion gauge with a Granville
Phillips controller. Partial pressures were corrected for ion sensitivities and systematic
error with an MKS baratron capacitance manometer in the 1 x 10'5 Torr range.
Pressure calibrations varied only slightly within the series of ruthenocene derivatives.
Ruthenocene Derivatives and Reference Compounds
Organic reference compounds, ruthenocene, and ferrocene were purchased from
Aldrich Chemicals. Ethylferrocene, manganocene, and nickelocene were purchased
from Strem Chemicals. All ETE reference compounds were used without further
purification.
Compounds 2, 5, 10, 11, and 12 were prepared by using (r|5-
per*amethylcyclopentadienyl)ruthenium(IIT) chloride oligomer, (Cp*RuCl2)x, as
described by Gassman and Winter.91 All compounds were obtained in yields of 40%
or greater and characterized by !H NMR (CD3C1 solvent) and mass spectrometry.
Resublimation of 11 was necessary prior to mass spectral studies to remove trace
impurities of unreacted fluorenyl.
Many of the ruthenocene derivative studied by electrochemistry and mass
spectrometry were donated by several researchers. Samples of Cp*(C5(CF3)4H)Ru,
Cp*(C5(CF3)4OSiEt3)Ru, and Cp*(N02Cp)Ru were donated by Dr. Mark Burk of E. I.
du Pont de Nermours in Wilmington, Delaware.20 Samples of Cp*(TMSCp)Ru,

147
(TMSCp)2Ru, IndjRu, were donated by Dr. Allen Siedle from 3M Corporate Research
Laboratories in St. Paul, Minnesota. Additionally, Dr. Siedle also provided
purified samples of the zirconocene derivatives for the hydrogenolysis studies.
Professor Russell Hughes of Dartmouth College donated a sample of Cp*(C5F5)Ru.113
Professor Charles Winter of Wayne State University donated a sample of
Cp*(C5Br5)Ru.114
Cyclic Voltammetry Studies
Cyclic voltammetry studies were performed with a PAR systems (Models
173/175). A platinum button working electrode and a Ag/AgCl reference electrode
were used. Acetonitrile (Fisher HPLC grade) was purified by shaking the solvent with
CaH2 and filtering. The solvent was then distilled from P205 (5g/100 ml) onto CaH0
and then redistilled from calcium hydride immediately prior to use. Pure CH2C12
(Fisher HPLC grade) was obtained by shaking the solvent with concentrated sulfuric
acid, followed by an aqueous solution of Na^C^, dried with anhydrous CaCl2,
filtered over neutral alumina, and stored over P205, finally CH2C12 was obtained from
distillation with P205 immediately prior to use. The electrolyte Bu4NPF6 (Aldrich)
was re crystallized from ethanol/acetone three times, washed with dry ethanol and dried
in a vacuum oven at 100 °C for ca. 24 hr. Electrolytic solutions were freshly prepared
prior to all voltammetry studies. Tetrabutylammonium tetrakis[3,5-
bis(trifluoromethyl)phenyl]borate, TBA+TFPB\ was prepared and purified by the
method of Nishida and Kobayashi.125

CHAPTER 5
GAS-PHASE IONIZATION ENERGETICS, THERMOCHEMISTRY, AND
ELECTRON-TRANSFER KINETICS OF DECAMETHYLMETALLOCENES,
CHROMOCENE, AND COBALTOCENE
Introduction
Many studies concerning the reactivity of metallocene complexes involve
variations of the substituents on the r|5-cyclopentadienide ligand (Cp) and substitution
of ri5-pentamethylcyclopentadienide (Cp*) for Cp is one common ligand
variation. ’ The rj -pentamethylcyclopentadienide ligand leads to changes in
electronic effects and steric congestion with respect to ri5-cyclopentadienide. These
changes in ancillary ligand character in many cases lead to desirable alterations of the
structure, reactivity, and stability for transition metal complexes.105 Furthermore, the
development of a convenient synthesis126 and commercial availability of
pentamethylcyclopentadiene has lead to increased study of permethylated metallocene
complexes.
Electronic effects due to permethylation of Cp ligands of metallocenes can be
assessed by comparing the oxidation/reduction potentials of the Cp and Cp* metal
complexes and solution electrochemical potentials,34,107 and gas-phase vertical
ionization energies35,37 have been determined for a number of Cp* complexes.
However, as discussed earlier, if electrochemical oxidation/reduction potentials for an
148

149
organometallic complex are irreversible, the measured value will not be a reliable
indication of the thermodynamic potential. In addition, the intrinsic effect of
permethylation on the oxidation energeucs is modified by the differential solvation
energies of the redox couple. Vertical ionization energies, measured by photoelectron
spectroscopy are only accurate measure the adiabatic ionization potential, A//¡ 0°, if the
equilibrium geometries of the ion and the neutral are similar. * Large structural
differences between the neutral and the ion may lead to broad peaks, as observed for
OC 0*7
the decamethylmetallocenes, and an adiabatic ionization potential significantly
different from the measured vertical ionization potential.
Gas-phase electron-transfer equilibrium (ETE) techniques have been used
extensively to determine thermal free energies of ionization, AGj°, for metal
complexes16'20 and organic compounds.39-42 Application of ETE methods has been
previously reported for a number of metallocenes and substituted ferrocenes with AG¡°
values in the 6-7 eV range.16,19 Fourier transform ion cyclotron resonance mass
spectrometry,44-47 FTMS, was used to determine thermal free energies of ionization
for several decamethylmetallocenes, cobaltocene, and chromocene, which all have
AG° values in the 5-6 eV range. The AG ° values reponed are for the oxidation
process shown by equation 5-1, where L = Cp or Cp*. Free energies of ionization
were determined from the measured equilibrium constant for the ETE reaction and are
anchored to the estimated value of AG ° for bis(benzene)chromium(0), I^Cr,
L2M (g) -> l^M4- (g) + e
(5-1)

150
which was estimated by a combination of photoelectron spectroscopy results and
statistical mechanical analyses.
Through the application of thermochemical cycles,16,18'20,84 the heterolytic and
homolytic bond disruption enthalpies for M-Cp cleavage of Cp2Cr+/0 and Cp2Co+/0
have been derived. Additionally, by combining AG° values with reported data for
cobaltocene and chromocene and measured E^ values for the Cp*2M complexes,89,105
differential solvation energies, AAGsolv°, have been estimated. The Bom model for
estimating AAGsolv° values for M4^0 couples has been previously applied to
metallocene couples and other metal complexes.54 The AAGsolv° values for Cp2M+/,°
and Cp*2M+/0 couples are determined here through thermochemical cycles and the
values are compared to values estimated by the Bom model.
The rates of gas-phase electron-transfer reactions of several metallocenes
couples have been reported previously,14 and rate constants for the forward ET process
have been determined for several of the reactions in this study. For most reactions,
the approach to equilibrium was monitored and the forward and reverse rate constants
were determined. The forward electron-transfer reaction is written, as shown in
equation 5-2, so that the ETE reaction proceeds in an exoergic direction.
Cp*2M + R+ —> R + Cp*2M+ (5-2)
Overall, the forward rate constants, kf, show some dependence upon the driving force
of the electron-transfer reaction.

151
Gas-Phase Electron-Transfer Equilibrium Studies
Methods used for studying gas-phase electron-transfer equilibrium reactions
have been described.16'20,39'42 The general reaction studied is shown in equation 5-3,
where L^M denotes a metallocene and R is a reference compound with known a AG°
value. From the measured for the ETE reaction, the reaction free energy, AGct°,
L2M + R+ - R + L2M+ (5-3)
can be determined. Since AGet° is equal to AAG ° for the two compounds, AGiT° of
the metallocene can be derived provided AGiT° of R is known (equation 5-4).
AGel° = AGfiLjM) - AGj°(R) (5-4)
Because of the low values of the ionization potentials for the metallocenes
studied in this chapter, suitable organic reference compounds for ETE were
unavailable; organic reference compounds used in previous ETE studies have low AGj°
values approaching 140 kcal mol'1. Accordingly, AG¡° values reported here are
anchored to the AG ° of Bz2Cr = 125.6 ± 1.0 kcal mol'1 at 350 K. The ionization
potential of Bz^r was determined through high resolution photoelectron spectroscopy
and assumptions necessary to obtain AG¡ 35o0(Bz2Cr) from the vIP are described later.
Figure 5-1 is an equilibrium ladder exhibiting all ETE reactions investigated in
this chapter. All AG¡° values lie adjacent to the complex formula, and free energy

152
changes, AG°eU350, for specific reactions couples are adjacent to the arrows. Free
energies of ionization energies along with accompanying spectroscopic vIP data37,38
are presented in Table 5-1. Also in Table 5-1 are AAG¡° values corresponding to the
difference in AG¿°(Cp*2M) relative to AGj°(Cp2M). Error limits for the AGet° values
are estimated at 1.5 kcal mol'1, due largely to errors in the measured pressures of the
neutral reagent gases. Each reaction was repeated at least three times and examined
from both endoergic and exoergic directions to insure that the equilibrium constants
for the electron-transfer reactions were not dependent upon the direction of approach
to equilibrium. Several checks were performed to test internal consistency and in all
cases AGet° values were consistent within the established limits of ± 1.5 kcal mol'1.
Electrochemical Studies for Some Decamethylmetallocenes
The electrochemistry of several Cp*2M+7° couples (M = Mn, Fe, Ni, and Ru)
were examined. Table 5-2 lists the electrode potentials, of the
permethylmetallocenes as measured by cyclic voltammetry at a platinum button
working electrode and a KC1 saturated Ag/AgCl reference electrode in acetonitrile
containing 0.1 M t-butylammonium hexafluorophosphate, Bu4NPF6. The E^ of
Cp*2Ru has been previously reported in methylene chloride containing 0.1 M
t-butylammonium perchlorate,91 TBAP, and is included in Table 5-2 along with our
Ei¿ value measured in CH2C12 containing 0.1 M Bu4NPF6. All measured E^ values
are reported as reversible potentials and are consistent with literature values.

153
Cp2Mn
(ethylCp)2N¡
142.5
138.2
i
L
4.6
i
.
r
1.8
r
\ 0.8
137.9 Cp'2Ru
136.4 Cp‘2Os
123.5 Cp2Co
Figure 5-1 Electron-transfer equilibrium ladder for several decamethylmetallocenes.
Values for individual ETE reactions are adjacent to arrows and free
energies of ionization, AG¡° values, are adjacent to the complex
formulas. Free energy values are reported at 350 K. All values
anchored to the AG¡° value of Bz2Cr except Cp*2Os and Cp*2Ru which
are anchored to N,N-dimethylaniline.

154
Table 5-1. Ionization Energy Data for Some Metallocenes and
Decamethylmetallocenes.
Cp2M
A Gioa
vIP*
-A(AGi°)a’b
Cp*2Mn
121.6
122.9C
21.0
Cp*2Fe
126.7
135.6°
26.4
Cp*2Ni
121.2
134.2C
22.7
Cp*2Ru
137.9d
-
26.7
Cp*2Os
136.4
-
24.2
Cp2Cr
127.5
131.4e
Cp2Co
123.5
128.0°
a. Units are kcal mol-1.
b. Difference in the ionization energy of Cp*2M and Cp2M. Values of AGi°(Cp2M)
taken from reference 16.
c. See reference 37.
d. See reference 20.
e. See reference 35.

155
Table 5-2 Electrochemical E^ Data and Differential Solvation Free Energies of
Some Cp*2M+/0 and Cp2M+/0 Couples.
Cp2M+/0
Eft (solv)
F 15
AGj°(soln)c
-aag°so1vc
^Fc
Cp2Cr
-0.67d (CH3CN)
92
36
2
Cp2Co
-0.94d (CH3CN)
86
38
0
Cp*2Cr
-1.14e (CH3CN)
82
25f
13f
Cp*2Mn
-0.56s (CH3CN)
-0.64
93
28
6
Cp*2Fe
-0.12e (CH3CN)
-0.20
103
24
10
Cp*2Co
-1.47e (CH3CN)
74
24f
10f
Cp*2Ni
-0.65e (CH3CN)
-0.69
92
29
5
Cp*2Ru
(CH3CN)
0.41
116
22
12
Cp*2Ru
0.58h (CH2C12)
0.55
120
18
16
Cp*2Os
(CH2C12)
0.35
115
21
13
a. Values reported in volts using 0.1 M Bu4NBF4 as supporting electrolyte against
SCE, except chromocene and cobaltocene in 0.1 M Bu4NPF6 against SCE.
b. Eft values determined in this work were measured in 0.1 M Bu4NPF6 in CFL,Cl->
with ferrocene as an internal standard and then referenced against SCE.
c. Units are kcal mol'1. Estimated error limits ± 2 kcal mol'1.
d. See reference 89.
e. See reference 105.
f. Estimated from PES data and reported against SHE.
g. See reference 127.
h. See reference 91.

156
Potentials measured in methylene chloride were observed to be anodically shifted
relative to potentials measured in acetonitrile since CH2C12 is a waeker dielectric
medium. Values of were measured with ferrocene as an internal standard and
referenced against SCE and £,¿(Fc+/0).103
Bis(benzene)Chromium as a Reference Compound for
Electron-Transfer Equilibrium Investigations
Most AG ° values in this work are anchored to an estimated value for
AG¡°(Bz2Cr), as shown in Figure 5-1. In this section, the assumptions used to
estimate this anchor value are discussed. The photoelectron spectrum of Bz^r
(Figure 5-2) was measured for this study by Professor D. L. Lichtenberger and
coworker.
Vertical ionization energies measured by photoelectron spectroscopy are
approximately equal to adiabatic IP values if the difference between the ground state
geometries of the ion and the neutral is negligible.27 Figure 5-2 is a high resolution
photoelectron spectrum of bis(benzene)chromium in the valence ionization region
(5.0 - 7.5 eV). The very narrow first ionization band at 5.47 eV indicates that
structural difference of the ion ground state with respect to the neutral is very small.
Potential well diagrams for this process are represented by Figure 1-1B.
Ionization entropies for metal complexes have been estimated previously by
using statistical mechanics.16,74 The value of AS° for Cp2Fe+/0 calculated from
spectroscopic data is ~5 cal mol^K'1 at 350 K, and the positive value of AS° is
primarily due to changes in the vibrational and electronic entropies for the ferrocene

157
couple.16 Translational and rotational entropy changes were found to be only 0.1 cal
mol^K'1 at 350 K, and ASvib° accounted for nearly half of the total AS¡°. Sufficient
spectroscopic data for bis(benzene)chromium and its cation do not exist to allow for
detailed statistical mechanical analyses, however several important assumptions can be
made in estimating a value of ASi°(Bz2Cr). First, differences in the translational and
rotational entropies for Bz^r470 will be negligible since changes in the molecular
mass and moments of inertia for the molecule and the ion will be small.71 As inferred
from the photoelectron spectrum, changes in the Cr-Bz metal-ligand bond length are
minimal, therefore moments of inertia will not change significantly upon oxidation.
The statistical mechanics study of Cp2Fe+/0 couple revealed that vibrational frequency
shifts isolated at the Cp rings (i.e. ring skeletal vibrations) were minor components of
A5vib°, but shifts in metal-ligand vibrational frequencies (i.e. asymmetric and
symmetric Fe-Cp stretch and Fe-Cp bends) significantly contributed to ASvib°.16 Since
vibrational frequencies change little between Bz2Cr and Bz2Cr+,71 A5vib°(Bz2Cr+/0) <*
0 cal mol^K'1 can be inferred.
The most significant estimated contribution to AS^0 for Bz^r is A5elec°. The
ground state configurations for Bz^Cr and its cation are JAlg and 2Alg respectively.35
The electronic entropy change can then be determined from the change in electronic
spin degeneracy given by ASelec° = R(la g 5 where g = 2. Therefore, a value for
A5i°(Bz2Cr+/0) is estimated to be 1.4 cal mol^K'1 (AS ° ~ ASelec°) and at 350 K,
TAS ° = 0.5 kcal mol'1.

158
Ionization Energy (eV)
7 6
Figure 5-2
High resolution He (I) photoelectron spectrum of bis(benzene)chromium
in the valence ionization region. Spectrum recorded for this study by
Professor D. L. Lichtenberger and workers.

159
Changes in the integrated heat capacities for the Bz^Cr^couple are also
expected to be small. A statistical mechanical analysis of the change in the integrated
heat capacities for Cp2Fe+/0 couple indicates that ACp° contributions are negligible and
that AH} T° values from 300 - 600 K are expected to be equal to the 0° within ±
0.9 kcal mol'1.16 The difference in integrated heat capacities for benzene and the
benzene cation determined from statistical mechanics are also estimated to be
negligible (ca. 0.3 kcal mol'1 from 0 to 350 K), thus A//i0° closely approximates
A//¡ 350°. Similarly, the difference in the integrated heat capacities for the Bz2Cr+/0
couple, from 0 K to 350 K, are not expected to be much larger than 1 kcal mol’1 and
aLP = AH{ 0° - A//¡ 35q°. Finally, AG¡ 35Q0(Bz2Cr) = 125.6 ± 1.0 kcal mol'1 can be
derived from the estimated AH ° and AS° values at 350 K.
Free Energies of Ionization for Some Decamethylmetallocenes and
Comparison to Photoelectron Spectroscopy Results
As expected from previous studies, methylation of the cyclopentadienyl rings
lowers the ionization potential with respect to the parent metallocenes.35,37 In a
classical model, the polarizability and inductive effects of the methyl groups stabilize
the metallocene cation leading to a less endoergic ionization energy.2,5 Structural
evidence for the increased electron donating ability of pentamethylcyclopentadienyl
ligand is clearly observed by a comparison of ruthenocene and Cp*2Ru.91 Crystal
structures of Cp2Ru and Cp*2Ru indicate that the M-Cp distance is 0.08 Á larger for
ruthenocene than decamethylruthenocene.91 The smaller M-Cp* distance for Cp*2Ru
increased Cp*-Cp* repulsions, can be rationalized by more electron rich Cp* rings

160
being better donors to the Ru(II) metal center (i.e Cp* is a stronger Lewis base the
cyclopentadienyl).
Generally, AG ° values for the Cp*2M compounds are approximately 24 ± 3
kcal mol"1 lower than their Cp analogues (Table 5-1). Similarly, the values for the
oxidation potentials of the decamethyl derivatives (Table 5-2) are cathodically shifted
by 0.65 ±0.13V(15±3 kcal mol'1).105 The effect of alkyl substitution upon Cp2M
ionization energies is expected to be different for each metal due to several factors
including the energy of the metal-based HOMO and different electronic
degeneracies,35 but the relative constant value of the shift in the AG ° values upon
substitution of Cp* for Cp indicates that the differences are not large as when metal
center is varied.
Although the magnitudes of the effect of ten methyl groups on the AG¡° values
for the metallocenes are somewhat different, seen in the A(AG¡°) values in Table 5-1,
present results are consistent with previous AG ° values for alkylated metallocenes
(e.g. (C5H4R)CpFe and (C5H4R)2Ni)). Comparison of AGj° values for ferrocene16 and
nickelocene19 (Figure 5-3) and their Cp* analogues indicates that ferrocene ionization
energy is moderately more sensitive to alkylation than nickelocene. The derived p
values for the two plots shown in Figure 5-3 are 57 ± 1 kcal mol'1 and 49 ± 1 kcal
mol'1 for ferrocene16 and nickelocene19 respectively. Predicted AG ° values for
Cp*2Fe and Cp*2Ni based on previously published Taft analyses for ferrocene and
nickelocene alkylated derivatives are within ± 1 kcal mol'1 of the experimentally
determined values in this work. Plots of alkylnickelocene and alkylferrocene

Ionization Free Energy(kcal/mol)
161
Figure 5-3 Plots of AG¡° values for alkyl ferrocene and nickelocene complexes
versus Taft C[ parameters. Best fit lines are drawn for ferrocene (r =
0.999) and nickelocene (r = 0.999). Predicted values are within ± 1
kcal mol'1 of experimental results. Asterisk represents AG¡°(Cp*2M)
values predicted values based on Taft analyses of alkylmetallocene data
excluding the Cp* derivatives (i.e. data from Chapters 2 and 3)

162
derivatives with Taft ctj parameters,5,52,53 (Figure 5-3) show strong correlation
(R = 0.999 for both Ni and Fe). Therefore, although AAG¡°(Cp*2M/Cp2M) values for
M = Fe and Ni are different, alkyl parameterization schemes demonstrate that
AGi°(Cp*2M) values are consistent with ionization free energy data for other alkyl
substituents.
The vertical ionization energies from PES for Cp*2Mn, Cp*2Fe, and Cp*2Ni
have been determined by Cauletti et al. The assigned electronic configurations for
i n or o-t
Cp*2Fe ( Alg), and Cp*2Ni ( E2g) are the same as the simple metallocene. ’
However, Cp*2Mn exists as a low spin 2E2g complex (alg2e2g3)37 rather than a high-
spin 6Alg complex as observed for manganocene.79 The PES of Cp2Mn is extremely
broad do to significant geometry changes for cation with respect to the neutral.37 In
comparison, the PES of Cp*2Mn is substantially different from manganocene and the
first ionization band is quite narrow relative to manganocene.37 The AGj° value in
this work is in agreement with the estimated adiabatic ionization potential (122
kcal mol'1) determined from the onset of the first PES manifold.
The first ionization bands in the Cp*2Fe and Cp*2Ni photoelectron spectra
have broader bands than are observed for the parent metallocenes and as a result,
relaxation energies, determiend from equation 5-5, are ~12 kcal mol'1 for these
complexes.35,37
Et = vIP - AG¡°
(5-5)

163
However, the relaxation energy for Cp*2Mn is only 2 kcal mol'1 which is considerably
lower than £r(Cp2Mn) =16.8 kcal mol'1.
The AG° value for Cp*2Ru has been discussed in Chapter 4. The AG° for
decamethylosmocene is 136.4 ± 2 kcal mol'1 which is 24.2 kcal mol'1 smaller than the
AGj° value for Cp2Os. The trend in the AG¡° values for the Cp*2M complexes, where
M = Fe, Ru, Os, is consistent with the ordering of the AG ° values for the parent
complexes.16
Free Energies of Ionization for Chromocene and Cobaltocene
Because of their low IP values, we have been unable to report AG ° values of
cobaltocene and chromocene until now. Estimates of thermochemical properties of the
ions in previous publications were based on vertical ionization energies from PES.16
The AGj° values shown in Figure 5-1 are in agreement with vertical ionization data (£r
1 7*7 70
= 5 kcal moT ). ’ Cauletti et al. measured the vertical ionization energies by
photoelectron spectroscopy of both chromocene and cobaltocene and the values are
0*7
5.71 ±0.1 and 5.56 ±0.1 eV respectively. The difference in the vertical ionization
energies and the AGj° values for the two compounds is consistent for the two studies.
The ground state configuration for chromocene is 3Elg and the low value of AG° with
respect to the ferrocene can attributed to the removal an essentially non-bonding
unpaired electron from a 16 e' open-shelled compound.37,38 Since the molecular
orbital scheme for cobaltocene (a^e^e^1) has one unpaired electron residing in an

164
antibonding elg orbital, the low ionization potential relative to other metallocenes is
easily rationalized.35,37,38
Bond Disruption Enthalpies for Chromocene and Cobaltocene
Thermochemical cycles that incorporate AG ° data have been used in previous
chapters to derive bond disruption enthalpies and differential solvation free energies.
The assumptions required to combine free energy of ionization data referenced at 350
K with A//vap 29g° for metal sublimations8 and A//¡ 0° bare metal ionization energies8
have been discussed earlier and will not be described here. The value for AH° in
equations 5-6 and 5-7 represents twice the average homolytic, A//hom°, (equation 6)
and heterolytic, A//het°, (equation 7) bond disruption enthalpies for metallocenes and
metallocenium ions.
Cp2M(g) -> 2Cp(g) + M°(g)
(5-6a)
Cp2M+(g) 2Cp(g) + M+(g)
(5-6b)
Cp2M(g) -4 2Cp(g) + M2+(g)
(5-7a)
Cp2M+(g) -> 2Cp-(g) + M3+(g)
(5-7b)
From the heats of formation for the respective ions and neutrals (estimated from
vertical ionization energies) and A//f° data for Cp2Cr and Cp2Co, average values for
A//het° and A//hom° were derived and are given in Table 5-3. Error limits for
heterolytic M-Cp bond disruptions enthalpies are larger than homolytic disruption

165
Table 5-3 Average Bond Disruption Enthalpies for Chromocene and Cobaltocene.
A/4et°(Cp2M3+y
A//„e,°(Cp2M2+)a
^hom°(Cp2M+)a
^hon,°(Cp2M0)»
Cp2Cr 647 ± 5
354 ± 4
91 ± 3
11 ±2
Cp2Co 646 ± 5
394 ± 4
101 ± 3
12 ±2
a. Units are kcal mol'1. Auxiliary data used to determine bond disruption enthalpies
were taken from references 8 and 27.

166
enthalpies due to the inclusion of extra thermochemical data necessary to estimate the
AHhc° quantities.
Values for A//het° and A//hom° for chromocene and cobaltocene are consistent
with other first transition row metallocenes. Although AG ° values for chromocene
and cobaltocene are much lower than other first transition row metallocenes, bond
disruption enthalpies are not significantly different. In the thermochemical cycles used
here, bond disruption enthalpies are dependent upon ionization energetics data for both
the complexes and the corresponding bare metal ions. For example, A//het°(Cp2Co+)
is 53 kcal mol'1 larger than A//het°(Cp2Fe+) primarily because A//f°[Co(3+)(g)] >
A//f°[Fe(3+)(g)] by 85 kcal mol'1.8 In terms of the thermochemical cycles, AH °
values for the bare ions compensate for the low AGj° values for Cp2Cr and Cp2Co
relative to the other metallocenes.
Estimates of A//hel° and AHhom° for the decamethylmetallocenes could not be
made due to lack of thermochemical data for the heats of formation of the Cp*2M
complexes and pentamethylcyclopentadienide (Cp*).
Evaluation of the Solvation Energetics for Decamethylmetallocenes
Chromocene, and Cobaltocene
Differential solvatipn energies, AAGsolv° for metallocene redox couples were
determined from comparison of AG¡°(g) values with AG^Csoln) values measured at
298 K. Assumptions necessary for combining free energies and AG¡°(soln) values
have been discussed previously.16’19’20,54,84 Values of for the metallocenes and
decamethylmetallocenes are used to derive AG°(soln) values as described in previous

167
chapters. Corrections for liquid junction potentials have not been made for values.
Entropy and heat capacity changes from 298 to 350 K are small for the metallocene
redox couples and will only lead to minor errors.54,84 The stationary electron
convention is assumed for both gas-phase and solution redox couples, although near
298 K both the stationary and thermal electron convention yield essentially the same
AG ° values.88
Data used to calculate AG¡°(soln) and AAGsolv° values are shown in Table 5-2.
The literature E^ values for the decamethyl complexes were obtained from several
sources91,105,127 and different reaction conditions, therefore we redetermined the £j¿
values for a more consistent comparison of Cp2*Mn+/0, Cp*2Fe+/0, Cp*2Ni+/,°, and
Cp*2Ru+/0, electrode potentials under uniform conditions (Table 5-2). Differential
solvation free energies for the decamethyl complexes were then determined using our
Evalues. Data for the E^ of Cp*2Ru measured in CH2C12 is included in Table 5-2
for comparison along with estimates of AAGsolv° for chromium and cobalt decamethyl-
metallocene derivatives derived from vIP data.35
Generally, the AAGsolv° values for Cp*2M+/0 couples are 12 ± 2 kcal mol'1
smaller than the parent Cp2M+/0 couples. This difference is of course the origin of the
differences in the shifts in electrode potentials and the gas-phase ionization energies
for the decamethyl derivatives relative to their parent metallocenes (15 ± 3 kcal mol'1
versus 24 ± 3 kcal mol'1, respectively). Figure 5-4 is a plot comparing gas-phase and
solution AG¡° values. Estimates of AG°(g) for Cp*2Cr and Cp*2Co were obtained
from PES dataJJ and AG¡°(soln) values were determined from potentials measured

Free Energy of Ionization (kcal/mol)
168
Figure 5-4 Plot demonstrating periodic trend of ionization energies for the first
transition row decamethylmetallocenes. Gas-phase data include AG°
values determined in this work (filled squares) and AG ° values
estimated from PES results (triangles). Solution AG¡° data were
determined from Evalues measured in CH3CN/O.I M Bu4NPF6 except
for Cp*2Cr and Cp*2Co taken from reference 105.

169
by Robbins et al.105 As was observed for the first transition row metallocenes, the
trends in both the gas-phase data and solution are very similar.
Included in Table 5-2, are estimates of differential solvation energies relative to
AAGSoiv°(Cp2^e+/(^- The electrochemistry of ferrocene is reversible in a wide variety
solvent systems and values for £o34’103 aG¡°16 for ferrocene have been wen
established resulting in a reliable estimate for AAGsolv°(Cp2Fe+/0) in CH3CN. In light
of this, relative AAGsolv° values can be estimated for complexes that have reversible or
quasi-reversible electrochemical responses in limited electrochemical conditions.
Additionally, the formal potentials relative to Ev4(Cp2Fe+/(>) are constant in various
conditions,103 therefore relative AAGsolv° values relative to ferrocene may be
determined in several solvents. The term 5pc is used in place A(AAGsolv°) to denote
solvation energetics relative to ferrocene (equation 5-8). Positive values of 5pc
indicate that solvent stabilization of an M+/0 couple is decreased relative to Cp2Fe+/0.
Spc = MGsoiv°(M+/0) - AACsolv°(Cp2Fe+/0) (5-8)
A negative 6pc value indicates that solvation energetics for a reaction couple are
increased (more exoergic) relative to the ferrocene/ferrocenium couple either due to
poor solvation of the complex neutral or a strongly solvated complex ion.
For the decamethylmetallocenes, 6pc values are all positive. The size and
lipophilicity of the methyl groups decreases the solvation of the Cp*2M+/0 couple
compared to the parent metallocenes. Further, as charge-to-solvent separation

170
increases, polarization of the surrounding solvent will decrease. The methyl groups
function as a solvent sphere around the parent metallocene complex and the second
solvent layer (the solvent itself) stabilizes the cation to a lesser extent than the methyl
groups. For example, the A(AG°) for Cp2Fe and Cp*2Fe is -26 kcal mol'1 and
AAGSoiv0(Cp*2Fe) is -24 kcal mol'1. The effect is more apparent for ruthenocene
where A(AG¡°) = -27 kcal mol'1 and AAGsolv°(Cp*2Ru+/0) in the lower dielectric
methylene chloride is only -17 kcal mol'1 .
The Bom equation (equation 5-9) has been used successfully for predicting ion
solvation energies for Cp2M+/0 couples16,19,54 but apparently does not predict AGel°
values accurately for Cp*2M complexes. The Bom equation is used to estimate the
AGel°= (-166z2/reff)( 1-1/D) kcal mol'1 (5-9)
change in electrostatic free energy (AGel° = AAGsolv°) when a charge z on a sphere of
radius reff is transferred from a vacuum to a sphere of equivalent radius in medium of
dielectric constant D.54 The parent metallocenes have quasi-spherical, compact
molecular structures and the decamethyl derivatives are less spherical than their Cp2M
analogues. For example, the maximum molecular radius for decamethylferrocene is
estimated to be ~5 Á,128 and the estimated AGel° = -32 kcal mol'1 (equation 5-9)
which can be compared to -24 kcal mol'1 estimated from experimental data.
Alternatively, from the estimated AAGsolv°(Cp*2Fe+/0) = -24 kcal mol'1, the Bom
model yields a thermochemical radius of 7.3 Á for Cp*2Fe.

171
The AAGsolv° values for Cp2Cr and Cp2Co are in agreement with other first
transition row metallocenes.16 A comparison of AAGsolv° indicates that a common
value of -37 ± 2 kcal mol'1 is obtained for Cr, Fe, Co, Ni, and Ru metallocenes.
Electron-Transfer Kinetics
Forward electron-transfer reaction rate constants from ETE studies were
determined from the approach to equilibrium. Electron-transfer reaction rate constants
were measured in order to test the consistency of determined ETE reaction free
energies, AGet°, and establish that random errors in our experiment do not exist.
Several electron-transfer reactions were observed in which ETE data could not be
obtained because the approach to equilibrium was hampered by inefficient reaction
rates compared to ion loss from the ion trap. Electron-transfer reaction rate constants
for selected reaction couples are presented in Table 5-4. A detailed discussion for
determining second-order rate constants for gas-phase electron-transfer reactions has
been given elsewhere,14 and similar methods have been used here.
The barrier for electron-transfer process involving decamethylmetallocenes is
expected to be similar to that for simple metallocenes. Nevertheless, reaction rates
involving the decamethylmetallocenes were nearly an order of magnitude slower than
for the metallocenes14 and reaction efficiencies were only 2-5% (an estimate of
kL, the Langevin collision rate constant, is 1.1 x 10"9 cm3 molec'V1 for the Cp*2M
complexes).

Table 5-4
Electron-Transfer Rate Constants for Some Metallocenes,
UM/ + L2Mb -> L.M, + L2Mb+
172
L,Ma
AGetoa
1
Cp*2Mn
Cp*2Ni
2.5 x 10'12
0.002
2
Cp2Cr
Cp*2Fe
0.8
6.7 x 10'11
0.06
3
Cp*2Fe
Bz^Cr
1.1
5.1 x 10'11
0.05
4
Cp2Co
Cp*2Mn
2.0
3.6 x 10'11
0.03
5
B^Cr
Cp*2Mn
~4
6.7 x 10'11
0.06
6
B^Cr
Cp*2Ni
4.4
6.5 x 10'11
0.06
7
Cp2Cr
Cp*2Ni
~7
7.4 x 10'11
0.07
8
Cp*2Fe
Cp*2Ni
~6
4.6 x 10'11
0.04
9
B^Cr
Cp2Co
2.1
5.2 x 10'10
0.52d
10
Cp*2Ni
Cp*2Ni
0
5.3 x 10'10
0.48
11
Cp2Ru
Cp*2Ni
43
3.8 x 10'10
0.35
a. Units are kcal mol'1.
b. Units are cm3 molec*1 s’1.
c. kL = 1.1 x 10'9 cm3 molec'1 s'1 unless noted otherwise. See reference 14.
d. kL estimated to be 1.0 x 10'9 cm3 molec'1 s'1. See reference 14.

173
Electron transfer reaction involving Cp*2Ni were very slow and in most cases,
equilibrium reactions needed to be followed for >10 s. For the Cp*2Ni/Cp*2Mn
couple, electron-transfer equilibrium was never observed even with reaction times of
15-20 s. It was observed that the two ions were decaying at different reaction rates
which is indicative of a slow electron-transfer process.61,80 The estimated driving
force for the reaction (~0 kcal mol'1) may be too low to promote efficient electron-
transfer. The electron-transfer reaction for Cp*2Ni+/0 was studied and was anticipated
to be slow in view of the cross reaction involving the Cp*2Ni+/0 couple. The
theoretical maximum rate constant for a metallocene self-exchange or thermoneutral
reaction given is 5 x 10'10 cm3 molec'V1 corresponding to kJkL = 0.5.80 However,
the self-exchange rate constant for Cp*2Ni Iq = 5.3 ± 1.0 x 10'10 cm3 molec'V1. For
comparison The reported self-exchange reaction for Cp*2Mn is 3.1 x 10" cm
molec'V1, eff = 0.28.14 For the case of the Cp*2Ni/Cp*2Fe couple, the rate constant
is an order of magnitude larger than the Ni/Mn couple. The driving force for the
Ni/Fe couple, reaction 8 in Table 5-4, is ca. -6 kcal mol"1, indicating that the some
driving force dependence on ET reaction efficiency. The driving force for the
electron-transfer reaction for Cp*2Ni with ruthenocene is -43 kcal mol'1, well above
the estimated activation barrier for the ET reaction, and the measured rate constant for
the reaction (kf = 3.8 ± 1.0 x 10"10 cm3 molec'V1) is in accord with other ET rate
constants involving ruthenocene.14 It is therefore possible that the anomalously fast
self-exchange for Cp*2Ni+/0 is in fact proceeding through an alternative reaction

174
pathway for exchange of the isotope label used to monitor the reaction (ligand
exchange, for example).
Conclusions
Free energies of ionization for Cp*2Mn, Cp*2Fe, Cp*2Ni, Cp*2Ru and Cp*2Os
determined from gas-phase electron-transfer equilibrium techniques were found to be
consistent with previous ETE studies photoelectron spectroscopy results. Additionally,
AG¡° values for Cp*2Fe and Cp*2Ni were found to be consistent with AG¡° data for
ir 10
other alkylated metallocene complexes. The AG ° value for manganocene is only
2 kcal mol"1 lower than the its reported vIP35 indicating that the Cp*2Mn+/0 couple is
structurally more similar than the Cp2Mn+/0 couple. Larger relaxation energies were
found for Cp*2Ni and Cp*2Fe.
Differential solvation free energies for chromocene and cobaltocene are in
agreement with other first transition row metallocenes.16 Thus, solvent interactions for
the first transition row metallocenes are expected to be dependent upon complex size
as inferred from the Bom model. An average value for AAGsolv° = 26 ± 3 kcal mol"1
was found for the decamethylmetallocenes, excluding ruthenocene. However, the Bom
model predicts values of AGel° that are larger than found experimentally.
Electron-transfer reaction rate constants involving decamethyl complexes
proceed at ~2-5% of the Langevin collision rate. Formation of a transition-state
complex with inappropriate orbital overlap may inhibit electron transfer. Although the
self-exchange reaction for Cp*2Ni was not inefficient, the electron-transfer reactions of

175
Cp*-,Ni with other metallocenes was very slow unless the driving force of the reaction
was substantial.
Experimental Methods
Electron-Transfer Equilibrium Investigations
Electron-transfer equilibrium studies were performed by using Fourier
transform ion cyclotron resonance mass spectrometry. The spectrometer utilized in the
present studies is equipped with a 3T superconducting magnet and controlled by an
Ionspec data station. Details of the mass spectrometer and electron-transfer
experimental procedures have been described previously.
Bis(benzene)chromium, cobaltocene, and chromocene were sublimed into the
FTMS high vacuum chamber through a heated Varian precision leak valve.
Decamethylmetallocenes were introduced into the vacuum chamber by using a heated
solids probe positioned adjacent to the reaction cell.80 When the ETE of two Cp*2M
complexes was investigated, the metallocene with the lower AG¡° (therefore requiring a
lower neutral vapor pressure for equilibrium studies) was introduced through a heated
leak valve and the second compound was sublimed from a heated solids probe.
Typical reaction pressures were between 10 and 10 Torr for all ETE studies.
The temperature of the reaction cell was 350 K as measured by an Omega
RTD thin film detector. Ions were produced through electron impact at 9-12 eV with
a 20-30 ms beam event. Ions were thermalized through ion-molecule collisions (~30
collisions s'1) prior to detection. Through the use of ion ejections, fragment ions could

176
be removed from the cell before the ET reaction. Prior to the reaction period, one of
the parent ions was ejected from the reaction cell and the population change of both
parent ions with time was observed at set time intervals.
Partial pressures of the parent neutrals were measured directly with an ion
gauge and then calibrated by using a Baratron capacitance manometer in the 10'5 Ton-
range. Pressure correction were extrapolated to experimental conditions in the 10'7-
KT6 Torr range. Details concerning partial pressure measurements and pressure
calibrations for our 3T FTMS system have been discussed in detail elsewhere.94
For kinetic studies, pressure corrections and normalization for all molecular
isotopes was performed prior to calculation of reaction rate constants. Corrections to
account for diffusive loss of ions from the reaction cell were not performed. A home¬
made reaction cell with cell dimensions of 1.88" x 1.88" x 3.00" was used for the
majority of ETE and kinetic experiments. Ions could be trapped for -5 s without
signr ant loss of ion signal. Additionally, since more ions can be stored in the larger
cell, reactions could be followed for longer reaction times thus distinguishing from a
slow electron-transfer reaction or an electron-transfer equilibrium.
Cyclic Voltammetry Studies
Cyclic voltammetry studies were performed with a PAR systems (Models 173/175). A
platinum button working electrode and a Ag/AgCl reference electrode were used.
Solvents used for electrochemistry were stored over molecular sieves for several days
prior to purification. Pure dry acetonitrile and methylene chloride were prepared prior

177
to electrochemical studies. Electrolytes were crystallized several times from
ethanol/acetone mixtures and dried in a vacuum oven. Electrolytic solutions were
freshly prepared prior to all voltammetry studies.
Compounds
Decamethylnickelocene, decamethylmanganocene, decamethylosmocene,
chromocene, cobaltocene and bis(benzene)chromium were purchased from Strem
Chemicals and used without further purification except bis(benzene)chromium which
was resublimed prior to use. Decamethylnickelocene is thermally unstable and after
prolonged storage (months) in an inert atmosphere required recrystallization.
Decamethylferrocene129 and decamethylruthenocene91 were prepared according to
literature procedures and their purity was evaluated by mass spectrometry and *H
NMR.

CHAPTER 6
OVERVIEW OF EXPERIMENTAL METHODS AND PROCEDURES
The procedures for studying electron-transfer equilibrium reactions by using
FTMS have been described in the previous chapters. A general background to the
principles and design of the Fourier transform mass spectrometer, in addition to further
experimental considerations, is given in this chapter. Also, the utility of FTMS for
investigating the thermodynamics and reactivity of inorganic complexes is discussed.
Fourier Transform Ion Cyclotron Resonance Mass Spectrometry
The fundamental principles of Fourier transform ion cyclotron resonance mass
spectrometry, FTMS, are described by the elementary laws of electromagnetism.45 In
a homogeneous magnetic field, with magnetic field strength B, an ion of charge q and
mass m will be subject to a force perpendicular to the direction of the ion motion.
The magnetic force, called the Lorentz force, acting on the charged particle is
described by equation 6-1, where v represents the ion velocity.
F = q(v x B) (6-1)
The Lorentz force confines the ion to a helical path, with a radius r proportional to the
velocity of the ion.45 The ion rotates at a characteristic frequency called the
178

179
cyclotron frequency, given by equation 6-2, such that each ion with a unique mass-to-
charge ratio, mlz, will have a characteristic cyclotron frequency. Determination of the
cyclotron frequency for an ion leads directly to its mlz value.
to = qB/2nm (6-2)
Through the use of excitation and detection operations, different masses for an
ensemble of ions may be determined.44'47 The measurement of co of an ion in order
to determine mlz central to Fourier transform ion cyclotron mass spectrometry.45
Trapping Ions for Mass Analysis
To prevent ions from traveling along a helical path and being lost prior to
detection, ions are produced between two trapping plates. A schematic representation
of the sample region, the ion trap is shown in Figure 6-1. The trapping plates are
located at the ends of the ion trap. An electrical potential is applied, perpendicular to
the magnetic field, to the trapping plates which are in the XY-plane of the ions’
helical motions. For positive ions, the trapping plates typically carry a +1 V potential
and for negative ions, the trapping plates typically carry a -1 V potential. The
trapping potential causes ions moving toward the end of the ion trap to be repelled
back into the center of the ion trap. Thus, ions are constrained in the XY-plane by the
magnetic field and along the z-axis by the trapping plates.

180
Trap Rate
Figure 6-1
Orthorhombic ion trap used in a Fourier transform ion cyclotron
resonance mass spectrometer

181
Ion Excitation and Detection
The ion trap is the region of the mass spectrometer in which all ion
manipulations occur. As shown in Figure 6-1, the trap plates are perpendicular to the
excitation and detection plates. Two excitation plate and two detection plates are
required for ion detection. The two sets of excitation and detection plates (Figure 6-1)
are positioned parallel to the magnetic field.
All ions are produced at random time intervals and have random velocities,
even ions with the same m/z value.44 Ions in a random phase can not produce a
macroscopic signal for detection since their incoherent motion will average-out to zero
net signal.44 In order to detect the ions, an external energy pulse, in the radio¬
frequency range, is applied at the excitation plates. The radio-frequency, rf, pulse is
swept throughout the range of all cyclotron frequencies of interest.45 Ions absorb the
energy from the oscillating electric field and spiral outward into larger orbits.45 The
absorbed kinetic energy of the ions transferred from a radio-frequency pulse is related
to the cyclotron radius by equation 6-3. The process of transferring energy from an
K.E. = IntKtí^r1 (6-3)
external source to ion in an orbiting path is called ion cyclotron resonance.45
As the ions absorb external energy from the applied rf-pulse, all ions of a given
mlz will from a coherent ion packet with a unique cyclotron frequency46 (equation
6-2). The orbit of the frequency is larger than the initial orbit. The ion packet will

182
induce an image current on the detect plates and the image current of each individual
ion with a unique cyclotron frequency is measured by the detect plates. The image
current is converted to a voltage, amplified and stored in a computer for Fourier
transformations of the time-domain signal into a frequency signal. The sequence of
exciting ions into ion packets and then detecting their image current can be repeated
numerous times to generate an average signal.
Ion Formation and Mass Selection
Ions are formed by electron impact from an electron beam. Typical beam
voltages for experiments in this work were from 9 to 15 V. Ion formation is also
controlled by the duration and current of the electron beam. Beam currents are set by
a potentiometer and are measured by a collector at the end of the ion trap opposed to
the electron beam. The beam event varied for these studies, however a average beam
time of ca. 20 ms is sufficient to ionize metallocene complexes in large ion
concentrations.
During the beam event, ions are produced with a range of mlz values, i.e.
unwanted molecular and fragment ions. It is therefore necessary to isolate the ions of
interest in order to unambiguously follow ETE reactions or ion-molecule reactions.
All unwanted ions are removed from the ion trap by ion ejections. A high energy rf-
pulse is applied to the ion of interest causing the ion to spiral outward and crash into
the cell plates.46 By means of several ion ejection pulses, specific ions can be
isolated. Ion ejection pulses require adjusting the energy of the ejection pulse, the

183
pulse duration and the mass window for ejection (the sweep width). Ejection energies
ranged from 10 to 20 volts. Ion ejection times of ~10 ms were used for most
ejections, however longer pulse widths were required for broad mass windows.
For electron-transfer equilibrium studies, the two molecular ions of interest are
isolated by a series of ejections. Then an additional ion ejection is set for one of the
ions of the ETE couple. In this way, the growth of a reactant ion to the other reactant
ion can be monitored over time. For example, if the reaction of interest is A + B+ =
B + A+ and ion A+ is ejected, the growth of A+ is monitored relative to the
concentration of ion B+ over time until equilibrium of the ion populations of A+/B+ is
observed (see Figure 2-1).
Measurement of Equilibrium Constants
An advantage of monitoring electron-transfer equilibrium reactions by using
FTMS, aside the obvious ability to trap ions for long periods of time, is that the
coefficients needed to determine an equilibrium constant are all measured directly.
For comparison, partial pressures in a high pressure mass spectrometer are determined
from concentrations of the neutral reagents.40"43 Direct pressure measurements for
FTMS are made with an uncorrected ion gauge and are corrected for sensitivity and
system factors after ETE investigations
Partial pressures are measured remote to the ion trap by an ion gauge as shown
in Figure 6-2. During ETE studies, the HVC diffusion pump is throttled back to ca.
25% of full pumping capacity. This allows for a dynamic pressure equilibrium to be

184
established throughout the vacuum chamber. However, there may exist a pressure
differential in the vacuum chamber resulting in a disparate pressure at the ion trap
relative to the pressure at the ion gauge, i.e. systematic error. Also, the ion gauge
may give biased pressure readings for different neutral gases (ion sensitivities). Ion
sensitivities of the ion gauge are dependent on the ionization potential and the size of
the neutral. System factors are dependent on the rate of diffusion of the neutrals
through the FTMS high vacuum chamber (HVC). Pressure corrections therefore entail
corrections of both factors. Pressures are calibrated by comparing ion gauge pressures
to pressures measured at a Baratron capacitance manometer in the 10'5 Torr range and
extrapolated back to experimental conditions. A Baratron capacitance manometer is
used to correct ion gauge pressures because of the inherent biased of the ion gauge.
The Baratron is insensitive to ionization cross sections and therefore is thought to give
unbiased pressure readings. Details concerning the pressure corrections for
metallocene complexes and aniline derivatives have been presented
previously.13,1A 16,94
For the metallocenes and metallocene derivatives, ion gauge sensitivities are
essentially the same within ± 20%. System factors, the rate of flow (or resistance to
flow) have been measured for several metallocenes. System factors are also consistent
for the metallocenes.13,16 For example, the system factors for both ferrocene and
nickelocene are 1.6 ± 0.20 (dimensionless factor, corrected for ion gauge sensitivity),
meaning the pressure at the ion trap is ~1.6 times higher than the pressure at the ion

185
Figure 6-2 Schematic representation of a Fourier transform ion cyclotron resonance
mass spectrometer.

186
gauge during experimental conditions. For comparison, the system factor for N,N’-
diethlytoluidine is 1.4 ± 0.1.
Sensitivities for the organometallic complexes studied in this work do not vary
significantly. The average sensitivity for the Cp*2M complexes, where M = Mn, Fe,
Ni, and Ru, is 1.2 ± 0.2. Similarly, the average sensitivity for Cp2Fe and several
ferrocene derivatives is 1.3 ± 0.2. Generally, the ion sensitivities for the metallocene
complexes have been found to be equivalent. Furthermore, the ion sensitivities of the
organic reference compounds are also equivalent to the metallocenes. The sensitivity
of azulene is 1.2 ± 0.1 and the sensitivity of DET is 1.1 ±0.1. Generally, the overall
pressure differential at the ion gauge and the ion trap for the metal complexes and the
organic reference compounds are equivalent.
Corrections for ion sensitivities involve normalization for all isotopes of the
ions involved in the ETE work. If only one isotope (the most abundant isotope, for
example) is monitored during the course of an ETE study, the balance of the minor
isotopes must be accounted. Isotopic corrections were performed by calculating the
isotopic distribution of an ion and then adjusting the measured ion intensities of the
ions followed during the ETE experiment to include 100% of the natural isotopic
distribution. It is important to note that partial pressure and ion intensity ratios are
used for the determination of equilibrium constants. Therefore, any systematic errors
(assuming such errors are comparable for both species involved) will cancel in the
determination of an equilibrium constant.

187
Temperature Dependence Studies
The vast majority of temperature dependence studies of electron-transfer
equilibrium reactions have been performed by using PHPMS.39-43 For FTMS, the
technique has been scarcely been applied,16'39 and most FTMS ETE data have been
acquired at 350 k.13,15’16,19’20’39’59 In order to assess the validity of investigating the
temperature dependence of ETE reactions by using FTMS, and thus determine reaction
enthalpies and entropies, a model reaction was initially investigated. The required
model reaction must have well-defined thermodynamic parameters that can be
critically compared to experimental results and must have an electron-transfer
equilibrium constant that is accessible by FTMS (< 100).60 The equilibrium reaction
6-4 was investigated as a function of temperature from 350 to 500 K.
CO+ + Kr = Kr+ + CO (6-4)
The temperature of the ion trap was heated by a custom built heater described in
Chapter 2. In addition to heating the ion trap, the vacuum chamber was heated in
order to minimize radiative cooling of the ion trap. The procedure described in
Chapter 2 was developed from the methods that evolved from the study of the CO/Kr
reaction couple.
For the model reaction, the heat capacities and entropies of all species were
calculated by statistical mechanics55 from known spectroscopic data for the CO+^° and
Kr+/0 couples in the gas-phase.7 A van’t Hoff plot for the electron-transfer

188
equilibrium CO/Kr couple is shown in Figure 6-3. The experimental derived
parameters are A//et° = +0.3 ± 0.8 kcal mol'1 and ASet° = +3.1 ± 1.4 cal mol'1 K'1
reported at the 95% confidence limit. The thermodynamic parameters derived from
the statistical mechanical analyses are A//et° = -0.23 ± 0.01 kcal mol"1 and ASet° =
+1.42 ± 0.06 cal mol'1 K'1. The enthalpy change is within the experimental limits for
ETE studies which is usually ±1.5 kcal mol'1.16'20’39"43 The error in the entropies is
probably due to errors in the measured partial pressures of the neutral gases. Given
the error limits for reponed reaction entropies (ca. ± 50%)42 determined from
temperature dependent ETE studies, the overall error in the experimental and
theoretical ASc° values is acceptable. Thus, the application of FTMS to determine
reaction entropies and enthalpies is validated through the relative success of the CO/Kr
ETE couple.
Application of FTMS for the Study of Metal Complexes
There are several advantages of using FTMS for investigations of
organometallic complexes. The advantage of low background pressures in the HVC
region makes it possible to sample complexes with low volatilities, which includes
most of the metallocene complexes studied in this work. Background pressures in the
mass spectrometer used in these studies were typically less than 2 x 10'8 Torr.
However, even at these low pressures, it is still possible to generate ion signals. Ion
signals have been observed with background pressures in the 10'11 Torr regime.47

Ln (Keq)
189
1000/T
Figure 6-3 van’t Hoff plot of the CO/Kr electron-transfer equilibrium reaction.

190
The melting point of Cp*2Ni is ca. 300 °C at STP,105 however the sample was readily
introduced into the mass spectrometer by a heated solids probe with T^obe ~ 80 °C.
Thus, the need to heat samples to 200-300 °C (which often leads to decomposition of
the complex) in order to achieve significant vapor pressures can be avoided.
Therefore, the low working pressure in the FTMS vacuum region facilitates the study
of inorganics and organometallics that have relatively small vapor pressures.
Further, only small amounts of sample are required for an experiment because
of the low background pressures in the FTMS. A sample of less than 1 mg can
realistically last for several hours on an unheated solids probe (vide infra). For
organometallic complexes which are either cosdy to synthesize and/or difficult to
isolate in high yields, this is a desirable quality. Additionally, oxygen and water
sensitive samples can easily be sampled into the mass spectrometer and decomposition
of these unstable complexes can be hindered by the low background pressure of the
system.

CHAPTER 7
SUMMARY
Thermal free energies of ionization for some prototypical organometallic
complexes have been measured by using Fourier transform ion cyclotron resonance
mass spectrometry. Temperature dependence studies of electron-transfer equilibrium
reactions were performed in order to measure reaction enthalpies and entropies.
Comparison of the experimental ASet° value for the Cp2Fe/DET couple and the value
predicted by statistical mechanical analyses indicates that the experimental reaction
entropy for the couple may be too large. Through a combination of spectroscopic data
and statistical mechanical analyses, the origins of the ionization entropy for the
Cp2Fe+/0 couple have been determined. Vibrational entropy changes account for over
50% of the total A5j°(Cp2Fe). Although the A5j° values for ferrocene in the 298 to
600 K range were found to be relatively large (~5 cal mol^K'1), the experimentally
derived value is about twice as large.
The AGj° values have been used to derive estimates of the average heterolytic
and homolytic M-Cp bond disruption enthalpies for the first transition-row
metallocenes and metallocenium ion through the application of thermochemical cycles.
Thermochemical cycles which include AG° values in the gas phase and
solution have been used to derive estimates of differential solvation energies,
AAGsolv°, for numerous metallocene complexes. Although solvation energetics are
191

192
difficult to measure for individual ions, AAGsolv° values can be used to understand
relative solvation energetics for metal redox couples. Variations in the values of
AAGsoiv° for the metallocenes have been shown to be primarily dependent on both size
of the metal complex and not necessarily the net charge or the metal center. For
example, the AAGsolv° values for the Cp2Ni+/0 and the Cp2Ni0/' couple are equivalent
within experimental error.19 Further the AAGsolv° value for Cp*2Ru is nearly one-half
AAGsolv°(Cp2Fe) which is consistent with the relative molecular volumes for the metal
complexes (the volume of Cp2Fe is -50% the volume of Cp*2Ru). Overall, the
AAGSoiv° values for the metallocenes studied in this work are consistent with AGel°
values estimated by the Bom solvation model.16,19
Correlations of AG¡° values for alkylated ferrocene and nickelocene derivatives
demonstrate that alkyl substituent parameters can be applied to predict thermodynamic
values for organometallic complexes. The application of various models5,53 allows for
the separation of various electron effects. The contribution of inductive and
polarization effects for the alkylnickelocenes is equally important in the stabilization of
(RCp)2Ni+ complexes. For comparison, polarization effects (P) for aliphatic alcohols
are more important in the stabilization of ROH2+ and RO' species than inductive
effects (7).102
The application of the alkyl Taft parameters failed to correlate with AGj°
values for several ruthenocene derivatives studied in this work. Therefore a
parameterization scheme, y parameters, based on the overall electronic nature of Cp
ligands was developed. The utility of Cp y parameters for correlating thermochemical

193
parameters has been applied to XPS data.91,108 The application of these parameters
for predicting reactivity of metal complexes bearing Cp derivative ligands is of great
interest and studies have been designed and performed with this goal in mind. The
ligand parameters are potentially useful in the areas of homogeneous catalysis and
advanced material science.
The information presented in this work was determined with the idea of
developing a foundation for organometallic thermochemistry. As stated earlier,
thermochemical information is fundamentally important to chemistry.
Thermochemistry allows chemists to better understand chemical transformations and
subsequently attenuate, adjust, and further refine chemical reactivity of specific
systems. More than a compilation of data, the values derived in this work have been
critically analyzed and should be useful to other workers in the area of inorganic
chemistry.

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BIOGRAPHICAL SKETCH
The author was bom in Brooklyn, New York, in 1965 and was raised in the
Long Island town of Kings Park. He graduated from Kings Park High School in June
of 1983 and left the small north shore community in August of 1983 in order to attend
college at the State University of New York in Oneonta, New York. Because of his
strong interest in science, he chose to study chemistry at the state college and
graduated in 1987 with a B.S. degree in chemistry. In addition to fostering and
developing academic interests at school, he learned a great deal of life and formed
several lifelong friendships.
After Oneonta, he enrolled in the University of Florida to pursue an advanced
degree in inorganic chemistry. He received his master’s in the summer of 1990. He
later completed his requirements for a Doctor of Philosophy in Inorganic Chemistry in
the Spring of 1993 under the guidance of Professor David E. Richardson. The author
has accepted a postdoctoral appointment with Professor Helmut Schwarz of the
Technische Universitat Berlin in Germany and will begin work in the summer of 1993.
203

I certify that I have read this study and that in my opinion it conforms to
acceptable standards of scholarly presentation and is fully adequate, in scope and
quality, as a dissertation for the degree of Doctor of Philosophy.
David E. Richardson, Chair
Associate Professor of Chemistry
I certify that I have read this study and that in my opinion it conforms to
acceptable standards of scholarly presentation and is fully adequate, in scope and
quality, as a dissertation for the degree of Doctor of Philosophy.
£
V
. Eyler ^
ssor 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
Russell S. Drago
Graduate Research Professor of Chemistry
DoctoiX)f Philosophy.
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.
Gus J. Paleñik
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.
James W. Dufty
Professor of Physics
This dissertation was submitted to the Graduate Faculty of the Department of
Chemistry in the College of Liberal Arts and Sciences and to the Graduate School and
was accepted as partial fulfillment of the requirements for the degree of Doctor of
Philosophy.
May 1993
Dean, Graduate School

UNIVERSITY OF FLORIDA
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