Ionization, bonding, and solvation energetics of organometallic complexes


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Ionization, bonding, and solvation energetics of organometallic complexes
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xiv, 203 leaves : ill. ; 29 cm.
Ryan, Matthew Francis, 1965-
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Subjects / Keywords:
Organometallic chemistry   ( lcsh )
Metal complexes   ( lcsh )
Metallocenes   ( lcsh )
Chemistry thesis Ph. D
Dissertations, Academic -- Chemistry -- UF
bibliography   ( marcgt )
non-fiction   ( marcgt )


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






To my Grandparents


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.


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

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

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

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


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


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

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

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

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


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


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


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



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


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.


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


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.


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

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


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


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


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.



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+



2.0 3.0
Reaction Time




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.




.,. 4*

+ NFc

m +
m ---- DET
m W

i- -



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


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 160.1
0 NMe2159.4

)NEt2 156.3

164.6- Cp2Ru

160.6 Cp20s




153.1 Cp2Fe
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- .


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.


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




, 3.0-




Figure 2-3


A. EtFc/Fc


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 =

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

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


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


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

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



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


Cl 00


\c so

C tr)

(Sl '/n

in kn
'f*) '/n
1r C-
C4 t





























0 0 0
Vc c 0
e e .
o o

Ut- (S



C o

- II



- cr w

0 0
00 00

C -

o cc


0 0

a E -


,- 8

e 0 ^a

LIa 0S03

'i 00
It 'IR

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

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


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

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


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


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



u 145


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


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)


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


M (g)

Cp2M (s)



M3+ (g) + 2Cp- (g) + e-


+ 2Cp- (g)



p2M (soln)


M+ (g) + 2Cp- (g) + e-


Cp2M+ (g) + e-


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




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.


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]


Electrochemical E Data and Differential Solvation Energies for Some
Metallocenes Cp2M+/o couples.












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)


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.























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







o 120

8 100'


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).



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


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


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


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.


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


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

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.



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


A- I




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



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---



141.5 EthylNc

139.6 (MethylCp)2Ni

138.2 (EthylCp)2Ni

136.4 (t-butylCp)2Ni


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.



CO 17.3


21.2 t-butyl

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.


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.









1.5 kcal mol1.









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


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


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

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)

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

w -145-

73 -150

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

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
















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


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





w -145-


"E -150-

-1 .60

Figure 3-5

t-Bu E

R=R'= Et --10

S\ R=R'=Me "
R=H;R'=Et u,,

-.-- -15 u-
SR= R'=H -


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.