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Experimental and theoretical investigation of the reactivity of partially fluorinated radicals

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Experimental and theoretical investigation of the reactivity of partially fluorinated radicals
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Includes bibliographical references (leaves 175-184).
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EXPERIMENTAL AND THEORETICAL INVESTIGATION OF THE REACTIVITY
OF PARTIALLY FLUORINATED RADICALS













By

MICHAEL DAVID BARTBERGER


A DISSERTATION PRESENTED TO THE GRADUATE SCHOOL
OF THE UNIVERSITY OF FLORIDA IN PARTIAL FULFILLMENT
OF THE REQUIREMENTS FOR THE DEGREE OF
DOCTOR OF PHILOSOPHY

UNIVERSITY OF FLORIDA


1998














ACKNOWLEDGEMENTS


Among the great number of individuals with whom I have interacted throughout

the course of my education at the University of Florida and elsewhere, I wish to express

my sincere appreciation to the special few that have motivated, challenged, and inspired

me.

I extend my deepest gratitude to Prof. William R. Dolbier, Jr., an outstanding

scientist and truly exceptional educator, for his excellent guidance, support, and

friendship throughout the course of my graduate career. My appreciation for the

knowledge he has shared with me, as well as his patience and level of understanding,

particularly during periods of difficulty and stress, can not be overstated.

I wish to thank the two finest classroom instructors I have ever had-my first

college level chemistry teacher, Dr. Jeanette Madea, for her profound influence in my

decision to pursue a career in the chemical sciences, and Prof. Seth Elsheimer, for my

initial exposure to the fascinating area of organofluorine chemistry in 1990 and his

friendship thereafter.

I am indebted to Dr. Max Muir for introducing me to computational chemistry.

The experience I have gained in the use of molecular orbital methods as a tool for the

understanding of chemical reactivity is due entirely to him. Special thanks go to Prof.

Benjamin Horenstein for his helpful discussions and generosity with regard to

computational resources.

I thank my colleagues, past and present, in both the Dolbier group and the

Department of Chemistry as a whole. A few bear special mention-Dr. Keith Palmer, for

his friendship and advice during my first year in the group; Dr. Xiao Xin Rong and He-Qi








Pan, for their camaraderie and early assistance with radical kinetics; Dr. Conrad

Burkholder, for numerous stimulating discussions; and Dr. Henryk Koroniak, Michelle

Fletcher, Lian Luo, Feng Tian, and Kevin Ley for their friendship (and tolerance!)

throughout the course of my stay in the department.

I wish to thank my graduate committee, particularly the "organic" portion thereof--

Profs. Merle Battiste and Kirk Schanze, for their advice and encouragement. Also,

special thanks go to Prof. R. J. Bartlett for taking seriously my interest in theoretical

methodology and the invitations to participate in his workshops on Applied Molecular

Orbital Theory.

I am especially grateful to my very best friend, Cynthia Dawn Zook, for her

unrelenting moral support and encouragement over the last several years. Finally, I

wish to acknowledge my parents, George Charles and Beverly Jean, for instilling in me

the work ethic which has likely had as much to do with the successful completion of this

work as any of the chemistry I ever learned.














TABLE OF CONTENTS



ACKNOW LEDGEMENTS ............................................................................................... ii

ABSTRACT.................................................................................................................... vi

CHAPTER

1 AN OVERVIEW OF ORGANIC FREE RADICAL REACTIONS ................. ........ 1

Introduction .............................................................................................................. 1
Radical Chain Processes .......................................................................................... 2
Hydrogen Atom Abstraction Reactions...................................................................... 4
Intermolecular Radical Addition Reactions ................................................................ 9
Intramolecular Addition Reactions: Radical Cyclizations.......................................... 13
Methods for Determination of Organic Radical Kinetics........................................... 22
Conclusion .............................................................................................................. 26

2 THE FLUORINE SUBSTITUENT IN ORGANIC SYSTEMS................................. 28

Introduction ............................................................................................................. 28
Structure, Bonding, and Reactivity in Saturated Systems........................................ 29
Structure, Bonding, and Reactivity in Unsaturated Systems.................................... 31
Fluorine Non-Bonded Interactions in Reactive Intermediates.................................. 33
Fluorine Steric Effects .............................................................................................35
The Fluorine Substituent in Free Radicals............................................................... 37
Organofluorine Radical Reactivity ........................................................................... 40
Conclusion .............................................................................................................. 48

3 THE REACTIVITY OF PARTIALLY FLUORINATED RADICALS IN
INTERMOLECULAR ADDITION AND HYDROGEN ABSTRACTION
REACTIONS .......................................................................... .......... .. ............. 49

Introduction ............................................................................................................. 49
Precursor Syntheses and Competitive Kinetic Studies ............................................ 50
Discussion ............................................................................................................... 61
Conclusion .............................................................................................................. 76

4 THE REACTIVITY OF PARTIALLY FLUORINATED RADICALS IN
INTRAMOLECULAR CYCLIZATION REACTIONS ..................... ............................ 77

Introduction ............................................................................................................. 77
Precursor Syntheses and Competitive Kinetic Studies............................................ 78
Discussion ............................................................................................................... 89
Conclusion ............................................................................................................ 100










5 EXPERIM ENTAL................................................................................................... 102

General Methods- Experimental............................................................................ 102
General Methods- Theoretical ............................................................................... 103
Synthetic Procedures ............................................................................................ 103
Com petitive Kinetic Procedures............................................................................. 129

APPENDIX A: SELECTED 19F NM R SPECTRA.......................................................... 136

APPENDIX B: B3LYP/6-31G(d) TOTAL AND ZERO-POINT ENERGIES FOR
DATA IN TABLES 3-3 AND 3-4............................................................................. 172

REFERENCES............................................................................................................ 175

BIOGRAPHICAL SKETCH .......................................................................................... 185














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

EXPERIMENTAL AND THEORETICAL INVESTIGATION OF THE REACTIVITY
OF PARTIALLY FLUORINATED RADICALS

By

Michael David Bartberger

May 1998

Chairman: William R. Dolbier, Jr.
Major Department: Chemistry

The reactivities of a series of partially-fluorinated radicals towards intermolecular

addition, hydrogen abstraction, and intramolecular cyclization have been investigated.

Based on competitive kinetic techniques and absolute rate contents for addition of these

radicals to styrene obtained by laser flash photolysis, absolute rate constants for

abstraction of hydrogen from tributylstannane have been determined for 1,1-

difluoroalkyl, 2,2-difluoroalkyl, 1,1,2,2-tetrafluoroalkyl, 3-perfluoroalkyl, and

pentafluoroethyl radicals. Fluorination at the 3-position of an alkyl radical was found to

exert a negligible effect on the kinetics of hydrogen abstraction. All other systems

exhibit rate enhancements relative to non-fluorinated analogues, the magnitudes of

which are dependent upon the degree and location of fluorine substitution. A parallel

computational study was performed utilizing density functional calculations, providing

estimates of carbon-carbon and carbon-hydrogen bond dissociation energies (BDEs) for

hydrofluorocarbons. The observed kinetic enhancements were attributed to a

combination of structural, charge transfer, and enthalpic effects, due to the pyramidal

nature of 1-fluoralkyl radicals, increased electrophilic character induced by successive

fluorination, and thermodynamics of carbon-carbon and carbon-hydrogen bond








formation. From the computation of partial atomic charges in fluoroalkanes, the

contrasting effect of 1-fluorination on carbon-carbon and carbon-hydrogen BDEs and

the consistent strengthening effect of such substitution at the 2-position have been

explained on the basis of Coulombic interactions.

Based on the rate constants obtained for hydrogen abstraction, absolute rate

constants for 5-exo and 6-endo intramolecular cyclization for a series of partially

fluorinated 5-hexenyl radicals have been obtained. These observed rates of cyclization

may be rationalized by the same combination of effects influencing their bimolecular

addition reactions. In some cases, the rate of 6-endo closure is dramatically

accelerated relative to the parent hydrocarbon without the introduction of reversibility of

ring closure.














CHAPTER 1

AN OVERVIEW OF ORGANIC FREE RADICAL REACTIONS


Introduction


The discovery of the first free radical, triphenylmethyl, by Moses Gomberg1 in

1900 initiated considerable effort directed toward the understanding of radical reactivity.

However, only after a series of pioneering investigations undertaken more than thirty

years later were the primary mechanistic pathways available to organic free radicals well

elucidated.2- These studies, most notably those of Kharasch et al.,47 demonstrated that

most radical processes can be expressed in terms of a small number of elementary

steps, or variations thereof, as shown below in Figure 1-1.8


A" + B A-B coupling / homolysis (1-1)

A* + B-D A-B + D substitution (SH2) (1-2)

A* + B=D A-B-D addition / p-fission (1-3)

A' + e A- ; electron transfer (1-4)

A' e -- A+

Figure 1-1. The Elementary Mechanistic Pathways of Free Radical Reactions.


Despite these breakthroughs, organic free radical reactions continued for years

to be regarded as unpredictable, unselective, and in general inadaptable to synthetic

application. Fortunately, subsequent kinetic and thermochemical studies have served to

uncover the factors governing the reactivity of organic radicals, and consequently in

recent years sentiment toward the utility of free radicals in synthesis has drastically

changed. Indeed, the number of elegant works in the literature based on radical








mediated transformations is a testament to their applicability in the construction of

natural products and other complex synthetic targets.915 It is the purpose of this

introductory chapter to acquaint the reader with the fundamental types of free radical

processes which occur in organic systems, as well as to provide an overview of the

wealth of physical studies which have given rise to the current level of understanding of

organic radical reactivity.


Radical Chain Processes


Most free radical reactions occur via a sequence of chain events, propagated by

intermediate steps during the course of the reaction. An example illustrating a

competition between two potential pathways is provided in Figure 1-2.


Initiation: In-In 2 In (1-5)

In* + M-H In-H + M" (1-6)

Propagation: M" + R-X -- M-X + R" (1-7)
kH
R" + M-H H-- R-H + M' (1-8)
kr
R" k -- R'" (1-9)

Propagation: R" + M-H -- R'-H + M- (1-10)

Figure 1-2. Radical Chain Process Involving Competition Between Rearrangement
versus Hydrogen Atom Transfer from a Donor Molecule M-H.


Homolysis of an initiator, typically accomplished by thermal or photochemical

means, provides a source of (often metal centered) radicals M" (equation 1-6) from

which intermediate radicals R are generated by reaction with a suitable precursor R-X

(equation 1-7). This species encounters one of two fates: trapping, in this case by

hydrogen atom donor, to yield R-H (equation 1-8) or transformation via a unimolecular or

bimolecular process (equation 1-9) to form radical R", itself then trapped producing








R'-H. In either case, additional metal radicals are formed and the chain process

continued via the propagation steps given in equations 1-7, 1-8 and 1-10.

The distribution of products R-H and R'-H is governed by the relative propensity

of R* toward rearrangement versus trapping (that is, kr and kH,) the latter dependent on

both the nature of R" and the type of trapping agent employed. In systems where

trapping is fast relative to rearrangement (kH >> kr), the partitioning radical R is

converted to R-H with little or no rearranged product. However, if kH and kr are of

comparable magnitude, product mixtures result. An understanding of the reactivity of a

radical intermediate toward potential competing processes is therefore essential for the

design of useful kinetic experiments, as well as for the development of effective synthetic

strategies.

For an efficient chain process, it is necessary that the propagation steps are

rapid relative to chain termination steps, thereby maintaining a low but constant

concentration of radical intermediates. Besides the obvious practical benefit (higher

product yields) resulting from such a condition, the occurrence of undesired chain

termination side reactions such as disproportionation and radical-radical coupling,

possibly complicating kinetic analyses, is minimized. In many cases, this may be

achieved by judicious selection of the type and concentrations of precursor R-X and

trapping agent.

This procedure enjoys wide application in both kinetic and synthetic studies

requiring the controlled generation of radical intermediates. One of its variants, likely the

most commonly used procedure for the indirect (competitive) determination of the rates

of organic radical reactions, is based on the trialkylstannane reduction of an alkyl halide

(the 'Tin Hydride Method").16-18 Other modifications of this general procedure exist,

accommodating a variety of radical precursors and trapping agents; a discussion of time-

resolved and competitive techniques utilized in radical kinetic measurements is provided

later in the chapter.








Finally, it is important to note the implication of kinetic control in the above

discussion. That is, that the distribution of products R-H and R'-H may be ascribed to

the relative values of kH and kr hinges on the absence of any thermodynamic

equilibration of products under the reaction conditions. This is of vital importance in the

design and interpretation of competitive kinetic studies and is discussed in detail in

Chapter 3.

Radical reactivity is dependent on the "complex interplay" of thermodynamic,

steric, and polar considerations.19 The relationship between enthalpies of activation and

heats of reaction, the basis of the thermochemical kinetic approach of Benson,20 was

recognized early on and holds for a number of radical addition and substitution

reactions, where the order of reactivity often parallels exothermicity.21-23 This relation

has led to such overgeneralizations as "radical reactions follow the most exothermic

available pathway" or ". afford the most stable possible product."8 However, reaction

thermochemistry is not the sole, nor even predominant decisive factor in the outcome of

radical reactions. Nonbonding interactions and the electronic influence of substituents in

ground and transition states (which may be rationalized in terms of Frontier Molecular

Orbital (FMO) theory)24'25 will also play a role. A discussion of the combination of these

effects as manifested in hydrogen atom abstractions and inter- and intramolecular

additions, the most commonly occurring and well-characterized reactions of organic free

radicals, will now be presented.


Hvdrogen Atom Abstraction Reactions


The vast majority of free radical applications involve the use of an organometallic

hydride of the type R3M-H (most commonly, where M = Sn, Si, or Ge) as a hydrogen

atom donor and chain propagation agent, the properties of which have been the focus of

extensive investigation by kineticists. Metal-hydrogen bond dissociation energies

(BDEs) along with activation parameters and associated absolute rate constants for





5

hydrogen atom transfer to n-alkyl hydrocarbon radicals by a series of donors R3M-H

have been determined and are provided below in Table 1-1 .16, 26-32

Table 1-1. Bond Dissociation Energies with Activation Parameters and Rate Constants
for Hydrogen Atom Transfer to Hydrocarbon Radicals by R3M-H.

R3M-H BDE. kcal mol1 log A E, kcal mol'1 k. _,10 M s' (298 K)

nBu3SnH 73.7 9.06 3.65 2.3

(TMS)3SiH 79.0 8.86 4.47 0.38

nBu3GeH 82.6 8.44 4.70 0.093

(TMS)2Si(CH3)H 82.9 8.89 5.98 0.032

Et3SiH 90.1 8.66 7.98 0.00064


Analysis of the data demonstrates that for hydrogen atom abstraction by

structurally similar radicals from this series of donors, a direct relation holds between the

rate of transfer and the strength of the metal-hydrogen bond being broken. This is

depicted graphically in Figure 1-3. In addition, it is noted that in each case the pre-

exponential term in the Arrhenius relation remains relatively constant. Thus, the rate

variations within the series are due almost entirely to differences in activation energies.


16 -
16m = -0.50403
S14- b =52.313
4r2 = 0.96858

120-


8 -
6 '-

72 74 76 78 80 82 84 86 88 90 92

BDE, kcal mol"1

Figure 1-3. Plot of In kH for Alkyl Radicals versus M-H Bond Dissociation Energies for
Hydrogen Atom Donors R3M-H in Table 1-1.








However, as previously mentioned, relative thermodynamics is not the only factor

which influences the kinetics of hydrogen atom transfer. The fast donor thiophenol

(PhSH) reacts with primary alkyl radicals with a rate constant of 1.36 x 108 M"' s1 at

298 K,33 and has been employed as a trapping agent in competitive kinetic studies

involving strained or otherwise highly reactive radicals with rearrangement rates upward

of 101" s1 and thus with lifetimes on the picosecond timescale.34 This enhanced rate of

transfer, not commensurate with its S-H BDE of 82.0 kcal mol1,35 gives rise to a severe

deviation from the plot in Figure 1-3 and indicates the presence of other influences.

Table 1-2. Absolute Rate Constants for Hydrogen Atom Transfer to tert-Butoxyl
Radicals by R3M-H.

R3M-H kH. 106 M1 s-' (300 K)

nBu3SnH 220

(TMS)3SiH 110

nBu3GeH 80

Et3SiH 5.7


Further evidence may be found in the rates of hydrogen atom transfer to tert-

butoxyl radicals by the same series of hydrogen atom donors, provided in Table 1-2 and

illustrated graphically in Figure 1-4.3637 It is observed that for tert-butoxyl radicals, rates

of hydrogen abstraction are at least two orders of magnitude greater than those of their

n-alkyl counterparts. Although the relative strengths of the newly formed C-H or 0-H

bonds will certainly play a role, the difference in BDE between tBuO-H and n-alkyl C-H

bonds (105 and 100 kcal mol1, respectively)38 is not sufficient to explain the increase in

reactivity, especially in light of the fact that such rapid hydrogen atom abstractions

should proceed with early transition states.36

At this time, the absolute rates of reduction of tert-butoxyl radicals by thiophenol

have yet to be determined. However, a series of competition studies by Hartung and








Gallou39 involving 4-pentenyl-l-oxy radicals and utilizing naphthalene 2-thiol (NpSH) as

a trapping agent have determined a ratio [ kH (NpSH) / kH (nBu3SnH) ] of 1.4. By

comparison, n-alkyl radicals afford the ratio [ kH (PhSH) / kH (nBu3SnH) ] = 59.1.

Although a leveling effect may be partly responsible for the compressed ratio of rates for

tert-butoxyl radicals (which are indeed within an order of magnitude of the diffusion-

controlled limit)40 it is logical to assume based on the aforementioned examples that

hydrogen abstraction reactivity will be governed to some extent by factors other than

simple relative BDE values of the donor species.



2Om = -0.22212
19- b =35.942
2'~~r ="S^r 0.91532
0 18

17

16-

15 1 T -------------
72 74 76 78 80 82 84 86 88 90 92
BDE, kcal" mol[1

Figure 1-4. Plot of In kH for tert-Butoxyl Radicals vs. M-H Bond Dissociation Energies for
Hydrogen Atom Donors R3M-H in Table 1-2.


Chatgilialoglu et al.3637 have attributed such differences in reactivity to a

polarized, or charge separated, transition state of the type depicted in Figure 1-5. Here

it can be seen that in the case of an electropositive metal hydride donor, hydrogen atom

transfer to alkoxyl radicals (b) is facilitated by greater stabilization of partial negative

charge on oxygen relative to carbon, with a resultant decrease in activation barrier.


*+ 6- 8* 8-
(a) R3M ----- H ----- R (b) R3M ----- H --- OR

Figure 1-5. Charge Polarized Transition State for Hydrogen Abstraction from R3M-H by
(a) Alkyl and (b) Alkoxyl radicals.








In the case of thiol donors, the opposite situation ensues. The greater

electronegativity of sulfur relative to tin (or other metal atom) gives rise to a reversal in

the transition state charge distribution. This arrangement, involving a partially negatively

charged sulfur atom, is better suited to the more nucleophilic alkyl radical, where in the

alkoxyl case a less- or non-polarized transition state results. Such a mismatch in the

latter is partially responsible for the decrease in rate enhancement for hydrogen

abstraction from thiols by alkoxyl radicals, relative to their alkyl analogues.


Frontier Molecular Orbital Theory of Atom Abstraction Reactions


FMO theory provides a satisfying rationale for the kinetic characteristics of

hydrogen abstraction reactions of free radicals. In general terms, radicals are species

containing an unpaired electron in a singly occupied molecular orbital (SOMO), which in

the ground state of the radical is its highest occupied orbital. According to the FMO

concept, during the course of the reaction this SOMO will interact with both the highest

occupied (HOMO) and lowest unoccupied (LUMO) molecular orbitals of the donor

molecule. Such interactions between these "frontier molecular orbitals"24 are not

necessarily equal. The extent of SOMO-HOMO and SOMO-LUMO interaction is

governed by their energy values, the strongest interaction occurring between orbitals

closest in energy.

It is these values, influenced by atom type as well as neighboring substituents,

from which the relative descriptors such as "nucleophilic" and "electrophilic" are derived

and provide the basis for the previously described concept of "polar factors." Electron

donating substituents generally serve to raise both HOMO and LUMO energies, with

electron withdrawing groups resulting in lowering. Radicals possessing a low energy

SOMO will display electrophilic character, whereas a higher energy SOMO gives rise to

a more nucleophilic species. Figure 1-6 depicts the FMO interactions between radical

and donor in each of these cases.








During the abstraction process, the primary interaction involves the radical

SOMO and the a and C* orbitals of the donor M-H bond. The antibonding a* orbitals of

the donor are typically quite high in energy, and thus in atom abstraction reactions the

SOMO-HOMO interaction dominates.



SOMO ,

SSOMOM
-- HOMO ', ^-- HOMC


(a) (b) *

Figure 1-6. FMO Diagram Illustrating the SOMO-HOMO Interaction Between (a) a
Nucleophilic Alkyl and (b) an Electrophilic Alkoxyl radical.


Here it is seen that the lower-energy SOMO of the alkoxyl radical (ca. -12 eV, as

determined from ionization potential measurements)25 leads to a reinforced interaction

with the donor HOMO (case b). This greater stabilizing interaction results in a lowered

activation barrier and hence a more facile transfer reaction, compared to the more

nucleophilic alkyl radical, (case a) whose SOMO energies range from -6.9 to -9.8 eV.25

Some of the most striking examples of such "polar" factors involve systems

where fluorine substitution has taken place at, or adjacent to, the radical center. This is

elaborated upon in Chapters 2 and 3, where the effects of fluorination on the structure

and reactivity of free radicals are discussed and compelling evidence provided based on

kinetic studies of hydrogen transfer to such partially and fully fluorinated alkyl radicals.

Intermolecular Radical Addition Reactions


Over the past twenty years, the intermolecular addition reactions of free radicals

(as well as their intramolecular cyclization counterparts) have become an important








addition to the arsenal of C-C bond formation methods available to the synthetic organic

chemist. Their mild means of generation from a variety of precursors and tolerance for a

wide variety of functional groups provide distinct advantages over ionic processes.

Alkyl radical additions to carbon-carbon double bonds are highly exothermic, as a

new a bond is formed at the expense of a n bond (in the case of methyl radical addition

to ethylene (Figure 1-7), AHrxn= ca. -22.6 kcal mol-1).41'42 In accordance with the

Hammond postulate,43 such additions should proceed via early transition states, with low

barriers of activation. This is indeed the case, as supported by a wealth of both

experimental44 and theoretical4,1'4549 data.


CH; + CH2=CH2 CH3CH2CH2"

Ea = 7.9 kcal molr1
AHrxn = -22.6 kcal mol-r1

Figure 1-7. Addition of Methyl Radical to Ethylene, Yielding n-Propyl Radical.
Experimental Activation Energies and Heats of Reaction are Shown.


FMO Theory of Radical Additions


Quantum mechanical molecular orbital calculations at levels of ab initio theory

ranging from UHF to UQCISD(T) and varying basis sizes from 3-21G to 6-311G(2df,p)

are consistent in their characterization of the transition structure for the above reaction




G I::pyr(CH3)= 101.9
\

2.246 A 109.1o 173.40



S1.382 A
154.7

Figure 1-8. UHF/6-31G(d) Transition Structure and Relevant Geometrical Parameters
for Addition of Methyl Radical to Ethylene.









S



SOMO ,I
a


- LUMO


-- HOMO


SOMO--


LUMC



' -
i
I

,'4| HOMC


(a) (b) -

Figure 1-9. FMO Diagram for Addition of (a) Nucleophilic and (b) Electrophilic Radicals
to Alkenes.

Table 1-3. Some Relative Rates of Addition of Methyl and tert-Butyl Radicals to Alkenes
CH2=CHX.


x

H

CH3

OCH2CH3

F

Cl

CN


KreL I (3

1

0.7


KeL( ^^03 3C

1

0.74

0.31


13.2

1920


a Data for ethyl radical.


(Figure 1-8), which possesses an incipient C-C bond distance of ca. 2.23 2.27 A. The

C-C-C attack angle of 109.1 is rationalized in FMO terms based on a primary interaction

between the radical SOMO and the LUMO of the alkene. It is in such reactions with high

exothermicities and early transition states that FMO interactions are most

substantial.2450 This postulate enjoys experimental support; for the t-butyl radical, a

correlation exists between rates of addition to alkenes and the experimentally








determined electron affinities of the latter.21'51 Such an FMO interaction for nucleophilic

radicals is shown in Figure 1-9 (case a) and is influenced by substituents on both radical

and olefin, a raising of radical SOMO and/or lowering of alkene MO energies

strengthening the SOMO-LUMO interaction and enhancing the rate of addition as seen

from the data in Table 1-3.21,52 Here, the greater nucleophilic character of t-butyl relative

to methyl is evident from its enhanced rate of addition to olefins bearing electron

withdrawing groups.

As previously discussed, electronegative substituents at the radical center which

substantially lower its SOMO energy will impart electrophilic character and reinforce the

transition state SOMO-HOMO interaction (Figure 1-9, case b). Indeed, it has been

shown that rates of addition of dicyanomethyl53 and perfluorinated54 radicals correlate

with the ionization potentials of the substrate alkenes. The intermediate behavior of
"ambiphilic" radicals, such as malonyl and (tert-butoxycarbonyl)methyl, has also been

documented, these species yielding "U"-shaped correlations between rates of addition

and alkene IP and EA values.5557

A more thorough presentation of kinetic results obtained to date for the addition

of partially and fully fluorinated radicals to alkenes is given in Chapters 2 and 3.


Steric Effects; Regiochemical Preferences in Addition


Competition studies on both nucleophilic and electrophilic radicals have provided

for some generalizations in terms of the regiochemical preference for addition to

unsymmetrically substituted olefins.52 The preferred orientation of radical addition

occurs to the unsubstituted end of the double bond, attributed to steric repulsion but also

influenced by the effect of substituents on the coefficients of the HOMO and LUMO of

the alkene. Such FMO effects can be the decisive factor in polysubstituted olefins if

steric effects are in opposition. Strongly spin delocalizing substituents on the alkene

reinforce such sterically induced regiochemical preferences and exert slight rate








enhancing effects; however, as such additions occur through early transition states the
effect of exothermicity on the kinetics of addition should be minimal.

The concepts introduced in the aforementioned discussion on intermolecular
radical additions extend to their intramolecular cyclization analogues, an overview of
which will now be presented.

Intramolecular Addition Reactions: Radical Cyclizations

The intramolecular addition reactions of alkenyl radicals enjoy a strong foothold

among the available strategies for the construction of cyclic organic molecules. In
addition to their synthetic utility, the kinetic, regioelectronic and stereoelectronic
characteristics of radical cyclization reactions as a function of substituent continue to fuel
an abundance of fundamental physical organic structure-reactivity investigations, more

than thirty years after the first report of the archetypal radical ring closure, cyclization of
hex-5-en-1-yl radical 1 (Figure 1-10).58



250 C
S6 + 0

(98%) (2%)
1 2 3
Figure 1-10. Cyclization of Hex-5-en-1-yl Radical 1 to Cyclopentylcarbinyl (2) and
Cyclohexyl (3) Radicals. At 250 C, 5-exo Closure Dominates 49: 1.

The most striking aspect of this reaction lies in the preferred regiochemistry of
addition. In the case of the parent hydrocarbon, 5-exo59 cyclization dominates (Eact [5-
exo] = 6.8 kcal mol1, Eac [6-endo] = 8.5 kcal mol-1)27'60 yielding the less

thermodynamically stable primary cyclopentylcarbinyl radical 2. This finding has
provided the driving force for a number of experimental61 and theoretical62"68








investigations geared toward the understanding of radical cyclization regiochemistry in

the hydrocarbon and related substituted systems.

Early explanations,69 later advanced by semiempirical techniques,65 attributed

this result to a less negative entropy of activation for cyclization to 2. Although

experiment demonstrates this to be true, the difference (AS1,5 AS*,6 = 2.8 eu)60 is

insufficient to completely account for the observed regiochemistry; the preferred mode of

cyclization resulting primarily from enthalpic (AH*1,6 AH*1,5 = 1.7 kcal mol1) rather than

entropic factors. Ab initio computations66 lend support to this conclusion.

Transition structures for 5-exo and 6-endo cyclization of 1 have been located

using a variety of theoretical treatments. The UHF/6-31G(d) structures leading to 2 and

3 are shown in Figures 1-11 and 1-12.





^ Ifc-JflL e0 = 109.70 ^ s
2.186 A ',



Figure 1-11. Two Views of the UHF/6-31G(d) "5-exo-chair" Transition Structure for
Cyclization Hex-5-en-l-yl Radical 1 to Cyclopentylcarbinyl Radical 2.





098.40 '





Figure 1-12. Two Views of the UHF/6-31G(d) "6-endo-chair" Transition Structure for
Cyclization of Hex-5-en-1-yl Radical I to Cyclohexyl radical 3.







Spellmeyer and Houk66 have also postulated additional "boat-like" transition
structures on the basis of molecular mechanics calculations parameterized by ab initio
results of model systems. Inclusion of these "boat-like" structures as viable competing
pathways was found to be necessary for the accurate prediction of regio- and
stereoselectivities in cyclizations of alkyl substituted and heteroalkenyl radicals. The
existence of such structures is corroborated by higher level ab initio treatments
performed as part of the present study and are shown in Figures 1-13 and 1-14.




O~~rM^^ Sj~s0 9108.20 d ~e

rK:is) 2.192 A



Figure 1-13. Two Views of the UHF/6-31G(d) "5-exo-boaf Transition Structure for
Cyclization of Hex-5-en-l-yl Radical 1 to Cyclopentylcarbinyl Radical 2.





0 =100.00 "
J^T \2.245 A %Am. r



Figure 1-14. Two Views of the UHF/6-31G(d) "6-endo-twist-boaf' Transition Structure
for Cyclization of Hex-5-en-l-yl Radical 1 to Cyclohexyl Radical 3.

Upon inspection of the forming bond lengths and angles of the 5-exo structure in
Figure 1-11, its similarity to the transition structure for addition of methyl radical to
ethylene is readily apparent. The C-C-C angle of attack, 109.7, is practically identical to








that in Figure 1-8 and fits the requirement for overlap of the radical SOMO with the n*

orbital of the alkene moiety. From Figure 1-12 it is observed that this angle is

significantly reduced (98.4) in the 6-endo approach. Thus the required disposition of

centers for optimal FMO overlap is more readily achieved in the 5-exo transition

structure, leading to the kinetically preferred cyclization product.


Substituent Effects on the Kinetics and Regiochemistry of 5-Hexenyl Cyclizations


Radical 1 cyclizes to cyclopentylmethyl radical 2 with a rate constant (kcs) of

2.3 x 105 s' at 25 C.27 In the parent hydrocarbon, 6-endo closure competes to a very

minor extent (kc6 = ca. 4.7 x 103 s1). However, it will be shown that the rates and

regiochemical preferences can be substantially affected by substitution at both the

radical center and terminal alkene.


Alkyl Substitution: Steric Effects


Beckwith et al. have reported rate constants for a number of alkyl-substituted

hexenyl radicals.70,71 The rates of 5-exo closure as a function of gem-dialkyl substitution

on the aliphatic portion of the hexenyl chain is shown in Table 1-4.

It is seen that substitution at the radical center has a nearly negligible effect on

the rate of cyclization, due to offsetting polar and steric considerations. Conversely, a

significant (> 10-fold) rate enhancement is observed with internal substitution (systems

6, 8, and 9), accelerated by relief of steric compression between alkyl groups during ring

formation (the "Thorpe-lngold", or "gem-dimethyl" effect).72

Intermediate kinetic behavior would be expected of monosubstituted 5-hexenyl

systems. The data in Table 1-5 show this to be the case. In addition, a stereochemical

preference for cis- or trans-dimethylcyclopentanes, depending on the location of the

substituent on the chain, is observed. This has been rationalized by Beckwith et al.71







Table 1-4. 5-exo Cyclization Rate Constants for gem-Disubstituted 5-Hexenyl Radical
Derivatives.

Cyclization Reaction kc5, 105 s1 (298 K)



2.3a
1 2




4 5
*

^ 36c

6 7

= __52c



8 7

= 32c
N*


9 5
a Reference 27. b Reference 28. c Reference 71.


based on the cyclization transition structure depicted in Figure 1-11. Although "early" in
terms of the forming C-C bond, the overall orientation of atoms in this structure is quite
product-like. According to this rationale, substituents on the aliphatic fragment are
likened to those in chair cyclohexane, which then occupy an equatorial position in the
chair transition structure. Minor products are assumed to derive from the occupation of








axial positions. The latter has been disputed by Spellmeyer and Houk,6 whose model

indicates that such secondary products originate from an equatorial disposition of

substituents in the boat-like transition structure of Figurel-12, rather than from an axial

orientation in the chair.

Table 1-5. Cis- and Trans- 5-exo Cyclization Rate Constants for Monosubstituted 5-
Hexenyl Systems. Rate Constants are for 298 K.

Cyclization Reaction kc5 (cis), 105 s"1 kc5 (trans), 105 s-1



+ 1.18 0.42a
10 cis-11 trans-11



+ 2.4" b 4 b

12 cis-13 trans-13

'N +
U Q + 7 7Ob 2.4


14 cis-13 trans-13


+ 0.75b 3.60

15 cis-11 trans-11
a Reference 73. b Reference 71.


In the above examples, 5-exo products are formed either predominantly (> 97%)

or exclusively. Substitution at the vinyl group leads to marked changes in regiochemical

ratios as indicated by the data in Table 1-6. Replacement of hydrogen by methyl (16) or

isopropyl (19) at C5 results in preferential formation 6-membered rings 18 and 21.

Inspection of the data indicates that this shift in regiochemistry is not due to a significant







Table 1-6. Rate Constants at 338 K for 5-exo and 6-endo Cyclization for Vinyl-
substituted 5-Hexenyl Radicals.

Cyclization Reaction kc5, 105 s1 kc6, 105 s1


+ )9.4 0.19 ab
1 2 3
wI d

0.21b 0. O37b

16 17 18


+ 0 0.21b 0.66b

19 20 21


+ 1. 0.19b

22 23 24


S+ 22b < 0.1b

25 26 27
a Calculated from the Arrhenius parameters given in reference 27. b Reference 70.

extent to rate enhancement for 6-endo closure, but rather a substantial (44-fold)
retardation of 5-membered ring formation due to a combination of 1,5 steric hindrance
and back strain engendered at Cs upon adaptation of sp3 character. Substitution at both
C5 and C6 again favors 5-exo closure, the rate of which decreased relative to the parent
system. Disubstitution at C6 (25) gives rise to a slight (2.3-fold) rate enhancement for 5-
exo closure, sufficiently explained on thermodynamic grounds, which dominates 6-endo








cyclization by a factor of at least 220. The above data indicates that the kinetic and

regiochemical characteristics of alkyl-substituted 5-hexenyl cyclizations may be

sufficiently rationalized by steric considerations. The effects of substitution by

conjugating, heteroatom-containing groups is outlined below.


FMO Considerations


Studies of 5-hexenyl systems bearing "polar" subsitutents have been

investigated by Newcomb.16'74-77 In line with those of intermolecular radical additions,

the kinetics of radical ring closure will be influenced by the impact of substitutents on the

SOMO-HOMO and SOMO-LUMO interaction in the cyclization transition state.


Table 1-7. 5-exo Cyclization Rate Constants For oa-Donor- and a-Acceptor-substituted
6,6-Diphenyl-5-hexenyl radicals. Rate Constants are for 298 K.

X Y
SX Y Ph
5-exo
Ph &" Ph

Ph

System X Y kc5. 107 s1 (298 K)a

28-> 29 H H 4

30- 31 H CH3 2

32-> 33 CH3 CH3 1

34- 35 H OCH3 4

36-> 37 H CO2CH2CH3 3.7

38 -39 CH3 CO2CH2CH3 0.04

40 41 CH3 CN 0.03
a Calculated from the Arrhenius parameters provided in reference 76.


As seen from the data in Table 1-7, only a very minor effect is exerted by either

a-donor or a-acceptor substituents, relative to parent system 28. The marked decrease








in rate for 38 and 40 is attributed to an increase in activation energy due to enforced

planarity at the radical site induced by the Tc-delocalizing substituents CO2CH2CH3 and

CN.51'78


Table 1-8. Absolute Rate Constants at 298 K for 5-exo Cyclization of 5-Hexenyl
Systems Bearing Vinylic Donor and Acceptor Substituents.


Cyclization Reaction


kc5, 105 s-1 (298 K)


H3CO


42




NC 44

44


2

*OCH3



43



* CN


45


680b


CN NCj OCH3
H3CON ___ A
L^ .--^ 1000b


46 47
a Reference 27. b Calculated from the Arrhenius parameters provided in reference 74.


Substitution at the vinyl terminus, especially by strong resonance-withdrawing

groups can significantly accelerate the rate of ring closure. Although possessing a

radical stabilizing group, methoxy analog 42 enjoys only a very minor increase in rate.

The donor substituent raises the energies of the frontier orbitals, increasing the SOMO-

HOMO interaction but widening the SOMO-LUMO energy gap, the latter more important








for relatively nucleophilic alkyl radicals (Figure 1-9). Substituents which serve to lower

the FMO energies should reinforce the SOMO-LUMO interaction, leading to rate

enhancement. The nearly 300-fold increase resulting from cyano substitution (44)
reflects such an effect. The slight rate increase 46 relative to 45 has been explained on

the basis of the suggested slight extra "push-pull," or "captodative" stabilization
manifested in donor-acceptor disubstituted systems.79'80

The importance of kinetic control was previously mentioned. Care must be taken
in assessing the potential for reversibility in such intramolecular additions, which may
obscure the effect of steric and/or polar influences on reaction kinetics. This is

demonstrated in the 5-exo:6-endo product ratios of highly stabilized systems 48 and
5 .81,82

Table 1-9. Product Ratios for 5-Hexenyl Cyclization Reactions Under Full or Partial
Thermodyamic Control.
Cyclization Reaction % 5-exo % 6-endo

.*^ 'rPh Ph^V

h <22 > 78
48 49 50

NC CO2Et CN CN
-U CO + L C02 16 84
C02Et

51 52 53

Given the data and discussion provided in the above sections, a review of the
direct and competitive techniques utilized in the determination of rates of organic radical

reactions is now in order.

Methods for Determination of Organic Radical Kinetics

The development of indirect competitive methods, in conjunction with laser flash
photolytic generation and time-resolved detection of transient intermediates, has greatly








expanded the dynamic range available for the measurement of radical reactions,

especially those at the upper end of the kinetic scale. Such advances have provided for
the use of a variety of precursors and the accurate determination of rate constants for

reactions approaching the diffusion-controlled limit in solution.16

Laser Flash Photolysis: Direct Measurement of Addition Rates

In the time-resolved laser flash photolysis method, described in detail in the

literature,83 radicals are generated from precursors possessing a suitable chromophore

by a laser pulse of the appropriate wavelength. Alkyl iodides, diacyl peroxides, and the

0-acylthiohydroxamic esters of Barton et al.84 are most commonly utilized in this regard

(Figure 1-15).

hv
(a) R-l ,-1 R


| R' R hv ||0 -C02
(b) R- 0 2 R1 O fast 2 R
0
O O-

( N hv 0 -CO2
(C) R 0 R 0 fast
s
Figure 1-15. Laser Flash Photolytic Generation of Radicals R" from (a) Iodide, (b) Diacyl
Peroxide, and (c) 0-Acylthiohydroxamic Ester Precursors.

These radicals so generated undergo further reaction, usually bimolecular
addition or unimolecular cyclization to a (typically phenyl-substituted) double bond (a

styrene in the case of bimolecular additions). The increase in the characteristic

absorption of this intermediate benzyl radical (Xmax 320 nm) is then followed in a time-

dependent manner by UV-visible spectroscopy (Figure 1-16). For bimolecular additions,

this experimental growth curve is fit to the expression in Equation 1-11, yielding absolute








rate of addition kwd. With such data in hand, this addition can now serve as a competing

basis reaction for the determination of rates of other transformations involving the same

or structurally similar radical.


kobs =ko + kadd [alkene] (1-11)


-\ R

hv -G /
R-X hvo R --- / ax ca. 320 nm
kadd 1 G
(monitor)

Figure 1-16. LFP Generation of R* and Detection of Transient Benzyl Radical Adduct
for Determination of Absolute Rate of Addition kad.


Indirect Methods (Competitive Techniques)


Indirect kinetic methods involve the partitioning of an intermediate between two

competing pathways, one with a known rate constant and the other whose rate constant

is to be determined. Post facto product analyses, typically by chromatographic or

spectroscopic means, provide a ratio of rate constants from which the new kinetic value

is obtained. Such radical kinetic measurements usually involve competition between two

bimolecular reactions or a bimolecular reaction competing with unimolecular

rearrangement; examples of both instances are provided in Figures 1-17 and 1-18.

Determination of rates of hydrogen abstraction by this method involves the

generation of R" in the presence of two trapping agents; in Figure 1-17, styrene and the

hydrogen atom donor. Both traps are typically present in excess to ensure pseudo-first

order behavior. Radical R* may undergo addition to styrene (with known rate constant

kadd) forming the intermediate benzyl radical, itself trapped with excess hydrogen donor,

yielding closed shell product with a rate which is kinetically unimportant provided the

addition reaction is irreversible. Alternatively, R* is trapped directly by hydrogen atom

donor with a rate constant kH, which may be obtained from the pseudo-first order relation







[reduced] [kH ] [ R1 [M-H ]
[ adduct ] [ kadd ] [ R I [ CH2=CHR']

where [reduced] and [adduct] are the final product concentrations, [M-H] is the
concentration of hydrogen atom donor, and [CH2=CHR'] the concentration of alkene trap.
For accurate kinetic determinations, a series of runs is performed where trap
concentrations are varied, (again, maintaining at least a five-fold excess) a plot of
product ratios versus that of trapping agents yielding the ratio kH I kay.

(vary) VR R


k /add --- ^{S--- ~


\ M-H (vary)
G M-H

R "----- ^ --H
GJG


H RH

Figure 1-17. Competition Between Bimolecular Addition to an Alkene with Rate
Constant kad and Bimolecular Trapping by Hydrogen Atom Donor with Rate Constant
kH.

Determination of rates of cyclization are performed in a similar manner, using
hydrogen atom abstraction (with the known value of kH) as the competitive basis
reaction. Unimolecular rearrangement competes with bimolecular trapping with an
excess of hydrogen atom donor (Figure 1-18) to yield intermediate cyclic radicals, further
trapped to form characterizable products.
The ratio of products of cyclization versus hydrogen abstraction are obtained
from the pseudo-first order relation in Equation 1-13, a plot of the ratio of products of
hydrogen abstraction to those of cyclization as a function of trapping agent concentration
affording ratios kcs I kH and kc6 I kH.
Finally, the importance of an efficient chain process should be reemphasized.
High conversions of precursors, although important for any radical reaction, are crucial in







competitive kinetic experiments. The reliability of data resulting from indirect methods is

directly dependent on high "mass balance" values, those of 90% or greater typically

being desired.


[reduced] [kH] [R] [M-H] (1-13)
[cyclized] [kcn ][R*]


M-H (vary)
kH RH



M-H
Ro
\ kc5 jL



\ l^ \ M -H1"
-k -6 0 0


Figure 1-18. Competition Between Bimolecular Hydrogen Atom Abstraction with Rate
Constant kad and Unimolecular 5-exo and 6-endo Cyclization with Rate Constants kc5
and kc6.

Such competitive processes have been employed extensively by the Dolbier
research group in the investigation of the rates of addition, hydrogen atom abstraction,

and cyclization of a variety of fluorinated open shell systems, providing the first

quantitative kinetic data for this class of reactive intermediates.54,85-91


Conclusion

The proceeding discussions have attempted to provide the reader with an
introduction to the chemistry of organic free radicals. Kinetic data for hydrogen

abstraction, addition, and cyclization reactions of hydrocarbon radicals, important
benchmarks for comparison of reactivity with other substituted systems, was provided.





27


Substituent effects, rationalized on the basis of a combination of thermodynamic, steric

and FMO considerations, were discussed.

The following chapter provides a review of the effects of fluorine substitution in

organic molecules, including fluorinated radicals. Previous research efforts in this area

by the Dolbier group are summarized, setting the stage for the presentation of results of

the current study.














CHAPTER 2

THE FLUORINE SUBSTITUENT IN ORGANIC SYSTEMS


Introduction

Incorporation of fluorine into organic molecules often imparts dramatic alterations

in structure and reactivity. These effects are induced by three major characteristics

inherent to the fluorine atom: extreme electronegativity, non-bonded electron pairs, and

relatively small size.

Fluorine possesses the highest electronegativity of all the elements, with a value

of 4.10 on the Pauling scale, compared to oxygen (3.50), chlorine (2.83), bromine (2.74),

carbon (2.50), and hydrogen (2.20).92 As a substituent in organic systems, this results in

strong inductive withdrawal of electron density through the a molecular framework and

highly polarized bonds with substantial ionic character.

Three non-bonding pairs of electrons in 2p orbitals similar in size to those of

other second-row elements provide for optimal overlap, and therefore an offsetting back

donation of electron density into the molecule to which it is bonded.

The accepted van der Waals radius of fluorine, 1.47 A, suggests minimal steric

impact in comparison with other halogens (chlorine, 1.73 A; bromine, 1.84 A; iodine,

2.01 A; carbon, 1.70 A; oxygen, 1.52 A; hydrogen, 1.20 A).93 This has allowed for the

complete replacement of hydrogen by fluorine in organic systems, a feat not possible to

such an extent with any other element.

The following sections, based on a number of excellent reviews,9497 provide an

introduction to the fascinating behavior exhibited by fluorinated stable molecules and

reactive intermediates due to a combination of the above effects.








Structure. Bonding, and Reactivity in Saturated Systems


The data provided in Tables 2-1 and 2-29'97 reveal a trend unique to fluorine

within the halogenated methanes. An incremental shortening of C-F interatomic

distances, with a resultant increase in bond dissociation energies, is observed as the

series is traversed. No such trend exists for any other member of the halomethane

family; on the contrary, it is seen from the data in Table 2-2 that such C-X BDE values

instead decrease with increasing halogen content. Strengthening of C-H bonds is also

observed within the fluoromethanes (CH3F, 101.3 kcal mol1; CH2F2, 103.2 kcal mol1;

CF3H, 106.7 kcal mol-1).95


Table 2-1. Carbon-Halogen Interatomic Distances (Angstroms) of Halomethanes

X CHa_ CH2X CHX3 CX.

F 1.385 1.357 1.332 1.319

Cl 1.781 1.772 1.758 1.767

Br 1.939 1.934 1.930 1.942

Table 2-2. Carbon-Halogen Bond Dissociation Energies (D, kcal mol') of
Halomethanes

X CH CHX CHXg CX9

F 108.3 119.5 127.5 130.5

Cl 82.9 81.0 77.7 72.9

Br 69.6 64 62 56.2

I 57.2 51.3 45.7 -

Data for geminally fluorinated ethanes parallel that of the methane series,

demonstrating a progressive strengthening and shortening of both C-C and C-F bonds

with increasing fluorination (Table 2-3). Conversely, vicinal fluorination gives rise to the

opposite effect on C-C bonds, a steady lengthening and weakening being observed.








A variety of hypotheses have been put forth to explain the observed trends. One

rationalization, invoked by Pauling98 and based on valence bond theory, involves "double

bond, no bond" resonance of the type depicted in Figure 2-1.

Table 2-3. Interatomic Distances and Dissociation Energies of Fluoroethanes

Ethane r (C-C). A Do (C-C). kcal mol-1 r (C-F). A Do (C-F). kcal mol-r1

CH3-CH3 1.532 90.4

CH3-CH2F 1.502 91.2 1.398 107.9

CH3-CHF2 1.498 95.6 1.343 Unknown

CH3-CF3 1.494 101.2 1.335 124.8

CH2F-CF3 1.501 94.6 Unknown 109.4 (CH2F)

CF3-CF3 1.545 98.7 Unknown 126.8

As the degree of geminal fluorination is increased, the number of such resonance

forms involving doubly bonded fluorine increases (0, 2, 6, and 12 in the case of CH3F,

CH2F2, CH3F, and CF4, respectively). This is supported by ab initio calculations at the

RHF/4-31G and 4-31G(d) levels,99-101 which illustrate back donation of electron density

from fluorine into the C-F o* orbitals. It is further observed that the overlap population

between the 2p orbitals of carbon and those of fluorine increases continually with

successive fluorination; in contrast, such carbon-chlorine overlap populations decrease

steadily from CH3CI to CCl4.

F F-
;L F ^ -----"-=

Figure 2-1. "Double Bond, No Bond" Resonance in Geminally Fluorinated Alkanes.


Alternative explanations based on hybridization schemes have also been
advanced. It is postulated that for electronegative elements bound to carbon,

rehybridization occurs causing an increase in the amount of p character directed toward








the substituent. Thus, in CH3F, the C-F bond possesses greater p character, with

greater s character in the C-H bonds. This rationale accounts not only for incremental

C-F bond strengthening, but also for the observed changes in geometry within the

fluoromethane series. Accumulation of p character in C-F bonds should lead to a

decrease in FCF bond angle, accompanied by HCH widening. This is consistent with

experimental observation (ZFCF in CF4, 109.5; CHF3, 108.7; CH2F2, 108.3;102'103 for

CH2F2, ZHCH = 113.70).103

Finally, a more recent argument has been advanced by Wiberg,104 on the basis

of Coulombic interactions between carbon and fluorine substituents. From charge-fitting

treatments based on calculated electrostatic potentials, a linear increase in positive

charge on carbon is observed, while the degree of negative charge on each of the

fluorine substituents remains quite constant. Thus, incremental fluorine substitution

strengthens not only new, but also previous, C-F bonds. This finding will be further

discussed in Chapter 3, as such ESP-derived charges calculated for larger fluorinated

systems as part of the present study were found to provide a cogent explanation for the

remarkable and contrasting effects of fluorination on the strengths of both C-C and a and

13 C-H bonds.


Structure. Bonding, and Reactivity in Unsaturated Systems


The most reliable structural and n-BDE data for the fluoroethylenes is provided in

Table 2-4.9'7 Vinylic fluorine substitution results in shorter C=C bonds than in the

parent hydrocarbon, and shorter C-F bonds than fluoroalkanes bearing the same

number of geminal or vicinal fluorine substituents. Ab initio investigations by Radom99

and others105107 attribute the C-F bond contraction to delocalization of fluorine 2p

electrons into the C-C n bond (depicted in resonance terms in Figure 2-2). Computed

atomic charges are consistent with this conclusion.








Table 2-4. Interatomic Distances, Angles, and n Bond Dissociation Energies of
Fluoroethenes.

CH2=CH2 CH2=CHF CH2=CFg CHF=CF2 CF2=CF,

r(C=C),A 1.339 1.333 1.316 1.309 1.311

r(C-F),A 1.348 1.324 1.336 1.319

ZHCH, deg. 117.8 114.7 119.3 -

ZHCF, deg. 111.3 114.0 -

ZFCF, deg. 109.7 109.1 112.6

7 D0, kcal mol' 63-64 Unknown 62.8 Unknown 52.3


The marked decrease in FCF bond angles has been rationalized by Bemrett108

and Kollman109 on the basis of hybridization arguments; Epiotis has advanced an

alternative explanation involving nonbonded attraction between fluorine atoms.110111'
+
>\ F F
F F

Figure 2-2. Fluorine 2p Electron Delocalization in Unsaturated Systems.


Photoelectron12 and electron attachment1"3 spectral data for the fluoroethylene

series are provided in Table 2-5. A significant lowering (over a range of 3.1 eV) of the a

MO energies is observed, with only a slight (ca. 0.3 eV) variation in the n MOs.

Stabilization of the a MOs is ascribed to extensive delocalization over the fluorine

substituents; such an interaction within the n system is diminished and counteracted by

strong C-F antibonding overlap.112 A steady increase in electron attachment energies

with successive fluorination can also be seen, attributed to destabilization of n* resulting

from an antibonding interaction with the fluorine 2p AOs.113

Heats of hydrogenation provided in Table 2-695 illustrate the reactivities of

fluorinated alkenes. In general, transformation of a polyfluorinated olefin into a saturated








Table 2-5. Vertical Ionization Potentials (I. P.) and Electron Attachment Energies (EA) for
the Fluoroethenes.

Ethene I.P., eV a I.P., eV E.A., eV

CH2=CH2 10.6 12.85 1.78

CH2=CHF 10.58 13.79 1.91

CH2=CF2 10.72 14.79 1.84

cis-CHF=CHF 10.43 13.97 2.18

trans-CHF=CHF 10.38 13.90 2.39

CF2=CHF 10.53 14.64 2.45

CF2=CF2 10.52 15.95 3.00

derivative is more exothermic than for the parent hydrocarbon. This arises from a

combination of the destabilizing effect of polyfluorination on double bonds and the

thermodynamic preference for gem-difluoro substitution at saturated carbon. The

deviation of CH2=CHF in Table 2-6 is explained by the preference of a single fluorine

substituent to reside at the vinylic position (Figure 2-3).114


Table 2-6. Heats of Hydrogenation of the Fluoroethenes.

Ethene AH (H,), kcal mol1

CH2=CH2 -32.6

CH2=CHF -29.7

CH2=CF2 -38.8

CF2=CF2 -45.7

Fluorine Non-Bonded Interactions in Reactive Intermediates


The r-donor ability of the fluorine substituent is reflected in its activating and

ortho, para-directing character in electrophilic aromatic substitution reactions,115

consistent with "3C NMR measurements of fluorobenzene, where shielding of these

positions is observed.1"6








H CH2F 12 H3C F 12 H3C H

H H AH = -3.34 kcal mol-1 H H AH = +0.65 kcal mol1- H F

54 Z-55 E-55


H3C F 12 H CF2H

H F AH =-2.5 kcal mol-1 H H

56 57

Figure 2-3. Thermodynamic Equilibria in Mono- and Difluoropropenes.


Delocalization of fluorine's nonbonded electrons into the vacant 2p orbital on

carbon more than compensates for its inductive withdrawal in a-fluoro carbocations,

resulting in net stabilization. In the gas phase, carbocation stability increases along the

series *CH3 < CF3 < *CH2F < CHF2 and *CH2CH3 << +CF2CH3 = +CHFCH3.117'118 The

+CF3 cation has been observed in the gas phase, with many others having been

successfully generated in solution.119121 In contrast, fluorination at the 1 position and

beyond destabilizes carbocations due to inductive effects; simple alkyl 13-fluoro

carbocations not benefiting from additional stabilizing factors have yet to be detected.

Electron pair repulsion in a-fluoro carbanions results in a strong preference for

pyramidal geometries, ab initio calculations122-124 predicting an FCF angle of ca. 99.5

and an inversion barrier of 119 kcal mol1 for CF3. Although fluorination does increase

C-H bond acidities in such pyramidal systems,125 a destabilizing effect is observed in

cases such as the 9-fluorofluorenyl anion (Figure 2-4, X = F) where coplanarity is forced

between the 2p orbitals on fluorine and the remainder of the R system.126 Fluorine

substitution in the p position stabilizes carbanions through a combination of inductive

and hyperconjugative effects, (Figure 2-5) the latter supported by X-ray crystallographic

data of perfluoroalkyl anion salts as well as through calculation.'22"127128











D X


58









Figure 2-4. Fluorine Destabilization in




F o


NaOCH3
CH3OH


H X


59
X kexc (reI

D 1

F 0.125

Cl 400

Br 700

Planar 9-Halofluorenyl Anions.





F


Figure 2-5. Negative Hyperconjugation in p-Fluorocarbanions.


Fluorine Steric Effects


The minimal spatial requirements of fluorine, the smallest non-hydrogenic

substituent, would imply a very minor steric impact on reaction kinetics and

thermochemistry. In most cases this is true; however, examples of steric inhibition in

reactions of fluoro-substituted systems do exist, typically in conformational and other

dynamic processes occurring via highly congested transition states. This is illustrated in

Figure 2-6;129'130 the meta ring flip in 61 (X = F) exhibiting the largest known rate

retardation induced by a single fluorine substituent.

The disparate behavior observed in the Cope rearrangements of d,I- and meso-

62 (Figure 2-7)131 provides a particularly striking example of a fluorine steric effect.

Transformation of d,l-62 to 63 proceeds via a typical chair-like transition structure, where

in meso-62 a "boat-like" structure is required for C1 C6 bond formation. The higher AH*








x x

q: )/


X = H, AG < 6 kcal morl'1 (340 K)
X = F, AG" = 11.1 kcal mol"1 (340 K)


kH / kF= 1011 (298 K)


Figure 2-6. Inhibition of Conformational Dynamics by Fluorine Substitution.


and positive ASt implies a dissociative, rather than concerted, transition state for

rearrangement of meso-62. This is induced by severe electrostatic repulsion between

the high charge densities of the terminal fluorine substituents, separated by less than the

sum of their van der Waals radii in the Cope transition state.132


H1 CF2
H CF2


H CF2
H OCF2


CF2
CF2


V F


AH* = 22.4 kcal molr1
AS* = -17.5 eu


F F

F F


CF2
CF2


meso-62 AHW = 49.5 kcal mol1 63
AS* = +8.1 eu
Figure 2-7. Chair- versus Boat-Constrained Cope Rearrangement Reactions of
Terminally Fluorinated Dienes d,I- and meso-62.


d,/-62








Steric effects are enhanced by perfluoroalkylation and branching. Cyclohexane
A values133 and modified Taft steric parameters1'34 demonstrate that CF3 is at least as
large as isopropyl; evidence exists135 to suggest that perfluoroisopropyl and tert-butyl are
comparable in size. The most remarkable example of the above affects is the existence
of perfluorinated radical 64, (Figure 2-8) found to be persistent by ESR even in the
presence of molecular oxygen.136 The astounding kinetic stability of 64 derives from
steric sheltering of the unpaired electron by the neighboring perfluoroalkyl groups.




2: 'F1 9

64
Figure 2-8. Scherer's Persistent Perfluoroalkyl Radical.

The Fluorine Substituent in Free Radicals

Early application of organofluorine radical chemistry was comprised of the chain-
mediated addition of polyhalomethanes and ethanes to olefins, first by Haszeldine137 and
soon thereafter by Tarrant.138'139 Subsequent relative rate studies by Stefani et al.140
followed by those of Tedder141'142 clearly demonstrate the contrasting behavior of
fluorinated and non-fluorinated radicals in their bimolecular additions to alkenes.
The fluorine substituent has a substantial effect on the structure of organic
radicals as well as dramatic, but comprehensible, alterations in hydrogen abstraction
and addition reactivity in comparison to hydrocarbon systems, resulting primarily from

fluorine's potent a-withdrawing character.

Structural Aspects

In contrast to the planar, n-type methyl radical, a-fluorination results in

increasingly pyramidal, a-type radicals, as indicated by electron paramagnetic








resonance measurements143 and ab initio theoretical studies.144147 Calculations by

Pasto at the UHF/4-31G level indicate barriers to inversion of 0.5, 6.8, and 25.1

kcal mol1 for CH2F, CHF2, and "CF3, respectively.147 Geometries of the fluoromethyl

radicals computed at the UHF/6-31G(d) level are provided in Figure 2-9.


I 90.0o I 101.1o I 105.80 I 107.60
II II





Figure 2-9. UHF/6-31G(d) Pyramidalization Angles of *CH3, "CH2F, "CHF2, and *CF3.


Inversion barriers in the series appear somewhat sensitive to the level of theory

employed and substantially increase with the inclusion of polarization functions. SCF

calculations by Dykstra145 employing a polarized double- basis result in an inversion

barrier of 33 kcal mol1 for "CF3; inclusion of electron correlation (QCISD(T)/

6-31G(d)//UHFI6-31G(d), present work) affords a value of 29.4 kcal mol-1. Such

successive deviation from planarity is due to a combination of effects; relief of It

repulsion between the singly-occupied orbital on carbon and the fluorine 2p electrons is

further reinforced by overlap between the carbon 2p and C-F antibonding orbitals. This

strong tendency for pyramidalization has been shown to be responsible for the low n

bond energy in tetrafluoroethylene (Tables 2-4 and 2-6).148

The structures of fluorinated C2 radicals have been theoretically probed by

Paddon-Row and Chen et al.149-152 Alkyl substitution induces slight pyramidalization as

observed in the ethyl radical, (Figure 2-10) due to hyperconjugation between the SOMO

and the staggered P3 C-H bond.153 Fluorination at the radical center exerts an effect

similar to that observed in the methyl series, while the structures of alkyl radicals are

found to be relatively insensitive to P-fluorination.









1 94.60 I 94.4
I I







105.7 / 105.9






Figure 2-10. UHF/6-31G(d) Pyramidalization Angles for Ethyl Radicals CH3CH2",
CF3CH2", CH3CF2", and CF3CF2.


Radical Stability


FMO theory dictates that for radicals bearing electronegative substituents with

lone pairs (F, OH, NH2, SH) an inductive, destabilizing influence exists, countered by

stabilization resulting from delocalization of the unpaired electron.154 Thus, in the a

sense, fluoroalkyl radicals are destabilized. Furthermore, the opposing i-stabilizing

effect of the fluorine lone pairs decreases with pyramidalization of the radical site, due to

diminished overlap with the 2p AO on carbon.

The progressive decrease in stability of alkyl radicals with a- or 3-fluorination has

been illustrated by Pasto,80147 in the form of calculated radical stabilization energies

(RSE) based on isodesmic reactions (Table 2-7). The aforementioned increase in C-H

bond dissociation energies along the fluoromethane series lends experimental support.

Although some degree of 3 C-F hyperconjugative interaction is observed in 2-

fluoro substituted radicals,152 such a contribution to overall radical stability is minor and

inductive destabilizing influences dominate, as indicated by experimental and theoretical

C-H BDEs for the 2-fluoroethanes (Table 2-8).








Table 2-7. Isodesmic Equation and Radical Stabilization Energies (RSE, kcal mol-1,
4-31G) for a- and -Substituted Systems. Positive Values Denote Radical Stabilization.

Xn'CH3-n + CH4 -* XnCH4-n + 'CH3 (2-1)

X RSE X RSE

F +1.64 CH3 +3.27

F2 +0.56 OH +5.73

F3 -4.21 OCH3 +5.30

CH2F +1.46 CN +5.34

CHF2 +0.16 NH2 +10.26

CF3 -1.34 *NH3 -4.07

SH +5.66

*SH2 -3.17

Table 2-8. Experimental and Theoretical C-H Bond Dissociation Energies (kcal mol'1)
of 2-Fluoroethanes.

CHCH,-H CH2FCH2-H CF2HCH2-H CFCH,-H
101.1ab 103.6' 106.7a.f

97.7c 99.6c 101.3c 102.0c

100.0d 104.3d

102.0e 104.le 105.9e 107.le
a Experimental Value; Reference 155. b Experimental Value, Reference 156.
c MP2/6-311G(d,p)//MP2/6-31G(d,p); Reference 157. d B3LYP/6-31G(d); Reference 91
and Present Study. e MP2/6-311+G(3df,2p)//MP2/6-31G(d); Reference 158.
f Experimental Value; Reference 159.

Organofluorine Radical Reactivity


As a result of the a-withdrawing character of the fluorine substituent and

interaction of the SOMO with C-F o* orbitals, fluoroalkyl radicals should possess lowered

SOMO energies and therefore exhibit enhanced SOMO-HOMO interactions in

comparison with reactions of their hydrocarbon counterparts. Experimental ionization








potential and electron affinity data, although sparse, has been compiled in a recent

review by Dolbier90 and demonstrates the greater absolute electronegativity of the

fluoroalkyl radicals. Calculated quantities, inferred from Koopmans' theorem160 or based

on radical-ion energy differences, follow the expected trend although quantitative

agreement is often lacking.161162 The combination of such FMO, geometric, and

enthalpy factors in hydrogen atom abstraction, intermolecular addition, and cyclization

reactions of fluoroalkyl radicals is now discussed.


Hvdroaen Atom Abstraction Reactions


A review by Tedder163 has underlined the importance of polar and enthalpic

factors in radical abstraction reactions. Activation parameters for abstraction by methyl

and trifluoromethyl radicals from a series of hydrogen donors are provided in Table 2-9.


Table 2-9. Arrhenius Parameters for Hydrogen Atom Abstraction by Methyl and
Trifluoromethyl Radicals.

CH3" CF3"
H-Donor Ea log A EAa oqg A

CH3-H 14.2 8.8 11.3 8.9

CH3CH2-H 11.8 8.8 6.9 8.4

(CH3)2CH-H 10.1 8.8 6.5 8.1

(CH3)3C-H 8.0 8.3 4.9 7.7

H-Cl 2.5b 5.Oc
a In kcal mol-1. b AHOXnn = -1 kcal mol-1. c AHOrxn = -3 kcal mol1.


In the first four examples, the decrease in activation barrier for both CH3* and

CF3 abstractions from alkanes are in line with the greater stability of the product radical;

in each case, the barrier to abstraction by CF3 is substantially lowered. In contrast,








abstraction of hydrogen atom from HCI by CF3* is much less facile, occurring with a

barrier twice that of CH3* in spite of its slightly greater exothermicity.

In collaboration with Lusztyk and Ingold at NRCC, LFP-determined rates of

perfluoroalkyl radical addition to a number of alkenes have been determined by the

Dolbier group.54'85'88'91 This has afforded, via competitive kinetic techniques, absolute

rate constants for hydrogen abstraction by the perfluoro-n-heptyl radical from a series of

donors, summarized in Table 2-10.54'86'87'89 For comparison, abstraction rate constants

for hydrocarbon n-alkyl radicals were provided in Table 1-1.

Table 2-10. Absolute Rate Constants for Hydrogen Atom Abstraction for Perfluoro-n-
heptyl Radicals. Rate Constants are at 303 K.

Et3SiH (TMS),Si(CH3)H nBu3GeH (TMS)SiH nBu3SnH PhSH

kH(n-C7F15"), 0.75 16 15 51 203 0.28
106 M1 s-1


For the first five donors in the series, a substantial rate enhancement is observed

over hydrocarbon radicals, ranging from 75-fold in the case of nBu3SnH to nearly 900-

fold in the case of Et3SiH, after slight temperature correction to 303 K. Although

hydrogen transfer to perfluoroalkyl radicals is more exothermic (see the previous

discussion on C-H BDEs) this is insufficient to account for a nearly three order of

magnitude difference in rate constants. Furthermore, thiophenol, an excellent donor to

hydrocarbon radicals, is found to suffer a greater than 400-fold decrease in transfer rate

to perfluoroalkyls.

These characteristics are explained by the ability of the radical-donor pair to

accommodate charge transfer interactions in the hydrogen transfer transition state, as

previously discussed in Chapter 1. The relatively electropositive donor agents

(stannanes, germanes, and silanes) lead to a more favorable polarity matchup with the

electronegative perfluoroalkyl radical than with the hydrocarbon. Conversely, transfer

from the more electronegative thiophenol results in a non-polarized or polarity-








mismatched transition state. The existence of such polar effects were confirmed by a

correlation between rates of hydrogen transfer to perfluoro-n-heptyl radicals by a series

of substituted thiophenols, versus their Hammett a+ constants (Figure 2-11).87 The

resulting p value of -0.56, when compared to that obtained in the case of tert-butoxyl

(-0.30)164 again reflects the high electrophilicity of perfluoroalkyl radicals.


0.6 -

0.4 -

0.2

0.0

-0.2 -

-0.4 7--
-1.00 -0.75 -0.50 -0.25


0.00 0.25 0.50 0.75 1.00


Figure 2-11. Hammett Plot for Hydrogen Abstraction from para-Substituted Thiophenols
by Perfluoro-n-heptyl Radical.


Intermolecular Addition Reactions


Relative rates of addition of small fluorocarbon radicals to fluorinated and non-

fluorinated olefins have been extensively investigated by Tedder and Walton,

culminating in a critical review in 1980.52 Table 2-11 illustrates the relative reactivity of


Table 2-11. Relative Rates of Addition of the Fluoromethyl Radicals to Ethene and
Tetrafluoroethene.


Radical

"CH3

"CH2F

*CHF2

"CF3


kdd (CF4)/ka (CdH4) (437 K)

9.5

3.4

1.1








the fluoromethyl radicals towards ethylene and tetrafluoroethylene. Additional studies,

utilizing a number of unsymmetrical methyl- and trifluoromethyl-substituted olefins,

solidified the ascription of relative rates and regiochemical preferences to a combination

of polar and steric influences.

Absolute rates of addition of perfluoroalkyl radicals to alkenes have been

determined by laser flash photolysis, a subset of the data acquired to date presented in

Table 2-12.90 The dramatic rate acceleration enjoyed by the perfluoroalkyl radicals

versus hydrocarbon n-alkyls is readily apparent, ranging from factors of 300 to 30,000 in

the case of the heptafluoropropyl radical addition to electron-rich alkenes. The rate of

addition to pentafluorostyrene, in contrast, is increased by only a factor of 42.


Table 2-12. Absolute Rate Constants for Addition of Perfluoroalkyl Radicals to Alkenes.

kdd, 106 M-1 s- (298 K)
Alkene C3F7 C7F15 "C8F17 CF CF3 RCH*

Styrene 43 46 46 79 53 0.12

a-Methylstyrene 78 89 94 87 0.059

13P-Methylstyrene 3.8 3.7 7.0 17

Pentafluorostyrene 13 23 26 0.31

4-Methylstyrene 61

4-Methoxystyrene 65

4-Chlorostyrene 36

4-(CF3)Styrene 35

1-Hexene 6.2 7.9 16 0.0002


Such enhancements may potentially be attributed to a combination of factors.

Relative reaction enthalpies should play a role, as a stronger C-C bond (from CH3-CH3

versus CF3-CH3 BDE data in Table 2-3, ca. 11 kcal mol-1) is formed upon perfluoroalkyl








addition. However, only slight increases (factors of 5-7) in addition rates of the

perfluoroalkyls to styrene versus 1-hexene are observed, despite the greater

exothermicity (ca. 16 kcal mol1) of the former. This demonstrates the relatively minor

importance of reaction enthalpy, in accord with the early transition states expected for

radical addition.

The pyramidal, a-character of the fluoroalkyl radicals should afford a kinetic

advantage (further discussed in Chapter 3) over the planar hydrocarbon, the LFP-

measured rate of addition of 1,1-difluoropentyl radical to styrene, 2.7 x 106 M'1 s1,88

giving rise to a 22-fold enhancement relative to n-alkyls.

The primary factor responsible for such striking increases in reactivity is believed

to be charge transfer influences (Figure 2-12) similar to those postulated for hydrogen

atom transfer reactions. The lowered SOMO energies of the perfluoroalkyl radicals and

resulting enhanced SOMO-HOMO interactions with alkenes in the addition transition

state leads to substantial rate enhancement. Supporting evidence, in the form of a

Hammett relation involving para-substituted styrenes and correlations between addition

rates and alkene ionization potentials, has been offered.54 Ab initio computations concur

with the experimental findings and are discussed in Chapter 3.


6-
CF3(CF2)nCF;
6+



Figure 2-12. Polarized Transition State for Addition of Perfluoroalkyl Radicals to
Alkenes.


Intramolecular Cyclization Reactions


Cyclopolymerization of fluorinated monomers has long been known as a means

for the generation of macromolecular materials with unique physical properties.165"168








However, despite the widespread popularity of the unimolecular 5-hexenyl radical

cyclization for the generation of five-membered rings, synthesis of fluorinated cyclic

products utilizing radical methodology has received limited attention.90 This is somewhat

surprising, in light of the current interest and demonstrated importance of fluorinated

analogues and mimics of pharmaceutical and agricultural agents.169

Until recently, no quantitative kinetic data existed for cyclization reactions of

fluorinated radicals. With the competitively-determined rate constants of hydrogen

abstraction by perfluoroalkyl radicals serving as basis reactions, rates of cyclization of a

number of fluorinated 5-hexenyl systems have now been determined,86,89,'170 examples of

which are provided in Figure 2-13.

Most obvious of the data is the remarkable rate acceleration in the 5-exo

cyclizations of 65 and 68, occurring with 163- and 41-fold increases relative to the parent

hydrocarbon I and ascribed to charge-transfer effects analogous to those occurring in

bimolecular additions. Consequently, with no such polarity matchup in 71 (which

involves cyclization of an electrophilic radical onto an electron-deficient double bond)

only a minor increase in rate is observed. The 5-exo cyclization rates of 71 and 74, in

line with those other hexenyl systems bearing fluorinated double bonds,170 demonstrate

the lack of kinetic impact of vinylic substitution.

Especially surprising is the degree to which 65 and 68 undergo 6-endo closure,

the former with a 1040-fold, and the latter a 700-fold acceleration relative to 1. This is

further discussed in Chapter 4, in which the reactivities of a series of lightly-fluorinated 5-

hexenyl systems are investigated.

In the proceeding discussions of hydrogen abstraction, addition, and cyclization

reactivity, the observed kinetic behavior was found to be due to a combination of

geometric and polar effects induced by polyfluorination. Related studies of partially

fluorinated alkyl radicals, addressed in the next two chapters, will aid in separating the

effects of fluorination at the a and p positions and beyond, providing insight into the








extent to which the effects of perfluorination on the reactivity of organic radicals are the

sum of their parts.


Table 2-13. Some Absolute Rate Constants for 5-exo and 6-endo
Fluorinated 5-Hexenyl Radicals. Rate Constants are for 303 K.


Cyclization Reaction


kc5, 105 s-1


Cyclization of


kc6, 105s1


F F2
C,-CF2
F2

65


^ CF2
F2CC.. CF2
F2


F
F F CF2
F2C...CF2
F2


F
F ,F.
F2C-cJ
F2
74


CF2
F2C-CF2

66


F2C CF2

F2C-CF2


F CF2
- F2C CF2
F2C-CF2


F CF2
F2C 2
F2C-6


75 76


0.05


440


+ 0

3



F2C...CF2
F2
67


+ r-CF2
F2C-.c.CF2
F2

70

F
F2C CF2
+ "- I I
F2C,.. CF2
F2

73

F
F2C CF2
+ C(F2


N/A


N/A


'U








Conclusion


An introduction to the general structural and reactivity characteristics imparted by

fluorine substitution in organic systems has been provided. Such substitution can either

be stabilizing or destabilizing, depending on the nature of the ground state molecule,

intermediate, or reaction in question.

A summary of the first absolute kinetic data obtained for reactions of fluorinated

radicals has been presented. Results of these studies, involving per- or otherwise highly

fluorinated systems, demonstrate the combined influence of polarity, structural, and

enthalpic factors. The following studies of the addition, hydrogen abstraction, and

cyclization kinetics of partially fluorinated alkyl radicals will serve to dissect the relative

magnitudes of these influences on organic radical reactivity.














CHAPTER 3

THE REACTIVITY OF PARTIALLY FLUORINATED RADICALS
IN INTERMOLECULAR ADDITION AND HYDROGEN ABSTRACTION REACTIONS

Introduction


Initial studies of the addition rates of some partially-fluorinated radicals to

alkenes88 demonstrated observable rate enhancements relative to n-alkyls, though not

nearly as great as those of perfluoroalkyl systems. Such reactivity derives from the

combination of structural and polar characteristics induced by fluorine substitution.

The current study extends the amount of absolute rate data assembled for the

addition of partially fluorinated radicals to olefins. In addition, through competitive kinetic

techniques, absolute rate constants for hydrogen abstraction from tri-n-butyltin hydride

(nBu3SnH) have been determined. Such kinetic data is necessary for the determination

of absolute rates of cyclization of partially fluorinated 5-hexenyl radicals, discussed in

Chapter 4, and allows for the partitioning of the gross reactivity characteristics of

perfluoralkyl radicals into the separate influences of a, P3, and y fluorination. This is

accomplished by use of the following systems: RCH2CH2CF2' (a,a-difluoro),

RCH2CF2CH2 (P3,P13-difluoro), RCH2CF2CF2* (a,cc,3,,3-tetrafluoro), CF3CF2 (a,a,p13,13,13-

pentafluoro), and RfCH2CH2 (y-perfluoro).

The existence of charge transfer stabilization in the transition states for

fluorinated radical addition to alkenes is corroborated by ab initio calculations, and the

thermodynamics of C-H and C-C bonding and radical stabilization in hydrofluorocarbons

rationalized on the basis of Coulombic interactions.








Precursor Syntheses and Competitive Kinetic Studies


In each of the competitive kinetic runs, hydrofluorocarbon radicals were

generated from bromide or iodide precursors by photoassisted C-X bond homolysis.

These radicals subsequently underwent competitive trapping with known, varying

concentrations of styrene or nBu3SnH, adjusted to ensure pseudo-first-order kinetic

behavior and to allow for accurately measurable amounts of trapping products, as

depicted in Figure 1-17 and in greater detail below.


1.,1-Difluorohex-l-vl Radical (77)


The synthesis of bromide precursor 80 was achieved in two steps (Figure 3-1) in

a straightforward manner. Copper(l)-mediated addition of dibromodifluoromethane to 1-

pentene (78), based on a modification by Gonzalez et al.171 of a procedure by Burton

and Kehoe172 afforded 1,3-dibromo-1,1-difluorohexane 79 in typical yield.

Regioselective displacement of the internal bromine was accomplished via sodium

borohydride reduction in DMSO, providing precursor 80 contaminated with a small

amount of overreduction product 81. Pure samples of each were obtained by

preparative GC separation, the former utilized in the competition run and the latter for

spectral comparison with kinetic NMR data.


CF2Br2 Br
s --CF2Br
(CH3)3COH, H2NCH2CH2OH
CuCI (cat.) (57.3%)
78 79


NaBH4
----- ^^^CF2Br + CF^*2Hl
DMSO
(62.1%)
80 81

Figure 3-1. Preparation of 1-Bromo-1,1-difluorohexane 80 and 1,1-Difluorohexane 81.









CF2Br hv
nBu3SnH
80 CAD6


nBu3SnH


kadd


F2


83

Figure 3-2. kH I kaw Competitive


77CF
L ~77


+ nBu3Sn"


81 F2H
81


F2
S ./ /^C. nBu3SnH
F 1 ______-- _

82


+ nBu3Sn"


Kinetic Scheme for 1,1-Difluorohex-l-yl Radical 77.


Photolysis of 80 as a C6D6 solution in the presence of an excess of nBu3SnH and

styrene (Figure 3-2) afforded intermediate radical 77. Subsequent entrapment by these

agents (both irreversible processes) yielded 81 and 82, respectively; the latter further

trapped by nBu3SnH to yield 3,3-difluoro-l-phenyloctane 83. Throughout the course of

the reaction, nBu3Sn* radicals are generated to propagate the chain process via

abstraction of halogen from precursor 80.

Product ratios for varied concentrations of trapping agents were determined by

19F NMR analysis according to the pseudo-first-order relation in Equation 3-1,


[81]
[83]


[kH] [ 77] [nBu3SnH]
[kadd] [ 77] [ C6H5CH=CH2 ]


(3-1)


a plot of product ratios obtained for each data point versus that of trapping agents

affording the ratio kH I kadd. The stability of trapped products under the reaction

conditions and lack of appreciable side reactions were demonstrated by the high








conversion of precursor 80 to 81 and 83 versus an internal standard of a,a,a,-

trifluorotoluene, (< -63.24) indicating in turn the high efficiency of the chain process and

reliability of the kinetic results. A partial 19F NMR spectrum of the first of six data points

is provided in Figure 3-3, a doublet of triplets (-CF2H, 4 -116.0) observed for 81 versus

an overlapping triplet of triplets at -99.1 (-CF2-) for 83. Full kinetic data and yields are

given in Table 3-1 below, the plot of which located in Figure 3-4. The slope of the line

(3.39 0.02) in conjunction with the known absolute rate constant for addition of













.. .... .. ... .. ......) ---.--.-.-.. .

-98 -99 -100 -101 -102 -103 -104 -105 -106 -107 -108 -109 -110 -111 -112 -113 -114 -115 -116 -117 -118

Figure 3-3. Partial 19F NMR Spectrum of Data Point 1 for kH I kaW Competition of 1,1-
Difluorohex-1-yl Radical 77.

Table 3-1. Competitive Kinetic Data for kH I kadd Competition of 1,1-Difluorohex-l-yl
Radical 77.

[801 CH5CH=CH2 1 [F nBuSnH 1 / [C6H5CH=CH2 1 f 811/ F 83 1 % Yield

0.094 2.01 0.719 2.30 95

0.094 1.81 0.847 2.73 95

0.094 1.61 1.01 3.26 96

0.094 1.41 1.21 3.93 97

0.094 1.21 1.49 4.88 96

0.094 1.01 1.87 6.21 95






53


1,1-difluoropent-l-yl radical to styrene, 2.7 ( 0.5) x 106 M1 s-1, resulted in an absolute

rate constant kH of 9.1 ( 1.7) x 106 M1 s1. It should be noted that the accuracy of such

derived kH values can be no better than those reported in the LFP determinations of kadd,

the error estimates in the former reflecting both the least-squares fit of the line and

propagated error of the latter. Synthesis of styrene adduct 83 was performed by classic

means for characterization and spectral comparison (Figure 3-5).


0.50 0.75 1.00 1.25 1.50
[ nBu3SnH ] / [ C6H5CH=CH2 ]

Figure 3-4. Plot of the Data in Columns 3 and 4 of Table 3-1.


MgBr


Na2Cr2O7 / H2S04
Et20


F2


(68.0%/)


1.75


2.00


1. CH3(CH2)4CHO

2. H30*


0


89.0%)
87


Figure 3-5. Preparation of Styrene Adduct 83.


NN Br


Mg
EtBO


OH


(78.7%)
86


DAST
CH2CI2








2.2-Difluorohex-l-vl Radical (88)


In accordance with literature precedent,173 a-bromination of 2-hexanone in the

presence of urea in acetic acid selectively afforded 1-bromo isomer 90 in 72.2% yield.

Subsequent treatment with diethylaminosulfurtrifluoride (DAST) provided bromo

precursor 91, originally purified by preparative GC for use in the kinetic study. However,

a sluggish chain reaction (further hindered by the strong UV absorption of the excess

styrene present in the kinetic samples) led to the preparation of iodo precursor 92

(Figure 3-6) via Finkelstein transformation at elevated temperature.


0 Br2 0 DAST
CH3CO2H, H2NCONH2 "J -CH2Br CHCI3
(72.2%)
89 90
Nal
.'/^^ C-cCH2Br Nal C c.CH21
F2 (CH3)2CO F2
(59.7%) (87.5%)
91 92

Figure 3-6. Preparation of 1-lodo-2,2-Difluorohexane, (92) Precursor to 2,2-Difluorohex-
1-yl Radical 88.


10 -
9 Coefficients:
-" m = 27.3
M 8 b =-0.065
r 2 = 0.998
V) 7-

6-

5 I --- I I -
0.20 0.22 0.24 0.26 0.28 0.30 0.32 0.34

[ nBu3SnH ] / [ C6H5CH=CH2]

Figure 3-7. Plot for kH I kad Competition of 2,2-Difluorohex-l-yl Radical 88.








The competition plot for kH I kadd determination is found in Figure 3-7; raw data for

this and all remaining experiments in this chapter may be found in Chapter 5. With the

absolute rate constant for addition of 2,2-difluoropent-l-yl radical to styrene known from

LFP experiments, a kH value of 1.4 ( 0.5) x 107 M1 s1 was determined. Authentic

samples of 93 and 98 were prepared for spectral comparison and characterization as

illustrated below in Figure 3-8.


c-^_/I .^CH2Br
F2
91


nBu3SnH
AIBN
C6H6


" c .-C H 3
F2

93


N. Br


Mg
Et20


. YMgBr
( ,;5"-


1. CH3(CH2)3CHO

2. H3O+


OH
(71.1%)
96


DAST
CH2CI2


Na2Cr2O7 / H2S04
Et2O


0
(83.7%)
97


F(66.3%)
(66.3%)


Figure 3-8. Preparation of Hydrogen Abstraction Product 93 and Styrene Adduct 98.


1.1.2,2-Tetrafluorobut-l-vl (99) and 1.1,2.2-Tetrafluorohex-l-yl (100) Radicals


At the time of this study, no absolute rate constants for addition of a 1,1,2,2-

tetrafluorinated radical to alkenes had been determined. Thus, in order to obtain a kH

value for such a system, a precursor suitable for absolute kad measurements was first

required. Bromide precursors, although in most cases sufficient for competition









experiments, are ineffective under the LFP operating conditions used in our kad

determinations (308 nm excimer laser pulses) due to their relatively short wavelength

chromophore and small extinction coefficient (for 102, Ew = 37.3 M1 cm1, ;,ax = 218

nm, cyclohexane solvent). The instability of 0-acylthiohydroxamic esters of the

perfluoroalkanoic acids has been noted by Barton.174 With neither these nor diacyl

peroxide precursors lending themselves to isolation and / or shipment to the NRCC in

Canada, the synthesis of a suitable iodide precursor was undertaken.

1-Bromo-1,1,2,2-tetrafluorohexane (102, Figure 3-9) was prepared in one step

from 6-bromo-5,5,6,6-tetrafluorohex-1-ene (supplied by Halocarbons, Inc.) and its

transformation to the corresponding iodide or carboxylic acid (the latter of which could be

converted to the iodide via Hunsdieker methodology) attempted under a variety of

conditions (Figure 3-10).

1. BH3 Me2S
,- c,.-CF2Br 2. C5HlCO2H Nc--CF2Br
F2 Tetraglyme F2
(79.6%)
101 102

Figure 3-9. Preparation of 1-Bromo-1,1,2,2-tetrafluorohexane (102).


Although the conversion of perfluoroalkyl iodides to their lighter analogues is

known in the literature,175176 downward transhalogenation of perfluoroaliphatic halides is

exceedingly difficult. Indeed, all attempts at conversion of 102 to the corresponding

iodide were met with failure. Perfluoroalkyl Grignard reagents, generated either directly

(and in low yield) or via transmetallation by an alkylmagnesium halide, utilize iodide

starting materials.177178 Investigations of perfluoroalkylzinc halides by Miller179 resulted

in a similar conclusion; perfluoro-n-propyl iodide may be converted to the organozinc

reagent in ca. 60-80% yield (as determined by aqueous hydrolysis or capture with

halogen electrophiles) after a brief induction period. On the contrary, reaction of

heptafluoro-n-propyl bromide with zinc in 1,2-dimethoxyethane afforded no product after









1) Mg
2) 12
Et2O


N.R.


1) t-BuLi, -80C
2) 12
SBEt20O

1) t-BuLi, -100C
2) 12
Et20

1) MeLi, -100C
2) 12


Et20

1) MeLi, -100C
2) CO2
Et2O


ICF2CF2I

tBuOH, H2NCH2CH2OH
CuCI (cat.)


Et3N 3HF / NIS
CH2CI2 0 C


F

F
103


It

if

to




It


N.R.


N.R.


103

Figure 3-10. Attempted Preparation of lodo Analogue of 102.


73 hours at 900 C; a 60% yield of the zinc derivative was obtained (inferred via

hydrolysis) after a period of 1.5 months. Attempted lithiation of 102 at low temperature

resulted in the formation of p-fluoride elimination product 1,1,2-trifluoro-l-hexene, (103)

as identified by its 19F NMR spectrum. Additionally, neither Cu(l)-induced addition of

1,2-diiodotetrafluoroethane to 1-pentene nor iodofluorination of 103 utilizing the

triethylamine trihydrofluoride / N-halosuccinimide methodology of Alvernhe et al.180 were

successful.


/~c/\.CF2Br
F2
102


,,-- c.CF2Br
F2
102


F

F
F


\


L








Synthesis of a C4 iodide was achieved through modification of a DuPont literature

procedure,181 whereby 1,4-diiodo-1,1,2,2-tetrafluorobutane was produced in 40.1% yield

via direct thermal addition of 1,2-diiodotetrafluoroethane to ethylene (Figure 3-11). DBU-

induced elimination of hydrogen iodide in ether followed by diimide hydrogenation with

hydrazine-hydrogen peroxide in methanol afforded tetrafluoroiodo precursor 107 in 2.8%

overall yield after preparative GC purification, which was sent to the NRCC for absolute

kadd measurements.


ICF2CF21

104


HN=NH
CH3OH


H2C=CH2
A


I c .CF21
F2
(40.1%)
105


DBU
Et20


'-c.CF21
F2
(66.0%)
106


c---CXF21
F2
(10.7%)


Figure 3-11. Preparation of 1,1,2,2-Tetrafluoro-l -iodobutane (107).


0.4 0.5 0.6 0.7 0.8 0.9 1.0 1.1 1.2 1.3 1.4 1.5
[ nBu3SnH ] / [ C6H5CH=CH2 ]

Figure 3-12. Plot for kH I kaw Competition of 1,1,2,2-Tetrafluorohex-l-yl Radical 100.


Bromide 102 was utilized in the kH I kadd competition, the kinetic plot for which

given in Figure 3-12. An authentic sample of styrene addition product 110 was prepared





59


by slow syringe pump addition of nBu3SnH to a heated, irradiated solution of 102 and

styrene in benzene (Figure 3-13).


,c,.CF2Br
F2
102


nBu3SnH
SC6H6


C61H5CH=CH2
nBu3SnH
C6H6


nBu3SnH


,--,,c-CF2H
F2
108

F2


109

F2

F2


110

Figure 3-13. Preparation of Hydrogen Abstraction Product 108 and Styrene Adduct 110.


2-f[Perfluorohexylleth-l-vi Radical (111)


2-[Perfluorohexyl]-1l-iodoethane 112 was provided as a gift from Prof. Neil Brace.

The kH I kadd competition plot is found in Figure 3-14; hydrogen abstraction product

1-[perfluorohexyl]ethane (113) and styrene adduct 1-[perfluorohexyl]-4-phenylbutane

(117) were prepared as shown in Figure 3-15.


7-
6 Coefficients:
% m = 16.0
5 b =-0.371
r 2 = 0.999
qE4-

3-

2 1I I
0.20 0.25 0.30 0.35 0.40 0.45
[nBu3SnH ] / [ C6H5CH=CH2 ]

Figure 3-14. Plot for kH I kadd Competition of 2-[Perfluorohexyl]eth-1-yl Radical 111.









C6F13CH-2CH21

112


nBu3SnH
A
C6H6


+ N


115


C6F13 -


(70.0%)
117


C6F13-CH2CH3

113


Et3B (cat.)_


C6F13

(82.8%)
116


Figure 3-15. Preparation of Hydrogen Atom Abstraction Product 113 and Styrene
Adduct117.

Pentafluoroethyl Radical (118)


lodopentafluoroethane (119) was obtained from PCR, Inc. Due to the high

volatility of both this precursor (bp 12-13 C) and hydrogen atom abstraction product

pentafluoroethane (120, bp -48.5 C) 119 was handled as a solution in degassed C6D6.

The kH I kadd competition experiment (Figure 3-16) was performed in tubes which were

quickly flame-sealed upon injection of an aliquot of the chilled precursor stock solution.


0.40 0.45 0.50 0.55 0.60 0.65 0.70 0.75 0.80
[ nBu3SnH ] / [ C6H5CH=CH2]
Figure 3-16. Plot for kH / kadd Competition of Pentafluoroethyl Radical 118.


C6F131


114


NaBH4
DMSO








Samples of 120 and 122 were prepared under free radical conditions for characterization

purposes as shown in Figure 3-17.


CF3CF2H
hv
nBu3SnH


119 C6eH5CH=CH2 F2 ]
L 121




nBu3SnH F3C2'
C61-16 121

F2
nBu3SnH F3cC- -


122

Figure 3-17. Preparation of Hydrogen Atom Abstraction Product 120 and Styrene
Adduct 122.

Discussion

Absolute rate constants for addition and hydrogen atom abstraction for systems

studied in this chapter are provided in Table 3-2. For comparison, data for hydrocarbon
(n-pentyl) and perfluoroalkyl (perfluoro-n-heptyl) radicals are included.

It is seen from the data that the reactivity trends for partially fluorinated radical

additions to styrene are generally adhered to in hydrogen abstractions. However, the

actual rate constants for the latter are observed to differ (on the average, by a factor of
11) from those for addition, and span a narrower range. The decrease in abstraction
rate ratios as a function of substitution is due to the proximity (within an order of
magnitude) to diffusion control for the more reactive radicals 100, 118, and 127.

C-H and C-C bond dissociation energies of the fluoroalkanes should reflect the
relative thermochemistry of hydrogen atom abstraction and addition to alkenes by their
respective radicals. However, very few experimentally determined BDE values for such








Table 3-2. Absolute Rate Constants for Hydrogen Abstraction from Tributyltin Hydride
and Addition to Styrene by Partially Fluorinated Radicals. Rate Constants are for 298 K.
Radical k (M-1 s-1) kk (M-1 s1) kH

RCH2CH2 1.2 xlO 105 b 1 2.4 X106 c 1
(123) a

RCH2CF2" 2.7 (0.5) x 106 e 22.5 9.1 ( 1.7) x 106 f 3.8
(124, 77) d

RCF2CH2" 5.2(1.8)x 105 e 4.3 1.4 (0.5) x 107 f 5.8
(125,88) g

RCF2CF2" 2.0 (0.1) x 107' 167 9.2 (0.8) x 107 38
(99,100) h

RfCH2CH2" 1.3(0.2) x 105 e 1.1 2.1 (0.3) x 106' 0.9
(126,111)'

CF3CF2= 7.9 (1.0) x 107 658 3.2 ( 0.3) x 108 133
(118)

C7F15* 4.6 (0.6) x 107J 383 2.0 ( 0.3) x 108j 83
(127)
a R = C3H7. b Reference 182, After Modification for Temperature and Other Factors in
Table III of Reference 183. c Reference 27. d For kadd Experiment, R = C3H7 (124); for
kH Experiment, R = C4H9 (77). e Reference 88. Reference 91 and Present Study. g For
kadd Experiment, R = C3H7 (125); for kH Experiment, R = C4H9 (88). h For kadd
Experiment, R = C2H5 (99); for kH Experiment, R = C4H9 (100). For kadd Experiment,
Rf = C4F9 (126); for kH Experiment, Rf = C6F13 (111). 1 Reference 54.

systems have been reported. Theoretical studies by Boyd157'184 at the MP2 level have

provided reasonably accurate C-H and C-C BDEs for C2 hydrofluorocarbons. A more

recent investigation based on isodesmic reactions by Marshall and Schwartz,158

published after the completion of the present study, has yielded C-H BDEs of

appreciably high quality for some linear (C2) and branched (up to C4) polyfluoroalkanes.

Out of interest in determining additional C-H and C-C BDEs for larger (through C4)

fluorinated n-alkyls, the geometries of a series of partially fluorinated ethanes, propanes,

and butanes, along with their respective radicals generated from terminal C-H or C-CH3

bond cleavage, were optimized at the hybrid density functional level. This DFT method

was chosen due to its implicit consideration of electron correlation at only slightly greater








computational expense than that of Hartree-Fock theory. Utilizing the three-parameter

exchange functional of Becke'85 and the correlation functional of Lee, Yang, and Parr186

with the 6-31G(d) basis, bond dissociation energies obtained in this "direct" fashion were

found to be, in the cases where such values are known, within 1-3 kcal mol1 of those

determined experimentally and in comparable or better agreement with experiment than

the MP2/6-31 1 G(d,p) values obtained by Boyd for C2 systems. Tabulated experimental

(where known) and calculated C-H and C-C BDE values are provided below in Tables

3-3 and 3-4, respectively.


Table 3-3. Theoretical and Experimental C-H Bond Dissociation Energies.

C-H Bond Calculated BDE. kcal mol1 a Experiment

CH3CH2-H 100.0 101.1 + 1 e
97.7 c
102.0 d

CH3CF2-H 97.4 b 99.5 2.5f
97.0 c

CF3CH2-H 104.3 b 106.7 1'
102.0 c
107.1 d

CF3CF2-H 99.5 b 102.7 0.5
99.7 c
104.6 d

CH3CH2CH2-H 100.3 b 100.4 0.6 e

CH3CH2CF2-H 97.7b

CH3CF2CH2-H 103.1 b

CH3CF2CF2-H 100.1 b

CF3CH2CH2-H 101.4 b

CF3CF2CH2-H 103.8 b
a Reported as Do (298.15 K). b B3LYP/6-31G(d); Reference 91 and Present Study.
c MP2/6-311 G(d,p)//MP2/6-31 G(d,p); Reference 157. d MP2/6-311 +G(3df,2p)//MP2/
6-31G(d); Reference 158. e Reference 187. f Reference 95.








Table 3-4. Experimental and Theoretical C-C Bond Dissociation Energies.

C-C Bond Calculated BDE. kcal mol1 a Experiment

CH3-CH3 89.4 b 90.4 0.2 d
90.6 c

CF3-CH3 99.6 b 101.2 1.1 d
103.3 c

CH3CH2-CH3 86.3 b

CH3CF2-CH3 91.4 b

CF3CH2-CH3 91.4 b

CF3CF2-CH3 95.5b

CH3CH2CH2-CH3 86.7 b

CH3CH2CF2-CH3 91.6 b

CH3CF2CH2-CH3 89.9 b

CH3CF2CF2-CH3 95.4 b

CF3CH2CH2-CH3 87.8 b
a Reported as Do (298.15 K). b B3LYP/6-31G(d); Reference 91 and Present Study.
c MP2/6-311G(d,p)//MP2/6-31G(d,p); Reference 157. d Reference 95.


Inspection of the data reveals interesting trends within both the C-H and C-C

BDE series. From Table 3-3, it is observed that a-fluorination results in a weakening of

C-H bonds, on the order of 1-3 kcal mol1, as predicted by the various levels of theory

and supported by experiment in the case of ethane versus 1,1-difluoroethane.

Conversely, P-fluoro substitution results in a 3-5 kcal molr1 increase in terminal C-H

BDEs. Furthermore, strengthening of C-C bonds (Table 3-4) is observed for all systems

examined, whether substituted at the a, P3, or even (albeit diminished) Y position. An

explanation for this behavior, consistent with the given BDE and other thermochemical

data, is provided later in the chapter.

With the relative thermodynamics of C-H and C-C bond formation investigated,

attention is turned to polarity effects. As mentioned previously, rate constants for








addition of the perfluoroalkyl radicals to subtituted styrenes (Table 2-12) are observed to

increase with decreasing ionization potential (styrene, IP 8.43 eV; a-methylstyrene, IP

8.19 eV; pentafluorostyrene, IP 9.20 eV).188 In contrast, 1,1-difluoropentyl radical (124)

was found to add with rates equal, within experimental error, for all three olefins.88 From

this observation, along with derived absolute electronegativities for the radicals CH3*

(4.96), CH2F" (4.73), CHF2" (4.91), and CF3" (5.74) it was concluded that a-fluoro

substitution alone does not impart electrophilicity to an alkyl radical, and may instead

give rise to nucleophilic behavior.88

A series of theoretical studies by Wong et al.48'49,189,190 have assessed the degree

of charge transfer interaction in the transition states for methyl, hydroxymethyl,

cyanomethyl and tert-butyl radical additions to a series of monosubstituted olefins.

Based in part on the computation of partial charges, it was concluded that the addition of

methyl radical was governed primarily by enthalpic effects, with no evidence for

nucleophilic character arising from either Mulliken or Bader-based charge-fitting

schemes. Hydroxymethyl and tert-butyl were found to be nucleophilic, with cyanomethyl

exhibiting substantial electrophilicity.

In order to determine such tendencies for the fluorinated radicals under

investigation, transition structures for the addition of hydrocarbon and fluoro-substituted

ethyl radicals to ethylene and C, of propene have been located at the UHF/6-31G(d)

level. Partial atomic charges were then computed, based on fits to the electrostatic

potential at points selected according the Merz-Kollman-Singh scheme.91'192

Previous studies by Houk et al. utilizing the 3-21G basis found that addition of

ethyl radical to ethylene occurs preferentially via a gauche conformation, with an

incipient C-C-C-C dihedral angle of ca. 600.46 Such gauche and anti structures

computed at the UHF/6-31(d) level are depicted in Figure 3-18, the former ca. 0.1

kcal mol'1 lower in energy than that of (b). Inclusion of electron correlation at the spin-








projected PMP2/6-311G(d,p)//UHF/6-31G(d) level increases this energy difference to ca.

0.3 kcal mol1.


103.80 \%
2.227 A \ 1100
3)o


103.3
\
2.231 A 109.90
\ 109n9


Figure 3-18. UHF/6-31G(d) (a) Gauche and (b) Anti Transition Structures for Addition of
Ethyl Radical to Ethylene. Relevant Geometrical Parameters are Shown.


Consistent with UHF/3-21G results, the preferred mode of addition of the ethyl

and radical to C1 of propene involves a gauche arrangement of the radical with the

alkene C=C i bond and a transoid orientation of the methyl group of the radical with

respect to that of the olefinic C2 carbon (Figure 3-19). This 'gauche-transoid' structure

lies ca. 0.1 and 0.4 kcal mol1 below the 'anti' and 'gauche-cisoid' conformers,

respectively, at the UHF/6-31(d) level after zero-point energy correction. Inclusion of

correlation effects (PMP2/6-311G(d,p)//UHF/6-31G(d)) yields energy differences of 0.3

and 0.4 kcal mol1. These orientation preferences extend to the fluoroethyl series, as

seen in Tables 3-5 and 3-6.


, 2.224 A

( a.


2.228 A

(b) CJ* ^


\ 2.219 A


(c) J


NJ 0 0
Figure 3-19. UHF/6-31G(d) (a) 'Gauche-transoid', (b) 'Anti', and (c) 'Gauche-cisoid'
Transition Structures for Addition of Ethyl Radical to C1 of Propene.








Table 3-5. Geometric Parameters and Total and Relative Energies of Transition
Structures for Addition of Fluorinated Methyl and Ethyl Radicals to Ethylene.
Radical r (C-C) r (C=C) z C-C-C E_ ZPE _
(Al LAl (Deg.) (Deg.) (au) a

CH3" 2.246 1.382 109.1 101.9 -117.575692 0.089165 -
(-118.045817)

CF3" 2.300 1.372 106.6 108.6 -414.156036 0.067955 -
(-415.285304)

CH3CH2 a 2.227 1.384 110.0 103.8 -156.612999 0.120578 0.0
(-157.243689) (0.0)

CH3CH2 b 2.231 1.383 109.9 103.3 -156.612807 0.120455 0.05
(-157.243039) (0.34)

CH3CF2 a 2.235 1.378 110.0 108.6 -354.333278 0.105597 0.0
(-355.404529) (0.0)

CH3CF2 b 2.245 1.378 106.3 108.2 -354.332326 0.105470 0.53
(-355.403222) (0.75)

CF3CH2 a 2.258 1.380 109.0 104.1 -453.202939 0.097486 0.0
(-454.490483) (0.0)

CF3CH2 b 2.258 1.380 107.4 103.5 -453.203135 0.097452 -0.14
(-454.490294) (0.10)

CF3CF2 a 2.275 1.375 108.2 108.6 -650.904439 0.081977 0.0
(-652.632413) (0.0)
CF3CF2 b 2.280 1.375 104.7 108.1 -650.904343 0.081948 0.04
(-652.632131) (0.16)
a Gauche. b Anti. c Degree of Radical Pyramidalization. d UHF/6-31G(d) Values;
PMP2/6-311G(d,p)//UHF/6-31G(d) Values in Parentheses. e Relative Conformer Energy
Differences, kcal mol-1. UHF/6-31G(d); PMP2/6-311G(d,p)IIUHFI6-31G(d) Values in
Parentheses.


Inspection of the data reveals the similarity in both angle of attack and forming

C-C bond length, regardless of either the nature of the radical or conformation of the

transition structure. Incipient bond lengths range from ca. 2.22 to 2.30 A, generally

slightly longer for additions of the fluorinated members of the series. These somewhat

earlier transition states are also reflected in the shorter olefinic C=C bonds, 1.382 A in

the case of methyl radical addition to ethylene versus 1.372 A for trifluoromethyl, in turn






68


Table 3-6. Geometric Parameters and Total and Relative Energies of Transition
Structures for Addition of Fluorinated Methyl and Ethyl Radicals to C1 of Propene.


Radical


CH3"


CF3"


CH3CH2 a


CH3CH2" b


CH3CH2


CH3CF2" a


CH3CF2" b


CH3CF2" c


CF3CH2 a


CF3CH2" b


CF3CH2" c


CF3CF2" a


CF3CF2" b


CF3CF2 c


r (C-C)
(A)

2.243


2.297


2.224


2.219


2.228


2.233


2.228


2.242


2.254


2.248


2.257


2.274


2.266


2.279


r(C=C-)


1.383


1.372


1.385


1.385


1.384


1.379


1.380


1.379


1.381


1.381


1.381


1.375


1.375


1.375


z C-C-C
(Deg.)

109.3


106.4


110.2


111.3


110.1


109.5


110.7


106.0


109.3


110.9


107.7


108.0


110.1


104.6


a Gauche-transoid. b Gauche-cisoid. c Anti.d Degree of Radical Pyramidalization.
e UHF/6-31G(d) Values; PMP2/6-311 IG(d,p)//UHF/6-31G(d) Values in Parentheses.
'Relative Conformer Energy Differences, kcal mol'. UHF/6-31G(d); PMP2/
6-311G(d,p)//UHF/6-31G(d) Values in Parentheses.


(Deg.)

102.0


108.9


103.9


104.2


103.3


108.7


108.9


108.5


104.2


104.6


103.6


108.7


108.9


108.3


(au)

-156.614381
(-157.244918)

-453.195586
(-454.485598)

-195.651571
(-196.442757)

-195.651075
(-196.442411)

-195.651399
(-196.442053)

-393.372286
(-394.604285)

-393.371834
(-394.603661)

-393.371641
(-394.603448)

-492.242374
(-493.691081)

-492.242279
(-493.691288)

-492.242269
(-493.690171)

-689.944220
(-691.833274)

-689.944106
(-691.833236)

-689.944068
(-691.832854)


ZPE
(au)

0.119634


0.098359


0.150984


0.151060


0.150878


0.135951


0.135994


0.135892


0.127933


0.128087


0.127869


0.112383


0.112473


0.112356


E^







0.0
(0.0)

0.35
(0.26)

0.05
(0.38)

0.0
(0.0)

0.31
(0.42)

0.37
(0.49)

0.0
(0.0)

0.15
(-0.04)

0.03
(0.54)

0.0
(0.0)

0.12
(0.07)

0.08
(0.25)









in accord with the greater exothermicity for the latter (AErxn = -22.35 kcal molr1 versus

-34.52 kcal mol-1, respectively, at the [QCISD(T)/6-311G(d,p)]'/UHF/6-31(d) level, and

C-C BDE data in Table 3-4).

Gauche-transoid addition of ethyl radical to C, of propene occurs via a transition

structure with a forming C-C interatomic distance of 2.224 A and and a C=C bond length

of 1.385 A, in comparison with 2.274 A and 1.375 A for addition of pentafluoroethyl along

the same trajectory. Attack angles appear slightly smaller for additions of the fluorinated

radicals, most notably in the case of CH3* versus CF3 and consistent with a reinforced

SOMO-HOMO interaction for the latter. However, such differences are barely

significant, and in the ethyl series appear to be influenced more by the conformation of

the transition structure than the nature of the attacking radical, in line with the previously

observed insensitivity of transition state geometry to additions of both nucleophilic and

electrophilic radicals.47'48

Of note is the degree of pyramidalization at the radical site in the addition

transition structure, ranging from 102-105 for radicals of the RCH2* type and 108-109

for a-fluorinated species. Considering the pyramidal nature of the ground states of the

latter as well (Figures 2-9 and 2-10) it follows that x-fluoroalkyl radicals enjoy a kinetic

advantage over their hydrocarbon analogues in that little or no additional bending is

necessary to accommodate the addition transition structure. The energetic cost of

pyramidalization of the methyl and tert-butyl radicals to the same extent as required for

their addition to ethylene has been computed at 1.5 and 1.6 kcal molr', respectively, at

the RMP2/6-31G(d)ilUHF/6-31G(d) level.49

Calculated degrees of charge transfer (CT) between radical and olefin moieties

of the addition transition structures are provided in Table 3-7. Where applicable, lowest

energy conformations (gauche in the case of ethyl radical additions to ethylene, gauche-








transoid for additions to C1 of propene) were used in the determination of the

electrostatic potential-derived charges.


Table 3-7. Calculated Charge Transfer Data (Electrons) for Transition Structures of
Hydrocarbon and Fluorinated Methyl and Ethyl Radical Addition to Ethylene and C1 of
Propene.

Radical Ethylene a Propene b

CH3" -0.019 -0.004

CF3" -0.013 -0.006

CH3CH2" +0.036 +0.052

CH3CF2" +0.030 +0.052

CF3CH2" -0.047 -0.037

CF3CF2" -0.045 -0.034

Note: Derived from UHF/6-31G(d) Electrostatic Potentials. Negative Values Denote
Electron Transfer from Alkene to Radical. a Gauche Transition Structure. b Gauche-
transoid Transition Structure.


Consistent with previous investigation, addition of CH3* to ethylene involves only

a slight degree of charge transfer (-0.019 e) from olefin to radical (Mulliken and Bader

analyses yield values of -0.017 and -0.011, respectively)48 and even less so for addition

to propene. Somewhat surprisingly, CF3" addition is also predicted to occur without

appreciable polarization.

Along the ethyl series, such interactions appear more clearly defined. Ethyl

radical addition to both ethylene and propene involves a shift of electron density from the

radical to the alkene (0.036 and 0.052 e, respectively) well in accord with the expected

nucleophilicity of the alkyl radicals. CH3CF2* is predicted to be nucleophilic as well,

exhibiting transition state polar characteristics very similar to those of its hydrocarbon

counterpart and consistent with the experimentally deduced non-electrophilicity of

the a-fluoro radicals.








A striking reversal in these trends occurs upon fluorination at the p3 carbon atom,

regardless of the nature of the radical site itself. Here it is seen that both CF3CH2* and

CF3CF2 exhibit substantial electrophilicty, with ca. 0.04 0.05 units of electron density

transferred from the alkene to the radical center. To place such values into some

degree of perspective, addition of the strongly electrophilic cyanomethyl radical to C2 of

electron-rich vinylamine is predicted to occur with a transfer of ca. 0.11 electrons from

CH2=CHNH2 to "CH2CN.48

It is especially noteworthy that the degree of CT in the case of CF3CH2" and

CF3CF2" addition is practically unaffected by fluorination at the a carbon (-0.047 versus

-0.045 and -0.037 versus -0.034, respectively). This, along with the demonstrated lack

of kinetic impact of fluorine substitution at the y position (Table 3-2) leads to the

conclusion that the electrophilic character of the perfluoroalkyl radicals derives

exclusively from substitution at the 2-position.

With geometric, enthalpic, and polar considerations for hydrofluorocarbon

radicals having been addressed, the influences of each of these effects on determined

kadd and kH values are now discussed.


a .a-Difluoroalkvl Radicals (77, 124)


The 1,1-difluoroalkyl radicals, as mentioned previously, benefit from

pyramidalization at their radical site, leading to a more facile adoption of the transition

structure for addition or hydrogen atom abstraction. However, such an advantage is

counteracted by the experimentally and theoretically demonstrated lack of electrophilicity

for such species. Moreover, the terminal C-H bond weakening effect of gem-difluoro

substitution (2 3 kcal mol1, Table 3-3) leads to a slight thermodynamic disadvantage

for hydrogen abstraction by the corresponding radical relative to the hydrocarbon. Thus,

the 3.8-fold rate enhancement enjoyed by 77 may be completely ascribed to the a-type,








pyramidal nature of its ground state, attenuated by enthalpic and polarity factors working

in opposition. In contrast, the strengthening effect of gem-difluorination on C-C bonds

compliments that of pyramidal geometry, leading to a more substantial (22.5-fold) rate

enhancement for the addition of 124 to styrene versus hydrocarbon 123.


B.3-Difluoroalkvl Radicals (88. 125)


The rate enhancements for addition (4.3) and hydrogen abstraction (5.8)

observed for 2,2-difluoroalk-l-yl radicals are due to a complimentary combination of

polar and enthalpic effects. The terminal C-H bond in 2,2-difluoropropane is predicted to

be 2.8 kcal mol' stronger than that of propane itself; similarly, gem-difluorination at C2

of butane leads to a 3.2 kcal mol"1 strengthening of its C3-C4 bond. In spite of these

favorable considerations, the near-planar ground state geometry of 88 results in a

modest net rate acceleration for hydrogen abstraction from nBu3SnH. This also

functions to oppose the CT and enthalpic advantages present in the addition reaction of

125, giving rise to only a slight rate increase relative to n-alkyls and certainly diminished

in comparison with that enjoyed by 124.


-y-Fluorinated Radicals (111, 126)


Due to the near-planar geometric character of 3-fluoroalk-l-yls and the lack of

effect (ca. 1 kcal mol1) on terminal C-H and C-C BDE values, fluorination beyond two

carbon atoms removed from the radical site exhibits a negligible effect on the rates of

both addition and hydrogen transfer. The reactivities of 111 and 126 are found to be,

within experimental error, identical to those of the corresponding hydrocarbon.


.ac.13.-Tetrafluoroalkvl Radicals (99.100) and Pentafluoroethyl Radical (118)


Radicals substituted at the a and P3 positions benefit from both pyramidal

geometries and electrophilic character. Since the degree of pyramidalization of 1,1-








difluoroalkyl radicals remains constant regardless of substitution at the 3- and further

positions, geometrically induced influences on the reactivities of such polyfluorinated

radicals are expected to be uniform. Consequently, rate enhancements for radicals of

the type RCH2CF2CF2, RfCF2CF2CF2", and CF3CF2" versus RCH2CH2CF2" (for kH: 38,

83, and 133 versus 3.8; for kadd: 167, 383 and 658 versus 22.5; all relative to n-alkyls)

derive from either an increasing degree of transition state charge transfer stabilization,

increasingly greater exothermicity of reaction, or a combination of both. The relevant

C-H BDE data in Table 3-3 yields no direct correlation between reaction rate and

enthalpy for the polyfluorinated radical series, with values of 97.7, 100.1, and 99.5

kcal mol' corresponding to hydrogen abstraction by radicals 77, 100, and 118.

Similarly, terminal C-C BDEs of 91.6, 95.4, and 95.5 kcal mol1, equated with the

additions of 124, 99, and 118, illustrate that although P-fluoro substitution should lead to

rate enhancement on thermochemical grounds, such an effect does not account for the

incremental acceleration across the series.

With the degree of radical electrophilicity related to its substitution at the 2-

position and the potential for additional 13 C-F a* delocalization made possible by the

"extra" fluorine substituent in CF3CF2* relative to RCH2CF2CF2" and RfCF2CF2CF2", it

follows that the increasing resonance and inductive withdrawal ability of these groups

relative to RCH2CH2CF2* sufficiently explain both the enhanced reactivity of these

radicals as a whole, as well as the observed trend.


Fluorine Substituent Effects on Bond Dissociation Energies; Coulombic Interactions


As previously discussed, substitution by fluorine in hydrocarbons gives rise to

nearly additive and sometimes opposite effects on C-H and C-C homolytic bond

dissociation energies. For example, the aforementioned 1-3 kcal mol' weakening

effect of a,a-difluoro substitution and the 3-5 kcal mol-1 strengthening of terminal C-H








bonds by p-fluorination lead to near cancellation in the case of pentafluoroethane and

1,1,2,2-tetrafluoropropane, yielding BDE values very near those of the parent

hydrocarbon (Table 3-3.) Furthermore, the 4.9 kcal mol1 strengthening brought about

by substitution at the breaking C-C bond is reinforced by an additional 3-5 kcal mol"1

upon further fluorine incorporation at the p position, leading to net increases of nearly 10

kcal mol-1 over the parent in the C2-C3 homolysis of 1,1,1,2,2-pentafluoropropane and

cleavage of the terminal C-C bond of 2,2,3,3-tetrafluorobutane.

The opposite effects of aa-difluoro substitution on C-H and C-C bond

dissociation energies bear special mention. Experimental BDE values in the

fluoromethanes are in accord with the general RSE expectations of Pasto (Table 2-7) in

that although substitution at a radical center by a single fluorine is stabilizing, its further

incorporation leads to a successive decrease in RSE, resulting in net destabilization for

CF3'. This is consistent with the incremental strengthening of C-F bonds along the

fluoromethane series, leading to a C-H BDE in CF3H which is 1.9 kcal mol1 stronger

than that of methane itself (methane BDE, 104.8 kcal morl1; see discussion below Table

2-2) and the comparatively weaker C-H bonds in CHsF and CH2F2.

The considerable stability of the 2,2-difluoroalkanes relative to their 1,1-difluoro

isomers, demonstrated by the isodesmic reaction in Equation 3-2 (calculated from

B3LYP/6-31G(d) total energies and zero-point corrections) provides the underlying

reason for why the stability trends observed above do not extend to C-C bonds.


CH3CF2CH3 + CH3CH3 CH3CF2H + CH3CH2CH3 (3-2)

AErxn = + 7.8 kcal mol1

In Chapter 2, the Wiberg rationale of electrostatic attraction for the incremental

strengthening and shortening of C-F bonds in the fluoromethanes was introduced.

Similarly, it is found that such a Coulombic-based argument sufficiently explains the

observed effects of fluorine substitution on C-H and C-C bond dissociation energies.








Atomic charges for select hydrofluorocarbons based on the B3LYP/6-31G(d)

electrostatic potential (Merz-Kollman radii) are provided in Table 3-8.

The dipolar nature of the C-C bond in 1,1,1-trifluoroethane and its resultant

increase in BDE relative to ethane and hexafluoroethane (Table 2-3) was first postulated

by Rodgers.16 Such stabilization due to increased C-C bond ionicity is seen in

Equations 3-3 (derived from experimental heats of formation188) and 3-4 (from B3LYP/6-

31G(d) total and zero-point energies).


Table 3-8. Atomic Charges in Hydrofluorocarbons, Based on B3LYP/6-31G(d) Density.

Hvdrofluorocarbon LF _Cg LC "HXHoi

CaH3Ca,,H3 -0.055 +0.018
Ethane

C,,H2F2 -0.200 +0.320 +0.041
Difluoromethane

CpH3CaF2H -0.228 +0.467 -0.386 +0.020 +0.110(2H)
1,1 -Difluoroethane +0.135 (1 H)

CH3C,,F2CpH3 -0.245 +0.631 -0.477 +0.133 (4H)
2,2-Difluoropropane +0.140 (2H)


CF3CF3 + CH3CH3 2 CH3CF3 (3-3)

AErxn = -16.9 kcal mol"1


CH3CH2CH2CH3 + CH3CF2CF2CH3 2 CH3CF2CH2CH3 (3-4)

AErxn = -5.0 kcal mol1-

The significant electrostatic attraction between adjacent carbon atoms in both

CH3CF2H and CH3CF2CH3 is readily apparent from the data in Table 3-8, providing an

explanation for the strengthening of these bonds relative to their hydrocarbon or

perfluorocarbon analogues. In addition, C-H repulsion in difluoromethane and 1,1-

difluoroethane is predicted, in accord with the observed weakening of these bonds

compared to those of methane and ethane. Conversely, the strong attraction between


I








the P carbon and hydrogen atoms of CH3CF2H (XC, -0.386; xHavg, +0.118) and

CH3CF2CH3 (.C, -0.477; XHavg, +0.135) is consistent with their greater theoretical and

experimental BDEs.


Conclusion


Based on time-resolved kwd measurements, absolute rate constants for

hydrogen abstraction from tri-n-butyltin hydride have been determined for a series

partially fluorinated radicals. The reactivities of such radicals towards nBu3SnH follow

those of addition to alkenes. The enhanced reactivity of a,a-difluoroalkyl radicals in

hydrogen abstraction reactions derives exclusively from their pyramidal geometry.

p-Fluorination leads to a favorable combination of polar and thermodynamic factors in

both addition and hydrogen transfer reactions, giving rise to the exceptional reactivity of

CF3CF2* and the perfluoroalkyl radicals as a whole. In Chapter 4, the kH values so

obtained are utilized in the determination of absolute rates of cyclization for partially

fluorinated 5-hexenyl radicals.

A self-consistent rationale for the impact of fluorine substitution on C-H and C-C

bond dissociation energies based on electrostatic considerations was offered, providing

new understanding of the thermochemistry of bonding and radical stabilization in

hydrofluorocarbons.














CHAPTER 4

THE REACTIVITY OF PARTIALLY FLUORINATED RADICALS
IN INTRAMOLECULAR CYCLIZATION REACTIONS


Introduction


The intramolecular addition reactions of 5-hexen-l-yl radicals continue to attract

the attention of synthetic and physical organic chemists alike. Such cyclizations to

(predominantly) 5-exo products have been utilized as probes for the detection of radical

intermediates and as basis reactions for the competitive determination of absolute

kinetic data for a number of free radical transformations.16 Rationalization of the rates,

and especially the regio- and stereochemistry, of intramolecular radical additions on the

basis of force field63'64'66'87 and molecular orbital62,65'68'68 techniques has proven to be one

of the greatest successes of theory in the prediction of organic reactivity. Due in no

small part to such structure-reactivity studies, application of free radical methodology to

the singular and tandem construction of 5-membered rings has been equally exploited,

providing for the assembly of functionalized organic systems under mild conditions, often

accomplished with a high degree of stereocontrol.11'15

Determination of absolute rates of cyclization of per- and other highly fluorinated

5-hexenyl systems86,89 170 have aided in solidifying the understanding of the effect of

fluorine substitution on the reactivity of organic radicals, though at the same time

generating a number of new questions, particularly with regard to cyclization

regiochemistry.

In order to examine the potentially more subtle influences of partial fluorination

on 5-hexenyl radical reactivity, and to obtain a set of data through which the effect of








incremental gem-difluoro substitution along the aliphatic portion of the 5-hexenyl chain

may be assessed, absolute rates of 5-exo and 6-endo cyclization for some partially

fluorinated 5-hexenyl radicals have been determined based on competitive kinetic

technique and the absolute rates of hydrogen abstraction obtained in Chapter 3.

Precursor Syntheses and Competitive Kinetic Studies


As in the bimolecular addition versus hydrogen abstraction competition studies,

bromide precursors were utilized in the generation of partially fluorinated 5-hexenyl

radicals. Photolysis by UV irradiation (Rayonet photoreactor) in the presence of known,

varying concentrations of hydrogen atom donor, carefully adjusted to ensure pseudo-first

order kinetic behavior and to allow for accurately measurable amounts of cyclization and

hydrogen abstraction products, provided the kinetic ratio kH I kcn. Absolute rate

constants for 5-exo and (where applicable) 6-endo cyclization were then determined

from the known value of hydrogen abstraction rate constant kH, illustrated in Figure 1-18

and in greater detail below.


1.1-Difluorohex-5-en-1-vyl Radical (128)


Synthesis of bromide 135 was achieved in six steps in ca. 14.5% overall yield,

starting from commercially available 3-buten-l-ol (129, Figure 4-1). Curiously, direct

addition of dibromodifluoromethane to 129 could not be induced, even through extended

reaction time at elevated temperatures. Although the presence of the alcohol

functionality in the alkene starting material would not have been expected to exhibit a

detrimental effect (in light of the hydroxylic nature of the ethanolamine / terft-butanol

cosolvent medium) protection of the hydroxyl moiety as its tert-butyldimethylsilyl ether

130 (TBDMSCI, imidazole in dimethylformamide) followed by dibromodifluoromethane

addition indeed afforded 1,3-dibromo-1,1-difluoro adduct 131 in good yield. Highly

selective displacement of the internal bromine yielded bromodifluoromethyl derivative








TBMSCI
//'\/OH ImH, DMF ^ OTBDMS
(89.3%)
129 130


CF2Br2
tBuOH, H2NCH2CH2OH
CuCI (cat.)


Br
BrF2C ^OTBD
SvOTBDI\
(72.4%)
131


BrF2C\7 O

(97.7%)
133


NaBH4
AIS DMSO





PCC DO BrF
CH2CI2


,/s\CF2Br

(48.1%)
135

Figure 4-1. Preparation of 6-Bromo-6,6-difluorohex-1-ene, Precursor to 1,1-Difluorohex-
5-en-1-yl Radical 128.


132 with virtually no overreduction product, as monitored via 1"F NMR through high

conversion of starting material. Lewis acid deprotection via the method of Cort193 and

subsequent pyridinium chlorochromate oxidation provided aldehyde 134, further

subjected to Wittig olefination to yield precursor 135.

Generation of 128 was achieved via irradiation of a solution of 135 in CeD6 in the

presence of excess nBu3SnH (Figure 4-2) and an internal standard of a,ca,a-

trifluorotoluene. Direct capture of 128 by hydrogen atom donor afforded reduction

product 6,6-difluorohex-l-ene 136, whereas intermediate 137, subsequently trapped by

nBu3SnH to yield spectroscopically observable cyclization product 138, was generated

via irreversible, unimolecular rearrangement with rate constant kc5 (no 6-endo cyclization

was observed, within NMR detection limits, (ca. 4%) for 128). During the course of the


FeCI3
CH3CN


r2C \/\ OTBDMS

(90.0%)
132


0
2 \AH
(53.1%)
134


Ph3P=CH2
THF








CF2Br hv ^[ ]
nBu3SnH
135 C6D6 128

nBu3SnH /,. /,\/CF2H + nBu3Sn.

S k136


F_ /nBu3SnH F
kc5 F] F

137 138
+ nBu3Sn

Figure 4-2. kH / kc Competitive Kinetic Scheme for 1,1-Difluorohex-5-en-1-yl Radical
128.

reaction, tributylstannyl radicals generated by transfer of hydrogen atom from nBu3SnH
to 128 and 137 served to propagate the chain process via bromine abstraction from 135.
Product ratios for varied concentrations of nBu3SnH were determined by 19F
NMR analysis according to the pseudo-first-order relation in Equation 4-1,

[136] [kH] [1281 [nBu3SnH (4-1)
[138] [kC5] [128]

a plot of which obtained for each data point versus nBu3SnH concentration providing the
ratio kH / kcs. Exceptionally clean spectra and high mass balances were obtained for
each kinetic point, indicating the efficiency of the radical chain process and reliability of
the obtained rate constant ratios. A partial 19F spectrum of the first of six data points is
provided in Figure 4-3, a doublet of triplets (-CF2H, 0 -116.2) observed for 136 versus

overlapping doublets of doublets of triplets at 4 -100.3 and -107.8 for the diastereotopic
-CF2- resonances of 138. Kinetic data and product yields are given in Table 4-1, a plot
of which found in Figure 4-4. The slope of the line (2.57 + 0.05) in conjunction with the
























-96 -98 -100 -102


-104 -106 -108 -110


-112 -114


-116 -118


Figure 4-3. Partial 19F NMR Spectrum of
Difluorohex-5-en-l-yl Radical 128.


Table 4-1. Competitive Kinetic Data for kH
Radical 77.


[1351

0.054

0.054

0.054

0.054

0.054

0.054


[nBu3SnH 1

0.673

0.807

0.942

1.08

1.21

1.35


Data Point 1 for kH I kc Competition of 1,1-



/ kc Competition of 1,1-Difluorohex-5-en-1-yl


[1361/f1381

1.53

1.91

2.28

2.57

2.93

3.29


% Yield

88

100

89

94

95

92


Coefficients:
m = 2.57
b =-0.175
r I = 0.999


0.6 0.7 0.8 0.9 1.0 1.1
[nBu3SnH]

Figure 4-4. Plot of the Data in Columns 2 and 3 of Table 4-1.


4.0 -
3.5-
3.0-
2.5 -
2.0 -
1.5-
in








known absolute rate constant for hydrogen atom abstraction from nBu3SnH by 1,1-

difluorohex-1-yl radical 77, 9.1 (+ 1.7) x 106 M"1 s'1, resulted in a kcs value of

3.5 ( 0.59) x 106 s1 for 5-exo closure of 128, with errors in kc reflecting both the least-

squares fit of the line and propagated error in kH. Syntheses of hydrogen atom transfer

and cyclization products 136 and 138 were performed as shown in Figure 4-5.


4 ,/CF2Br nBu3SnH ___/ /,/CF2H
AIBN
135 Mesitylene 136


0
DAST F
CH2C12 F

139 138

Figure 4-5. Preparation of Hydrogen Abstraction and 5-Exo Cyclization Products 136
and 138.


2.2-Difluorohex-5-en-1-vyl Radical (140)


Bromide 144 was obtained in a three-step synthesis starting from 1,2-epoxy-5-

hexene (Figure 4-6). Regiospecific ring opening by a Corey94 procedure afforded

bromohydrin 142, converted to the corresponding a-haloketone via Jones oxidation.

Treatment of 143 with DAST in dichloromethane afforded precursor 144 in 40.9% overall

yield, purified by preparative GC for competitive kinetic study.

In contrast to the virtually regiospecific 5-exo closure of 128, a broad singlet

resonance at -95.8, comprising approximately 9% of cyclized products, was observed

in the 1"F NMR spectra for the kH / kc competition of 140. This is attributed to competing

6-endo cyclization to 148 (Figure 4-7), the presence of which was confirmed by spectral

comparison with that of an authentic sample of 149.







KBr, CH3CO2H
THF / H20


/,,Br
OH


Na2Cr2O7 / H2S04
Et20


(88.1%)


141


DAST
CH2CI2


'tCF2CH2Br


(56.2%)


Figure 4-6. Preparation of 6-Bromo-5,5-difluorohex-1-ene, Precursor to 2,2-Difluorohex-
5-en-1-yl Radical 140.


,/xv, CF2CH2Br

144


hv
nBu3SnH
C6D6


[^CF

140


\/yCF2CH3
145


F ]
F


146


F F


nBu3SnH


148


Figure 4-7. kH I kc Competitive Kinetic
140.


Scheme for 2,2-Difluorohex-5-en-1 -yl Radical


0


0
CH2Br


(82.6%)


nBu3SnH


nBu3SnH


F
F


147


F6


149





84


Cyclizations of 1p,3-difluoroalkyl radicals have appeared in the synthetic

literature,195 utilized in the generation of alkoxy-substituted gem-difluorocyclopentane,

cyclohexane, and tetrahydropyran derivatives, though reported to undergo addition in an

exo-specific manner. However, the regiochemical behavior exhibited in the cyclization of

140 provided experimental verification of that previously predicted on the basis of ab

initio calculations, performed as part of the present study and elaborated upon in the

Discussion section of the chapter. Competition plots for kH / kc5 and kH I kce are found in

Figures 4-8 and 4-9, respectively. Preparation of hydrogen abstraction and 5-exo and 6-

endo cyclization products was performed as shown in Figure 4-10.


7
6 Coefficients:
I- m=12.5
S5 b =-0.924
r2 = 0.999
S4-3
3-

2 -1 I I I I11
0.25 0.30 0.35 0.40 0.45 0.50 0.55 0.60

[nBu3SnH]

Figure 4-8. Plot of kH I kc5 Competition of 2,2-Difluorohex-5-en-1-yl Radical 140.



60 coefficients:
0 m=132
50 b = -14.0 &
40 r 2 = 0.995
-" 40

S30 -

20 ------
0.25 0.30 0.35 0.40 0.45 0.50 0.55 0.60

[nBu3SnH]

Figure 4-9. Plot of kH / kc6 Competition of 2,2-Difluorohex-5-en-1-yl Radical 140.








F2CH2Br nBu3SnH M-W /VCF2CH3
AIBN
144 Mesitylene 145


0
DAST
CH2CI2
F

150 147

O F\ F

.^ DAST .
CH2CI2

151 149

Figure 4-10. Preparation of Hydrogen Abstraction and 5-Exo and 6-Endo Cyclization
Products 145,147, and 149.


1.1,2,2-Tetrafluorohex-5-en-1-vyl Radical (152)


Bromide 101 (Halocarbons, Inc.) served as the precursor to a,a,pj3-

tetrafluorinated radical 152. However, attempts at determination of accurate kH I kc

ratios using nBu3SnH as a trapping agent met with failure, leading primarily to reduction

product 153, with only minor amounts of 155 and 177 evident in the 19F NMR baseline

which could not be integrated accurately over a span of hydrogen atom donor

concentrations (Figure 4-11). In principle, lowering the concentration of both radical

precursor (typically in the 0.05 0.1 M range) and trapping agent (while still maintaining

pseudo-first order conditions) should effectively decrease the amount of reduction

product and allow for a greater degree of cyclization to be observed. However, too great

of a decrease in precursor concentration leads to decreased NMR signal to noise ratios,

the necessity of longer acquisition times per sample, and increased potential for the

introduction of systematic error.









hv F 1
^?c^-\ CFzBr ----- ",v.' c,.C F2
F2 nBu3SnH CF2
F2 F6D6
101 c6D 152

nBu3SnH ^..c-CF2H
kH F2

153



virtually no cyclization products
kc5 observed






kc6

Figure 4-11. Attempted kH / kc Competition of 1,1,2,2-Tetrafluorohex-5-en-1-yl Radical
152 with nBu3SnH as Trapping Agent.


As it was evident that any cyclization reaction of 152 occurred with a rate

constant too low to be competitive with transfer of hydrogen from nBu3SnH, attention

was turned to alternative trapping agents. With the rate of hydrogen atom transfer to

perfluoroalkyl radicals by a number of reducing agents having been accurately

determined (Table 2-10), it was decided to investigate the suitability of

tris(trimethylsilyl)silane ((TMS)3SiH) as a competitive trapping agent for the calibration of

cyclization rate constants for 152, due to its approximately four-fold decrease in

hydrogen transfer rate to perfluoroalkyls relative to nBu3SnH. For such a competition to

be of kinetic value, however, it was necessary to determine rate constant kH for

tetrafluoroalkyl radical 100 with (TMS)3SiH, using its known rate of addition to styrene as

a competing basis reaction. The plot for the kH ((TMS)3SiH) / kadd (styrene) competition

of 100 is provided in Figure 4-12.





87


With the kH ((TMS)3SiH) value of 1.8 ( 0.1) x 107 M1 s1 for 100 in hand (which,

along with its kH (nBu3SnH) of 9.2 ( 0.8) x 107 M1 s-', (Table 3-2) may be compared

with 5.1 x 107 M1 s1 and 2.0 x 108 M1 s1, respectively, for perfluoro-n-alkyl radicals;

Table 2-10) rate constants kc5 and kc6 for 152 were then determined (Figures 4-13 and

4-14.) Use of this slower hydrogen transfer agent allowed for sufficient competitive

(including significant 6-endo) cyclization such that accurate kH / kcn ratios could be

obtained. Isolation of products 153, 155, and 157 was achieved by slow syringe pump

addition of nBu3SnH to a heated, irradiated solution of 101 in mesitylene (Figure 4-15).


2.2 -
S2.0 Coefficients:
0 m=0.913
1.8 -
,e- b =0.189
1. r=0.999
Go
S1.4-
1.2 -
1 .0 1 I I I 1 I
0.9 1.0 1.1 1.2 1.3 1.4 1.5 1.6 1.7 1.8 1.9 2.0
[ (TMS)3SiH ] / [ C6H5CH=CH2]

Figure 4-12. Plot for kH ((TMS)3SiH) / kadd (Styrene) Competition of 1,1,2,2-
Tetrafluorohex-1-yl Radical 100.


2.5 -
S.. Coefficients:
u) 2.0 m=2.11
tOe b =-0.242 J
S 1.5 r =0.998

1.0

0 .5 1 1 1 1 1-1
0.5 0.6 0.7 0.8 0.9 1.0 1.1 1.2
[ (TMS)3SiH ]

Figure 4-13. Plot of kH / kc5 Competition of 1,1,2,2-Tetrafluorohex-5-en-1-yl Radical 152.




















0.6 0.7 0.8 0.9 1.0


[ (TMS)3SiH ]

Figure 4-14. Plot of kH I kc6 Competition of 1,1,2,2-Tetrafluorohex-5-en-1-yl Radical 152.


^, -^^ .^-CF2Br
F2
101


nBu3SnH
IDO
AIBN
Mesitylene


.. CF2
F2


nBu3SnH


,,,,s,/ CF2H
F2

153


F
FF
F


F


nBu3SnH


nBu3SnH


F
F
FF
F


F



157


Figure 4-15. Preparation of Hydrogen
Products 153,155, and 157.


Abstraction and 5-Exo and 6-Endo Cyclization


1.1 1.2








Discussion


Absolute rate constants of cyclization for radicals 128, 140, and 152 are given in

Table 4-2. For comparison, such kcs and (where applicable) kc6 values for parent

hydrocarbon 1 and fluorinated radicals 65 and 68 are also provided, the latter two

systems along with those of the current study found to give rise to the greatest impact on

cyclization kinetics and regiochemistry. Recent studies of 5-hexenyl systems bearing

vinylic fluorine substituents have demonstrated that the effect of such substitution is

relatively minor, with no 6-endo products observed within the detection limits imposed by

NMR analysis and kc5 (re) values with respect to 1 ranging from ca. 0.09 to 2.3.170


Cyclization Kinetics


Rates of intramolecular addition of partially-fluorinated radicals to alkenes should

be governed by the same combination of steric, polar, and thermodynamic factors which

influence the reactivity of their intermolecular counterparts. As seen by comparison of

the data in Tables 3-2 and 4-2, the reactivity characteristics of the above radicals in

unimolecular cyclization reactions, particularly 5-exo closure, generally reflect those

observed in bimolecular additions. This is logical in light of the similarity of their

transition structures, elaborated upon in Chapters 1 and 3.Y

The pyramidal nature of a,a-difluoroalkyl radicals, combined with the more

favorable thermodynamics of C-C bond formation involving fluorinated carbon (see

related discussions in Chapters 2 and 3, along with cyclization transition structures and

energies of reaction below) provide sufficient explanation for the 13-fold increase in rate

of 5-exo ring closure of 128 relative to 1. The factor of 22.5 observed for addition of 77

versus 124 to styrene is consistent with the observed cyclization rate ratios.

The increase in kcs of 4.1 enjoyed by 140 parallels that of bimolecular addition of

125 to styrene, (4.3) due to its increased electrophilicty over both hydrocarbon 1 and





90

Table 4-2. Absolute Rate Constants for 5-Exo and 6-Endo Cyclization of Partially
Fluorinated 5-Hexenyl Radicals. Rate Constants are for 303 K; Relative kcn Values in
Parentheses.


Cyclization Reaction


kc5, 105 s-1


k6, 105 s-1


0.05 a


+ 0

3


N/A b,c


0CF2


(13.0)


158


F2


cCF2
F2
156


11 (3.8) b
(4.1)


87 (t 4.1) b
(32.2)


1.1 (0.34)" b
(22)


19 ( 1.1) b
(380)


F2C-.. -CF2
F2


+ r CF2
F2C,. CF2
F2


440 ( 46) d
(163)


110 (1.7) d
(40.7)


8 Reference 16. b Current Study. c 6-Endo Cyclization Not Observed Within 19F NMR
Detection Limits (Approximately 4%). d Reference 170.


jF2 C


2


CF2


CF2

140


CF cF2
^CF2


6CF2

146



CF2
CF2


154


1 CFF2
` CCF2
F2


'ICF2*
F2C.cCF2
F2


0
FCCF2
F2C-CF2


F2C CF2
F2C-CF2


52 (6.4) d
(1040)


35 ( 4.4) d
(700)








a,a-difluorocarbon 128 and greater exothermicity of addition relative to n-alkyls, though

tempered by the effectively planar, n-nature of its radical center.

a,a,p3,13-Tetrafluorinated radical 152, as in the case of bimolecular additions,

benefits from favorable thermodynamics of addition as well as its electrophilicty and

G-character, leading to the 32-fold increase in kcs compared to parent 1. It should be

noted that such unimolecular cyclizations possess an inherent entropic advantage over

their bimolecular analogues, generally proceeding with log A values ca. 2 units larger

than those for the latter18 and resulting in a leveling of rate ratios relative to

intermolecular additions.

Upon additional fluorination of the aliphatic moiety of the 5-hexenyl chain (65, 68)

such radicals undertake perfluoroalkyl character, leading to further increase in reactivity

akin to that observed for C7Fs* (127, kadd (.e) = 383) versus CH3CH2CF2CF2" (99,

kadd (re1) = 167, relative to n-alkyl) in bimolecular additions to styrene. Geminal

difluorination at the allylic position (68) serves to diminish the transition state SOMO-

HOMO interaction, and hence kcs and kcm, relative to 65.


Cyclization Regiochemistrv


The significant degree to which 152, 65, 68, and even 140 undergo 6-endo

cyclization is particularly striking, with six-membered ring formation in 140 occurring with

a rate nearly half, and 65 and 68 more than an order of magnitude greater than, that of

5-exo closure for hydrocarbon 1. In comparison, 5-hexenyl systems bearing alkyl

substituents along the aliphatic fragment exhibit regiochemical profiles similar to that of

the unsubstituted parent.70'71

The question of potential reversibility in the above cyclizations has been

addressed, in light of the greater relative thermodynamic stability of secondary

cyclohexyl radicals. Upon independent generation of 5-exo adduct radical 69 from

precursor 1-(iodomethyl)-2,2,3,3,4,4,5,5-octafluorocyclopentane in the presence of








hydrogen atom donor triethylsilane in C6D6, the only product observed after complete

consumption of starting material was that resulting from direct capture of 69 by Et3SiH.170

The lack of 6-endo or ring-opened products originating from 69, coupled with the ab initio

predictions based on relative energies of cyclization transition structures described

below, demonstrates that the regiochemical characteristics of fluorinated 5-hexenyl

radical cyclizations are indeed kinetic in nature.

Of further note is that system 68, which undergoes the greatest percentage

(24.1%) of 6-endo closure (that is, exhibiting the least selectivity) is not the most

reactive. Hexafluoro system 65, though forming 67 with a rate constant 1.5 times that of

analogous closure of 68 to 70, does so only to an extent of 10.6% of total cyclized

products.

A combined ab initio I molecular mechanics approach has allowed for accurate

regiochemical predictions for a number of alkyl and heteroalkyl intramolecular radical

additions.66 In order to examine the effect of the degree and location of fluorine

substitution on transition structure geometry and energetic, as well as on activation

barriers and reaction enthalpy, the "chair-like" and "boat-like" 5-exo and 6-endo

cyclization transition structures for the parent hydrocarbon and various fluorinated

5-hexenyl systems, along with their respective open-chain radicals and products of

5- and 6-membered ring closure, have been investigated with ab initio techniques.

In accordance with a UMINDO/3 investigation of Bischof,62 the lowest energy

conformation of 5-hexenyl radical 1 was found to be an all-trans methylene chain in a

gauche orientation with the internal vinyl hydrogen (Figure 4-16.) Alignment of the singly

occupied orbital of 1 with the adjacent C-H bond was found to be slightly preferred (ca.

0.1 kcal morl1) over similar C-C alignment at the UHF/6-31G(d) + ZPE level.

From the calculated structures and energies of cyclization products

(cyclopentylmethyl and cyclohexyl radicals) it was possible to compute energies of

reaction for hydrocarbon 1 and its fluorinated analogues. Total energies of reactant and








product radicals and exothermicities of 5-exo and 6-endo cyclization for 1,128, 152, and

65 are provided below in Table 4-3.







r(C=C) = 1.318A


Figure 4-16. Lowest Energy Conformation of 5-Hexen-l-yl Radical 1; SOMO (On Right)
Aligned with Adjacent C-H Bond. UHF/6-31G(d) Optimized Geometry.


As expected from C-C BDE data, (Table 3-4) intermolecular additions of the

fluorinated species are, as a whole, more exothermic than for parent 1. However, no

direct correlation exists between either absolute rates of 5-exo and 6-endo addition or

relative percentage of 6-membered ring formation and its corresponding reaction

exothermicity. Although a steady increase in both kc5 and kce is observed along the

series (1 -> 128 --> 152 -> 65), both cyclizations of 65 are predicted to be less

exothermic than those of 152. Furthermore, relative enthalpies (AEn(15.s) AErxn(.e)) are

found to rise with the degree of fluorination, favoring 6-endo closure in consistent

manner for both levels of theory employed. This is at variance with the lesser extent of

6-endo closure in 65 compared to 152.

Total, zero-point, and relative energies along with pertinent geometrical

parameters for the UHF/6-31G(d) cyclization transition structures of 1, (depicted in

Figures 1-11 1-14) 128, 140, 152, and 65 are reported in Table 4-4. Although the

calculated energy differences between "chair" and "boat" forms of either 5-exo or 6-endo

transition structures are quite consistent among the theoretical methods, energies of the

"6-endo-chaif' and "6-endo-twist-boaf' structures relative to the "5-exo-chaie" and

"5-exo-boaf' appear to be overestimated at the PMP2/6-311G(d,p)//UHF/6-31G(d) level

compared to both UHF and QCISD(T) results. Bearing this in mind, relative transition




Full Text
EXPERIMENTAL AND THEORETICAL INVESTIGATION OF THE REACTIVITY
OF PARTIALLY FLUORINATED RADICALS
By
MICHAEL DAVID BARTBERGER
A DISSERTATION PRESENTED TO THE GRADUATE SCHOOL
OF THE UNIVERSITY OF FLORIDA IN PARTIAL FULFILLMENT
OF THE REQUIREMENTS FOR THE DEGREE OF
DOCTOR OF PHILOSOPHY
UNIVERSITY OF FLORIDA
1998

ACKNOWLEDGEMENTS
Among the great number of individuals with whom I have interacted throughout
the course of my education at the University of Florida and elsewhere, I wish to express
my sincere appreciation to the special few that have motivated, challenged, and inspired
me.
I extend my deepest gratitude to Prof. William R. Dolbier, Jr., an outstanding
scientist and truly exceptional educator, for his excellent guidance, support, and
friendship throughout the course of my graduate career. My appreciation for the
knowledge he has shared with me, as well as his patience and level of understanding,
particularly during periods of difficulty and stress, can not be overstated.
I wish to thank the two finest classroom instructors I have ever had--my first
college level chemistry teacher, Dr. Jeanette Madea, for her profound influence in my
decision to pursue a career in the chemical sciences, and Prof. Seth Elsheimer, for my
initial exposure to the fascinating area of organofluorine chemistry in 1990 and his
friendship thereafter.
I am indebted to Dr. Max Muir for introducing me to computational chemistry.
The experience I have gained in the use of molecular orbital methods as a tool for the
understanding of chemical reactivity is due entirely to him. Special thanks go to Prof.
Benjamin Horenstein for his helpful discussions and generosity with regard to
computational resources.
I thank my colleagues, past and present, in both the Dolbier group and the
Department of Chemistry as a whole. A few bear special mention--Dr. Keith Palmer, for
his friendship and advice during my first year in the group; Dr. Xiao Xin Rong and He-Qi

Pan, for their camaraderie and early assistance with radical kinetics; Dr. Conrad
Burkholder, for numerous stimulating discussions; and Dr. Henryk Koroniak, Michelle
Fletcher, Lian Luo, Feng Tian, and Kevin Ley for their friendship (and tolerance!)
throughout the course of my stay in the department.
I wish to thank my graduate committee, particularly the "organic" portion thereof-
Profs. Merle Battiste and Kirk Schanze, for their advice and encouragement. Also,
special thanks go to Prof. R. J. Bartlett for taking seriously my interest in theoretical
methodology and the invitations to participate in his workshops on Applied Molecular
Orbital Theory.
I am especially grateful to my very best friend, Cynthia Dawn Zook, for her
unrelenting moral support and encouragement over the last several years. Finally, I
wish to acknowledge my parents, George Charles and Beverly Jean, for instilling in me
the work ethic which has likely had as much to do with the successful completion of this
work as any of the chemistry I ever learned.

TABLE OF CONTENTS
ACKNOWLEDGEMENTS Ü
ABSTRACT vi
CHAPTER
1 AN OVERVIEW OF ORGANIC FREE RADICAL REACTIONS 1
Introduction 1
Radical Chain Processes 2
Hydrogen Atom Abstraction Reactions 4
Intermolecular Radical Addition Reactions 9
Intramolecular Addition Reactions: Radical Cyclizations 13
Methods for Determination of Organic Radical Kinetics 22
Conclusion 26
2 THE FLUORINE SUBSTITUENT IN ORGANIC SYSTEMS 28
Introduction 28
Structure, Bonding, and Reactivity in Saturated Systems 29
Structure, Bonding, and Reactivity in Unsaturated Systems 31
Fluorine Non-Bonded Interactions in Reactive Intermediates 33
Fluorine Steric Effects 35
The Fluorine Substituent in Free Radicals 37
Organofluorine Radical Reactivity 40
Conclusion 48
3 THE REACTIVITY OF PARTIALLY FLUORINATED RADICALS IN
INTERMOLECULAR ADDITION AND HYDROGEN ABSTRACTION
REACTIONS 49
Introduction 49
Precursor Syntheses and Competitive Kinetic Studies 50
Discussion 61
Conclusion 76
4 THE REACTIVITY OF PARTIALLY FLUORINATED RADICALS IN
INTRAMOLECULAR CYCLIZATION REACTIONS 77
Introduction 77
Precursor Syntheses and Competitive Kinetic Studies 78
Discussion 89
Conclusion 100
IV

5 EXPERIMENTAL
102
General Methods- Experimental 102
General Methods- Theoretical 103
Synthetic Procedures 103
Competitive Kinetic Procedures 129
APPENDIX A: SELECTED 19F NMR SPECTRA 136
APPENDIX B: B3LYP/6-31G(d) TOTAL AND ZERO-POINT ENERGIES FOR
DATA IN TABLES 3-3 AND 3-4 172
REFERENCES 175
BIOGRAPHICAL SKETCH 185
v

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
EXPERIMENTAL AND THEORETICAL INVESTIGATION OF THE REACTIVITY
OF PARTIALLY FLUORINATED RADICALS
By
Michael David Bartberger
May 1998
Chairman: William R. Dolbier, Jr.
Major Department: Chemistry
The reactivities of a series of partially-fluorinated radicals towards intermolecular
addition, hydrogen abstraction, and intramolecular cyclization have been investigated.
Based on competitive kinetic techniques and absolute rate contants for addition of these
radicals to styrene obtained by laser flash photolysis, absolute rate constants for
abstraction of hydrogen from tributylstannane have been determined for 1,1-
difluoroalkyl, 2,2-difluoroalkyl, 1,1,2,2-tetrafluoroalkyl, 3-perfluoroalkyl, and
pentafluoroethyl radicals. Fluorination at the 3-position of an alkyl radical was found to
exert a negligible effect on the kinetics of hydrogen abstraction. All other systems
exhibit rate enhancements relative to non-fluorinated analogues, the magnitudes of
which are dependent upon the degree and location of fluorine substitution. A parallel
computational study was performed utilizing density functional calculations, providing
estimates of carbon-carbon and carbon-hydrogen bond dissociation energies (BDEs) for
hydrofluorocarbons. The observed kinetic enhancements were attributed to a
combination of structural, charge transfer, and enthalpic effects, due to the pyramidal
nature of 1-fluoralkyl radicals, increased electrophilic character induced by successive
fluorination, and thermodynamics of carbon-carbon and carbon-hydrogen bond
VI

formation. From the computation of partial atomic charges in fluoroalkanes, the
contrasting effect of 1-fluorination on carbon-carbon and carbon-hydrogen BDEs and
the consistent strengthening effect of such substitution at the 2-position have been
explained on the basis of Coulombic interactions.
Based on the rate constants obtained for hydrogen abstraction, absolute rate
constants for 5-exo and 6-endo intramolecular cyclization for a series of partially
fluorinated 5-hexenyl radicals have been obtained. These observed rates of cyclization
may be rationalized by the same combination of effects influencing their bimolecular
addition reactions. In some cases, the rate of 6-endo closure is dramatically
accelerated relative to the parent hydrocarbon without the introduction of reversibility of
ring closure.

CHAPTER 1
AN OVERVIEW OF ORGANIC FREE RADICAL REACTIONS
Introduction
The discovery of the first free radical, triphenylmethyl, by Moses Gomberg1 in
1900 initiated considerable effort directed toward the understanding of radical reactivity.
However, only after a series of pioneering investigations undertaken more than thirty
years later were the primary mechanistic pathways available to organic free radicals well
elucidated.2'6 These studies, most notably those of Kharasch et a/.,4 7 demonstrated that
most radical processes can be expressed in terms of a small number of elementary
steps, or variations thereof, as shown below in Figure 1-1.8
A*
+
B-
A-B
coupling / homolysis
(1-1)
A*
+
B-D
A-B + D *
substitution (Sh2)
(1-2)
A •
+
B=D
A-B-D'
addition / p-fission
(1-3)
A*
+
e
A- ;
electron transfer
(1-4)
A ‘
_
e
A +
Figure 1-1. The Elementary Mechanistic Pathways of Free Radical Reactions.
Despite these breakthroughs, organic free radical reactions continued for years
to be regarded as unpredictable, unselective, and in general inadaptable to synthetic
application. Fortunately, subsequent kinetic and thermochemical studies have served to
uncover the factors governing the reactivity of organic radicals, and consequently in
recent years sentiment toward the utility of free radicals in synthesis has drastically
changed. Indeed, the number of elegant works in the literature based on radical
1

2
mediated transformations is a testament to their applicability in the construction of
natural products and other complex synthetic targets.9"15 It is the purpose of this
introductory chapter to acquaint the reader with the fundamental types of free radical
processes which occur in organic systems, as well as to provide an overview of the
wealth of physical studies which have given rise to the current level of understanding of
organic radical reactivity.
Radical Chain Processes
Most free radical reactions occur via a sequence of chain events, propagated by
intermediate steps during the course of the reaction. An example illustrating a
competition between two potential pathways is provided in Figure 1-2.
Initiation:
In' +
Propagation: M * +
R- +
Propagation: r* • +
In-In â–º 2 In*
M-H â–º In-H +
R-X â–º M-X +
M-H ^ R-H +
*r
R ‘ ► R' *
M-H â–º R'-H +
(1-5)
M’
(1-6)
R*
(1-7)
M *
(1-8)
(1-9)
M-
(1-10)
Figure 1-2. Radical Chain Process Involving Competition Between Rearrangement
versus Hydrogen Atom Transfer from a Donor Molecule M-H.
Homolysis of an initiator, typically accomplished by thermal or photochemical
means, provides a source of (often metal centered) radicals M’ (equation 1-6) from
which intermediate radicals R* are generated by reaction with a suitable precursor R-X
(equation 1-7). This species encounters one of two fates: trapping, in this case by
hydrogen atom donor, to yield R-H (equation 1-8) or transformation via a unimolecular or
bimolecular process (equation 1-9) to form radical R'\ itself then trapped producing

3
R'-H. In either case, additional metal radicals are formed and the chain process
continued via the propagation steps given in equations 1-7, 1-8 and 1-10.
The distribution of products R-H and R'-H is governed by the relative propensity
of R* toward rearrangement versus trapping (that is, kr and kH,) the latter dependent on
both the nature of R* and the type of trapping agent employed. In systems where
trapping is fast relative to rearrangement (kH » kr), the partitioning radical R* is
converted to R-H with little or no rearranged product. However, if kH and kr are of
comparable magnitude, product mixtures result. An understanding of the reactivity of a
radical intermediate toward potential competing processes is therefore essential for the
design of useful kinetic experiments, as well as for the development of effective synthetic
strategies.
For an efficient chain process, it is necessary that the propagation steps are
rapid relative to chain termination steps, thereby maintaining a low but constant
concentration of radical intermediates. Besides the obvious practical benefit (higher
product yields) resulting from such a condition, the occurrence of undesired chain
termination side reactions such as disproportionation and radical-radical coupling,
possibly complicating kinetic analyses, is minimized. In many cases, this may be
achieved by judicious selection of the type and concentrations of precursor R-X and
trapping agent.
This procedure enjoys wide application in both kinetic and synthetic studies
requiring the controlled generation of radical intermediates. One of its variants, likely the
most commonly used procedure for the indirect (competitive) determination of the rates
of organic radical reactions, is based on the trialkylstannane reduction of an alkyl halide
(the "Tin Hydride Method").1618 Other modifications of this general procedure exist,
accommodating a variety of radical precursors and trapping agents; a discussion of time-
resolved and competitive techniques utilized in radical kinetic measurements is provided
later in the chapter.

4
Finally, it is important to note the implication of kinetic control in the above
discussion. That is, that the distribution of products R-H and R'-H may be ascribed to
the relative values of kH and kr hinges on the absence of any thermodynamic
equilibration of products under the reaction conditions. This is of vital importance in the
design and interpretation of competitive kinetic studies and is discussed in detail in
Chapter 3.
Radical reactivity is dependent on the "complex interplay" of thermodynamic,
steric, and polar considerations.19 The relationship between enthalpies of activation and
heats of reaction, the basis of the thermochemical kinetic approach of Benson,20 was
recognized early on and holds for a number of radical addition and substitution
reactions, where the order of reactivity often parallels exothermicity.21'23 This relation
has led to such overgeneralizations as "radical reactions follow the most exothermic
available pathway" or". . . afford the most stable possible product."8 However, reaction
thermochemistry is not the sole, nor even predominant decisive factor in the outcome of
radical reactions. Nonbonding interactions and the electronic influence of substituents in
ground and transition states (which may be rationalized in terms of Frontier Molecular
Orbital (FMO) theory)24 25 will also play a role. A discussion of the combination of these
effects as manifested in hydrogen atom abstractions and inter- and intramolecular
additions, the most commonly occurring and well-characterized reactions of organic free
radicals, will now be presented.
Hydrogen Atom Abstraction Reactions
The vast majority of free radical applications involve the use of an organometallic
hydride of the type R3M-H (most commonly, where M = Sn, Si, or Ge) as a hydrogen
atom donor and chain propagation agent, the properties of which have been the focus of
extensive investigation by kineticists. Metal-hydrogen bond dissociation energies
(BDEs) along with activation parameters and associated absolute rate constants for

5
hydrogen atom transfer to /7-alkyl hydrocarbon radicals by a series of donors R3M-H
have been determined and are provided below in Table i-i 16 26 32
Table 1-1. Bond Dissociation Energies with Activation Parameters and Rate Constants
for Hydrogen Atom Transfer to Hydrocarbon Radicals by R3M-H.
R,M-H
BDE, kcal mol1
loq A
Ea. kcal mol'1
kH. 106 M'1 s'1 (298 K)
nBu3SnH
73.7
9.06
3.65
2.3
(TMS)3SiH
79.0
8.86
4.47
0.38
nBu3GeH
82.6
8.44
4.70
0.093
(TMS)2Si(CH3)H
82.9
8.89
5.98
0.032
Et3SiH
90.1
8.66
7.98
0.00064
Analysis of the data demonstrates that for hydrogen atom abstraction by
structurally similar radicals from this series of donors, a direct relation holds between the
rate of transfer and the strength of the metal-hydrogen bond being broken. This is
depicted graphically in Figure 1-3. In addition, it is noted that in each case the pre¬
exponential term in the Arrhenius relation remains relatively constant. Thus, the rate
variations within the series are due almost entirely to differences in activation energies.
Figure 1-3. Plot of In kH for Alkyl Radicals versus M-H Bond Dissociation Energies for
Hydrogen Atom Donors R3M-H in Table 1-1.

6
However, as previously mentioned, relative thermodynamics is not the only factor
which influences the kinetics of hydrogen atom transfer. The fast donor thiophenol
(PhSH) reacts with primary alkyl radicals with a rate constant of 1.36 x 108 M 1 s'1 at
298 K,33 and has been employed as a trapping agent in competitive kinetic studies
involving strained or otherwise highly reactive radicals with rearrangement rates upward
of 1011 s'1 and thus with lifetimes on the picosecond timescale.34 This enhanced rate of
transfer, not commensurate with its S-H BDE of 82.0 kcal mol'1,35 gives rise to a severe
deviation from the plot in Figure 1-3 and indicates the presence of other influences.
Table 1-2. Absolute Rate Constants for Hydrogen Atom Transfer to ferf-Butoxyl
Radicals by R3M-H.
R,M-H
kH. 106 M'1 s'
nBu3SnH
220
(TMS)3SiH
110
nBu3GeH
80
Et3SiH
5.7
Further evidence may be found in the rates of hydrogen atom transfer to tert-
butoxyl radicals by the same series of hydrogen atom donors, provided in Table 1-2 and
illustrated graphically in Figure 1-4.36 37 It is observed that for fe/t-butoxyl radicals, rates
of hydrogen abstraction are at least two orders of magnitude greater than those of their
n-alkyl counterparts. Although the relative strengths of the newly formed C-H or O-H
bonds will certainly play a role, the difference in BDE between fBuO-H and n-alkyl C-H
bonds (105 and 100 kcal mol'1, respectively)38 is not sufficient to explain the increase in
reactivity, especially in light of the fact that such rapid hydrogen atom abstractions
should proceed with early transition states.36
At this time, the absolute rates of reduction of terf-butoxyl radicals by thiophenol
have yet to be determined. However, a series of competition studies by Hartung and

7
Gallou39 involving 4-pentenyl-1-oxy radicals and utilizing naphthalene 2-thiol (NpSH) as
a trapping agent have determined a ratio [ kH (NpSH) / kH (nBu3SnH) ] of 1.4. By
comparison, n-alkyl radicals afford the ratio [ kH (PhSH) / kH (nBu3SnH) ] = 59.1.
Although a leveling effect may be partly responsible for the compressed ratio of rates for
tert-butoxyl radicals (which are indeed within an order of magnitude of the diffusion-
controlled limit)40 it is logical to assume based on the aforementioned examples that
hydrogen abstraction reactivity will be governed to some extent by factors other than
simple relative BDE values of the donor species.
Figure 1-4. Plot of In kH for terf-Butoxyl Radicals vs. M-H Bond Dissociation Energies for
Hydrogen Atom Donors R3M-H in Table 1-2.
Chatgilialoglu et al,3637 have attributed such differences in reactivity to a
polarized, or charge separated, transition state of the type depicted in Figure 1-5. Here
it can be seen that in the case of an electropositive metal hydride donor, hydrogen atom
transfer to alkoxyl radicals (b) is facilitated by greater stabilization of partial negative
charge on oxygen relative to carbon, with a resultant decrease in activation barrier.
8+ 8' 8+ S'
(a) R3M H R (b) R3M H OR
Figure 1-5. Charge Polarized Transition State for Hydrogen Abstraction from R3M-H by
(a) Alkyl and (b) Alkoxyl radicals.

8
In the case of thiol donors, the opposite situation ensues. The greater
electronegativity of sulfur relative to tin (or other metal atom) gives rise to a reversal in
the transition state charge distribution. This arrangement, involving a partially negatively
charged sulfur atom, is better suited to the more nucleophilic alkyl radical, where in the
alkoxyl case a less- or non-polarized transition state results. Such a mismatch in the
latter is partially responsible for the decrease in rate enhancement for hydrogen
abstraction from thiols by alkoxyl radicals, relative to their alkyl analogues.
Frontier Molecular Orbital Theory of Atom Abstraction Reactions
FMO theory provides a satisfying rationale for the kinetic characteristics of
hydrogen abstraction reactions of free radicals. In general terms, radicals are species
containing an unpaired electron in a singly occupied molecular orbital (SOMO), which in
the ground state of the radical is its highest occupied orbital. According to the FMO
concept, during the course of the reaction this SOMO will interact with both the highest
occupied (HOMO) and lowest unoccupied (LUMO) molecular orbitals of the donor
molecule. Such interactions between these "frontier molecular orbitals"24 are not
necessarily equal. The extent of SOMO-HOMO and SOMO-LUMO interaction is
governed by their energy values, the strongest interaction occurring between orbitals
closest in energy.
It is these values, influenced by atom type as well as neighboring substituents,
from which the relative descriptors such as "nucleophilic" and "electrophilic" are derived
and provide the basis for the previously described concept of "polar factors." Electron
donating substituents generally serve to raise both HOMO and LUMO energies, with
electron withdrawing groups resulting in lowering. Radicals possessing a low energy
SOMO will display electrophilic character, whereas a higher energy SOMO gives rise to
a more nucleophilic species. Figure 1-6 depicts the FMO interactions between radical
and donor in each of these cases.

9
During the abstraction process, the primary interaction involves the radical
SOMO and the o and a* orbitals of the donor M-H bond. The antibonding a* orbitals of
the donor are typically quite high in energy, and thus in atom abstraction reactions the
SOMO-HOMO interaction dominates.
SOMO -+-(
\ »
\ i
SOMO
,h4- homo
HK
(a) (b)
Figure 1-6. FMO Diagram Illustrating the SOMO-HOMO Interaction Between (a) a
Nucleophilic Alkyl and (b) an Electrophilic Alkoxyl radical.
Here it is seen that the lower-energy SOMO of the alkoxyl radical (ca. -12 eV, as
determined from ionization potential measurements)25 leads to a reinforced interaction
with the donor HOMO (case b). This greater stabilizing interaction results in a lowered
activation barrier and hence a more facile transfer reaction, compared to the more
nucleophilic alkyl radical, (case a) whose SOMO energies range from -6.9 to -9.8 eV.25
Some of the most striking examples of such "polar" factors involve systems
where fluorine substitution has taken place at, or adjacent to, the radical center. This is
elaborated upon in Chapters 2 and 3, where the effects of fluorination on the structure
and reactivity of free radicals are discussed and compelling evidence provided based on
kinetic studies of hydrogen transfer to such partially and fully fluorinated alkyl radicals.
Intermolecular Radical Addition Reactions
Over the past twenty years, the intermolecular addition reactions of free radicals
(as well as their intramolecular cyclization counterparts) have become an important

10
addition to the arsenal of C-C bond formation methods available to the synthetic organic
chemist. Their mild means of generation from a variety of precursors and tolerance for a
wide variety of functional groups provide distinct advantages over ionic processes.
Alkyl radical additions to carbon-carbon double bonds are highly exothermic, as a
new o bond is formed at the expense of a n bond (in the case of methyl radical addition
to ethylene (Figure 1-7), AHn Hammond postulate43 such additions should proceed via early transition states, with low
barriers of activation. This is indeed the case, as supported by a wealth of both
experimental44 and theoretical41 45-49 data.
CH3 + ch2=ch2 â–º ch3ch2ch2
Ea = 7.9 kcal mol'1
AHâ„¢ = -22.6 kcal mol'1
Figure 1-7. Addition of Methyl Radical to Ethylene, Yielding n-Propyl Radical.
Experimental Activation Energies and Heats of Reaction are Shown.
FMO Theory of Radical Additions
Quantum mechanical molecular orbital calculations at levels of ab initio theory
ranging from UHF to UQCISD(T) and varying basis sizes from 3-21G to 6-311G(2df,p)
are consistent in their characterization of the transition structure for the above reaction
154.7°
Figure 1-8. UHF/6-31G(d) Transition Structure and Relevant Geometrical Parameters
for Addition of Methyl Radical to Ethylene.

V
LUMO
LUMC
SOMO
HOMO
(a)
HOMC
Figure 1-9. FMO Diagram for Addition of (a) Nucleophilic and (b) Electrophilic Radicals
to Alkenes.
Table 1-3. Some Relative Rates of Addition of Methyl and teri-Butyl Radicals to Alkenes
CH2=CHX.
X
krel (CH3*)
krPi ((CH313CD
H
1
1
ch3
0.7
0.74
och2ch3
0.31
F
0.9
Cl
4.2
13.2
CN
343
1920
3 Data for ethyl radical.
(Figure 1-8), which possesses an incipient C-C bond distance of ca. 2.23 - 2.27 Á. The
C-C-C attack angle of 109.1° is rationalized in FMO terms based on a primary interaction
between the radical SOMO and the LUMO of the alkene. It is in such reactions with high
exothermicities and early transition states that FMO interactions are most
substantial.24 50 This postulate enjoys experimental support; for the f-butyl radical, a
correlation exists between rates of addition to alkenes and the experimentally

12
determined electron affinities of the latter.2151 Such an FMO interaction for nucleophilic
radicals is shown in Figure 1-9 (case a) and is influenced by substituents on both radical
and olefin, a raising of radical SOMO and/or lowering of alkene MO energies
strengthening the SOMO-LUMO interaction and enhancing the rate of addition as seen
from the data in Table 1-3.21,52 Here, the greater nucleophilic character of f-butyl relative
to methyl is evident from its enhanced rate of addition to olefins bearing electron
withdrawing groups.
As previously discussed, electronegative substituents at the radical center which
substantially lower its SOMO energy will impart electrophilic character and reinforce the
transition state SOMO-HOMO interaction (Figure 1-9, case b). Indeed, it has been
shown that rates of addition of dicyanomethyl53 and perfluorinated54 radicals correlate
with the ionization potentials of the substrate alkenes. The intermediate behavior of
"ambiphilic" radicals, such as malonyl and (tert-butoxycarbonyl)methyl, has also been
documented, these species yielding "U"-shaped correlations between rates of addition
and alkene IP and EA values.55'57
A more thorough presentation of kinetic results obtained to date for the addition
of partially and fully fluorinated radicals to alkenes is given in Chapters 2 and 3.
Steric Effects; Reqiochemical Preferences in Addition
Competition studies on both nucleophilic and electrophilic radicals have provided
for some generalizations in terms of the regiochemical preference for addition to
unsymmetrically substituted olefins.52 The preferred orientation of radical addition
occurs to the unsubstituted end of the double bond, attributed to steric repulsion but also
influenced by the effect of substituents on the coefficients of the HOMO and LUMO of
the alkene. Such FMO effects can be the decisive factor in polysubstituted olefins if
steric effects are in opposition. Strongly spin delocalizing substituents on the alkene
reinforce such sterically induced regiochemical preferences and exert slight rate

13
enhancing effects; however, as such additions occur through early transition states the
effect of exothermicity on the kinetics of addition should be minimal.
The concepts introduced in the aforementioned discussion on intermolecular
radical additions extend to their intramolecular cyclization analogues, an overview of
which will now be presented.
Intramolecular Addition Reactions: Radical Cvclizations
The intramolecular addition reactions of alkenyl radicals enjoy a strong foothold
among the available strategies for the construction of cyclic organic molecules. In
addition to their synthetic utility, the kinetic, regioelectronic and stereoelectronic
characteristics of radical cyclization reactions as a function of substituent continue to fuel
an abundance of fundamental physical organic structure-reactivity investigations, more
than thirty years after the first report of the archetypal radical ring closure, cyclization of
hex-5-en-1-yl radical 1 (Figure 1-10).58
(98%) (2%)
1 2 3
Figure 1-10. Cyclization of Hex-5-en-1-yl Radical 1 to Cyclopentylcarbinyl (2) and
Cyclohexyl (3) Radicals. At 25° C, 5-exo Closure Dominates 49:1.
The most striking aspect of this reaction lies in the preferred regiochemistry of
addition. In the case of the parent hydrocarbon, 5-exo59 cyclization dominates (Eact [5-
exo] = 6.8 kcal mol'1, Eact [6-endo] = 8.5 kcal mol'1)27 60 yielding the less
thermodynamically stable primary cyclopentylcarbinyl radical 2. This finding has
provided the driving force for a number of experimental61 and theoretical62-68

14
investigations geared toward the understanding of radical cyclization regiochemistry in
the hydrocarbon and related substituted systems.
Early explanations,69 later advanced by semiempirical techniques,65 attributed
this result to a less negative entropy of activation for cyclization to 2. Although
experiment demonstrates this to be true, the difference (AS*1i5 - AS*i,6 = 2.8 eu)60 is
insufficient to completely account for the observed regiochemistry; the preferred mode of
cyclization resulting primarily from enthalpic (AH*16 - AH*15 = 1.7 kcal mol'1) rather than
entropic factors. Ab initio computations66 lend support to this conclusion.
Transition structures for 5-exo and 6-endo cyclization of 1 have been located
using a variety of theoretical treatments. The UHF/6-31G(d) structures leading to 2 and
3 are shown in Figures 1-11 and 1-12.
Figure 1-11. Two Views of the UHF/6-31G(d) "5-exo-chair" Transition Structure for
Cyclization Hex-5-en-1-yl Radical 1 to Cyclopentylcarbinyl Radical 2.
Figure 1-12. Two Views of the UHF/6-31G(d) "6-endo-chair" Transition Structure for
Cyclization of Hex-5-en-1-yl Radical 1 to Cyclohexyl radical 3.

15
Spellmeyer and Houk66 have also postulated additional "boat-like" transition
structures on the basis of molecular mechanics calculations parameterized by ab initio
results of model systems. Inclusion of these "boat-like" structures as viable competing
pathways was found to be necessary for the accurate prediction of regio- and
stereoselectivities in cyclizations of alkyl substituted and heteroalkenyl radicals. The
existence of such structures is corroborated by higher level ab initio treatments
performed as part of the present study and are shown in Figures 1-13 and 1-14.
Figure 1-13. Two Views of the UHF/6-31G(d) "5-exo-boat Transition Structure for
Cyclization of Hex-5-en-1-yl Radical 1 to Cyclopentylcarbinyl Radical 2.
Figure 1-14. Two Views of the UHF/6-31G(d) "6-endo-twist-boaf Transition Structure
for Cyclization of Hex-5-en-1-yl Radical 1 to Cyclohexyl Radical 3.
Upon inspection of the forming bond lengths and angles of the 5-exo structure in
Figure 1-11, its similarity to the transition structure for addition of methyl radical to
ethylene is readily apparent. The C-C-C angle of attack, 109.7°, is practically identical to

16
that in Figure 1-8 and fits the requirement for overlap of the radical SOMO with the 71*
orbital of the alkene moiety. From Figure 1-12 it is observed that this angle is
significantly reduced (98.4°) in the 6-endo approach. Thus the required disposition of
centers for optimal FMO overlap is more readily achieved in the 5-exo transition
structure, leading to the kinetically preferred cyclization product.
Substituent Effects on the Kinetics and Reqiochemistry of 5-Hexenyl Cvclizations
Radical 1 cyclizes to cyclopentylmethyl radical 2 with a rate constant (kC5) of
2.3 x 105 s'1 at 25 °C.27 In the parent hydrocarbon, 6-endo closure competes to a very
minor extent (kCe = ca. 4.7 x 103 s'1). However, it will be shown that the rates and
regiochemical preferences can be substantially affected by substitution at both the
radical center and terminal alkene.
Alkyl Substitution; Steric Effects
Beckwith et al. have reported rate constants for a number of alkyl-substituted
hexenyl radicals.70 71 The rates of 5-exo closure as a function of gem-dialkyl substitution
on the aliphatic portion of the hexenyl chain is shown in Table 1-4.
It is seen that substitution at the radical center has a nearly negligible effect on
the rate of cyclization, due to offsetting polar and steric considerations. Conversely, a
significant (> 10-fold) rate enhancement is observed with internal substitution (systems
6, 8, and 9), accelerated by relief of steric compression between alkyl groups during ring
formation (the "Thorpe-lngold", or "gem-dimethyl" effect).72
Intermediate kinetic behavior would be expected of monosubstituted 5-hexenyl
systems. The data in Table 1-5 show this to be the case. In addition, a stereochemical
preference for cis- or frans-dimethylcyclopentanes, depending on the location of the
substituent on the chain, is observed. This has been rationalized by Beckwith et al71

17
Table 1-4. 5-exo Cyclization Rate Constants for gem-Disubstituted 5-Hexenyl Radical
Derivatives.
Cyclization Reaction kC5, 105 s'1 (298 K)
based on the cyclization transition structure depicted in Figure 1-11. Although "early" in
terms of the forming C-C bond, the overall orientation of atoms in this structure is quite
product-like. According to this rationale, substituents on the aliphatic fragment are
likened to those in chair cyclohexane, which then occupy an equatorial position in the
chair transition structure. Minor products are assumed to derive from the occupation of

18
axial positions. The latter has been disputed by Spellmeyer and Houk,66 whose model
indicates that such secondary products originate from an equatorial disposition of
substituents in the boat-like transition structure of Figure1-12, rather than from an axial
orientation in the chair.
Table 1-5. C/s- and Trans- 5-exo Cyclization Rate Constants for Monosubstituted 5-
Hexenyl Systems. Rate Constants are for 298 K.
Cyclization Reaction kC5 (c/s), 105 s'1 kC5 (trans), 105 s'1
14 c/s-13 frans-13
15 cis-11 transé 1
3 Reference 73. b Reference 71.
In the above examples, 5-exo products are formed either predominantly (> 97%)
or exclusively. Substitution at the vinyl group leads to marked changes in regiochemical
ratios as indicated by the data in Table 1-6. Replacement of hydrogen by methyl (16) or
isopropyl (19) at C5 results in preferential formation 6-membered rings 18 and 21.
Inspection of the data indicates that this shift in regiochemistry is not due to a significant

19
Table 1-6. Rate Constants at 338 K for 5-exo and 6-endo Cyclization for Vinyl-
substituted 5-Hexenyl Radicals.
Cyclization Reaction /cqs, 105s~1 /cqs, 105s'1
extent to rate enhancement for 6-endo closure, but rather a substantial (44-fold)
retardation of 5-membered ring formation due to a combination of 1,5 steric hindrance
and back strain engendered at C5 upon adaptation of sp3 character. Substitution at both
C5 and C6 again favors 5-exo closure, the rate of which decreased relative to the parent
system. Disubstitution at C6 (25) gives rise to a slight (2.3-fold) rate enhancement for 5-
exo closure, sufficiently explained on thermodynamic grounds, which dominates 6-endo

20
cyclization by a factor of at least 220. The above data indicates that the kinetic and
regiochemical characteristics of alkyl-substituted 5-hexenyl cyclizations may be
sufficiently rationalized by steric considerations. The effects of substitution by
conjugating, heteroatom-containing groups is outlined below.
FMO Considerations
Studies of 5-hexenyl systems bearing "polar" subsitutents have been
investigated by Newcomb.16 74'77 In line with those of intermolecular radical additions,
the kinetics of radical ring closure will be influenced by the impact of substitutents on the
SOMO-HOMO and SOMO-LUMO interaction in the cyclization transition state.
Table 1-7. 5-exo Cyclization Rate Constants For a-Donor- and a-Acceptor-substituted
6,6-Diphenyl-5-hexenyl radicals. Rate Constants are for 298 K.
X Y
•
5-exo
â–º
X Y Ph
0^'ph
Svstem
Ph
X
Y
kn*. 107 s'1 (298 I
28 -> 29
H
H
4
30 -â–º 31
H
ch3
2
32 -» 33
ch3
ch3
1
34 -> 35
H
och3
4
36 -» 37
H
co2ch2ch3
«3.7
38 -> 39
ch3
co2ch2ch3
0.04
40 -> 41
ch3
CN
0.03
3 Calculated from the Arrhenius parameters provided in reference 76.
As seen from the data in Table 1-7, only a very minor effect is exerted by either
a-donor or a-acceptor substituents, relative to parent system 28. The marked decrease

21
in rate for 38 and 40 is attributed to an increase in activation energy due to enforced
planarity at the radical site induced by the 7t-delocalizing substituents C02CH2CH3 and
CN.51,78
Table 1-8. Absolute Rate Constants at 298 K for 5-exo Cyclization of 5-Hexenyl
Systems Bearing Vinylic Donor and Acceptor Substituents.
Cyclization Reaction kc5, 105 s'1 (298 K)
a Reference 27. b Calculated from the Arrhenius parameters provided in reference 74.
Substitution at the vinyl terminus, especially by strong resonance-withdrawing
groups can significantly accelerate the rate of ring closure. Although possessing a
radical stabilizing group, methoxy analog 42 enjoys only a very minor increase in rate.
The donor substituent raises the energies of the frontier orbitals, increasing the SOMO-
HOMO interaction but widening the SOMO-LUMO energy gap, the latter more important

22
for relatively nucleophilic alkyl radicals (Figure 1-9). Substituents which serve to lower
the FMO energies should reinforce the SOMO-LUMO interaction, leading to rate
enhancement. The nearly 300-fold increase resulting from cyano substitution (44)
reflects such an effect. The slight rate increase 46 relative to 45 has been explained on
the basis of the suggested slight extra "push-pull," or "captodative" stabilization
manifested in donor-acceptor disubstituted systems.79 80
The importance of kinetic control was previously mentioned. Care must be taken
in assessing the potential for reversibility in such intramolecular additions, which may
obscure the effect of steric and/or polar influences on reaction kinetics. This is
demonstrated in the 5-exo:6-endo product ratios of highly stabilized systems 48 and
51.8182
Table 1-9. Product Ratios for 5-Hexenyl Cyclization Reactions Under Full or Partial
Thermodyamic Control.
Cyclization Reaction % 5-exo % 6-endo
Given the data and discussion provided in the above sections, a review of the
direct and competitive techniques utilized in the determination of rates of organic radical
reactions is now in order.
Methods for Determination of Organic Radical Kinetics
The development of indirect competitive methods, in conjunction with laser flash
photolytic generation and time-resolved detection of transient intermediates, has greatly

23
expanded the dynamic range available for the measurement of radical reactions,
especially those at the upper end of the kinetic scale. Such advances have provided for
the use of a variety of precursors and the accurate determination of rate constants for
reactions approaching the diffusion-controlled limit in solution.16
Laser Flash Photolysis; Direct Measurement of Addition Rates
In the time-resolved laser flash photolysis method, described in detail in the
literature,83 radicals are generated from precursors possessing a suitable chromophore
by a laser pulse of the appropriate wavelength. Alkyl iodides, diacyl peroxides, and the
O-acylthiohydroxamic esters of Barton et a/.84 are most commonly utilized in this regard
(Figure 1-15).
(a) R-l ► R •
(b)
O
.A
o'Y
o
hv
2
-C02
â–º 2 R *
fast
(c)
-co2
fast
R •
Figure 1-15. Laser Flash Photolytic Generation of Radicals R* from (a) Iodide, (b) Diacyl
Peroxide, and (c) O-Acylthiohydroxamic Ester Precursors.
These radicals so generated undergo further reaction, usually bimolecular
addition or unimolecular cyclization to a (typically phenyl-substituted) double bond (a
styrene in the case of bimolecular additions). The increase in the characteristic
absorption of this intermediate benzyl radical (^max * 320 nm) is then followed in a time-
dependent manner by UV-visible spectroscopy (Figure 1-16). For bimolecular additions,
this experimental growth curve is fit to the expression in Equation 1-11, yielding absolute

24
rate of addition kadd. With such data in hand, this addition can now serve as a competing
basis reaction for the determination of rates of other transformations involving the same
or structurally similar radical.
k0bs = k0 + ^0 [alkene] (1-11)
R-X
A.max ca. 320 nm
Figure 1-16. LFP Generation of R* and Detection of Transient Benzyl Radical Adduct
for Determination of Absolute Rate of Addition kaM.
Indirect Methods (Competitive Techniques)
Indirect kinetic methods involve the partitioning of an intermediate between two
competing pathways, one with a known rate constant and the other whose rate constant
is to be determined. Post facto product analyses, typically by chromatographic or
spectroscopic means, provide a ratio of rate constants from which the new kinetic value
is obtained. Such radical kinetic measurements usually involve competition between two
bimolecular reactions or a bimolecular reaction competing with unimolecular
rearrangement; examples of both instances are provided in Figures 1-17 and 1-18.
Determination of rates of hydrogen abstraction by this method involves the
generation of R* in the presence of two trapping agents; in Figure 1-17, styrene and the
hydrogen atom donor. Both traps are typically present in excess to ensure pseudo-first
order behavior. Radical R* may undergo addition to styrene (with known rate constant
kadd) forming the intermediate benzyl radical, itself trapped with excess hydrogen donor,
yielding closed shell product with a rate which is kinetically unimportant provided the
addition reaction is irreversible. Alternatively, R’ is trapped directly by hydrogen atom
donor with a rate constant /rH, which may be obtained from the pseudo-first order relation

25
[ reduced ]
[*H] [R*][M-H]
(1-12)
[ adduct ]
[*add] [R*][CH2=CHR']
where [reduced] and [adduct] are the final product concentrations, [M-H] is the
concentration of hydrogen atom donor, and [CH2=CHR'] the concentration of alkene trap.
For accurate kinetic determinations, a series of runs is performed where trap
concentrations are varied, (again, maintaining at least a five-fold excess) a plot of
product ratios versus that of trapping agents yielding the ratio kH / kadd.
R •
Figure 1-17. Competition Between Bimolecular Addition to an Alkene with Rate
Constant /cadd and Bimolecular Trapping by Hydrogen Atom Donor with Rate Constant
/(h.
Determination of rates of cyclization are performed in a similar manner, using
hydrogen atom abstraction (with the known value of kH) as the competitive basis
reaction. Unimolecular rearrangement competes with bimolecular trapping with an
excess of hydrogen atom donor (Figure 1-18) to yield intermediate cyclic radicals, further
trapped to form characterizable products.
The ratio of products of cyclization versus hydrogen abstraction are obtained
from the pseudo-first order relation in Equation 1-13, a plot of the ratio of products of
hydrogen abstraction to those of cyclization as a function of trapping agent concentration
affording ratios kC5 / kH and /cC6 / kH.
Finally, the importance of an efficient chain process should be reemphasized.
High conversions of precursors, although important for any radical reaction, are crucial in

26
competitive kinetic experiments. The reliability of data resulting from indirect methods is
directly dependent on high "mass balance" values, those of 90% or greater typically
being desired.
[reduced ]
[ cyclized ]
[kH] [r.] [M-H]
[ fcCn ] [ R #1
(1-13)
Figure 1-18. Competition Between Bimolecular Hydrogen Atom Abstraction with Rate
Constant kaM and Unimolecular 5-exo and 6-endo Cyclization with Rate Constants kCs
and kC6.
Such competitive processes have been employed extensively by the Dolbier
research group in the investigation of the rates of addition, hydrogen atom abstraction,
and cyclization of a variety of fluorinated open shell systems, providing the first
quantitative kinetic data for this class of reactive intermediates.54,85'91
Conclusion
The preceeding discussions have attempted to provide the reader with an
introduction to the chemistry of organic free radicals. Kinetic data for hydrogen
abstraction, addition, and cyclization reactions of hydrocarbon radicals, important
benchmarks for comparison of reactivity with other substituted systems, was provided.

27
Substituent effects, rationalized on the basis of a combination of thermodynamic, steric
and FMO considerations, were discussed.
The following chapter provides a review of the effects of fluorine substitution in
organic molecules, including fluorinated radicals. Previous research efforts in this area
by the Dolbier group are summarized, setting the stage for the presentation of results of
the current study.

CHAPTER 2
THE FLUORINE SUBSTITUENT IN ORGANIC SYSTEMS
Introduction
Incorporation of fluorine into organic molecules often imparts dramatic alterations
in structure and reactivity. These effects are induced by three major characteristics
inherent to the fluorine atom: extreme electronegativity, non-bonded electron pairs, and
relatively small size.
Fluorine possesses the highest electronegativity of all the elements, with a value
of 4.10 on the Pauling scale, compared to oxygen (3.50), chlorine (2.83), bromine (2.74),
carbon (2.50), and hydrogen (2.20).92 As a substituent in organic systems, this results in
strong inductive withdrawal of electron density through the a molecular framework and
highly polarized bonds with substantial ionic character.
Three non-bonding pairs of electrons in 2p orbitals similar in size to those of
other second-row elements provide for optimal overlap, and therefore an offsetting back
donation of electron density into the molecule to which it is bonded.
The accepted van der Waals radius of fluorine, 1.47 A, suggests minimal steric
impact in comparison with other halogens (chlorine, 1.73 A; bromine, 1.84 A; iodine,
2.01 A; carbon, 1.70 A; oxygen, 1.52 A; hydrogen, 1.20 A).93 This has allowed for the
complete replacement of hydrogen by fluorine in organic systems, a feat not possible to
such an extent with any other element.
The following sections, based on a number of excellent reviews,94'97 provide an
introduction to the fascinating behavior exhibited by fluorinated stable molecules and
reactive intermediates due to a combination of the above effects.
28

29
Structure, Bonding, and Reactivity in Saturated Systems
The data provided in Tables 2-1 and 2-295,97 reveal a trend unique to fluorine
within the halogenated methanes. An incremental shortening of C-F interatomic
distances, with a resultant increase in bond dissociation energies, is observed as the
series is traversed. No such trend exists for any other member of the halomethane
family; on the contrary, it is seen from the data in Table 2-2 that such C-X BDE values
instead decrease with increasing halogen content. Strengthening of C-H bonds is also
observed within the fluoromethanes (CH3F, 101.3 kcal mol"1; CH2F2, 103.2 kcal mol'1;
CF3H, 106.7 kcal mol"1).95
Table 2-1. Carbon-Halogen Interatomic Distances (Angstroms) of Halomethanes
X
CH3X
CH2X2
chx3
CX4
F
1.385
1.357
1.332
1.319
Cl
1.781
1.772
1.758
1.767
Br
1.939
1.934
1.930
1.942
Table 2-2. Carbon-Halogen Bond Dissociation Energies (D°, kcal mol"1) of
Halomethanes
X
ch3x
CH2X2
chx3
CX4
F
108.3
119.5
127.5
130.5
Cl
82.9
81.0
77.7
72.9
Br
69.6
64
62
56.2
I
57.2
51.3
45.7
_
Data for geminally fluorinated ethanes parallel that of the methane series,
demonstrating a progressive strengthening and shortening of both C-C and C-F bonds
with increasing fluorination (Table 2-3). Conversely, vicinal fluorination gives rise to the
opposite effect on C-C bonds, a steady lengthening and weakening being observed.

30
A variety of hypotheses have been put forth to explain the observed trends. One
rationalization, invoked by Pauling98 and based on valence bond theory, involves "double
bond, no bond" resonance of the type depicted in Figure 2-1.
Table 2-3. Interatomic Distances and Dissociation Energies of Fluoroethanes
Ethane
r (C-C). A
D° (C-C). kcal mol'1
r (C-F). A
D° (C-F). kcal mol'1
ch3-ch3
1.532
90.4
-
-
ch3-ch2f
1.502
91.2
1.398
107.9
ch3-chf2
1.498
95.6
1.343
Unknown
ch3-cf3
1.494
101.2
1.335
124.8
ch2f-cf3
1.501
94.6
Unknown
109.4 (CH2F)
cf3-cf3
1.545
98.7
Unknown
126.8
As the degree of geminal fluorination is increased, the number of such resonance
forms involving doubly bonded fluorine increases (0, 2, 6, and 12 in the case of CH3F,
CH2F2, CH3F, and CF4, respectively). This is supported by ab initio calculations at the
RHF/4-31G and 4-31 G(d) levels,99'101 which illustrate back donation of electron density
from fluorine into the C-F a* orbitals. It is further observed that the overlap population
between the 2p orbitals of carbon and those of fluorine increases continually with
successive fluorination; in contrast, such carbon-chlorine overlap populations decrease
steadily from CH3CI to CCI4.
F F
Figure 2-1. "Double Bond, No Bond" Resonance in Geminally Fluorinated Alkanes.
Alternative explanations based on hybridization schemes have also been
advanced. It is postulated that for electronegative elements bound to carbon,
rehybridization occurs causing an increase in the amount of p character directed toward

31
the substituent. Thus, in CH3F, the C-F bond possesses greater p character, with
greater s character in the C-H bonds. This rationale accounts not only for incremental
C-F bond strengthening, but also for the observed changes in geometry within the
fluoromethane series. Accumulation of p character in C-F bonds should lead to a
decrease in FCF bond angle, accompanied by HCH widening. This is consistent with
experimental observation (ZFCF in CF4, 109.5°; CHF3, 108.7°; CH2F2, 108.3°;102'103 for
CH2F2, ZHCH = 113.7°).103
Finally, a more recent argument has been advanced by Wiberg,104 on the basis
of Coulombic interactions between carbon and fluorine substituents. From charge-fitting
treatments based on calculated electrostatic potentials, a linear increase in positive
charge on carbon is observed, while the degree of negative charge on each of the
fluorine substituents remains quite constant. Thus, incremental fluorine substitution
strengthens not only new, but also previous, C-F bonds. This finding will be further
discussed in Chapter 3, as such ESP-derived charges calculated for larger fluorinated
systems as part of the present study were found to provide a cogent explanation for the
remarkable and contrasting effects of fluorination on the strengths of both C-C and a and
P C-H bonds.
Structure. Bonding, and Reactivity in Unsaturated Systems
The most reliable structural and 7T-BDE data for the fluoroethylenes is provided in
Table 2-4.95 9' Vinylic fluorine substitution results in shorter C=C bonds than in the
parent hydrocarbon, and shorter C-F bonds than fluoroalkanes bearing the same
number of geminal or vicinal fluorine substituents. Ab initio investigations by Radom"
and others105 107 attribute the C-F bond contraction to delocalization of fluorine 2p
electrons into the C-C n bond (depicted in resonance terms in Figure 2-2). Computed
atomic charges are consistent with this conclusion.

32
Table 2-4. Interatomic Distances, Angles, and n Bond Dissociation Energies of
Fluoroethenes.
fvU
X
o
II
X
o
CH,=CHF
CH,=CF,
CHF=CF,
O
m
-o
ii
o
T|
r (C=C), A
1.339
1.333
1.316
1.309
1.311
r (C-F), A
-
1.348
1.324
1.336
1.319
ZHCH, deg.
117.8
114.7
119.3
-
-
ZHCF, deg.
-
111.3
-
114.0
-
ZFCF, deg.
-
-
109.7
109.1
112.6
n D°, kcal mol'1
63-64
Unknown
62.8
Unknown
52.3
The marked decrease in FCF bond angles has been rationalized by Bernett108
and «oilman109 on the basis of hybridization arguments; Epiotis has advanced an
alternative explanation involving nonbonded attraction between fluorine atoms.110111
Figure 2-2. Fluorine 2p Electron Delocalization in Unsaturated Systems.
Photoelectron112 and electron attachment113 spectral data for the fluoroethylene
series are provided in Table 2-5. A significant lowering (over a range of 3.1 eV) of the a
MO energies is observed, with only a slight (ca. 0.3 eV) variation in the n MOs.
Stabilization of the a MOs is ascribed to extensive delocalization over the fluorine
substituents; such an interaction within the n system is diminished and counteracted by
strong C-F antibonding overlap.112 A steady increase in electron attachment energies
with successive fluorination can also be seen, attributed to destabilization of rc* resulting
from an antibonding interaction with the fluorine 2p AOs.113
Heats of hydrogenation provided in Table 2-695 illustrate the reactivities of
fluorinated alkenes. In general, transformation of a polyfluorinated olefin into a saturated

33
Table 2-5. Vertical Ionization Potentials (I.P.) and Electron Attachment Energies (EA) for
the Fluoroethenes.
Ethene
Ti I.P.. eV
a I.P., eV
E.A.. eV
ch2=ch2
10.6
12.85
1.78
ch2=chf
10.58
13.79
1.91
CM
u_
O
ii
CM
X
o
10.72
14.79
1.84
c/s-CHF=CHF
10.43
13.97
2.18
frans-CHF=CHF
10.38
13.90
2.39
cf2=chf
10.53
14.64
2.45
CM
U_
o
II
CM
LL
O
10.52
15.95
3.00
derivative is more exothermic than for the parent hydrocarbon. This arises from a
combination of the destabilizing effect of polyfluorination on double bonds and the
thermodynamic preference for gem-difluoro substitution at saturated carbon. The
deviation of CH2=CHF in Table 2-6 is explained by the preference of a single fluorine
substituent to reside at the vinylic position (Figure 2-3).114
Table 2-6. Heats of Hydrogenation of the Fluoroethenes.
Ethene
AH0 (HA. kcal
ch2=ch2
-32.6
ch2=chf
-29.7
ch2=cf2
-38.8
cf2=cf2
-45.7
Fluorine Non-Bonded Interactions in Reactive Intermediates
The 7r-donor ability of the fluorine substituent is reflected in its activating and
ortho, para-directing character in electrophilic aromatic substitution reactions,115
consistent with 13C NMR measurements of fluorobenzene, where shielding of these
positions is observed.116

34
H CH2F i2 h3c F l2 h3c H
M - H( - X
H H AH° = -3.34 kcal mol'1 H H AH0 = +0.65 kcal mol'1 H F
54 Z-55 E- 55
H3C F l2 H CF2H
M
AH0 = -2.5 kcal mol"1
>=K
H F
H H
56
57
Figure 2-3. Thermodynamic Equilibria in Mono- and Difluoropropenes.
Delocalization of fluorine's nonbonded electrons into the vacant 2p orbital on
carbon more than compensates for its inductive withdrawal in a-fluoro carbocations,
resulting in net stabilization. In the gas phase, carbocation stability increases along the
series +CH3 < +CF3 < +CH2F < +CHF2 and +CH2CH3 « +CF2CH3 = +CHFCH3.117,118 The
+CF3 cation has been observed in the gas phase, with many others having been
successfully generated in solution.119'121 In constrast, fluorination at the 0 position and
beyond destabilizes carbocations due to inductive effects; simple alkyl p-fluoro
carbocations not benefiting from additional stabilizing factors have yet to be detected.
Electron pair repulsion in a-fluoro carbanions results in a strong preference for
pyramidal geometries, ab initio calculations122'124 predicting an FCF angle of ca. 99.5°
and an inversion barrier of 119 kcal mol'1 for CF3. Although fluorination does increase
C-H bond acidities in such pyramidal systems,125 a destabilizing effect is observed in
cases such as the 9-fluorofluorenyl anion (Figure 2-4, X = F) where coplanarity is forced
between the 2p orbitals on fluorine and the remainder of the n system.126 Fluorine
substitution in the p position stabilizes carbanions through a combination of inductive
and hyperconjugative effects, (Figure 2-5) the latter supported by X-ray crystallographic
data of perfluoroalkyl anion salts as well as through calculation.122,127 128

35
D X
NaOCH3
CH3OH
H X
58
59
X kexc (rel)
D 1
F 0.125
Cl 400
Br 700
Figure 2-4. Fluorine Destabilization in Planar 9-Halofluorenyl Anions.
*0,
c-c^i
F 0
^c=c
Figure 2-5. Negative Hyperconjugation in p-Fluorocarbanions.
Fluorine Steric Effects
The minimal spatial requirements of fluorine, the smallest non-hydrogenic
substituent, would imply a very minor steric impact on reaction kinetics and
thermochemistry. In most cases this is true; however, examples of steric inhibition in
reactions of fluoro-substituted systems do exist, typically in conformational and other
dynamic processes occuring via highly congested transition states. This is illustrated in
Figure 2-6;129 130 the meta ring flip in 61 (X = F) exhibiting the largest known rate
retardation induced by a single fluorine substituent.
The disparate behavior observed in the Cope rearrangements of d,l- and meso-
62 (Figure 2-7)131 provides a particularly striking example of a fluorine steric effect.
Transformation of d,l-62 to 63 proceeds via a typical chair-like transition structure, where
in meso-62 a "boat-like" structure is required for Ct - C6 bond formation. The higher AH1

36
X X
X = H, AG* < 6 kcal mol'1 (340 K)
X = F, AG* =11.1 kcal mol'1 (340 K)
60
kHlkF = 1011 (298 K)
61
Figure 2-6. Inhibition of Conformational Dynamics by Fluorine Substitution.
and positive AS* implies a
dissociative, rather than concerted,
transition state for
rearrangement of meso-62.
This is induced by severe electrostatic repulsion between
the high charge densities of the terminal fluorine substituents, separated by less than the
sum of their van der Waals radii in the Cope transition state.132
°v,
H^ycF2
O'"'
d,l- 62
AH* = 22.4 kcal mol'1
63
AS* = -17.5 eu
r\
F r-
a
Uh2
- ifW
x"t;f2
1
or"'
meso-62
AH* = 49.5 kcal mol'1
63
AS* = +8.1 eu
Figure 2-7. Chair- versus Boat-Constrained Cope Rearrangement Reactions of
Terminally Fluorinated Dienes d,l- and meso-62.

37
Steric effects are enhanced by perfluoroalkylation and branching. Cyclohexane
A values133 and modified Taft steric parameters134 demonstrate that CF3 is at least as
large as isopropyl; evidence exists135 to suggest that perfluoroisopropyl and fe/f-butyl are
comparable in size. The most remarkable example of the above affects is the existence
of perfluorinated radical 64, (Figure 2-8) found to be persistent by ESR even in the
presence of molecular oxygen.136 The astounding kinetic stability of 64 derives from
steric sheltering of the unpaired electron by the neighboring perfluoroalkyl groups.
64
Figure 2-8. Scherer's Persistent Perfluoroalkyl Radical.
The Fluorine Substituent in Free Radicals
Early application of organofluorine radical chemistry was comprised of the chain-
mediated addition of polyhalomethanes and ethanes to olefins, first by Haszeldine137 and
soon thereafter by Tarrant.138 139 Subsequent relative rate studies by Stefani et a/.140
followed by those of Tedder141142 clearly demonstrate the contrasting behavior of
fluorinated and non-fluorinated radicals in their bimolecular additions to alkenes.
The fluorine substituent has a substantial effect on the structure of organic
radicals as well as dramatic, but comprehensible, alterations in hydrogen abstraction
and addition reactivity in comparison to hydrocarbon systems, resulting primarily from
fluorine's potent a-withdrawing character.
Structural Aspects
In contrast to the planar, n-type methyl radical, a-fluorination results in
increasingly pyramidal, a-type radicals, as indicated by electron paramagnetic

38
resonance measurements143 and ab initio theoretical studies.144'147 Calculations by
Pasto at the UHF/4-31G level indicate barriers to inversion of 0.5, 6.8, and 25.1
kcal mol'1 for *CH2F, *CHF2, and *CF3, respectively.147 Geometries of the fluoromethyl
radicals computed at the UHF/6-31G(d) level are provided in Figure 2-9.
Figure 2-9. UHF/6-31G(d) Pyramidalization Angles of *CH3, *CH2F, *CHF2, and *CF3.
Inversion barriers in the series appear somewhat sensitive to the level of theory
employed and substantially increase with the inclusion of polarization functions. SCF
calculations by Dykstra145 employing a polarized double-^ basis result in an inversion
barrier of 33 kcal mol'1 for *CF3; inclusion of electron correlation (QCISD(T)/
6-31G(d)//UHF/6-31G(d), present work) affords a value of 29.4 kcal mol'1. Such
successive deviation from planarity is due to a combination of effects; relief of In
repulsion between the singly-occupied orbital on carbon and the fluorine 2p electrons is
further reinforced by overlap between the carbon 2p and C-F antibonding orbitals. This
strong tendency for pyramidalization has been shown to be responsible for the low n
bond energy in tetrafluoroethylene (Tables 2-4 and 2-6).148
The structures of fluorinated C2 radicals have been theoretically probed by
Paddon-Row and Chen et a/.149'152 Alkyl substitution induces slight pyramidalization as
observed in the ethyl radical, (Figure 2-10) due to hyperconjugation between the SOMO
and the staggered p C-H bond.153 Fluorination at the radical center exerts an effect
similar to that observed in the methyl series, while the structures of alkyl radicals are
found to be relatively insensitive to p-fluorination.

39
Figure 2-10. UHF/6-31G(d) Pyramidalization Angles for Ethyl Radicals CH3CH2*,
CF3CH2*, CH3CF2*, and CF3CF2*.
Radical Stability
FMO theory dictates that for radicals bearing electronegative substituents with
lone pairs (F, OH, NH2, SH) an inductive, destabilizing influence exists, countered by
stabilization resulting from delocalization of the unpaired electron.154 Thus, in the a
sense, fluoroalkyl radicals are destabilized. Furthermore, the opposing n-stabilizing
effect of the fluorine lone pairs decreases with pyramidalization of the radical site, due to
diminished overlap with the 2p AO on carbon.
The progressive decrease in stability of alkyl radicals with a- or p-fluorination has
been illustrated by Pasto,80 147 in the form of calculated radical stabilization energies
(RSE) based on isodesmic reactions (Table 2-7). The aforementioned increase in C-H
bond dissociation energies along the fluoromethane series lends experimental support.
Although some degree of p C-F hyperconjugative interaction is observed in 2-
fluoro substituted radicals,152 such a contribution to overall radical stability is minor and
inductive destabilizing influences dominate, as indicated by experimental and theoretical
C-H BDEs for the 2-fluoroethanes (Table 2-8).

40
Table 2-7. Isodesmic Equation and Radical Stabilization Energies (RSE, kcal mol'1,
4-31G) for a- and p-Substituted Systems. Positive Values Denote Radical Stabilization.
Xn*
CH3-n + CH4
—> XnCH4_n +
*ch3
X
RSE
X
RSE
F
+ 1.64
ch3
+3.27
f2
+0.56
OH
+5.73
f3
-4.21
och3
+5.30
ch2f
+1.46
CN
+5.34
chf2
+0.16
nh2
+10.26
CO
LL
o
-1.34
+nh3
-4.07
SH
+5.66
+sh2
-3.17
Table 2-8. Experimental and Theoretical C-H Bond Dissociation Energies (kcal mol'1)
of 2-Fluoroethanes.
CH3CH?-H
CH9FCH9-H
CF,HCH,-H
CF3CH2-H
101.1a6
103.6f
106.7a,f
97.7C
99.6C
101.3°
102.0C
100.0d
104.3d
102.0e
104.1e
105.9®
107.1®
3 Experimental Value; Reference 155. b Experimental Value, Reference 156.
c MP2/6-311G(d,p)//MP2/6-31G(d,p); Reference 157. d B3LYP/6-31G(d); Reference 91
and Present Study.e MP2/6-311+G(3df,2p)//MP2/6-31G(d); Reference 158.
' Experimental Value; Reference 159.
Organofluorine Radical Reactivity
As a result of the a-withdrawing character of the fluorine substituent and
interaction of the SOMO with C-F a* orbitals, fluoroalkyl radicals should possess lowered
SOMO energies and therefore exhibit enhanced SOMO-HOMO interactions in
comparison with reactions of their hydrocarbon counterparts. Experimental ionization

41
potential and electron affinity data, although sparse, has been compiled in a recent
review by Dolbier90 and demonstrates the greater absolute electronegativity of the
fluoroalkyl radicals. Calculated quantities, inferred from Koopmans' theorem160 or based
on radical-ion energy differences, follow the expected trend although quantitative
agreement is often lacking.161162 The combination of such FMO, geometric, and
enthalpy factors in hydrogen atom abstraction, intermolecular addition, and cyclization
reactions of fluoroalkyl radicals is now discussed.
Hydrogen Atom Abstraction Reactions
A review by Tedder163 has underlined the importance of polar and enthalpic
factors in radical abstraction reactions. Activation parameters for abstraction by methyl
and trifluoromethyl radicals from a series of hydrogen donors are provided in Table 2-9.
Table 2-9. Arrhenius
Trifluoromethyl Radicals.
Parameters for Hydrogen Atom
Abstraction
by Methyl
ch3-
CF;
•
3
H-Donor
im
10)
0)
loq A
im
IO>
0)
loq A
ch3-h
14.2
8.8
11.3
8.9
ch3ch2-h
11.8
8.8
6.9
8.4
(CH3)2CH-H
10.1
8.8
6.5
8.1
(CH3)3C-H
8.0
8.3
4.9
7.7
H-CI
2.5b
5.0C
a In kcal mol'1. b AH0â„¢ =
- -1 kcal mol'1.
c AH0â„¢ = -3 kcal
mol'1.
In the first four examples, the decrease in activation barrier for both CH3* and
CF3* abstractions from alkanes are in line with the greater stability of the product radical;
in each case, the barrier to abstraction by CF3* is substantially lowered. In contrast,

42
abstraction of hydrogen atom from HCI by CF3* is much less facile, occuring with a
barrier twice that of CH3* in spite of its slightly greater exothermicity.
In collaboration with Lusztyk and Ingold at NRCC, LFP-determined rates of
perfluoroalkyl radical addition to a number of alkenes have been determined by the
Dolbier group.54,85,88,91 This has afforded, via competitive kinetic techniques, absolute
rate constants for hydrogen abstraction by the perfluoro-n-heptyl radical from a series of
donors, summarized in Table 2-1 0.54,86,87 89 For comparison, abstraction rate constants
for hydrocarbon n-alkyl radicals were provided in Table 1-1.
Table 2-10. Absolute Rate Constants for Hydrogen Atom Abstraction for Perfluoro-n-
heptyl Radicals. Rate Constants are at 303 K.
EtiSiH (TMS),Si(CHOH nBu.GeH (TMS),SiH nBu.SnH PhSH
kH( n-C7F15*), 0.75 16 15 51 203 0.28
106 M'1 s1
For the first five donors in the series, a substantial rate enhancement is observed
over hydrocarbon radicals, ranging from 75-fold in the case of nBu3SnH to nearly 900-
fold in the case of Et3SiH, after slight temperature correction to 303 K. Although
hydrogen transfer to perfluoroalkyl radicals is more exothermic (see the previous
discussion on C-H BDEs) this is insufficent to account for a nearly three order of
magnitude difference in rate constants. Furthermore, thiophenol, an excellent donor to
hydrocarbon radicals, is found to suffer a greater than 400-fold decrease in transfer rate
to perfluoroalkyls.
These characteristics are explained by the ability of the radical-donor pair to
accommodate charge transfer interactions in the hydrogen transfer transition state, as
previously discussed in Chapter 1. The relatively electropositive donor agents
(stannanes, germanes, and silanes) lead to a more favorable polarity matchup with the
electronegative perfluoroalkyl radical than with the hydrocarbon. Conversely, transfer
from the more electronegative thiophenol results in a non-polarized or polarity-

43
mismatched transition state. The existence of such polar effects were confirmed by a
correlation between rates of hydrogen transfer to perfluoro-n-heptyl radicals by a series
of substituted thiophenols, versus their Hammett a+ constants (Figure 2-11).87 The
resulting p value of -0.56, when compared to that obtained in the case of te/f-butoxyl
(-0.30)164 again reflects the high electrophilicity of perfluoroalkyl radicals.
Figure 2-11. Hammett Plot for Hydrogen Abstraction from para-Substituted Thiophenols
by Perfluoro-n-heptyl Radical.
Intermolecular Addition Reactions
Relative rates of addition of small fluorocarbon radicals to fluorinated and non-
fluorinated olefins have been extensively investigated by Tedder and Walton,
culminating in a critial review in 1980.52 Table 2-11 illustrates the relative reactivity of
Table 2-11. Relative Rates of Addition of the Fluoromethyl Radicals to Ethene and
Tetrafluoroethene.
Radical kartrt (C,FJ / kart* (C,HJ (437 K)
*CH3 9.5
*CH2F 3.4
*CHF2 1.1
*CF3
0.1

44
the fluoromethyl radicals towards ethylene and tetrafluoroethylene. Additional studies,
utilizing a number of unsymmetrical methyl- and trifluoromethyl-substituted olefins,
solidified the ascription of relative rates and regiochemical preferences to a combination
of polar and steric influences.
Absolute rates of addition of perfluoroalkyl radicals to alkenes have been
determined by laser flash photolysis, a subset of the data acquired to date presented in
Table 2-12.90 The dramatic rate acceleration enjoyed by the perfluoroalkyl radicals
versus hydrocarbon n-alkyls is readily apparent, ranging from factors of 300 to 30,000 in
the case of the heptafluoropropyl radical addition to electron-rich alkenes. The rate of
addition to pentafluorostyrene, in constrast, is increased by only a factor of 42.
Table 2-12. Absolute Rate Constants for Addition of Perfluoroalkyl Radicals to Alkenes.
kartri, 106 M'1 s1 (298 K)
Alkene
Id
rn
t-4
1CzFi5
Id
100
S'
*C,Fs
•cf3
RCH/
Styrene
43
46
46
79
53
0.12
a-Methylstyrene
78
89
94
87
0.059
p-Methylstyrene
3.8
3.7
7.0
17
Pentafluorostyrene
13
23
26
0.31
4-Methylstyrene
61
4-Methoxystyrene
65
4-Chlorostyrene
36
4-(CF3)Styrene
35
1-Hexene
6.2
7.9
16
0.0002
Such enhancements may potentially be attributed to a combination of factors.
Relative reaction enthalpies should play a role, as a stronger C-C bond (from CH3-CH3
versus CF3-CH3 BDE data in Table 2-3, ca. 11 kcal mol'1) is formed upon perfluoroalkyl

45
addition. However, only slight increases (factors of 5-7) in addition rates of the
perfluoroalkyls to styrene versus 1-hexene are observed, despite the greater
exothermicity (ca. 16 kcal mol'1) of the former. This demonstrates the relatively minor
importance of reaction enthalpy, in accord with the early transition states expected for
radical addition.
The pyramidal, a-character of the fluoroalkyl radicals should afford a kinetic
advantage (further discussed in Chapter 3) over the planar hydrocarbon, the LFP-
measured rate of addition of 1,1 -difluoropentyl radical to styrene, 2.7 x 106 M 1 s'1,88
giving rise to a 22-fold enhancement relative to n-alkyls.
The primary factor responsible for such striking increases in reactivity is believed
to be charge transfer influences (Figure 2-12) similar to those postulated for hydrogen
atom transfer reactions. The lowered SOMO energies of the perfluoroalkyl radicals and
resulting enhanced SOMO-HOMO interactions with alkenes in the addition transition
state leads to substantial rate enhancement. Supporting evidence, in the form of a
Hammett relation involving para-substituted styrenes and correlations between addition
rates and alkene ionization potentials, has been offered.54 Ab initio computations concur
with the experimental findings and are discussed in Chapter 3.
8-
CF3(CF2)nCF2*
8+
Figure 2-12. Polarized Transition State for Addition of Perfluoroalkyl Radicals to
Alkenes.
Intramolecular Cvclization Reactions
Cyclopolymerization of fluorinated monomers has long been known as a means
for the generation of macromolecular materials with unique physical properties.165'168

46
However, despite the widespread popularity of the unimolecular 5-hexenyl radical
cyclization for the generation of five-membered rings, synthesis of fluorinated cyclic
products utilizing radical methodology has received limited attention.90 This is somewhat
surprising, in light of the current interest and demonstrated importance of fluorinated
analogues and mimics of pharmaceutical and agricultural agents.169
Until recently, no quantitative kinetic data existed for cyclization reactions of
fluorinated radicals. With the competitively-determined rate constants of hydrogen
abstraction by perfluoroalkyl radicals serving as basis reactions, rates of cyclization of a
number of fluorinated 5-hexenyl systems have now been determined,86 89 170 examples of
which are provided in Figure 2-13.
Most obvious of the data is the remarkable rate acceleration in the 5-exo
cyclizations of 65 and 68, occurring with 163- and 41-fold increases relative to the parent
hydrocarbon 1 and ascribed to charge-transfer effects analogous to those occurring in
bimolecular additions. Consequently, with no such polarity matchup in 71 (which
involves cyclization of an electrophilic radical onto an electron-deficient double bond)
only a minor increase in rate is observed. The 5-exo cyclization rates of 71 and 74, in
line with those other hexenyl systems bearing fluorinated double bonds,170 demonstrate
the lack of kinetic impact of vinylic substitution.
Especially surprising is the degree to which 65 and 68 undergo 6-endo closure,
the former with a 1040-fold, and the latter a 700-fold acceleration relative to 1. This is
further discussed in Chapter 4, in which the reactivities of a series of lightly-fluorinated 5-
hexenyl systems are investigated.
In the preceeding discussions of hydrogen abstraction, addition, and cyclization
reactivity, the observed kinetic behavior was found to be due to a combination of
geometric and polar effects induced by polyfluorination. Related studies of partially
fluorinated alkyl radicals, addressed in the next two chapters, will aid in separating the
effects of fluorination at the a and p positions and beyond, providing insight into the

47
extent to which the effects of perfluorination on the reactivity of organic radicals are the
sum of their parts.
Table 2-13. Some Absolute Rate Constants for 5-exo and 6-endo Cyclization of
Fluorinated 5-Hexenyl Radicals. Rate Constants are for 303 K.
Cyclization Reaction kC5, 105s'1 /cC6, 105s'1
2.7
0.05
52
1 ?p2
f2c.c.cf2
f2
68
f2c cf2
f2c-cf2
69
CF,
f2c.c.cf2
f2
70
110
cf2
f2c.c.cf2
f2
71
VF2
f2c^cf2
f2c-cf2
72
F
j.
f2c cf2
f2c.c.cf2
f2
73
4.9
F .
F T
f2c. J
c
F2
Fn .CF,
F2C
F,C-
75
F
1.
f2c cf2
xf2
76
4.3
35
N/A
N/A
74

48
Conclusion
An introduction to the general structural and reactivity characteristics imparted by
fluorine substitution in organic systems has been provided. Such substitution can either
be stabilizing or destabilizing, depending on the nature of the ground state molecule,
intermediate, or reaction in question.
A summary of the first absolute kinetic data obtained for reactions of fluorinated
radicals has been presented. Results of these studies, involving per- or otherwise highly
fluorinated systems, demonstrate the combined influence of polarity, structural, and
enthalpic factors. The following studies of the addition, hydrogen abstraction, and
cyclization kinetics of partially fluorinated alkyl radicals will serve to dissect the relative
magnitudes of these influences on organic radical reactivity.

CHAPTER 3
THE REACTIVITY OF PARTIALLY FLUORINATED RADICALS
IN INTERMOLECULAR ADDITION AND HYDROGEN ABSTRACTION REACTIONS
Introduction
Initial studies of the addition rates of some partially-fluorinated radicals to
alkenes88 demonstrated observable rate enhancements relative to n-alkyls, though not
nearly as great as those of perfluoroalkyl systems. Such reactivity derives from the
combination of structural and polar characteristics induced by fluorine substitution.
The current study extends the amount of absolute rate data assembled for the
addition of partially fluorinated radicals to olefins. In addition, through competitive kinetic
techniques, absolute rate constants for hydrogen abstraction from tri-n-butyltin hydride
(nBu3SnH) have been determined. Such kinetic data is necessary for the determination
of absolute rates of cyclization of partially fluorinated 5-hexenyl radicals, discussed in
Chapter 4, and allows for the partitioning of the gross reactivity characteristics of
perfluoralkyl radicals into the separate influences of a, p, and y fluorination. This is
accomplished by use of the following systems: RCH2CH2CF2* (a,a-difluoro),
RCH2CF2CH2* (p.p-difluoro), RCH2CF2CF2* (a,a,p,p-tetrafluoro), CF3CF2‘ (a.a.p.p.p-
pentafluoro), and RfCH2CH2* (y-perfluoro).
The existence of charge transfer stabilization in the transition states for
fluorinated radical addition to alkenes is corroborated by ab initio calculations, and the
thermodynamics of C-H and C-C bonding and radical stablization in hydrofluorocarbons
rationalized on the basis of Coulombic interactions.
49

50
Precursor Syntheses and Competitive Kinetic Studies
In each of the competitive kinetic runs, hydrofluorocarbon radicals were
generated from bromide or iodide precursors by photoassisted C-X bond homolysis.
These radicals subsequently underwent competitive trapping with known, varying
concentrations of styrene or nBu3SnH, adjusted to ensure pseudo-first-order kinetic
behavior and to allow for accurately measurable amounts of trapping products, as
depicted in Figure 1-17 and in greater detail below.
1.1-Difluorohex-1-vl Radical (77)
The synthesis of bromide precursor 80 was achieved in two steps (Figure 3-1) in
a straightforward manner. Copper(l)-mediated addition of dibromodifluoromethane to 1-
pentene (78), based on a modification by Gonzalez et a/.171 of a procedure by Burton
and Kehoe172 afforded 1,3-dibromo-1,1-difluorohexane 79 in typical yield.
Regioselective displacement of the internal bromine was accomplished via sodium
borohydride reduction in DMSO, providing precursor 80 contaminated with a small
amount of overreduction product 81. Pure samples of each were obtained by
preparative GC separation, the former utilized in the competition run and the latter for
spectral comparison with kinetic NMR data.
78
CF2Br 2
â–º
(CH3)3COH, h2nch2ch2oh
CuCI (cat.)
(57.3%)
79
NaBH4
DMSO
(62.1%)
80 81
Figure 3-1. Preparation of 1-Bromo-1,1-difluorohexane 80 and 1,1-Difluorohexane 81.

51
hv
nBu3SnH
C6D6
77
Figure 3-2. kH / kaM Competitive Kinetic Scheme for 1,1-Difluorohex-1-yl Radical 77.
Photolysis of 80 as a C6D6 solution in the presence of an excess of nBu3SnH and
styrene (Figure 3-2) afforded intermediate radical 77. Subsequent entrapment by these
agents (both irreversible processes) yielded 81 and 82, respectively; the latter further
trapped by nBu3SnH to yield 3,3-difluoro-1-phenyloctane 83. Throughout the course of
the reaction, nBu3Sn* radicals are generated to propagate the chain process via
abstraction of halogen from precursor 80.
Product ratios for varied concentrations of trapping agents were determined by
19F NMR analysis according to the pseudo-first-order relation in Equation 3-1,
[ 81 ] _ [ *h ] [ 77 ] [ nBu3SnH ]
[ 83 ] ' [ /fadd ] [ 77 ] [ C6H5CH=CH2 ]
a plot of product ratios obtained for each data point versus that of trapping agents
affording the ratio kH / kaM. The stability of trapped products under the reaction
conditions and lack of appreciable side reactions were demonstrated by the high

52
conversion of precursor 80 to 81 and 83 versus an internal standard of a,a,a,-
trifluorotoluene, (<(> -63.24) indicating in turn the high efficiency of the chain process and
reliability of the kinetic results. A partial 19F NMR spectrum of the first of six data points
is provided in Figure 3-3, a doublet of triplets (-CF2H, 4> -116.0) observed for 81 versus
an overlapping triplet of triplets at -99.1 (-CF2-) for 83. Full kinetic data and yields are
given in Table 3-1 below, the plot of which located in Figure 3-4. The slope of the line
(3.39 ± 0.02) in conjunction with the known absolute rate constant for addition of
-98 -99 -100 -101 -102 -103 -104 -105 -106 -107 -108 -109 -110 -111 -112 -113 -114 -115 -116 -117 -118
Figure 3-3. Partial 19F NMR Spectrum of Data Point 1 for kH / kadd Competition of 1,1-
Difluorohex-1-yl Radical 77.
Table 3-1. Competitive Kinetic Data for kH I kadd Competition of 1,1-Difluorohex-1-yl
Radical 77.
L80J
f CrFUCH=CH, 1
f nBu.SnH 1 / rCRFUCH=CH, 1
r 81 1 / f 83 1
% Yield
0.094
2.01
0.719
2.30
95
0.094
1.81
0.847
2.73
95
0.094
1.61
1.01
3.26
96
0.094
1.41
1.21
3.93
97
0.094
1.21
1.49
4.88
96
0.094
1.01
1.87
6.21
95

53
1,1 -difluoropent-1-yl radical to styrene, 2.7 (± 0.5) x 106 M'1 s'1, resulted in an absolute
rate constant kH of 9.1 (± 1.7) x 10® M 1 s'1. It should be noted that the accuracy of such
derived kH values can be no better than those reported in the LFP determinations of kaM,
the error estimates in the former reflecting both the least-squares fit of the line and
propagated error of the latter. Synthesis of styrene adduct 83 was performed by classic
means for characterization and spectral comparison (Figure 3-5).
Figure 3-5. Preparation of Styrene Adduct 83.

54
2.2-Difluorohex-1-vl Radical (88)
In accordance with literature precedent,173 a-bromination of 2-hexanone in the
presence of urea in acetic acid selectively afforded 1-bromo isomer 90 in 72.2% yield.
Subsequent treatment with diethylaminosulfurtrifluoride (DAST) provided bromo
precursor 91, originally purified by preparative GC for use in the kinetic study. However,
a sluggish chain reaction (further hindered by the strong UV absorption of the excess
styrene present in the kinetic samples) led to the preparation of iodo precursor 92
(Figure 3-6) via Finkelstein transformation at elevated temperature.
Br2 9 DAST
^ â–º
CH3C02H, H2NCONH2 CH2Br CHCI3
(72.2%)
90
^c.CH2Br
F2
(59.7%)
91
Nal
(CH3)2CO
Figure 3-6. Preparation of 1-lodo-2,2-Difluorohexane, (92) Precursor to 2,2-Difluorohex-
1-yl Radical 88
Figure 3-7. Plot for kH / kadd Competition of 2,2-Difluorohex-1-yl Radical 88.

55
The competition plot for kH / kadd determination is found in Figure 3-7; raw data for
this and all remaining experiments in this chapter may be found in Chapter 5. With the
absolute rate constant for addition of 2,2-difluoropent-1-yl radical to styrene known from
LFP experiments, a kH value of 1.4 (± 0.5) x 107 M'1 s'1 was determined. Authentic
samples of 93 and 98 were prepared for spectral comparison and characterization as
illustrated below in Figure 3-8.
91
nBu3SnH
AIBN
c6h6
(66.3%)
98
Figure 3-8. Preparation of Hydrogen Abstraction Product 93 and Styrene Adduct 98.
1,1,2,2-Tetrafluorobut-1-vl (99) and 1,1,2,2-Tetrafluorohex-1-yl (100) Radicals
At the time of this study, no absolute rate constants for addition of a 1,1,2,2-
tetrafluorinated radical to alkenes had been determined. Thus, in order to obtain a kH
value for such a system, a precursor suitable for absolute /cadd measurements was first
required. Bromide precursors, although in most cases sufficient for competition

56
experiments, are ineffective under the LFP operating conditions used in our kaM
determinations (308 nm excimer laser pulses) due to their relatively short wavelength
chromophore and small extinction coefficient (for 102, £max = 37.3 M 1 cm'1, Xmax = 218
nm, cyclohexane solvent). The instability of O-acylthiohydroxamic esters of the
perfluoroalkanoic acids has been noted by Barton.1'4 With neither these nor diacyl
peroxide precursors lending themselves to isolation and / or shipment to the NRCC in
Canada, the synthesis of a suitable iodide precursor was undertaken.
1-Bromo-1,1,2,2-tetrafluorohexane (102, Figure 3-9) was prepared in one step
from 6-bromo-5,5,6,6-tetrafluorohex-1-ene (supplied by Halocarbons, Inc.) and its
transformation to the corresponding iodide or carboxylic acid (the latter of which could be
converted to the iodide via Hunsdieker methodology) attempted under a variety of
conditions (Figure 3-10).
1. BH3 Me2S
2. CsHnCOaH /^^^r-CF2Br
F2
Tetraglyme
(79.6%)
102
Figure 3-9. Preparation of 1-Bromo-1,1,2,2-tetrafluorohexane (102).
Although the conversion of perfluoroalkyl iodides to their lighter analogues is
known in the literature,175 176 downward transhalogenation of perfluoroaliphatic halides is
exceedingly difficult. Indeed, all attempts at conversion of 102 to the corresponding
iodide were met with failure. Perfluoroalkyl Grignard reagents, generated either directly
(and in low yield) or via transmetallation by an alkylmagnesium halide, utilize iodide
starting materials.177 178 Investigations of perfluoroalkylzinc halides by Miller179 resulted
in a similar conclusion; perfluoro-n-propyl iodide may be converted to the organozinc
reagent in ca. 60-80% yield (as determined by aqueous hydrolysis or capture with
halogen electrophiles) after a brief induction period. On the contrary, reaction of
heptafluoro-n-propyl bromide with zinc in 1,2-dimethoxyethane afforded no product after

57
C
F2
102
.CF2Br
"C'
f2
102
CF2Br
1) Mg
2) l2
Et,0
N. R.
78
icf2cf2i
tBuOH, H2NCH2CH2OH
CuCI (cat.)
N. R.
103
Et3N 3HF/NIS
CH2CI2,0° C
N. R.
Figure 3-10. Attempted Preparation of lodo Analogue of 102.
73 hours at 90° C; a 60% yield of the zinc derivative was obtained (inferred via
hydrolysis) after a period of 1.5 months. Attempted lithiation of 102 at low temperature
resulted in the formation of (3-fluoride elimination product 1,1,2-trifluoro-1 -hexene, (103)
as identified by its 19F NMR spectrum. Additionally, neither Cu(l)-induced addition of
1,2-diiodotetrafluoroethane to 1-pentene nor iodofluorination of 103 utilizing the
triethylamine trihydrofluoride / N-halosuccinimide methodology of Alvernhe et a/.180 were
successful.

58
Synthesis of a C4 iodide was achieved through modification of a DuPont literature
procedure,181 whereby 1,4-diiodo-1,1,2,2-tetrafluorobutane was produced in 40.1% yield
via direct thermal addition of 1,2-diiodotetrafluoroethane to ethylene (Figure 3-11). DBU-
induced elimination of hydrogen iodide in ether followed by diimide hydrogenation with
hydrazine-hydrogen peroxide in methanol afforded tetrafluoroiodo precursor 107 in 2.8%
overall yield after preparative GC purification, which was sent to the NRCC for absolute
kadd measurements.
ICF2CF2I
104
h2c=ch2
A
â–º
(40.1%)
105
DBU
â–º
Et20
^(rCF2'
F2
(66.0%)
106
HN=NH ^^CF2I
â–º F
CH3OH 2
(10.7%)
107
Figure 3-11. Preparation of 1,1,2,2-Tetrafluoro-1-iodobutane (107).
[ nBu3SnH ] / [ CgH5CH=CH2 ]
Figure 3-12. Plot for kH / kadd Competition of 1,1,2,2-Tetrafluorohex-1-yl Radical 100.
Bromide 102 was utilized in the kH / kadd competition, the kinetic plot for which
given in Figure 3-12. An authentic sample of styrene addition product 110 was prepared

59
by slow syringe pump addition of nBu3SnH to a heated, irradiated solution of 102 and
styrene in benzene (Figure 3-13).
nBu3SnH
C6H6
c6h5ch=ch2
nBu3SnH
C6H6
Figure 3-13. Preparation of Hydrogen Abstraction Product 108 and Styrene Adduct 110.
2-fPerfluorohexyl1eth-1-vl Radical (111)
2-[Perfluorohexyl]-1-iodoethane 112 was provided as a gift from Prof. Neil Brace.
The kH / kadd competition plot is found in Figure 3-14; hydrogen abstraction product
1-[perfluorohexyl]ethane (113) and styrene adduct 1-[perfluorohexyl]-4-phenylbutane
(117) were prepared as shown in Figure 3-15.
Figure 3-14. Plot for kH / kadd Competition of 2-[Perfluorohexyl]eth-1-yl Radical 111.

60
C6F13CH2CH2I
112
nBu3SnH
A
C6H6
CgF 13CH2CH3
113
(70.0%)
117
Figure 3-15. Preparation of Hydrogen Atom Abstraction Product 113 and Styrene
Adduct 117.
Pentafluoroethyl Radical (118)
lodopentafluoroethane (119) was obtained from PCR, Inc. Due to the high
volatility of both this precursor (bp 12-13° C) and hydrogen atom abstraction product
pentafluoroethane (120, bp -48.5° C) 119 was handled as a solution in degassed C6D6.
The kH / /radd competition experiment (Figure 3-16) was performed in tubes which were
quickly flame-sealed upon injection of an aliquot of the chilled precursor stock solution.
[ nBu3SnH ] / [ CgH5CH=CH2 ]
Figure 3-16. Plot for kH / /cadd Competition of Pentafluoroethyl Radical 118.

61
Samples of 120 and 122 were prepared under free radical conditions for characterization
purposes as shown in Figure 3-17.
CF3CF2l
119
hv
nBu3SnH
C6H5CH=CH2
nBu3SnH
CeH6
CF3CF2H
120
f3c
121
Figure 3-17. Preparation of Hydrogen Atom Abstraction Product 120 and Styrene
Adduct 122.
Discussion
Absolute rate constants for addition and hydrogen atom abstraction for systems
studied in this chapter are provided in Table 3-2. For comparison, data for hydrocarbon
(n-pentyl) and perfluoroalkyl (perfluoro-n-heptyl) radicals are included.
It is seen from the data that the reactivity trends for partially fluorinated radical
additions to styrene are generally adhered to in hydrogen abstractions. However, the
actual rate constants for the latter are observed to differ (on the average, by a factor of
11) from those for addition, and span a narrower range. The decrease in abstraction
rate ratios as a function of substitution is due to the proximity (within an order of
magnitude) to diffusion control for the more reactive radicals 100, 118, and 127.
C-H and C-C bond dissociation energies of the fluoroalkanes should reflect the
relative thermochemistry of hydrogen atom abstraction and addition to alkenes by their
respective radicals. However, very few experimentally determined BDE values for such

62
Table 3-2. Absolute Rate Constants for Hydrogen Abstraction from Tributyltin Hydride
and Addition to Styrene by Partially Fluorinated Radicals. Rate Constants are for 298 K.
Radical
k^ (M'1 s1)
kaMJrei)
/cH (M-1 s1)
ÍÍHJreí)
RCH2CH2*
(123)3
1.2 x 10s 5
1
2.4 x 106c
1
RCH2CF2*
(124, 77) d
2.7 (± 0.5) x 106 e
22.5
9.1 (± 1.7) x 106 f
3.8
rcf2ch2*
(125, 88) 9
5.2 (± 1.8) x 105 e
4.3
1.4 (± 0.5) x 107'
5.8
rcf2cf2*
(99, 100) h
2.0 (± 0.1) x 107f
167
9.2 (± 0.8) x 107 f
38
r,ch2ch2*
(126, 111)'
1.3 (± 0.2) x 105 e
1.1
2.1 (± 0.3) x 106 f
0.9
cf3cf2*
(118)
7.9 (± 1.0) x 107 f
658
3.2 (± 0.3) x 108'
133
c7f15*
(127)
4.6 (± 0.6) x 107y
383
2.0 (± 0.3) x 108y
83
a R = C3H7. b Reference 182, After Modification for Temperature and Other Factors in
Table III of Reference 183. c Reference 27. d For kadd Experiment, R = C3H7 (124); for
kH Experiment, R = C4H9 (77). e Reference 88. ' Reference 91 and Present Study. 9 For
kaM Experiment, R = C3H7 (125); for kH Experiment, R = C4H9 (88). h For kaM
Experiment, R = C2H5 (99); for kH Experiment, R = C4H9 (100). ' For kaM Experiment,
Rf = C4F9 (126); for kH Experiment, Rf = C6F13 (111). 1 Reference 54.
systems have been reported. Theoretical studies by Boyd157 184 at the MP2 level have
provided reasonably accurate C-H and C-C BDEs for C2 hydrofluorocarbons. A more
recent investigation based on isodesmic reactions by Marshall and Schwartz,158
published after the completion of the present study, has yielded C-H BDEs of
appreciably high quality for some linear (C2) and branched (up to C4) polyfluoroalkanes.
Out of interest in determining additional C-H and C-C BDEs for larger (through C4)
fluorinated n-alkyls, the geometries of a series of partially fluorinated ethanes, propanes,
and butanes, along with their respective radicals generated from terminal C-H or C-CH3
bond cleavage, were optimized at the hybrid density functional level. This DFT method
was chosen due to its implicit consideration of electron correlation at only slightly greater

63
computational expense than that of Hartree-Fock theory. Utilizing the three-parameter
exchange functional of Becke185 and the correlation functional of Lee, Yang, and Parr186
with the 6-31 G(d) basis, bond dissociation energies obtained in this "direct" fashion were
found to be, in the cases where such values are known, within 1-3 kcal mol'1 of those
determined experimentally and in comparable or better agreement with experiment than
the MP2/6-311G(d,p) values obtained by Boyd for C2 systems. Tabulated experimental
(where known) and calculated C-H and C-C BDE values are provided below in Tables
3-3 and 3-4, respectively.
Table 3-3. Theoretical and Experimental C-H Bond Dissociation Energies.
C-H Bond
Calculated BDE, kcal mol'1 a
Experiment
CH3CH2-H
100.06
97.7 c
102.0 d
101.1 ± 1 e
ch3cf2-h
97.4*
97.0C
99.5 + 2.5'
cf3ch2-h
104.3*
102.0 c
107.1 *
106.7 ± 1 1
cf3cf2-h
99.5*
99.7 c
104.6 d
102.7 ± 0.5 f
ch3ch2ch2-h
100.3*
100.4 ± 0.6 e
ch3ch2cf2-h
97.7*
ch3cf2ch2-h
103.1 *
ch3cf2cf2-h
100.1 *
cf3ch2ch2-h
101.4*
cf3cf2ch2-h
103.8*
a Reported as D0 (298.15 K). * B3LYP/6-31G(d); Reference 91 and Present Study.
c MP2/6-311G(d,p)//MP2/6-31G(d,p); Reference 157. * MP2/6-311+G(3df,2p)//MP2/
6-31 G(d); Reference 158. e Reference 187. Reference 95.

64
Table 3-4. Experimental and Theoretical C-C Bond Dissociation Energies.
C-C Bond
Calculated BDE, kcal mol1 a
Experiment
CH3-CH3
89.4 b
90.6 c
90.4 ± 0.2 d
CF3-CH3
99.6 6
103.3 c
101.2 ±1.1 ‘
CH3CH2-CH3
86.3 b
ch3cf2-ch3
91.4 b
cf3ch2-ch3
91.4 6
cf3cf2-ch3
95.5 6
ch3ch2ch2-ch3
86.7 b
ch3ch2cf2-ch3
91.6 6
ch3cf2ch2-ch3
89.9 b
ch3cf2cf2-ch3
95.4 6
cf3ch2ch2-ch3
87.8 b
a Reported as D0 (298.15 K). b B3LYP/6-31G(d); Reference 91 and Present Study.
c MP2/6-311 G(d,p)//MP2/6-31 G(d,p); Reference 157. d Reference 95.
Inspection of the data reveals interesting trends within both the C-H and C-C
BDE series. From Table 3-3, it is observed that a-fluorination results in a weakening of
C-H bonds, on the order of 1-3 kcal mol"1, as predicted by the various levels of theory
and supported by experiment in the case of ethane versus 1,1-difluoroethane.
Conversely, p-fluoro substitution results in a 3-5 kcal mol'1 increase in terminal C-H
BDEs. Furthermore, strengthening of C-C bonds (Table 3-4) is observed for all systems
examined, whether substituted at the a, p, or even (albeit diminished) y position. An
explanation for this behavior, consistent with the given BDE and other thermochemical
data, is provided later in the chapter.
With the relative thermodynamics of C-H and C-C bond formation investigated,
attention is turned to polarity effects. As mentioned previously, rate constants for

65
addition of the perfluoroalkyl radicals to subtituted styrenes (Table 2-12) are observed to
increase with decreasing ionization potential (styrene, IP 8.43 eV; a-methylstyrene, IP
8.19 eV; pentafluorostyrene, IP 9.20 eV).188 In contrast, 1,1 -difluoropentyl radical (124)
was found to add with rates equal, within experimental error, for all three olefins.88 From
this observation, along with derived absolute electronegativities for the radicals CH3*
(4.96), CH2F* (4.73), CHF2* (4.91), and CF3* (5.74) it was concluded that a-fluoro
substitution alone does not impart electrophilicity to an alkyl radical, and may instead
give rise to nucleophilic behavior.88
A series of theoretical studies by Wong et a/ 48'49'189 190 have assessed the degree
of charge transfer interaction in the transition states for methyl, hydroxymethyl,
cyanomethyl and ferf-butyl radical additions to a series of monosubstituted olefins.
Based in part on the computation of partial charges, it was concluded that the addition of
methyl radical was governed primarily by enthalpic effects, with no evidence for
nucleophilic character arising from either Mulliken or Bader-based charge-fitting
schemes. Hydroxymethyl and terf-butyl were found to be nucleophilic, with cyanomethyl
exhibiting substantial electrophilicity.
In order to determine such tendencies for the fluorinated radicals under
investigation, transition structures for the addition of hydrocarbon and fluoro-substituted
ethyl radicals to ethylene and of propene have been located at the UHF/6-31G(d)
level. Partial atomic charges were then computed, based on fits to the electrostatic
potential at points selected according the Merz-Kollman-Singh scheme.191192
Previous studies by Houk et at. utilizing the 3-21G basis found that addition of
ethyl radical to ethylene occurs preferentially via a gauche conformation, with an
incipient C-C-C-C dihedral angle of ca. 60°.46 Such gauche and anti structures
computed at the UHF/6-31(d) level are depicted in Figure 3-18, the former ca. 0.1
kcal mol'1 lower in energy than that of (b). Inclusion of electron correlation at the spin-

66
projected PMP2/6-311G(d,p)//UHF/6-31G(d) level increases this energy difference to ca.
0.3 kcal mol"1.
Figure 3-18. UHF/6-31G(d) (a) Gauche and (b) Anti Transition Structures for Addition of
Ethyl Radical to Ethylene. Relevant Geometrical Parameters are Shown.
Consistent with UHF/3-21G results, the preferred mode of addition of the ethyl
and radical to Ct of propene involves a gauche arrangement of the radical with the
alkene C=C n bond and a transoid orientation of the methyl group of the radical with
respect to that of the olefinic C2 carbon (Figure 3-19). This 'gauche-transoid' structure
lies ca. 0.1 and 0.4 kcal mol1 below the 'anti' and 'gauche-cisoid' conformers,
respectively, at the UHF/6-31(d) level after zero-point energy correction. Inclusion of
correlation effects (PMP2/6-311G(d,p)//UHF/6-31G(d)) yields energy differences of 0.3
and 0.4 kcal mol'1. These orientation preferences extend to the fluoroethyl series, as
seen in Tables 3-5 and 3-6.
Figure 3-19. UHF/6-31G(d) (a) 'Gauche-transoid', (b) 'Anti', and (c) 'Gauche-cisoid'
Transition Structures for Addition of Ethyl Radical to Ci of Propene.

67
Table 3-5. Geometric Parameters and Total and Relative Energies of Transition
Structures for Addition of Fluorinated Methyl and Ethyl Radicals to Ethylene.
Radical
r(C-C)
(A)
LiC=C)
íá)
Z C-C-C
(Peg)
^Dvr_
(DecL)
E,ot_d
iau)
ZPE
iau]
p e
Pre|
ch3*
2.246
1.382
109.1
101.9
-117.575692
(-118.045817)
0.089165
-
cf3*
2.300
1.372
106.6
108.6
-414.156036
(-415.285304)
0.067955
-
ch3ch2* 3
2.227
1.384
110.0
103.8
-156.612999
(-157.243689)
0.120578
0.0
(0.0)
ch3ch2* b
2.231
1.383
109.9
103.3
-156.612807
(-157.243039)
0.120455
0.05
(0.34)
CH3CF2* 3
2.235
1.378
110.0
108.6
-354.333278
(-355.404529)
0.105597
0.0
(0.0)
CH3CF2* b
2.245
1.378
106.3
108.2
-354.332326
(-355.403222)
0.105470
0.53
(0.75)
CF3CH2*3
2.258
1.380
109.0
104.1
-453.202939
(-454.490483)
0.097486
0.0
(0.0)
cf3ch2* 6
2.258
1.380
107.4
103.5
-453.203135
(-454.490294)
0.097452
-0.14
(0.10)
cf3cf2* 3
2.275
1.375
108.2
108.6
-650.904439
(-652.632413)
0.081977
0.0
(0.0)
cf3cf2* b
2.280
1.375
104.7
108.1
-650.904343
(-652.632131)
0.081948
0.04
(0.16)
3 Gauche. b Anti. c Degree of Radical Pyramidalization. d UHF/6-31G(d) Values;
PMP2/6-311G(d,p)//UHF/6-31G(d) Values in Parentheses. e Relative Conformer Energy
Differences, kcal mol’1. UHF/6-31G(d); PMP2/6-311G(d,p)//UHF/6-31G(d) Values in
Parentheses.
Inspection of the data reveals the similarity in both angle of attack and forming
C-C bond length, regardless of either the nature of the radical or conformation of the
transition structure. Incipient bond lengths range from ca. 2.22 to 2.30 A, generally
slightly longer for additions of the fluorinated members of the series. These somewhat
earlier transition states are also reflected in the shorter olefinic C=C bonds, 1.382 A in
the case of methyl radical addition to ethylene versus 1.372 A for trifluoromethyl, in turn

68
Table 3-6. Geometric Parameters and Total and Relative Energies of Transition
Structures for Addition of Fluorinated Methyl and Ethyl Radicals to Ci of Propene.
Radical
r(C-C)
(A)
r (C=C)
iAi
Z C-C-C
(Deg,)
^EYT—
(Deg.)
E,ot!
ÍM
ZPE
(au)
Erel_
ch3*
2.243
1.383
109.3
102.0
-156.614381
(-157.244918)
0.119634
-
cf3*
2.297
1.372
106.4
108.9
-453.195586
(-454.485598)
0.098359
-
ch3ch2* a
2.224
1.385
110.2
103.9
-195.651571
(-196.442757)
0.150984
0.0
(0.0)
CH3CH2*b
2.219
1.385
111.3
104.2
-195.651075
(-196.442411)
0.151060
0.35
(0.26)
CH3CH2* c
2.228
1.384
110.1
103.3
-195.651399
(-196.442053)
0.150878
0.05
(0.38)
ch3cf2* a
2.233
1.379
109.5
108.7
-393.372286
(-394.604285)
0.135951
0.0
(0.0)
ch3cf2* b
2.228
1.380
110.7
108.9
-393.371834
(-394.603661)
0.135994
0.31
(0.42)
ch3cf2* c
2.242
1.379
106.0
108.5
-393.371641
(-394.603448)
0.135892
0.37
(0.49)
cf3ch2* a
2.254
1.381
109.3
104.2
-492.242374
(-493.691081)
0.127933
0.0
(0.0)
CF3CH2’ b
2.248
1.381
110.9
104.6
-492.242279
(-493.691288)
0.128087
0.15
(-0.04)
CF3CH2’ c
2.257
1.381
107.7
103.6
-492.242269
(-493.690171)
0.127869
0.03
(0.54)
cf3cf2* a
2.274
1.375
108.0
108.7
-689.944220
(-691.833274)
0.112383
0.0
(0.0)
cf3cf2* b
2.266
1.375
110.1
108.9
-689.944106
(-691.833236)
0.112473
0.12
(0.07)
CF3CF2‘ c
2.279
1.375
104.6
108.3
-689.944068
(-691.832854)
0.112356
0.08
(0.25)
3 Gauche-transoid. b Gauche-cisoid. c Anti.d Degree of Radical Pyramidalization.
e UHF/6-31G(d) Values; PMP2/6-311G(d,p)//UHF/6-31G(d) Values in Parentheses.
' Relative Conformer Energy Differences, kcal mol'1. UHF/ 6-31G(d); PMP2/
6-311G(d,p)//UHF/6-31G(d) Values in Parentheses.

69
in accord with the greater exothermicity for the latter (AEâ„¢ = -22.35 kcal mol1 versus
-34.52 kcal mol'1, respectively, at the [QCISD(T)/6-311G(d,p)]7/UHF/6-31(d) level, and
C-C BDE data in Table 3-4).
Gauche-transoid addition of ethyl radical to Ci of propene occurs via a transition
structure with a forming C-C interatomic distance of 2.224 A and and a C=C bond length
of 1.385 A, in comparison with 2.274 A and 1.375 A for addition of pentafluoroethyl along
the same trajectory. Attack angles appear slightly smaller for additions of the fluorinated
radicals, most notably in the case of CH3* versus CF3* and consistent with a reinforced
SOMO-HOMO interaction for the latter. However, such differences are barely
significant, and in the ethyl series appear to be influenced more by the conformation of
the transition structure than the nature of the attacking radical, in line with the previously
observed insensitivity of transition state geometry to additions of both nucleophilic and
electrophilic radicals.47 48
Of note is the degree of pyramidalization at the radical site in the addition
transition structure, ranging from 102-105° for radicals of the RCH2* type and 108-109°
for a-fluorinated species. Considering the pyramidal nature of the ground states of the
latter as well (Figures 2-9 and 2-10) it follows that a-fluoroalkyl radicals enjoy a kinetic
advantage over their hydrocarbon analogues in that little or no additional bending is
necessary to accommodate the addition transition structure. The energetic cost of
pyramidalization of the methyl and fert-butyl radicals to the same extent as required for
their addition to ethylene has been computed at 1.5 and 1.6 kcal mol'1, respectively, at
the RMP2/6-31 G(d)//UHF/6-31 G(d) level.49
Calculated degrees of charge transfer (CT) between radical and olefin moieties
of the addition transition structures are provided in Table 3-7. Where applicable, lowest
energy conformations (gauche in the case of ethyl radical additions to ethylene, gauche-

70
transoid for additions to Ci of propene) were used in the determination of the
electrostatic potential-derived charges.
Table 3-7. Calculated Charge Transfer Data (Electrons) for Transition Structures of
Hydrocarbon and Fluorinated Methyl and Ethyl Radical Addition to Ethylene and Ci of
Propene.
Radical
Ethylene a
Prooene b
ch3*
-0.019
-0.004
cf3-
-0.013
-0.006
ch3ch2*
+0.036
+0.052
ch3cf2*
+0.030
+0.052
cf3ch2*
-0.047
-0.037
cf3cf2*
-0.045
-0.034
Note: Derived from UHF/6-31G(d) Electrostatic Potentials. Negative Values Denote
Electron Transfer from Alkene to Radical. 3 Gauche Transition Structure. b Gauche-
transoid Transition Structure.
Consistent with previous investigation, addition of CH3* to ethylene involves only
a slight degree of charge transfer (-0.019 e) from olefin to radical (Mulliken and Bader
analyses yield values of -0.017 and -0.011, respectively)48 and even less so for addition
to propene. Somewhat surprisingly, CF3* addition is also predicted to occur without
appreciable polarization.
Along the ethyl series, such interactions appear more clearly defined. Ethyl
radical addition to both ethylene and propene involves a shift of electron density from the
radical to the alkene (0.036 and 0.052 e, respectively) well in accord with the expected
nucleophilicity of the alkyl radicals. CH3CF2* is predicted to be nucleophilic as well,
exhibiting transition state polar characteristics very similar to those of its hydrocarbon
counterpart and consistent with the experimentally deduced non-electrophilicity of
the a-fluoro radicals.

71
A striking reversal in these trends occurs upon fluorination at the p carbon atom,
regardless of the nature of the radical site itself. Here it is seen that both CF3CH2* and
CF3CF2* exhibit substantial electrophilicty, with ca. 0.04 - 0.05 units of electron density
transferred from the alkene to the radical center. To place such values into some
degree of perspective, addition of the strongly electrophilic cyanomethyl radical to C2 of
electron-rich vinylamine is predicted to occur with a transfer of ca. 0.11 electrons from
CH2=CHNH2 to *CH2CN.48
It is especially noteworthy that the degree of CT in the case of CF3CH2* and
CF3CF2* addition is practically unaffected by fluorination at the a carbon (-0.047 versus
-0.045 and -0.037 versus -0.034, respectively). This, along with the demonstrated lack
of kinetic impact of fluorine substitution at the y position (Table 3-2) leads to the
conclusion that the electrophilic character of the perfluoroalkyl radicals derives
exclusively from substitution at the 2-position.
With geometric, enthalpic, and polar considerations for hydrofluorocarbon
radicals having been addressed, the influences of each of these effects on determined
kadd and kH values are now discussed.
a,a-Difluoroalkvl Radicals (77, 124)
The 1,1 -difluoroalkyl radicals, as mentioned previously, benefit from
pyramidalization at their radical site, leading to a more facile adoption of the transition
structure for addition or hydrogen atom abstraction. However, such an advantage is
counteracted by the experimentally and theoretically demonstrated lack of electrophilicity
for such species. Moreover, the terminal C-H bond weakening effect of gem-difluoro
substitution (2-3 kcal mol'1, Table 3-3) leads to a slight thermodynamic disadvantage
for hydrogen abstraction by the corresponding radical relative to the hydrocarbon. Thus,
the 3.8-fold rate enhancement enjoyed by 77 may be completely ascribed to the a-type,

72
pyramidal nature of its ground state, attenuated by enthalpic and polarity factors working
in opposition. In contrast, the strengthening effect of gem-difluorination on C-C bonds
compliments that of pyramidal geometry, leading to a more substantial (22.5-fold) rate
enhancement for the addition of 124 to styrene versus hydrocarbon 123.
B.B-Difluoroalkyl Radicals (88, 125)
The rate enhancements for addition (4.3) and hydrogen abstraction (5.8)
observed for 2,2-difluoroalk-1-yl radicals are due to a complimentary combination of
polar and enthalpic effects. The terminal C-H bond in 2,2-difluoropropane is predicted to
be 2.8 kcal mol'1 stronger than that of propane itself; similarly, gem-difluorination at C2
of butane leads to a 3.2 kcal mol'1 strengthening of its C3-C4 bond. In spite of these
favorable considerations, the near-planar ground state geometry of 88 results in a
modest net rate acceleration for hydrogen abstraction from nBu3SnH. This also
functions to oppose the CT and enthalpic advantages present in the addition reaction of
125, giving rise to only a slight rate increase relative to n-alkyls and certainly diminished
in comparison with that enjoyed by 124.
y-Fluorinated Radicals (111, 126)
Due to the near-planar geometric character of 3-fluoroalk-1-yls and the lack of
effect (ca. 1 kcal mol'1) on terminal C-H and C-C BDE values, fluorination beyond two
carbon atoms removed from the radical site exhibits a negligible effect on the rates of
both addition and hydrogen transfer. The reactivities of 111 and 126 are found to be,
within experimental error, identical to those of the corresponding hydrocarbon.
g.g.B.B-Tetrafluoroalkvl Radicals (99,100) and Pentafluoroethyl Radical (118)
Radicals substituted at the g and p positions benefit from both pyramidal
geometries and electrophilic character. Since the degree of pyramidalization of 1,1-

73
difluoroalkyl radicals remains constant regardless of substitution at the p- and further
positions, geometrically induced influences on the reactivities of such polyfluorinated
radicals are expected to be uniform. Consequently, rate enhancements for radicals of
the type RCH2CF2CF2\ RfCF2CF2CF2*, and CF3CF2* versus RCH2CH2CF2* (for kH\ 38,
83, and 133 versus 3.8; for kadd: 167, 383 and 658 versus 22.5; all relative to n-alkyls)
derive from either an increasing degree of transition state charge transfer stabilization,
increasingly greater exothermicity of reaction, or a combination of both. The relevant
C-H BDE data in Table 3-3 yields no direct correlation between reaction rate and
enthalpy for the polyfluorinated radical series, with values of 97.7, 100.1, and 99.5
kcal mol'1 corresponding to hydrogen abstraction by radicals 77, 100, and 118.
Similarly, terminal C-C BDEs of 91.6, 95.4, and 95.5 kcal mol'1, equated with the
additions of 124, 99, and 118, illustrate that although p-fluoro substitution should lead to
rate enhancement on thermochemical grounds, such an effect does not account for the
incremental acceleration across the series.
With the degree of radical electrophilicity related to its substitution at the 2-
position and the potential for additional p C-F a* delocalization made possible by the
"extra" fluorine substituent in CF3CF2* relative to RCH2CF2CF2* and RfCF2CF2CF2\ it
follows that the increasing resonance and inductive withdrawal ability of these groups
relative to RCH2CH2CF2* sufficiently explain both the enhanced reactivity of these
radicals as a whole, as well as the observed trend.
Fluorine Substituent Effects on Bond Dissociation Energies; Coulombic Interactions
As previously discussed, substitution by fluorine in hydrocarbons gives rise to
nearly additive and sometimes opposite effects on C-H and C-C homolytic bond
dissociation energies. For example, the aforementioned 1-3 kcal mol'1 weakening
effect of a,a-difluoro substitution and the 3-5 kcal mol1 strengthening of terminal C-H

74
bonds by p-fluorination lead to near cancellation in the case of pentafluoroethane and
1,1,2,2-tetrafluoropropane, yielding BDE values very near those of the parent
hydrocarbon (Table 3-3.) Furthermore, the 4.9 kcal mol1 strengthening brought about
by substitution at the breaking C-C bond is reinforced by an additional 3-5 kcal mol'1
upon further fluorine incorporation at the p position, leading to net increases of nearly 10
kcal mol'1 over the parent in the C2-C3 homolysis of 1,1,1,2,2-pentafluoropropane and
cleavage of the terminal C-C bond of 2,2,3,3-tetrafluorobutane.
The opposite effects of a.a-difluoro substitution on C-H and C-C bond
dissociation energies bear special mention. Experimental BDE values in the
fluoromethanes are in accord with the general RSE expectations of Pasto (Table 2-7) in
that although substitution at a radical center by a single fluorine is stabilizing, its further
incorporation leads to a successive decrease in RSE, resulting in net destabilization for
CF3*. This is consistent with the incremental strengthening of C-F bonds along the
fluoromethane series, leading to a C-H BDE in CF3H which is 1.9 kcal mol1 stronger
than that of methane itself (methane BDE, 104.8 kcal mol'1; see discussion below Table
2-2) and the comparatively weaker C-H bonds in CH3F and CH2F2.
The considerable stability of the 2,2-difluoroalkanes relative to their 1,1 -difluoro
isomers, demonstrated by the isodesmic reaction in Equation 3-2 (calculated from
B3LYP/6-31G(d) total energies and zero-point corrections) provides the underlying
reason for why the stability trends observed above do not extend to C-C bonds.
CH3CF2CH3 + CH3CH3 - CH3CF2H + CH3CH2CH3 (3-2)
AE^ = + 7.8 kcal mol'1
In Chapter 2, the Wiberg rationale of electrostatic attraction for the incremental
strengthening and shortening of C-F bonds in the fluoromethanes was introduced.
Similarly, it is found that such a Coulombic-based argument sufficiently explains the
observed effects of fluorine substitution on C-H and C-C bond dissociation energies.

75
Atomic charges for select hydrofluorocarbons based on the B3LYP/6-31G(d)
electrostatic potential (Merz-Kollman radii) are provided in Table 3-8.
The dipolar nature of the C-C bond in 1,1,1-trifluoroethane and its resultant
increase in BDE relative to ethane and hexafluoroethane (Table 2-3) was first postulated
by Rodgers.156 Such stabilization due to increased C-C bond ionicity is seen in
Equations 3-3 (derived from experimental heats of formation188) and 3-4 (from B3LYP/6-
31G(d) total and zero-point energies).
Table 3-8. Atomic Charges in Hydrofluorocarbons, Based on B3LYP/6-31G(d) Density.
Hvdrofluorocarbon
xE
XQa
lCE
Xtla
Xtle
CaH3CaH3
Ethane
-0.055
+0.018
CaH2F2
Difluoromethane
-0.200
+0.320
+0.041
CpH3CaF2H
1,1-Difluoroethane
-0.228
+0.467
-0.386
+0.020
+0.110 (2H)
+0.135 (1H)
C|iH3CaF2C3H3
2,2-Difluoropropane
-0.245
+0.631
-0.477
+0.133 (4H)
+0.140 (2H)
CF3CF3 + CH3CH3 —^ 2 CH3CF3 (3-3)
AErxn = -16.9 kcal mol'1
CH3CH2CH2CH3 + CH3CF2CF2CH3 ^ 2 CH3CF2CH2CH3 (3-4)
AE^ = -5.0 kcal mol'1
The significant electrostatic attraction between adjacent carbon atoms in both
CH3CF2H and CH3CF2CH3 is readily apparent from the data in Table 3-8, providing an
explanation for the strengthening of these bonds relative to their hydrocarbon or
perfluorocarbon analogues. In addition, C-H repulsion in difluoromethane and 1,1-
difluoroethane is predicted, in accord with the observed weakening of these bonds
compared to those of methane and ethane. Conversely, the strong attraction between

76
the p carbon and hydrogen atoms of CH3CF2H (xC, -0.386; +0.118) and
CH3CF2CH3 (xC, -0.477; xHavg, +0.135) is consistent with their greater theoretical and
experimental BDEs.
Conclusion
Based on time-resolved kaM measurements, absolute rate constants for
hydrogen abstraction from tri-n-butyltin hydride have been determined for a series
partially fluorinated radicals. The reactivities of such radicals towards nBu3SnH follow
those of addition to alkenes. The enhanced reactivity of a,a-difluoroalkyl radicals in
hydrogen abstraction reactions derives exclusively from their pyramidal geometry.
P-Fluorination leads to a favorable combination of polar and thermodynamic factors in
both addition and hydrogen transfer reactions, giving rise to the exceptional reactivity of
CF3CF2* and the perfluoroalkyl radicals as a whole. In Chapter 4, the kH values so
obtained are utilized in the determination of absolute rates of cyclization for partially
fluorinated 5-hexenyl radicals.
A self-consistent rationale for the impact of fluorine substitution on C-H and C-C
bond dissociation energies based on electrostatic considerations was offered, providing
new understanding of the thermochemistry of bonding and radical stabilization in
hydrofluorocarbons.

CHAPTER 4
THE REACTIVITY OF PARTIALLY FLUORINATED RADICALS
IN INTRAMOLECULAR CYCLIZATION REACTIONS
Introduction
The intramolecular addition reactions of 5-hexen-1-yl radicals continue to attract
the attention of synthetic and physical organic chemists alike. Such cyclizations to
(predominantly) 5-exo products have been utilized as probes for the detection of radical
intermediates and as basis reactions for the competitive determination of absolute
kinetic data for a number of free radical transformations.16 Rationalization of the rates,
and especially the regio- and stereochemistry, of intramolecular radical additions on the
basis of force field63 64 66 67 and molecular orbital62 65 66 68 techniques has proven to be one
of the greatest successes of theory in the prediction of organic reactivity. Due in no
small part to such structure-reactivity studies, application of free radical methodology to
the singular and tandem construction of 5-membered rings has been equally exploited,
providing for the assembly of functionalized organic systems under mild conditions, often
accomplished with a high degree of stereocontrol.11'15
Determination of absolute rates of cyclization of per- and other highly fluorinated
5-hexenyl systems86 89 170 have aided in solidifying the understanding of the effect of
fluorine substitution on the reactivity of organic radicals, though at the same time
generating a number of new questions, particularly with regard to cyclization
regiochemistry.
In order to examine the potentially more subtle influences of partial fluorination
on 5-hexenyl radical reactivity, and to obtain a set of data through which the effect of
77

78
incremental gem-difluoro substitution along the aliphatic portion of the 5-hexenyl chain
may be assessed, absolute rates of 5-exo and 6-endo cyclization for some partially
fluorinated 5-hexenyl radicals have been determined based on competitive kinetic
technique and the absolute rates of hydrogen abstraction obtained in Chapter 3.
Precursor Syntheses and Competitive Kinetic Studies
As in the bimolecular addition versus hydrogen abstraction competition studies,
bromide precursors were utilized in the generation of partially fluorinated 5-hexenyl
radicals. Photolysis by UV irradiation (Rayonet photoreactor) in the presence of known,
varying concentrations of hydrogen atom donor, carefully adjusted to ensure pseudo-first
order kinetic behavior and to allow for accurately measurable amounts of cyclization and
hydrogen abstraction products, provided the kinetic ratio kH / kCn. Absolute rate
constants for 5-exo and (where applicable) 6-endo cyclization were then determined
from the known value of hydrogen abstraction rate constant kH, illustrated in Figure 1-18
and in greater detail below.
1.1-Difluorohex-5-en-1-yl Radical (128)
Synthesis of bromide 135 was achieved in six steps in ca. 14.5% overall yield,
starting from commercially available 3-buten-1-ol (129, Figure 4-1). Curiously, direct
addition of dibromodifluoromethane to 129 could not be induced, even through extended
reaction time at elevated temperatures. Although the presence of the alcohol
functionality in the alkene starting material would not have been expected to exhibit a
detrimental effect (in light of the hydroxylic nature of the ethanolamine / ferf-butanol
cosolvent medium) protection of the hydroxyl moiety as its te/1-butyldimethylsilyl ether
130 (TBDMSCI, imidazole in dimethylformamide) followed by dibromodifluoromethane
addition indeed afforded 1,3-dibromo-1,1-difluoro adduct 131 in good yield. Highly
selective displacement of the internal bromine yielded bromodifluoromethyl derivative

79
129
TBMSCI
â–º
ImH, DMF
^^XOTBDMS
(89.3%)
130
CF2Br2
â–º
tBuOH, H2NCH2CH2OH
CuCI (cat.)
Br
BrF2C Jv y\
v V OTBDMS
(72.4%)
131
NaBH4
â–º
DMSO
BrF2C yv /x
v v OTBDMS
(90.0%)
132
FeCI3
CH3CN
BrF2C /\
^ V OH
(97.7%)
133
PCC
ch2ci2 ^
o
(53.1%)
134
Ph3P=CH2
THF
^^/CF2Br
(48.1%)
135
Figure 4-1. Preparation of 6-Bromo-6,6-difluorohex-1-ene, Precursor to 1,1-Difluorohex-
5-en-1-yl Radical 128
132 with virtually no overreduction product, as monitored via 19F NMR through high
conversion of starting material. Lewis acid deprotection via the method of Cort193 and
subsequent pyridinium chlorochromate oxidation provided aldehyde 134, further
subjected to Wittig olefination to yield precursor 135.
Generation of 128 was achieved via irradiation of a solution of 135 in C6D6 in the
presence of excess nBu3SnH (Figure 4-2) and an internal standard of a,a,a-
trifluorotoluene. Direct capture of 128 by hydrogen atom donor afforded reduction
product 6,6-difluorohex-1-ene 136, whereas intermediate 137, subsequently trapped by
nBu3SnH to yield spectroscopically observable cyclization product 138, was generated
via irreversible, unimolecular rearrangement with rate constant kc5 (no 6-endo cyclization
was observed, within NMR detection limits, (ca. 4%) for 128). During the course of the

80
¿^\/CF2Br
135
hv
nBu3SnH
CeDg
/'^/CF2'
128
+ nBu3Sn‘
nBu3SnH
137 138
+ nBu3Sn‘
Figure 4-2. kH / kc Competitive Kinetic Scheme for 1,1-Difluorohex-5-en-1-yl Radical
128
reaction, tributylstannyl radicals generated by transfer of hydrogen atom from nBu3SnH
to 128 and 137 served to propagate the chain process via bromine abstraction from 135.
Product ratios for varied concentrations of nBu3SnH were determined by 19F
NMR analysis according to the pseudo-first-order relation in Equation 4-1,
[136] _ [*h] [128 ] [ nBu3SnH ]
[138] ' [ Arcs I 1128 ]
a plot of which obtained for each data point versus nBu3SnH concentration providing the
ratio kH / /cC5. Exceptionally clean spectra and high mass balances were obtained for
each kinetic point, indicating the efficiency of the radical chain process and reliability of
the obtained rate constant ratios. A partial 19F spectrum of the first of six data points is
provided in Figure 4-3, a doublet of triplets (-CF2H, <(> -116.2) observed for 136 versus
overlapping doublets of doublets of triplets at <(> -100.3 and -107.8 for the diastereotopic
-CF2- resonances of 138. Kinetic data and product yields are given in Table 4-1, a plot
of which found in Figure 4-4. The slope of the line (2.57 ± 0.05) in conjunction with the

81
-96
-98
-100 -102 -104 -106 -108 -110
-112
-114 -116
-118
Figure 4-3. Partial 19F NMR Spectrum of Data Point 1 for kH / kc Competition of 1,1-
Difluorohex-5-en-1-yl Radical 128.
Table 4-1. Competitive Kinetic Data for kH / kc Competition of 1,1-Difluorohex-5-en-1-yl
Radical 77.
L135J
Í nBusSnH 1
i 1361/f1381
% Yield
0.054
0.673
1.53
88
0.054
0.807
1.91
100
0.054
0.942
2.28
89
0.054
1.08
2.57
94
0.054
1.21
2.93
95
0.054
1.35
3.29
92
Figure 4-4. Plot of the Data in Columns 2 and 3 of Table 4-1.

82
known absolute rate constant for hydrogen atom abstraction from nBu3SnH by 1,1-
difluorohex-1-yl radical 77, 9.1 (± 1.7) x 106 M1 s'1, resulted in a kCs value of
3.5 (± 0.59) x 106 s'1 for 5-exo closure of 128, with errors in kc reflecting both the least-
squares fit of the line and propagated error in kH. Syntheses of hydrogen atom transfer
and cyclization products 136 and 138 were performed as shown in Figure 4-5.
,^WCF2Br
135
nBu3SnH
AIBN
Mesitylene
O
DAST
CH2CI2
139 138
Figure 4-5. Preparation of Hydrogen Abstraction and 5-Exo Cyclization Products 136
and 138
2,2-Difluorohex-5-en-1-vl Radical (140)
Bromide 144 was obtained in a three-step synthesis starting from 1,2-epoxy-5-
hexene (Figure 4-6). Regiospecific ring opening by a Corey194 procedure afforded
bromohydrin 142, converted to the corresponding a-haloketone via Jones oxidation.
Treatment of 143 with DAST in dichloromethane afforded precursor 144 in 40.9% overall
yield, purified by preparative GC for competitive kinetic study.
In contrast to the virtually regiospecific 5-exo closure of 128, a broad singlet
resonance at 4» -95.8, comprising approximately 9% of cyclized products, was observed
in the 19F NMR spectra for the kH / kc competition of 140. This is attributed to competing
6-endo cyclization to 148 (Figure 4-7), the presence of which was confirmed by spectral
comparison with that of an authentic sample of 149.

83
141
KBr, CH3C02H
â–º
thf/h2o
(88.1%)
142
Na2Cr207 / H2SO4
Et20
DAST
â–º
ch2ci2
'^^x^/CF2CH2Br
(82.6%) (56.2%)
143 144
Figure 4-6. Preparation of 6-Bromo-5,5-difluorohex-1-ene, Precursor to 2,2-Difluorohex-
5-en-1-yl Radical 140.
^/\^CF2CH2Br
144
hv
nBu3SnH
CgDg
\/\v/cf2ch2 •
140
148
149
Figure 4-7. kH / kc Competitive Kinetic Scheme for 2,2-Difluorohex-5-en-1-yl Radical
140

84
Cyclizations of p,p-difluoroalkyl radicals have appeared in the synthetic
literature,195 utilized in the generation of alkoxy-substituted gem-difluorocyclopentane,
cyclohexane, and tetrahydropyran derivatives, though reported to undergo addition in an
exo-specific manner. However, the regiochemical behavior exhibited in the cyclization of
140 provided experimental verification of that previously predicted on the basis of ab
initio calculations, performed as part of the present study and elaborated upon in the
Discussion section of the chapter. Competition plots for kH / kC5 and kH / kCe are found in
Figures 4-8 and 4-9, respectively. Preparation of hydrogen abstraction and 5-exo and 6-
endo cyclization products was performed as shown in Figure 4-10.
[nBugSnH]
Figure 4-8. Plot of kH / kC5 Competition of 2,2-Difluorohex-5-en-1-yl Radical 140.
[nBugSnH ]
Figure 4-9. Plot of kH I kCe Competition of 2,2-Difluorohex-5-en-1-yl Radical 140.

85
<^/^/CF2CH2Br
144
nBu3SnH
AIBN
Mesitylene
^/\yCF2CH3
145
Figure 4-10. Preparation of Hydrogen Abstraction and 5-Exo and 6-Endo Cyclization
Products 145, 147, and 149.
1,1,2,2-Tetrafluorohex-5-en-1-vl Radical (152)
Bromide 101 (Halocarbons, Inc.) served as the precursor to a,a,p,p-
tetrafluorinated radical 152. However, attempts at determination of accurate kH / kc
ratios using nBu3SnH as a trapping agent met with failure, leading primarily to reduction
product 153, with only minor amounts of 155 and 177 evident in the 19F NMR baseline
which could not be integrated accurately over a span of hydrogen atom donor
concentrations (Figure 4-11). In principle, lowering the concentration of both radical
precursor (typically in the 0.05 - 0.1 M range) and trapping agent (while still maintaining
pseudo-first order conditions) should effectively decrease the amount of reduction
product and allow for a greater degree of cyclization to be observed. However, too great
of a decrease in precursor concentration leads to decreased NMR signal to noise ratios,
the necessity of longer acquisition times per sample, and increased potential for the
introduction of systematic error.

86
CF2Br
hv
nBu3SnH
CgDe
101
152
nBu3SnH
.cf2h
virtually no cyclization products
observed
153
Figure 4-11. Attempted kH / kc Competition of 1,1,2,2-Tetrafluorohex-5-en-1-yl Radical
152 with nBu3SnH as Trapping Agent.
As it was evident that any cyclization reaction of 152 occurred with a rate
constant too low to be competitive with transfer of hydrogen from nBu3SnH, attention
was turned to alternative trapping agents. With the rate of hydrogen atom transfer to
perfluoroalkyl radicals by a number of reducing agents having been accurately
determined (Table 2-10), it was decided to investigate the suitability of
tris(trimethylsilyl)silane ((TMS)3SiH) as a competitive trapping agent for the calibration of
cyclization rate constants for 152, due to its approximately four-fold decrease in
hydrogen transfer rate to perfluoroalkyls relative to nBu3SnH. For such a competition to
be of kinetic value, however, it was necessary to determine rate constant kH for
tetrafluoroalkyl radical 100 with (TMS)3SiH, using its known rate of addition to styrene as
a competing basis reaction. The plot for the kH ((TMS)3SiH) / kadd (styrene) competition
of 100 is provided in Figure 4-12.

87
With the kH ((TMS)3SiH) value of 1.8 (± 0.1) x 107 M'1 s1 for 100 in hand (which,
along with its kH (nBu3SnH) of 9.2 (± 0.8) x 107 M 1 s’1, (Table 3-2) may be compared
with 5.1 x 107 M 1 s'1 and 2.0 x 108 M'1 s'1, respectively, for perfluoro-n-alkyl radicals;
Table 2-10) rate constants kCb and kC6 for 152 were then determined (Figures 4-13 and
4-14.) Use of this slower hydrogen transfer agent allowed for sufficient competitive
(including significant 6-endo) cyclization such that accurate kH / kCn ratios could be
obtained. Isolation of products 153, 155, and 157 was achieved by slow syringe pump
addition of nBu3SnH to a heated, irradiated solution of 101 in mesitylene (Figure 4-15).
Figure 4-12. Plot for kH ((TMS)3SiH) / kadd (Styrene) Competition of 1,1,2,2-
Tetrafluorohex-1-yl Radical 100.
[ (TMS)3SiH ]
Figure 4-13. Plot of kH / kc5 Competition of 1,1,2,2-Tetrafluorohex-5-en-1-yl Radical 152.

[153]/[157 ]
88
[ (TMS)3SiH ]
Figure 4-14. Plot of kH / kce Competition of 1,1,2,2-Tetrafluorohex-5-en-1-yl Radical 152.
,CF2Br
101
nBu3SnH
AIBN
Mesitylene
152
156
157
Figure 4-15. Preparation of Hydrogen Abstraction and 5-Exo and 6-Endo Cyclization
Products 153,155, and 157.

89
Discussion
Absolute rate constants of cyclization for radicals 128, 140, and 152 are given in
Table 4-2. For comparison, such kc5 and (where applicable) kce values for parent
hydrocarbon 1 and fluorinated radicals 65 and 68 are also provided, the latter two
systems along with those of the current study found to give rise to the greatest impact on
cyclization kinetics and regiochemistry. Recent studies of 5-hexenyl systems bearing
vinylic fluorine substituents have demonstrated that the effect of such substitution is
relatively minor, with no 6-endo products observed within the detection limits imposed by
NMR analysis and AcC5 (rei) values with respect to 1 ranging from ca. 0.09 to 2.3.170
Cyclization Kinetics
Rates of intramolecular addition of partially-fluorinated radicals to alkenes should
be governed by the same combination of steric, polar, and thermodynamic factors which
influence the reactivity of their intermolecular counterparts. As seen by comparison of
the data in Tables 3-2 and 4-2, the reactivity characteristics of the above radicals in
unimolecular cyclization reactions, particularly 5-exo closure, generally reflect those
observed in bimolecular additions. This is logical in light of the similarity of their
transition structures, elaborated upon in Chapters 1 and 3.“
The pyramidal nature of a,a-difluoroalkyl radicals, combined with the more
favorable thermodynamics of C-C bond formation involving fluorinated carbon (see
related discussions in Chapters 2 and 3, along with cyclization transition structures and
energies of reaction below) provide sufficient explanation for the 13-fold increase in rate
of 5-exo ring closure of 128 relative to 1. The factor of 22.5 observed for addition of 77
versus 124 to styrene is consistent with the observed cyclization rate ratios.
The increase in kc5 of 4.1 enjoyed by 140 parallels that of bimolecular addition of
125 to styrene, (4.3) due to its increased electrophilicty over both hydrocarbon 1 and

90
Table 4-2. Absolute Rate Constants for 5-Exo and 6-Endo Cyclization of Partially
Fluorinated 5-Hexenyl Radicals. Rate Constants are for 303 K; Relative kCn Values in
Parentheses.
Cyclization Reaction kcs, 105s'1 *C6- 105s'1
2.7 a
0.05
(1)
(1)
128
137 158
35 (± 5.9) b N/A b'c
(13.0)
140
11 (± 3.8) b 1.1 (± 0.34) b
(4.1) (22)
152
154
87 (± 4.1) b 19 (± 1.1) b
kC'CF2 (32.2) (380)
F2
156
440 (± 46) d 52 (±6.4) d
(163) (1040)
F2C.
c
F2
68
cf2
I
cf2
f2c cf2
f2c-cf2
69
cf2
f2c.c.cf2
f2
70
110 (±1.7)
(40.7)
35 (± 4.4)
(700)
a Reference 16. b Current Study. c 6-Endo Cyclization Not Observed Within 19F NMR
Detection Limits (Approximately 4%). d Reference 170.

91
a,a-difluorocarbon 128 and greater exothermicity of addition relative to n-alkyls, though
tempered by the effectively planar, 71-nature of its radical center.
a,a,p,p-Tetrafluorinated radical 152, as in the case of bimolecular additions,
benefits from favorable thermodynamics of addition as well as its electrophilicty and
a-character, leading to the 32-fold increase in kc5 compared to parent 1. It should be
noted that such unimolecular cyclizations possess an inherent entropic advantage over
their bimolecular analogues, generally proceeding with log A values ca. 2 units larger
than those for the latter16 and resulting in a leveling of rate ratios relative to
intermolecular additions.
Upon additional fluorination of the aliphatic moiety of the 5-hexenyl chain (65, 68)
such radicals undertake perfluoroalkyl character, leading to further increase in reactivity
akin to that observed for C7Fi5* (127, kaM (re() = 383) versus CH3CH2CF2CF2* (99,
kadd (rei) = 167, relative to /7-alkyl) in bimolecular additions to styrene. Geminal
difluorination at the allylic position (68) serves to diminish the transition state SOMO-
HOMO interaction, and hence kC5 and kce, relative to 65.
Cvclization Reqiochemistrv
The significant degree to which 152, 65, 68, and even 140 undergo 6-endo
cyclization is particularly striking, with six-membered ring formation in 140 occuring with
a rate nearly half, and 65 and 68 more than an order of magnitude greater than, that of
5-exo closure for hydrocarbon 1. In comparison, 5-hexenyl systems bearing alkyl
substituents along the aliphatic fragment exhibit regiochemical profiles similar to that of
the unsubstituted parent.70,71
The question of potential reversibility in the above cyclizations has been
addressed, in light of the greater relative thermodynamic stability of secondary
cyclohexyl radicals. Upon independent generation of 5-exo adduct radical 69 from
precursor 1-(iodomethyl)-2,2,3,3,4,4,5,5-octafluorocyclopentane in the presence of

92
hydrogen atom donor triethylsilane in C6D6, the only product observed after complete
consumption of starting material was that resulting from direct capture of 69 by Et3SiH.170
The lack of 6-endo or ring-opened products originating from 69, coupled with the ab initio
predictions based on relative energies of cyclization transition structures described
below, demonstrates that the regiochemical characteristics of fluorinated 5-hexenyl
radical cyclizations are indeed kinetic in nature.
Of further note is that system 68, which undergoes the greatest percentage
(24.1%) of 6-endo closure (that is, exhibiting the least selectivity) is not the most
reactive. Hexafluoro system 65, though forming 67 with a rate constant 1.5 times that of
analogous closure of 68 to 70, does so only to an extent of 10.6% of total cyclized
products.
A combined ab initio / molecular mechanics approach has allowed for accurate
regiochemical predictions for a number of alkyl and heteroalkyl intramolecular radical
additions.66 In order to examine the effect of the degree and location of fluorine
substitution on transition structure geometry and energetics, as well as on activation
barriers and reaction enthalpy, the "chair-like" and "boat-like" 5-exo and 6-endo
cyclization transition structures for the parent hydrocarbon and various fluorinated
5-hexenyl systems, along with their respective open-chain radicals and products of
5- and 6-membered ring closure, have been investigated with ab initio techniques.
In accordance with a UMINDO/3 investgation of Bischof,62 the lowest energy
conformation of 5-hexenyl radical 1 was found to be an all-trans methylene chain in a
gauche orientation with the internal vinyl hydrogen (Figure 4-16.) Alignment of the singly
occupied orbital of 1 with the adjacent C-H bond was found to be slightly preferred (ca.
0.1 kcal mol’1) over similar C-C alignment at the UHF/6-31G(d) + ZPE level.
From the calculated structures and energies of cyclization products
(cyclopentylmethyl and cyclohexyl radicals) it was possible to compute energies of
reaction for hydrocarbon 1 and its fluorinated analogues. Total energies of reactant and

93
product radicals and exothermicities of 5-exo and 6-endo cyclization for 1, 128, 152, and
65 are provided below in Table 4-3.
Figure 4-16. Lowest Energy Conformation of 5-Hexen-1-yl Radical 1; SOMO (On Right)
Aligned with Adjacent C-H Bond. UHF/6-31G(d) Optimized Geometry.
As expected from C-C BDE data, (Table 3-4) intermolecular additions of the
fluorinated species are, as a whole, more exothermic than for parent 1. However, no
direct correlation exists between either absolute rates of 5-exo and 6-endo addition or
relative percentage of 6-membered ring formation and its corresponding reaction
exothermicity. Although a steady increase in both kC5 and kCe is observed along the
series (1 -> 128 -> 152 -> 65), both cyclizations of 65 are predicted to be less
exothermic than those of 152. Furthermore, relative enthalpies (AErxn(i,5) - AEâ„¢^)) are
found to rise with the degree of fluorination, favoring 6-endo closure in consistent
manner for both levels of theory employed. This is at variance with the lesser extent of
6-endo closure in 65 compared to 152.
Total, zero-point, and relative energies along with pertinent geometrical
parameters for the UHF/6-31G(d) cyclization transition structures of 1, (depicted in
Figures 1-11 - 1-14) 128, 140, 152, and 65 are reported in Table 4-4. Although the
calculated energy differences between "chair" and "boat" forms of either 5-exo or 6-endo
transition structures are quite consistent among the theoretical methods, energies of the
"6-endo-chair1' and "6-endo-twist-boat' structures relative to the "5-exo-chair" and
"5-exo-boat appear to be overestimated at the PMP2/6-311G(d,p)//UHF/6-31G(d) level
compared to both UHF and QCISD(T) results. Bearing this in mind, relative transition

94
Table 4-3. Total and Zero-Point Energies and Energies of Reaction for 5-Exo and
6-Endo Cyclizations of Hydrocarbon and Partially Fluorinated 5-Hexenyl Radicals.
Radical
F a
ZPE b
AEnm (5-Exo) C
AEâ„¢ (6-Endo)c
AAErxn
1
-233.543768
(-234.461575)
[-234.391386]
0.161154
-13.20
(-18.58)
[-15.90]
-19.45
(-23.44)
[-21.22]
6.25
(4.86)
[5.32]
2
-233.567930
(-234.494315)
[-234.419913]
0.164664
3
-233.579940
(-234.504105)
[-234.431547]
0.166947
128
-431.260147
(-432.620036)
0.148540
-19.91
(-23.09)
-27.50
(-29.51)
7.59
(6.42)
137
-431.291745
(-432.656712)
0.148402
158
-431.305809
(-432.668897)
0.150592
152
-628.970751
(-630.769253)
0.132048
-20.29
(-23.38)
-29.15
(-31.30)
8.86
(7.92)
154
-629.003100
(-630.806518)
0.132061
156
-629.018995
(-630.820937)
0.134061
65
-826.678248
(-828.915986)
0.115172
-18.68
(-21.77)
-28.22
(-30.36)
9.54
(8.59)
66
-826.708181
(-828.950843)
0.115350
67
-826.724889
(-828.966024)
0.117035
3 In Hartrees. UHF/6-31G(d); PMP2/6-311G(d,p)//UHF/6-31G(d) Values in Parentheses,
QCISD(T)/6-31G(d)//UHF/6-31G(d) Values in Brackets. 6 In Hartrees, from UHF/
6-31 G(d) Vibrational Frequencies. c In kcal mol'1; UHF/6-31G(d); PMP2/
6-311G(d,p)//UHF/6-31G(d) Values in Parentheses, [QCISD(T)/6-311G(d,p)]7/UHF/
6-31 (d) Values in Brackets. d (AErxn(15) - AE^ 6)) in kcal mol'1; UHF/6-31G(d); PMP2/
6-311G(d,p)//UHF/6-31G(d) Values in Parentheses, [QCISD(T)/6-311G(d,p)]7/UHF/
6-31 (d) Values in Brackets.

95
Table 4-4. UHF/6-31G(d) Geometric Parameters and Energies of Transition Structures
for Hydrocarbon and Fluorocarbon 5-Hexenyl Cyclizations.
Radical
r(C-C)
(A)
r(C=C)
(A)
Z C-C-C
(Deg.)
Eisu
(au)
ZPE
(au)
AS
(eu)
P f
£reL
1 a
2.186
1.393
109.7
-233.525546
(-234.455098)
[-234.379299]
0.161472
78.11
0
^b
2.192
1.394
108.2
-233.522908
(-234.452664)
[-234.376772]
0.161317
78.91
1.57
(1.44)
[1.38]
1 c
2.260
1.384
98.4
-233.521897
(-234.450180)
[-234.375531]
0.161801
76.91
2.24
(3.27)
[2.47]
1d
2.245
1.386
100.0
-233.517681
(-234.446462)
[-234.371537]
0.161736
77.77
5.08
(5.56)
[4.81]
128 3
2.192
1.387
108.9
-431.244513
(-432.614627)
0.146381
86.27
0
128 ü
2.199
1.388
107.0
-431.241713
(-432.611948)
0.146261
86.97
1.69
(1.61)
128 c
2.259
1.383
96.8
-431.241187
(-432.609955)
0.146926
85.37
2.39
(3.24)
128 d
2.240
1.385
99.9
-431.237289
(-432.606471)
0.146800
85.89
4.76
(5.35)
140 a
2.190
1.391
109.2
-431.248280
(-432.615856)
0.145212
85.54
0
140 b
2.193
1.393
107.4
-431.245575
(-432.613199)
0.145080
86.64
1.62
(1.59)
140 c
2.272
1.382
98.0
-431.247082
(-432.613844)
0.145752
83.98
1.05
(1.57)
140 d
2.257
1.384
99.6
-431.241993
(-432.608898)
0.145540
84.99
4.13
(4.55)
152 3
2.193
1.386
108.6
-628.954014
(-630.762361)
0.129951
93.47
0
152 b
2.195
1.387
106.5
-628.951440
(-630.759854)
0.129840
94.44
1.55
(1.51)
152 c
2.269
1.381
96.3
-628.953289
(-630.760695)
0.130521
92.33
0.77
(1.36)

96
Table 4-4-- continued
Radical
r (C-C)
(A)
r (C=C)
(A)
Z C-C-C
(Deg)
Etot_!
(au)
ZPE
(au)
AS
ieu]
P f
^rel_
152 d
2.249
1.383
100.0
-628.947730
(-630.754984)
0.130365
78.91
4.17
(4.86)
65 a
2.195
1.387
107.9
-826.661679
(-828.909230)
0.113158
76.91
0
65 b
2.199
1.389
106.3
-826.659545
(-828.907252)
0.113017
77.77
1.26
(1.16)
65 c
2.271
1.381
96.8
-826.660359
(-828.906567)
0.113677
86.27
1.12
(1.96)
65 d
2.252
1.383
99.3
-826.656303
(-828.902416)
0.113613
86.97
3.62
(4.53)
a 5-Exo-chair. b 5-Exo-boat. c 6-Endo-chair. d 6-Endo-twist-boat. e UHF/6-31G(d);
PMP2/6-311G(d,p)//UHF/6-31G(d) Values in Parentheses, QCISD(T)/6-31G(d)//UHF/
6-31 G(d) Values in Brackets. Mn kcal mol'1; UHF/6-31G(d); PMP2/6-311G(d,p)//
UHF/6-31G(d) Values in Parentheses, [QCISD(T)/6-311 G(d.p)]7/UHF/6-31 (d) Values in
Brackets.
state free energies were obtained for each structure based on the UHF/6-31G(d)
energies and entropies, and regiochemical ratios determined based on a Boltzmann
distribution including all four transition structures.
In a similar manner, the above structures along with those of 1,1,2,2,3,3,4,4-
octafluoro radical 68, perfluorinated system 71, and partially fluorinated 74 were also
investigated with the UHF/4-31G model, this smaller basis set implemented due to the
size of the latter three systems. Relative free energies for all systems investigated are
provided in Table 4-5, with resultant predicted and experimental 5-exo : 6-endo ratios for
each cyclization given in Table 4-6.
Agreement between the computed and observed values is quite remarkable,
even with the relatively small 4-31G basis, the largest deviation from experiment
approximately eight percent. From these results, in can be inferred that the
regioselectivities for a variety of radical cyclizations (hydrocarbon, fluorocarbon, or
otherwise) can be predicted with a reasonably high degree of confidence at minimal

97
Table 4-5. Relative Free
Energies of
Hydrocarbon and
Fluorocarbon Transition
Structures (kcal
mol'1) at 303 K.
Radical
5-Exo-chair
5-Exo-boat
6-Endo-chair
6-Endo-twist-boat
1
0.0
1.22
3.10
5.33
(0.0)
(1.33)
(2.84)
(5.19)
128
0.0
1.44
2.91
4.94
(0.0)
(1.48)
(2.67)
(4.88)
140
0.0
1.23
1.35
4.28
(0.0)
(1.29)
(1.53)
(4.30)
152
0.0
1.26
1.07
4.57
(0.0)
(1.26)
(1.12)
(4.37)
65
0.0
1.15
1.68
4.02
(0.0)
(1.22)
(1.45)
(3.96)
68
0.0
0.72
0.83
4.01
(0.0)
71
0.0
0.61
2.95
6.66
(0.0)
74
0.0
1.14
2.10
6.57
(0.0)
Note: UHF/4-31G; UHF/6-31G(d) Values in Parentheses.
computational expense and without the introduction of experimentally-based
parameters. This is due in part to the fact that the degree of spin contamination in each
of these structures, though rather high (with values ranging in magnitude from
approximately 1.0 to 1.1) is fairly consistent. This is not true, however, of the ground
state starting structures, which adversely affects computed activation barriers as
discussed below.
The experimental activation barriers for 5-exo and 6-endo closure of 1 are 6.8
and 8.5 kcal mol"1, respectively.63 64 Calculated energies of activation for the parent
hydrocarbon and number of fluorinated analogues are found in Table 4-7.
It can be seen from the results of 1 that absolute barriers for radical cyclization
reactions are extremely difficult to model theoretically. Even with the inclusion of

98
Table 4-6. Computed and Experimental Regiochemical Ratios at 303 K for Hydrocarbon
and Fluorocarbon 5-Hexenyl Radical Cyclizations.
Radical 5-Exo : 6-Endo Ratio 8
Predicted
Observed
1
100 : 0
(99.3 : 0.7)
98 : 2
128
100 : 0
(98.9: 1.1)
> 96 : < 4
140
91.3 : 8.7
(93.3 : 6.7)
90.9 : 9.1
152
86.9 : 13.1
(87.8 : 12.2)
82.1 : 17.9
65
94.8 : 5.2
(92.6 : 7.4)
89.4 : 10.6
68
83.8 : 16.2
75.9: 24.1
71
99.5 : 0.5
> 96 : < 4
74
97.4 : 2.6
> 96 : < 4
a UHF/4-31G; UHF/6-31G(d) Values in Parentheses.
electron correlation, accuracy leaves quite a bit to be desired. These results are in
accord with the those of Wong and Radom,41 whose systematic investigation of
intermolecular additions of small model radicals to olefins has demonstrated that high
levels of theory are required to accurately reproduce (and predict) activation barriers for
radical reactions. This is due in part to the differing degree of spin contamination in the
starting radicals and transition structures as mentioned previously, the former
possessing "acceptable" values in the range of 0.76 - 0.79, versus that of 0.75 for
a "pure" doublet radical. Such differential spin contamination has a disastrous effect on
activation barriers computed using perturbation theory. For example, UMP2/6-
311G(d,p)//UHF/6-31G(d) yields values of 13.5 and 17.3 kcal mol'1, respectively, for 1;
calculations at the UMP4SDQ/6-311G(d,p) level on UHF/6-31G(d) geometries fare only
slightly better with values of 12.1 and 15.3, only a modest improvement over UMP2/

99
Table 4-7. Calculated Activation Barriers for 5-Exo and 6-Endo Cyclizations of
Hydrocarbon and Fluorocarbon 5-Hexen-1-yl Radicals.
Radical
UHF/4-31G
UHF/6-31G(d)
PMP2/6-311G(d.D)//UHF/6-31G(d)
1
5-exo: 10.67
5-exo : 11.61
5-exo : 4.24
6-endo: 13.38
6-endo : 14.09
6-endo: 7.51
128
5-exo: 6.10
5-exo: 8.60
5-exo: 2.18
6-endo: 8.75
6-endo: 10.99
6-endo: 5.42
140
5-exo: 6.16
5-exo: 9.33
5-exo: 3.15
6-endo : 6.89
6-endo: 10.10
6-endo: 4.51
152
5-exo: 6.56
5-exo : 9.27
5-exo: 3.11
6-endo : 7.93
6-endo : 10.39
6-endo: 5.07
65
5-exo : 8.86
6-endo: 9.41
Note: In kcal mol'1. Includes 0.8929 ZPE Correction.
6-311G(d,p) and still worse than the Hartree-Fock results. The slow convergence of the
UMPn series in systems suffering from severe spin contamination has been noted by
Nobes et a/.196
The spin-projected PMP2/6-311G(d,p) values, though faring somewhat better in
the case of 1, do not reproduce the correct trends for either 5-exo or 6-endo reactivity
along the fluorinated series. Methods of high accuracy such as coupled cluster (CC)
and quadratic configuration interaction (QCI) have also been shown to be robust against
spin contamination197 198 but at significant (and, for larger systems, prohibitive) expense.
Results of QCISD(T) barrier calculations for 1 and 128 are provided in Table 4-8, the
former requiring 7-10 and the latter 28-35 hours of CPU time and approximately 2-4 GB
of disk space on the Cray C90 at the San Diego Supercomputer Center for each
calculation of the energies of starting species and transition structures. As expected,
however, the agreement between theory and experiment for 1 is much improved at this
level, especially with partial inclusion of basis extension effects by evaluation at the
PMP2/6-311G(d,p) level.

100
Table 4-8. Activation Barriers for 5-Exo and 6-Endo Cyclization of 1 and 128 at the
QCISD(T) Level.
Radical
QCISDm/6-31G(d)
fQCISDfn/6-311G(d,D)T a
1
5-exo : 7.76
6-endo: 10.31
5-exo: 7.30
6-endo: 9.77
128
5-exo : 4.98
6-endo: 7.21
5-exo : 5.21
6-endo : 7.75
Note: UHF/6-31G(d) Geometries, kcal mol'1. Includes 0.8929 ZPE Correction.
3 E[QCISD(T)/6-311G(d,p)]- * E[QCISD(T)/6-311G(d,p)] = E[QCISD(T)/6-31G(d)]
E(PMP2/6-311G(d,p)) - E(PMP2/6-31G(d)).
With such wide variations in variations in regiochemistry observed in the
cyclizations of 1, it is difficult to provide a rationale which satisfactorily accounts for all of
the data. This is especially so given the insensitivity of cyclization transition structure
geometry to substitution (Table 4-4 demonstrates the similarity in geometric trends to
those of intermolecular additions). However, the predictive value of the ab initio results
is encouraging, allowing for the potential design of precursors with desired regiochemical
behavior in cyclization and cyclopolymerization reactions.
Conclusion
Based on absolute rate constants of hydrogen abstraction from tributylstannane
and fr/s(trimethylsilyl)silane by partially fluorinated alkyl radicals, absolute rate constants
for 5-exo and 6-endo intramolecular addition reactions of partially fluorinated 5-hexen-1-
yl radicals have been obtained through competitive kinetic methods. Regardless of
either the degree or location of fluorine substitution on the 5-hexenyl chain, closure
proceeds predominantly in an exo fashion, although in some cases substantial
competing 6-endo cyclization is observed.
The kinetics of cyclization are found to follow those observed in the bimolecular
addition of partially fluorinated radicals to alkenes. The pyramidal nature of a,a-
difluoroalkyl radicals in combination with increased electrophilic character as function of

101
incremental fluorine substitution are responsible for their enhanced reactivity over
hydrocarbon analogues.
Consistent with high level computational studies on model radical addition
reactions, it is found that absolute activation barriers for intramolecular addition reactions
of 5-hexenyl radicals prove extremely difficult to model theoretically. Predictions of
cyclization regiochemistry based on relative ab initio transition state free energies may
be made with a reliable degree of accuracy.

CHAPTER 5
EXPERIMENTAL
General Methods- Experimental
NMR nuclear magnetic resonance (NMR) spectra (300 MHz, 75 MHz and 282
MHz for 1H, 13C, and 19F respectively) are reported in parts per million ppm downfield (8)
versus tetramethylsilane (TMS) for 1H and 13C, and in ppm upfield ((|)) versus CFCI3 for
19F. All NMR spectra were recorded on Varian VXR-300 or Gemini-300 spectrometers.
Preparative gas chromatographic (GC) separations were carried out with a 20
foot x 0.25 inch copper column packed with 20% SE-30 on Chromosorb P and
performed on a Varian Aerograph A-90 gas chromatograph equipped with a thermal
conductivity detector.
High resolution mass spectra were obtained on a Finnegan MAT-95
spectrometer.
Ultraviolet (UV) spectra were obtained on a Perkin-Elmer Lambda 9 UVA/IS/NIR
spectrophotometer.
All reagents, unless otherwise specified, were purchased from Aldrich, Fisher,
PCR, or Acros, and used as received. Styrene (Fisher) was freed from inhibitor by
passage through a column of neutral alumina. Dichloromethane was distilled from
calcium hydride and used immediately. Diethyl ether and tetrahydrofuran (THF) were
distilled from sodium benzophenone ketyl and used immediately. Benzene was distilled
from lithium aluminum hydride and stored over 4 Á molecular sieves. Chloroform,
dimethylsulfoxide, and tetraglyme were commercial anhydrous grade. All reactions were
performed under an inert atmosphere of argon.
102

103
General Methods- Theoretical
All ab initio calculations were performed with the Gaussian92 and Gaussian94
program systems199 on Cray Y-MP 4/32 or Cray C90 supercomputers, an IBM RS/6000
SP2 cluster, or Intel-based PCs. Geometry optimizations were performed using
standard gradient techniques. All stationary points were characterized by harmonic
frequency analysis, minima and transition structures giving rise to zero or exactly one
negative eigenvalue, respectively, in the second derivative matrix. All post-Hartree-Fock
calculations were performed with core orbitals frozen. Hartree-Fock frequencies and
resultant zero-point energies have been scaled by a factor of 0.8929.
Density functional theory calculations were performed with the Gaussian94
program on an IBM RS/6000 SP2 cluster or Intel-based PCs, implementing the
(75,302)p pruned integration grid. Zero-point energy corrections were scaled by 0.9806
as suggested by Scott and Radom and by Bauschlicher and Partridge.200 201 For bond
dissociation energy calculations, a thermal correction of ART (for C-C BDEs) or 2.5R7
(for C-H BDEs) has been applied as recommended by Hehre et al.202 Where applicable,
reported BDEs correspond to those resulting from the lowest electronic energy
conformer of the closed-shell species and / or radical.
Synthetic Procedures
Preparation of 1,3-Dibromo-1.1-difluorohexane (79)
A Carius tube of approximately 200 ml. capacity equipped with a small magnetic
stir bar was charged with 0.14 g (1.42 x 10'3 mol) cuprous chloride, 4.36 g (7.12 x 10'2
mol) ethanolamine, 12 mL tert-butanol, 10.0 g (1.42 x 10'1 mol) 1-pentene, and 59.84 g
(2.85 x 10'1 mol) dibromodifluoromethane. The tube was flushed with nitrogen and
flame-sealed; upon swirling, a deep blue coloration was observed. The tube was
immersed halfway into a silicon oil bath preheated at 85° C and allowed to stir for 48
hours, during which time the coloration turned from deep blue to olive green to brown.

104
(Caution: this procedure should be performed behind a safety shield). The tube was
cooled in an ice bath, opened, and the contents transferred to a 250-mL Erlenmeyer
flask (at this point, unreacted dibromodifluoromethane may be recovered by distillation)
and the tube rinsed with three 50 ml. portions of hexanes. All organic material (which
consisted of a cloudy yellow-green supernatant and a brown resin) was filtered through
50 mL of silica gel, which was rinsed with two additional 50 ml. portions of hexanes. The
resulting colorless filtrate was concentrated by rotary evaporation and subject to reduced
pressure fractional distillation through a 15 cm Vigreux column. A total of 22.79 g
(57.3%) 79 was obtained as a colorless liquid, bp 80-85° C / 25 mm Hg.
1.3-Dibromo-1,1 -difluorohexane (79): 1H NMR. 8 0.96 (3H, t, 3JHh = 7.42 Hz),
1.50 (2H, m), 1.86 (2H, m), 3.01 (2H, m), 4.24 (1H, m); 13C NMR: 8 13.2, 20.4, 40.5,
46.6, 52.7 (t, 2Jcf = 21.8 Hz), 120.6 (t, 1JCF = 306.9 Hz); 19F NMR: -43.2 (m); HRMS for
C6H10F2Br2: calc. 277.9117, calc. (M-Br) 198.9934, found 198.9983; CHN for C6H10F2Br2:
calc. 25.74% C, 3.60% H, found 25.53% C, 3.43% H.
Preparation of 1-Bromo-1,1 -difluorohexane (80) and 1,1-Difluorohexane (81)
A 250 mL three-necked round-bottomed flask equipped with an ice-water
condenser, argon inlet, and magnetic stir bar was charged with 20.0 g (7.14 x 10'2 mol)
1,3-dibromo-1,1 -difluorohexane (79) dissolved in 100 mL anhydrous dimethylsulfoxide.
A total of 10.8 g (2.85 x 10"1 mol) of sodium borohydride was then added in small
portions with vigorous stirring over the course of 1 hour, during which time the flask
became warm and a semisolid gel was observed to form. After the addition was
complete, the bath temperature was raised to 70° C over the course of 1 hour and
heating continued for an additional 6 hours (19F NMR analysis of a small aliquot of the
reaction mixture at this time showed complete consumption of starting material.) The
flask was cooled to room temperature, the contents transferred to a 1 L Erlenmeyer
flask, and the reaction quenched with chips of ice. The resulting mixture was carefully

105
acidified with concentrated aqueous HCI, 100 ml. ether was added, and the aqueous /
DMSO layer extracted with three 50 mL portions of ether. The combined ether layers
were washed with three 25 mL portions of water, dried over MgS04, and subject to
ambient pressure fractional distillation through a 15 cm Vigreux column. After
concentration in this way, 8.91 g (62.1%) 80 (contaminated with a small amount of 81)
was obtained as a colorless liquid, bp 125-128° C. Preparative GC separation afforded
analytically pure samples of each.
1 -Bromo-1,1 -difluorohexane (801: 1H NMR: 5 0.92 (3H, t, 3JHH = 7.5 Hz), 1.35
(4H, m), 1.62 (2H, m), 2.33 (2H, m); 13C NMR: 5 13.8, 22.3, 23.6, 30.6, 44.3 (t, 2JCF =
21.1 Hz), 123.3 (t, 1JCF = 303.5 Hz); 19F NMR. <|> -43.9 (t, 3JFH = 14.7 Hz); HRMS for
C6H11F2Br: calc. 200.0012, calc. (M-Br) 121.0829, found 121.0832; CHN for C6H11F2Br:
calc. 35.84% C, 5.51% H, found 35.47% C, 5.54% H.
1.1-Difluorohexane (81): 1H NMR: 8 0.91 (3H, t, 3JHH = 6.9 Hz), 1.34 (4H, m), 1.45
(2H, m), 1.80 (2H, m), 5.79 (1H, tt, 3JHH = 4.5 Hz, 2JHF = 57 Hz); 13C NMR: 8 13.8, 21.8 (t,
3Jcf = 5.5 Hz), 22.4, 31.2, 34.1 (t, 2JCF = 20.6 Hz), 117.5 (t, 1JCF = 237.4 Hz); 19F NMR: <(>
-116.3 (dt, 3Jfh = 17.1 Hz, 2JFH = 56.2 Hz); HRMS for C6H12F2: calc. 122.0907, calc.
(M-H) 121.0829, found 121.0825.
Preparation of 1-Phenyloctan-3-ol (86)
Into a flame-dried 300 mL three-necked round-bottomed flask equipped with a
125 mL pressure-equalizing addition funnel, condenser with argon inlet, and magnetic
stir bar was dispensed 3.94 g (1.62 x 10'1 mol) Mg turnings and a small crystal of iodine.
(2-Bromoethyl)benzene (10.0 g, 5.40 x 102 mol) was dissolved in 100 mL of anhydrous
ether and added to the funnel. After addition of a small amount of solution and initiation
of the reaction, addition was continued at a rate that maintained gentle reflux, after which
time the mixure was refluxed for an additional 2 hours. The flask was then cooled to
room temperature and 5.41 g (5.40 x 10"2 mol) of hexanal was dissolved in 50 mL

106
anhydrous ether, added to the addition funnel and dispensed dropwise at a rate that
maintained gentle reflux. The mixture was refluxed for an additional 2 hours, cooled to
room temperature, and quenched with a mixture of 100 g of ice and 50 mL concentrated
aqueous HCI. The layers were separated, the aqueous layer extracted with three 50 mL
portions of ether, and the combined organic extracts washed twice with 20 mL 5%
aqueous NaHC03 and once with 20 mL of water. The organic layer was dried over
MgS04, the solvent rotary evaporated, and the resulting liquid subjected to reduced
pressure fractional distillation through a 15 cm Vigreux column. 8.75 g (78.7%) 86 was
obtained as a colorless liquid, bp 148-151° C / 0.1 mm Hg.
1-Phenvloctan-3-ol (86): 1H NMR: 8 0.94 (3H, t, 3JHH = 5.7 Hz), 1.34 (4H, m), 1.51
(4H, m), 1.81 (2H, m), 2.76 (2H, m), 3.67 (1H, m), 7.28 (5H, m); 13C NMR: 8 14.0, 22.6,
25.2, 31.8, 32.0, 37.5, 39.0, 71.4, 125.7, 128.3, 128.4, 142.2; HRMS for C14H220: calc.
206.1671, found 206.1672.
Preparation of 1-Phenyloctan-3-one (87)
Into a 100 mL round-bottomed flask was placed 3.8 g (1.84 x 10'2 mol) of
1-phenyloctan-3-ol (86) dissolved in 10 mL ether. Jones' reagent (8.5 mL, previously
prepared with of 5.0 g Na2Cr207 and 3.65 mL concentrated H2S04 diluted with water to a
total volume of 25 mL) was added dropwise to the alcohol solution with stirring, and
allowed to react for 4 hours at room temperature. The dark green reaction mixture was
diluted with 20 mL ether and 20 mL water. The aqueous layer was extracted with three
20 mL portions of ether, and the combined organic extracts washed twice with 20 mL
saturated NaHC03, once with 20 mL brine, and dried over MgS04. Rotary evaporation
of the solvent followed by reduced pressure fractional distillation through a 10 cm
Vigreux column afforded 3.34 g (89.0%) 87, bp 108-111° C / 0.35 mm Hg.
1 -Phenvloctan-3-one (87): 1H NMR: 8 0.87 (3H, t, 3JHH = 6.9 Hz), 1.26 (4H, m),
1.55 (2H, overlapping tt, 3JHH = 7.5 Hz), 2.37 (2H, t, 3JHH = 7.2 Hz), 2.72 (2H, m), 2.89

107
(2H, m), 7.23 (5H, m); 13C NMR: 5 13.9, 22.4, 23.5, 29.8, 31.4, 43.0, 44.2, 126.0, 128.3,
128.4, 141.2, 210.3; HRMS for C14H2oO: calc. 204.1514, found 204.1532; CHN for
C14H20O: calc. 82.30% C, 9.87% H, found 82.53% C, 10.14% H.
Preparation of 3,3-Difluoro-1-phenyloctane (83)
Into a 100 mL three-necked round-bottomed flask equipped with stir bar, rubber
septum, and condenser with argon inlet was placed 2.0 g (9.80 x 10'3 mol) of 1-
phenyloctan-3-one (87) dissolved in 15 mL anhydrous CH2CI2. Into the reaction mixture
was slowly injected 1.74 g (1.08 x 10'2 mol) diethylaminosulfurtrifluoride (DAST) with
stirring, during which time the flask became slightly warm. The mixture was heated at
reflux for 72 hours, over which time a dark amber coloration was observed. After
cooling, the reaction mixture was carefully decanted onto 50 mL of ice water and diluted
with 10 mL CH2CI2. The organic layer was separated and the aqueous layer extracted
with three 20 mL portions of CH2CI2. The combined organic extracts were washed once
with 10 mL of 10% aqueous NaHC03 and once with 10 mL water, dried over MgS04,
and the solvent rotary evaporated. Fractional reduced pressure distillation of the
resulting liquid afforded 1.66 g (68.0%) 83, bp 100-105° C / 0.2 mm Hg.
3,3-Difluoro-1 -phenvloctane (83): 1H NMR: 8 0.92 (3H, t, 3JHH = 6.3 Hz), 1.33
(4H, m), 1.48 (2H, m), 1.87 (2H, m), 2.13 (2H, m), 2.82 (2H, m), 7.27 (5H, m); 13C NMR:
5 13.9, 22.0, 22.4, 28.5, 31.5, 36.5 (t, 2JCF = 25.1 Hz), 38.2 (t, 2JCF = 25.5 Hz), 124.8 (t,
1JCF = 239.4 Hz), 126.1, 128.3, 128.5, 140.8); 19F NMR: <|> -99.1 (overlapping tt, 3JFH =
17.1 Hz); HRMS for C14H20F2: calc. 226.1533, found 226.1529; CHN for C14H20F2: calc.
74.30% C, 8.91% H, found 74.14% C, 9.15% H.
Preparation of 1-Bromohexan-2-one (90)
Into a 250 mL three-necked round-bottomed flask equipped with an argon inlet,
magnetic stir bar, and 100 mL pressure-equalizing addition funnel was placed 10.0 g
(9.98 x 10'2 mol) of 2-hexanone, 40 mL glacial acetic acid, and 9.75 g (1.62 x 10'1 mol)

108
urea. The flask was cooled to 0° C and 17.15 g (1.07 x 10'1 mol) of bromine was
introduced dropwise into the reaction mixture over the course of 1 hour. After the
addition was complete, the cooling bath was removed and the reaction mixture allowed
to warm to room temperature and stirred overnight. The contents were then transferred
to a separatory funnel, diluted with 200 ml_ water, and extracted with four 50 mL portions
of CH2CI2. The combined organic extracts were washed with three 20 mL portions of
10% aqueous NaHC03, dried over MgS04, and the solvent rotary evaporated.
Fractional reduced pressure distillation through a 15 cm Vigreux column yielded 12.91 g
(72.2%) 90, contaminated with ca. 6% (by 1H NMR) of the 3-bromo isomer, bp 85-88° C /
15 mm Hg. Column chromatography (silica gel, 20% ethyl acetate in hexanes) of a
small sample yielded pure 90, which was used in characterization.
1 -Bromohexan-2-one (90): 1H NMR: 8 0.91 (3H, t, 3JHH = 7.5 Hz), 1.33 (2H,
overlapping tt, 3JHh = 7.5 Hz), 1.60 (2H, overlapping tt, 3JHh = 7.5 Hz), 2.64 (2H, t, 3JHh =
7.2 Hz), 3.88 (2H, s); 13C NMR: 5 13.7, 22.1, 25.9, 34.3, 39.5, 202.2; HRMS for
CeHuBrO: calc. 177.9993, found 177.9934; CHN for C6HnBrO: calc. 40.25% C, 6.19%
H, found 40.10% C, 6.19% H. (Caution: a-bromoketones are powerful lachrymators).
Preparation of 1-Bromo-2,2-difluorohexane (91)
A 100 mL three-necked round-bottomed flask equipped with magnetic stir bar,
condenser with argon inlet, and rubber septum was charged with 5.0 g (2.79 x 10'2 mol)
1-bromohexan-2-one (90) dissolved in 20 mL anhydrous CHCI3. DAST (9.15 g,
5.68 x 10'2 mol) was slowly injected into the reaction mixture, which was then heated at
50° C for 48 hours. The flask was cooled and the contents carefully dispensed into 50
mL of ice water. The layers were separated, the aqueous layer extracted with three 5
mL portions of CH2CI2, and the combined organic extracts washed twice with 10 mL
portions of 10% aqueous NaHC03 and once with 10 mL of water. Drying over MgS04,
evaporation of the halogenated solvents by ambient-pressure distillation through a 10

109
cm Vigreux column, and fractional, ambient pressure distillation afforded 3.35 g (59.7%)
91 as a colorless liquid, bp 120-122° C.
1-Bromo-2,2-difluorohexane (91): 1H NMR: 5 0.94 (3H, t, 3JHH = 7.2 Hz), 1.42
(4H, m), 2.02 (2H, m), 3.51 (2H, t, 3JHF = 13.2 Hz); 13C NMR: 13.7, 22.3, 24.1, 31.4 (t,
2Jcf = 34.1 Hz), 34.2 (t, 2JCF = 24.0 Hz), 121.5 (t, 1JCF = 240.9 Hz); 19F NMR: <|> -99.2 (m);
HRMS for CgHuFzBr: calc. 200.0012, calc. (M-H) 198.9934, found 198.9974; CHN for
C6HnF2Br: calc. 35.84% C, 5.51% H, found 36.00% C, 5.48% H.
Preparation of 1-lodo-2,2-difluorohexane (92)
1-Bromo-2,2-difluorohexane (91) (2.0 g, 9.95 x 10'3 mol) was placed into a thick-
walled Carius tube of approximately 150 mL capacity, along with a small magnetic stir
bar. 100 mL of a hot saturated solution of sodium iodide in acetone was added to the
tube, which was then cooled in a dry ice-isopropanol slush and flame-sealed. After
warming to room temperature, the tube was immersed halfway in an oil bath atop a
magnetic stirrer. The mixture was stirred at 85-90° C for 96 hours (a safety shield is
recommended) at which time the tube was cooled, opened, and the contents transferred
to a 250 mL round-bottomed flask. Most of the solvent was removed by rotary
evaporation, and the remaining residue taken up in a mixture of 50 mL of water and 50
mL of ether. The layers were separated, the aqueous layer extracted with three 50 mL
portions of ether, and the combined organic extracts washed with 20 mL of water.
Drying over MgS04, rotary evaporation of the solvent, and fractional reduced-pressure
distillation afforded 2.16 g (87.5%) of very pure 92, bp 101-102° C / 68 mm Hg.
1 -lodo-2.2-difluorohexane (92): 1H NMR: 8 0.93 (3H, t, 3JHH = 7.2 Hz); 1.41 (4H,
m), 2.06 (2H, m), 3.40 (2H, t, 3JHF = 14.4 Hz); 13C NMR: 5 3.95 (t, 2JCF = 31.5 Hz), 13.7,
22.3, 24.4, 35.0 (t, 2JCF = 24.5 Hz), 121.1 (t, 1JCF = 240.8 Hz); 19F NMR: <|) -94.9
(overlapping tt, 3JFH = 17.1 Hz); HRMS for C6HnF2l: calc. 247.9833, found 247.9862;
CHN for CeHuFzl: calc. 29.05% C, 4.47% H, found 28.82% C, 4.48% H.

110
Preparation of 2,2-Difluorohexane (93)
1-Bromo-2,2-difluorohexane (91) (1.0 g, 4.97 x 10'3 mol) was dissolved in 5 mL
benzene in a 25 mL round-bottomed flask equipped with septum-capped sidearm inlet
and magnetic stir bar, and attached to a small distillation apparatus equipped with with
ice water condenser and fractionating column. The bath temperature was raised to
60° C and 1.60 g (5.50 x 10‘3 mol) nBu3SnH and 0.01 g (6.09 x 10‘5 mol) 2,2-
azobisisobutyronitrile (AIBN) dissolved in 0.5 mL benzene was added slowly via syringe
through the rubber septum. When the addition was complete, the mixture was stirred for
an additional 30 minutes and the bath temperature quickly raised to 150° C. All volatile
material was flash distilled and subjected to preparative GC separation which afforded
analytically pure 93 as a colorless liquid.
2,2-Difluorohexane (93): 1H NMR: 5 0.92 (3H, t, 3JHH = 7.2 Hz), 1.42 (4H, m), 1.58
(3H, t, 3Jhf = 18.3 Hz), 1.83 (2H, m); 13C NMR: 5 13.8, 22.5, 23.2 (t, 2JCF = 28.1 Hz), 24.9
(t, 3Jcf = 4.5 Hz), 37.7 (t, 2JCF = 25.1 Hz), 124.5 (t, 1JCF = 236.4 Hz); 19F NMR: -90.9
(m); HRMS forC6H12F2: calc. 122.0907, calc. (M-HF) 102.0845, found 102.0822.
Preparation of 1-Phenyloctan-4-ol (96)
In a manner similar to that of the preparation of 86, 10.0 g (5.02 x 10'2 mol) of
1-bromo-3-phenylpropane dissolved in 100 mL anhydrous ether was added to 3.67 g
(1.51 x 10‘1 mol) Mg turnings to which a crystal of iodine had been added. Subsequent
addition of 4.54 g (5.27 x 10'2 mol) of valeraldehyde dissolved in 50 mL anhydrous ether
followed by workup in the usual way afforded 7.36 g (71.1%) 96, bp 102-103° C / 0.05
mm Hg.
1-Phenvloctan-4-ol (96): 1H NMR: 8 0.91 (3H, t, 3JHH = 6.6 Hz), 1.52 (10H,
overlapping m), 2.65 (2H, t, 3JHH = 7.5 Hz), 3.62 (1H, m), 7.24 (5H, m); 13C NMR: 14.0,
22.7, 27.4, 27.8, 35.9, 37.0, 37.1, 71.7, 125.7, 128.2, 128.3, 142.4; HRMS for C14H220;

111
calc. 206.1671, found 206.1671; CHN for C14H220: calc. 81.50% C, 10.75% H, found
81.72% C, 10.92% H.
Preparation of 1-Phenvloctan-4-one (97)
In a manner similar to that of the preparation of 87, 15 mL of Jones’ reagent was
added to 5.0 g (2.42 x 10‘2 mol) 1-phenyloctan-4-ol (96) in 25 mL ether. Workup in the
usual way followed by distillation at reduced pressure yielded 4.14 g (83.7%) 97, bp 123-
125° C / 0.4 mm Hg.
1 -Phenvloctan-4-one (97): 1H NMR: 8 0.96 (3H, t, 3JHH = 7.2 Hz), 1.36 (2H, m),
1.60 (2H, m), 1.97 (2H, overlapping tt, 3JHH = 7.5 Hz), 2.46 (4H, m), 2.69 (2H, t, 3JHH =
7.2 Hz), 7.30 (5H, m); 13C NMR: 13.8, 22.3, 25.2, 25.8, 35.0, 41.8, 42.5, 125.8, 128.3,
128.4, 141.6, 211.0; HRMS for C14H20O: calc. 204.1514, found 204.1576; CHN C14H20O:
calc. 82.30% C, 9.87% H, found 82.40% C, 9.83% H.
Preparation of 4,4-Difluoro-1-phenyloctane (98)
In a manner similar to the preparation of 83, 2.5 g (1.22 x 10'2 mol)
1-phenyloctan-4-one (97) and 3.93 g (2.44 x 10'2 mol) DAST in 25 mL CH2CI2 were
refluxed for 48 hours. Workup in the usual way afforded 1.83 g (66.3%) 98, bp
85-88° C/0.1 mm Hg.
4,4-Difluoro-1-phenvloctane (98): 1H NMR: 5 0.89 (3H, t, 3JHH = 6.9 Hz), 1.34 (4H,
m), 1.80 (6H, m), 2.63 (2H, t, 3JHH = 6.9 Hz), 7.22 (5H, m); 13C NMR: 5 13.8, 22.5, 24.0
(t, 3Jcf = 4.5 Hz), 24.38 (t, 3JCF = 4.7 Hz), 35.4, 35.7 (t, 2JCF = 26.0 Hz), 36.1 (t, 2JCF =
25.1 Hz), 125.2 (t, 1JCF = 238.9 Hz), 125.9, 128.4 (2C, overlapping), 141.5; 19F NMR: 98.3 (overlapping tt, 3JFH = 17.1 Hz); HRMS for C14H20F2: calc. 226.1533, found
226.1533; CHN for C14H20F2: calc. 74.30% C, 8.91% H, found 74.40%C, 8.85% H.

112
Preparation of 1-Bromo-1,1.2.2-tetrafluorohexane (102)
Into a 250-mL three-necked round-bottomed flask equipped with a magnetic stir
bar, rubber septum, and condenser equipped with an argon inlet was added 36 mL
tetraglyme and 30.0 g (1.28 x 10'1 mol) of 6-bromo-5,5,6,6-tetrafluorohex-1-ene (101). A
2.0 M solution of borane dimethyl sulfide in diethyl ether (24 mL, 4.80 x 10'2 mol) was
slowly injected into the flask through the septum. Some bubbling was evident and the
flask became slightly warm. This mixture was stirred for two hours at room temperature
then heated at reflux with stirring overnight. The flask was then cooled to room
temperature and 64 mL (59.3 g, 5.11 x 101 mol) of hexanoic acid was slowly injected
into the reaction mixture with stirring. Vigorous bubbling was evident and the flask
became warm. Stirring was continued for 2 hours at room temperature then at reflux
overnight. The mixture was distilled through a 15 cm Vigreux column at ambient
pressure until the head temperature reached 140° C, at which time distillation ceased.
The distillate was diluted with 50 mL of ether, washed with two 10 mL portions of
saturated aqueous sodium bicarbonate and two 10 mL portions of water, dried, and the
resulting solution distilled at ambient pressure. After removal of ether and residual
dimethyl sulfide in this way, 24.16 g (79.6%) 102 was obtained as a colorless liquid
boiling at 122-124° C.
1-Bromo-1,1,2.2-tetrafluorohexane (102): 1H NMR: 8 0.95 (3H, t, 3JHH = 7.2 Hz),
1.41 (2H, m), 1.59 (2H, m), 2.07 (2H, m); 13C NMR: 8 13.7, 22.3, 22.6, 30.1 (t, 2JCF =
22.5 Hz), 117.5 (tt, 2Jcf = 31.1 Hz, 1JCF = 251.8 Hz,), 118.0 (tt, 2JCF = 39.6 Hz, 1JCF =
309.5 Hz,); 19F NMR: <|> -65.9 (2F, s), -112.6 (2F, t, 3JFH = 19.5 Hz); HRMS for C6H9F4Br:
calc. 235.9823, calc. (M+H) 236.9902, found 236.9738; CHN for C6H9F4Br: calc. 30.40%
C, 3.83% H, found 30.72% C, 3.46% H.

113
Attempted Transhaloqenation Reactions of 102 via Lithium-Halogen Exchange and
Formation of 1,1,2-Trifluoro-1-hexene (103)
In each instance, 1.0 g (4.22 x 10'3 mol) 102 was dissolved in 25 mL dry diethyl
ether and cooled in a bath of liquid nitrogen / pentane. A slight excess (1.1 equivalents;
2.2 in the case of f-butyllithium) of the desired alkyllithium as a solution in diethyl ether
was added very slowly to the reaction mixture via syringe. An solution of excess iodine
in diethyl ether was then added dropwise to the mixture and allowed to warm to room
temperature and stir overnight in each case. 19F NMR analysis showed in all cases
quantitative formation of 103. Use of carbon dioxide as an alternative electrophile was
not successful; p-fluoride elimination of the lithiated species was too rapid, even at -100°
C, to be successfully trapped.
1.1.2-Trifluoro-1-hexene (103): 19F NMR (CDCI3): 5 -106.9 (1F, dd, 2JFF = 90 Hz,
3Jff = 32 Hz), -125.9 (1F, dd, 2JFF = 90 Hz, 3JFF = 112 Hz), -174 .8 (1F, dtd, 3JFF = 112
Hz, 3Jff = 32 Hz, 3Jfh = 22 Hz).
Attempted lodofluorination of 103.
To a solution of 1,1,2-trifluoro-1 -hexene (103) (containing 9.04 x 10'3 mol as
judged by a 19F NMR standard of hexafluorobenzene) in 10 mL of diethyl ether at 0° C
was added 3.7 mL (3.65 g, 2.27 x 10‘2 mol) of triethylamine trihydrofluoride. N-
iodosuccinimide (2.24 g, 9.94 x 103 mol) was added in portions, the mixture warmed to
room temperature, and stirred overnight. 19F NMR analysis of the mixture showed no
change in the spectrum from that of the starting material. The mixture was then heated
at reflux for an additional 24 hours, at which time 19F NMR analysis indicated no
reaction.
Preparation of 1.1.2,2-Tetrafluoro-1,4-diiodobutane (105)
Into a stainless steel pressure reactor of approximately 700 mL capacity was
added 100 g (2.83 x 10'1 mol) of 1,2-diiodotetrafluoroethane and 0.5 g (3.67 x 10'3 mol)

114
of d-limonene. The reactor was tightly sealed, connected to a vacuum line, immersed in
a liquid nitrogen bath, and subjected to three successive freeze-pump-thaw cycles to
remove oxygen. A total of 15.86 g (5.65 x 10'1 mol) of ethylene was transferred into the
bomb, which was then sealed and placed into a heating manifold. (Caution: the
following is performed behind a safety shield.) The thermostat was then set at 210° C,
and with constant stirring the reaction mixture was heated for 8 hours. After cooling in
an ice bath, venting, and opening, the contents (a dark violet liquid and dark solid) were
diluted with 50 mL CHCI3, transferred to a 250 mL Erlenmeyer flask, and chilled in a
freezer at -20° C, which caused precipitation of a large amount of dark solid. The
mixture was filtered, the solid washed with three 20 mL portions of cold CHCI3, and the
combined organic liquids rotary evaporated at room temperature. The remaining liquid
was fractionally distilled at reduced pressure through a 15 cm Vigreux column. A total of
43.36 g (40.1%) 105 was obtained as a violet liquid, bp 82-85° C / 22 mm Hg.
1,1,2.2-Tetrafluoro-1.4-diiodobutane (105): 1H NMR: 8 2.70 (2H, m), 3.24 (2H,
m); 13C NMR: 8 -10.5, 34.7 (t, 2JCF = 22.7 Hz), 96.7 (tt, 2JCF = 42.8 Hz, 1JCF = 315.9 Hz),
116.5 (tt, 2Jcf = 30.9 Hz, 1JCF = 253.1 Hz); 19F NMR: (2F, tt, 3Jff = 4.9 Hz, 3Jfh = 17.0 Hz).
Preparation of 3,3,4,4-Tetrafluoro-4-iodobut-1-ene (106)
Into a 250 mL three-necked round-bottomed flask equipped with magnetic stir
bar, argon inlet, and pressure-equalizing addition funnel was added 30.69 g (8.03 x 10‘2
mol) 1,1,2,2-tetrafluoro-1,4-diiodobutane (105) dissolved in 60 mL anhydrous ether.
DBU (26.88 g, 1.77 x 10'1 mol) dissolved in 50 mL anydrous ether was added dropwise
at room temperature and allowed to stir for an additional 6 hours; at this time, 19F NMR
analysis of a small aliquot of the reaction mixture showed complete consumption of
starting material. The mixture was poured into 50 mL of 5% aqueous HCI, the layers
separated, and the aqueous phase extracted with three 10 mL portions of ether. The

115
combined organic fractions were washed once with 20 ml. of saturated NaHC03
solution, once with 10 mL of water, dried over MgS04, and the ether carefully removed
by gentle, ambient pressure distillation through a 15 cm Vigreux column. Upon removal
of most of the solvent in this way, the bath temperature was increased and 13.46 g
(66.0%) 106 was collected over a range of 90-92° C.
3.3,4,4-Tetrafluoro-4-iodobut-1-ene (106): 1H NMR: 5 5.83 (1H, m), 5.99 (2H, m);
13C NMR: 5 97.3 (tt, 2JCF = 44.1 Hz, 1JCF = 316.5 Hz), 113.5 (tt, 2JCF = 30.1 Hz, 1JCF =
249.9 Hz), 124.2 (t, 2JCF = 26.1 Hz), 126.0 (t, 3JCF = 8.6 Hz); 19F NMR: -60.7 (2F, t, 3JFH
= 7.3 Hz),-108.7 (2F, m).
Preparation of 1,1,2,2-Tetrafluoro-1-iodobutane (107)
Into a 250 mL three-necked round-bottomed flask equipped with reflux
condenser, pressure-equalizing addition funnel and magnetic stir bar was added 10.0 g
(3.94 x 10'2 mol) of 3,3,4,4-tetrafluoro-4-iodobut-1-ene (106) dissolved in 50 mL of dry
methanol, 3.20 g (9.98 x 10'2 mol) anhydrous hydrazine, and 0.1 g (1.01 x 10'3 mol)
cuprous chloride. The flask was cooled to 0° C, and 14.2 g of a 30% aqueous solution of
hydrogen peroxide was delivered dropwise over a period of 20 minutes, over which time
the color of the reaction mixture changed to powder blue then to a dark amber as the
H202 addition neared completion. The cooling bath was then removed, the mixture
allowed to stir at room temperature for an additional 2 hours, and the contents poured
into a solution of 1 mL of concentrated HCI in 200 mL of water. The mixture was
extracted with six 10 mL portions of 1,2-dichlorobenzene, dried over Na2S04, and
transferred to a 100 mL flask attached to a 15 cm Vigreux column. A vacuum adapter
attached to small trap (immersed in a dry ice-isopropanol slush) was attached, the
pressure lowered to 100 mm Hg, and the flask heated with vigorous stirring until the
contents began to boil. At this time, it was observed that approximately 1 mL of material

116
had accumulated in the trap, which was subjected to preparative GC separation. A total
of 1.08 g (10.7%) 107 was collected as a colorless liquid.
1.1.2.2-Tetrafluoro-1 -iodobutane (107): 1H NMR: 8 1.15 (3H, t, 3JHH = 7.5 Hz),
2.11 (2H, m); 13C NMR: 5.05, 23.0 (t, 2JCF = 23.6 Hz); 98.5 (tt, 2JCF = 43.5 Hz, 1JCF =
315.9 Hz); 117.5 (tt, 2JCF = 29.6 Hz, 1JCF = 250.9 Hz); 19F NMR: -59.9 (2F, t, 3JFF = 4.9
Hz), -110.4 (2F, tt, 3Jff = 4.9 Hz, 3JFH = 17.1 Hz).
Preparation of 1,1,2,2-Tetrafluorohexane (108)
1-Bromo-1,1,2,2-tetrafluorohexane (102) (2.0 g, 8.44 x 10'3 mol) was dissolved in
5 ml. of benzene in a 25 ml_ round-bottomed flask equipped with septum-capped
sidearm inlet and stir bar, and attached to a small distillation apparatus equipped with ice
water condenser and fractionating column. The bath temperature was raised to 60° C
and 2.95 g (1.01 x 10'2 mol) nBu3SnH and 0.01 g (6.09 x 10"5 mol) AIBN dissolved in 0.5
mL benzene was added slowly via syringe through the rubber septum. When the
addition was complete, the mixture was stirred for an additional 30 minutes and the bath
temperature quickly raised to 150° C. All volatile material was flash distilled and
subjected to preparative GC separation, which afforded pure 108 as a colorless liquid.
1.1.2.2-Tetrafluorohexane (108): 1H NMR: 8 0.95 (3H, t, 3JHH = 7.5 Hz), 1.40 (2H,
m), 1.55 (2H, m), 1.95 (2H, m), 5.70 (1H, tt, 3JHF = 3.0 Hz, 2JHF = 54.0 Hz); 13C NMR:
13.6, 22.4 (2C, overlapping), 29.6 (t, 2JCF = 22.5 Hz), 110.4 (tt, 2JCF = 41.1 Hz, 1JCF =
248.4 Hz), 118.1 (tt, 2Jcf = 28.5 Hz, 1JCF = 243.8 Hz); 19F NMR: <|> -116.8 (2F, t, 3JFH =
19.5 Hz), -136.1 (2F, d, 2JFH = 53.7 Hz); HRMS for C6H10F4: calc. 158.0719, calc. (M-H)
157.0640, found 157.0648.
Preparation of 3.3,4,4-Tetrafluoro-1-phenyloctane (110)
1-Bromo-1,1,2,2-tetrafluorohexane (102) (2.0 g, 8.44 x 10‘3 mol) and 1.76 g
(1.69 x 10'2 mol) styrene dissolved in 25 mL benzene were added to a 100 mL three¬
necked round-bottomed flask fitted with rubber septum, condenser with argon inlet, and

117
magnetic stirrer. A total of 4.90 g (1.68 x 10"2 mol) nBu3SnH and 0.05 g (3.04 x 10'4 mol)
AIBN were dissolved in 5 mL benzene and taken up in a syringe. A 150 W flood lamp
was placed at a distance of approximately 15 cm from the flask, and with irradiation (at
this distance, sufficient heat was generated to cause the solvent to reflux as well) the
nBu3SnH solution was delivered to the reaction mixture via syringe pump over a 36 hour
period. After rotary evaporation at elevated temperature, column chromatography (silica
gel, hexanes) afforded four fractions of 110 free from organotin contaminants.
3,3,4,4-Tetrafluoro-1-phenyloctane (110): 1H NMR: 6 0.96 (3H, t, 3JHh = 7.2 Hz),
1.41 (2H, m), 1.57 (2H, m), 2.02 (2H, m), 2.31 (2H, m), 2.90 (2H, m), 7.28 (5H, m) 13C
NMR: 1.0, 13.8, 22.5, 26.9, 29.7 (t, 2JCF = 22.7 Hz), 32.1 (t, 2JCF = 22.8 Hz), 118.8 (tt,
2Jcf = 37.6 Hz, 1JCF = 250.4 Hz); 119.3 (tt, 2JCF = 34.1 Hz, 1JCF = 245.9 Hz), 126.3, 128.3,
128.6, 140.3); 19F NMR: <|> -116.0 (2F, m), -116.4 (2F, m); HRMS for C14H18F4: calc.
262.1345, found 262.1307.
Preparation of 1-fPerfluorohexyllethane (113)
Into a 25 mL round-bottomed flask equipped with septum-capped sidearm inlet
and magnetic stir bar was placed 4.22 g (8.90 x 10'3 mol) 2-[perfluorohexyl]-1-
iodoethane (112.) The flask was attached to a distillation apparatus equipped with a
small fractionating column and ice water condenser. The bath temperature was raised
to 80° C and 3.11 g (1.07 x 10‘2 mol) nBu3SnH was slowly injected into the flask with
stirring. After 30 minutes at this temperature, the bath temperature was raised to 150° C
and the product distilled over a range of 81-82° C. Preparative GC purification afforded
pure 113.
1 -[Perfluorohexvllethane (1131: 1H NMR: 6 1.14 (3H, m), 2.10 (2H, m); 19F NMR:
-81.5 (3F, t, 3Jff = 9.8 Hz), -117.0 (2F, m), -122.5 (2F, br s), -123.4 (2F, br m), -124.2
(2F, br m), -126.8 (2F, m); HRMS for C8H5F13: calc. 348.0183, calc. (M-F) 329.0200,
found 329.0280.

118
Preparation of 2-lodo-1-fperfluorohexvn-4-phenylbutane (116)
Into a 50 mL three-necked round-bottomed flask equipped with argon inlet,
rubber septum and stir bar was added 0.5 g (3.78 x 10'3 mol) 4-phenyl-1-butene and
2.02 g (4.53 x 10'3 mol) perfluorohexyl iodide dissolved in 20 mL of hexanes. A 1.0 M
solution of triethylborane in hexanes (0.4 mL, 4 x 10"4 mol) was slowly injected through
the septum, and the reaction allowed to stir for 6 hours at room temperature. The
mixture was washed twice with 10 mL water, the solvent rotary evaporated, and the
remaining liquid subject to reduced pressure fractional distillation. A total of 1.81 g
(82.8%) 116 was collected as a light violet liquid, bp 125-127° C /1 mm Hg.
2-lodo-1-rperfluorohexvn-4-phenvlbutane (116): 1H NMR: 8 2.12 (2H, m), 2.75
(2H, m), 2.92 (2H, m), 4.27 (1H, m), 7.27 (5H, m); 19F NMR: <)> -81.3 (3F, t, 3JFF = 9.8 Hz),
-111.8 (1F, dm, 2Jff = 275.9 Hz), -115.2 (1F, dm, 2JFF = 267.3 Hz), -122.3 (2F, br s),
-123.3 (2F, s), -124.2 (2F, br s), -126.7 (2F, m); HRMS for C16H12F13I: calc. 577.9776,
found 577.9815.
Preparation 1-fPerfluorohexvl1-4-phenvlbutane (117)
Into a 50 mL three-necked round-bottomed flask equipped with argon inlet and
magnetic stirrer was added 1.5 g (2.59 x 10‘3 mol) 2-iodo-1-[perfluorohexyl]-4-
phenylbutane (116) dissolved in 10 mL DMSO. The bath temperature was raised to
70° C and 0.39 g (1.03 x 10"2 mol) sodium borohydride was added in portions to the
reaction mixure. Stirring was continued for an additional 6 hours, at which time the
mixture was poured into 50 mL of ice water, and carefully acidified with 6 M HCI. 10 mL
ether was added, the layers separated, and the aqueous layer extracted with three 10
mL portions of ether. The combined organic layers were washed once with 10 mL 5%
aqueous NaHC03 and twice with 10 mL of water, dried over MgS04, and the solvent
rotary evaporated. Reduced pressure fractional distillation afforded 0.82 g (70.0 %) 117,
bp 83-85° C / 0.25 mm Hg.

119
1-ÍPerfluorohexvll-4-phenvlbutane (117): 1H NMR: 5 1.67 (4H, m), 2.06 (2H, m),
2.64 (2H, t, 3Jhh = 6.9 Hz), 7.24 (5H, m); 19F NMR: <|> -81.4 (3F, t, 3JFF = 9.8 Hz), -114.9
(2F, m), -122.5 (2F, br s), -123.4 (2F, s), -124.1 (2F, s), -126.7 (2F, m); HRMS for
Ci6H13F13: cale. 452.0810, found 452.0797.
Preparation of Pentafluoroethane (120)
lodopentafluoroethane (2.16 g, 8.78 x 10'3 mol) was condensed into a Pyrex tube
equipped with a Rotaflo stopcock, rubber septum, and immersed in a dry ice-isopropanol
slush. In one portion, 3.07 g (1.05 x 10‘2 mol) nBu3SnH was injected into the tube and
the stopcock closed. The tube was then subjected to photolysis in a Rayonet reactor for
30 minutes at room temperature with periodic shaking. The tube was cooled to -20° C in
a dry ice-isopropanol bath, a rubber hose connected to the Rotaflo tube and to a trap
immersed in a liquid nitrogen-ether slush, and the stopcock opened. After trap-to-trap
transfer in this way, an NMR tube containing ca. 1 mL CDCI3, capped with a rubber
septum, and immersed in a liquid-nitrogen-ether slush was connected via cannula. The
trap was warmed to -20° C as a sample of 120 collected in the tube, which was flame
sealed and taken for NMR analysis.
Pentafluoroethane (120): 1H NMR: 5 5.88 (tq, 3JHF = 2.4 Hz, 2JHF = 52.2 Hz); 19F
NMR: -86.1 (3F, s), -138.5 (2F, d, 2JFH = 51.3 Hz).
Preparation of 1.1,1.2,2-Pentafluoro-4-phenylbutane (122)
lodopentafluoroethane (5.47 g, 2.22 x 10"2 mol) of was condensed into a Pyrex
Rotaflo tube containing 2.12 g (2.04 x 10'2 mol) styrene, 5.92 g (2.03 x 10'2 mol)
nBu3SnH, 5 mL benzene and a small sir bar. The stopcock was closed and the tube
irradiated with a 150 W flood lamp with stirring for 72 hours. Volatiles were removed by
rotary evaporation and the residue subject to flash vaccuum distillation at 100 mm Hg
(bath temperature 150° C) during which time the distillate temperature reached 69° C.
Preparative GC purification of the distillate yielded an analytically pure sample of 122.

120
1,1,1.2,2-Pentafluoro-4-phenvlbutane (122): 1H NMR: 5 2.38 (2H, m), 2.95 (2H,
m), 7.31 (5H, m); 13C NMR: 6 26.6, 32.7 (t, 2JCF = 21.5 Hz), 115.5 (tq, 2JCF = 37.6 Hz, 1JCF
= 250.4 Hz), 119.3 (qt, 2JCF = 36.1 Hz, 1JCF = 283.4 Hz), 126.7, 128.2, 128.8, 139.2; 19F
NMR: (j) -85.9 (3F, s), -119.1 (2F, t, 3JFH = 19.5 Hz); HRMS for C10H9F5: calc. 224.0624,
found 224.0670; CHN for C10H9F5: calc. 53.58% C, 4.05% H, found 53.42%C, 3.97% H.
Preparation of 1-(te/t-Butvldimethylsiloxyl)-3-butene (130)
Into a 250-mL three-necked round-bottomed flask equipped with a condenser
and argon inlet was placed 9.60 g (1.33 x 10'1 mol) of 3-buten-1-ol, 20 ml. DMF, 24.1 g
(1.60 x 10'1 mol) fert-butyldimethylsilyl chloride, and 22.7 g (3.33 x 10‘1 mol) imidazole.
This was stirred for 48 hours at room temperature under an argon atmosphere. The
contents of the flask were then poured into 250 ml. of pentane, and washed with three
50 mL portions of water followed by three 50 mL portions of saturated aqueous sodium
chloride. The organic phase was dried, the solvent rotary evaporated, and the resulting
liquid subjected to fractional reduced pressure distillation through a 15 cm vigreaux
column. A total of 22.08 g (89.3%) of pure 130 was obtained in 4 fractions as a colorless
liquid boiling at 102-105° C / 75 mm Hg.
1 -(tert-Butvldimethvlsiloxvl)-3-butene (130): 1H NMR (CDCI3): 5 0.05 (6H, s), 0.90
(9H, s), 2.27 (2H, dt, 3JHH = 7 Hz, 3JHH = 3 Hz), 3.66 (2H, t, 3JHH = 7 Hz), 5.02 (1H, m),
5.10 (1H, m), 5.81 (1H, m); 13C NMR (CDCI3): 5 -5.27, 18.3, 25.9, 37.5, 62.8, 116.2,
135.4; HRMS for Ci0H22SiO: calc. 186.1440, found (M+H) 187.1561; CHN for C10H22SiO:
calc. 64.45% C, 11.90% H, found 64.32%C, 12.03 % H.
Preparation of 1,3-Dibromo-5-(tert-butvldimethvlsiloxyl)-1,1-difluoropentane (131)
Cuprous chloride (0.10 g, 1.25 x 10'3 mol), along with 12.5 mL terf-butanol,
3.83 g (6.26 x 10'2 mol) ethanolamine, 23.30 g (1.25 x 10'1 mol) 1 -(tert-
butyldimethylsiloxyl)-3-butene (130), and 52.60 g (2.51 x 10‘1 mol) CF2Br2 were added

121
to a Carius tube. A small stir bar was added and the tube flame-sealed. After stirring at
85° C for 96 hours {performed behind a safety shield) the tube was cooled in an ice bath,
opened, and the contents transferred to a 500 mL Erlenmeyer flask. The tube was
rinsed with four 50 mL portions of hexanes, and the combined organic material filtered
through a 50 mL pad of silica gel, which was rinsed with three additional 50 mL portions
of hexanes. Rotary evaporation of the solvent afforded a colorless liquid judged by 1H
NMR to contain some unreacted starting material, 2.78 g of which was successfully
recovered by reduced pressure distillation at 102-105° C / 75 mm Hg. High vacuum was
then applied and a total of 35.83 g (72.4%, 82.0% based on consumed 130) 131 was
obtained as a colorless liquid boiling at 75-79° C / 0.09 mm Hg.
1.3-Dibromo-5-(tert-butvldimethvlsiloxvl)-1,1-difluoropentane (131): 1H NMR: 5
0.07 (3H, s), 0.08 (3H, s), 0.90 (9H, s), 1.95 (1H, m), 2.15 (1H, m), 3.08 (2H, m), 3.80
(2H, m), 4.46 (1H, m); 13C NMR: 5 -5.5, 18.2, 25.9, 41.3, 43.8, 52.9 (t, 2JCF = 19 Hz),
60.2, 120.64 (t, 1 JCf = 306 Hz); 19F NMR: <)> -43.1 (m); HRMS for Ci1H22SiOF2Br2: calc.
393.9774, calc. (M-f-C4H9) 336.9070, found 336.905; CHN for C11H22SiOF2Br2: calc.
33.35% C, 5.60% H, found 33.62% C, 5.62% H.
Preparation of 1-Bromo-5-(tert-butvldimethylsiloxvl)-1,1-difluoropentane (132)
1.3-Dibromo-5-(terf-butyldimethylsiloxyl)-1,1-difluoropentane (131) (30.15 g, 7.60
x 102 mol) was dissolved in 150 mL of dry DMSO in a 500 mL three-necked round-
bottomed flask equipped with an argon inlet and strong magnetic stir bar. Sodium
borohydride (11.5 g, 3.04 x 10'1 mol) was then added in portions with vigorous stirring.
After the addition was complete, the temperature was raised to 70-75° C over the course
of one hour and stirring continued for an additional 6 hours, at which time analysis of the
reaction mixture by 19F NMR demonstrated complete consumption of starting material.
The flask was cooled and carefully quenched with ca. 100 g of ice, and the contents
carefully acidified with concentrated hydrochloric acid and transferred to a 1 liter

122
separatory funnel. After extraction with three 100 mL portions of ether, the combined
extracts were washed with two 25 mL portions of water, dried over MgS04, and rotary
evaporated. The remaining liquid was distilled at reduced pressure through a 15 cm
vigreaux column, affording 21.70 g (90.0%) 132 as a colorless liquid, bp 108-111° C /10
mm Hg.
1 -Bromo-5-(te/f-butvldimethvlsiloxyl)-1,1 -difluoropentane (132): 1H NMR: 5 0.05
(6H, s), 0.90 (9H, s), 1.63 (4H, m), 2.38 (2H, m), 3.63 (2H, t, 3JHH = 6 Hz); 13C NMR: 8
18.3, 20.7, 25.9 (2C, overlapping), 31.4, 44.1 (t, 2JCF = 22.5 Hz), 62.4, 123.2 (t, 1JCF =
304 Hz); 19F NMR: <|> -44.0 (m); HRMS for C11H23SiOF2Br: calc. 316.0669, found (M+H)
317.0630.
Preparation of 5-Bromo-5,5-difluoropentan-1-ol (133)
Into a 250-mL round-bottomed flask was placed 15.04 g (4.74 x 102 mol)
1-bromo-5-(terf-butyldimethylsiloxyl)-1,1 -difluoropentane (132) along with 41 mL of
acetonitrile. To this was slowly added with stirring 7.70 g (4.75 x 10'2 mol) ferric chloride.
The reaction mixture turned a brick-red color and became slightly warm. The reaction
was allowed to stir for 3 hours at room temperature, at which time the contents of the
flask were poured into 250 mL of water and 100 mL of chloroform was added. The
chloroform layer was drained and the aqueous layer extracted with three 50 mL portions
of chloroform. These combined extracts were washed twice with 25 mL of water, dried
over MgS04 and the solvent rotary evaporated. A total of 9.41 g (97.7%) 132 was
collected by fractional reduced pressure distillation, bp 80-82° C /10 mm Hg.
5-Bromo-5,5-difluoropentan-1-ol (133): 1H NMR: 8 1.39 (1H, s), 1.69 (4H, m),
2.40 (2H, m), 3.69 (2H, t, 3JHH = 6 Hz); 13C NMR: 8 20.5, 31.3, 44.0 (t, 2JCF = 21.5 Hz),
62.2, 123.0 (t, 1JCF = 303.5 Hz); 19F NMR: <|> -44.1 (t, 3JFH = 14.7 Hz); HRMS for
C5H9F2BrO: calc. 201.9804, found (M+H) 202.9957; CHN for C5H9F2BrO: calc. 29.58%
C, 4.47% H, found 29.71% C, 4.46% H.

123
Preparation of 5-Bromo-5,5-difluoropentanal (134)
Into a 250-mL round-bottomed flask equipped with a magnetic stir bar was added
8.5 g (4.18 x 10'2 mol) 5-bromo-5,5-difluoropentan-1-ol (133) dissolved in 85 ml. of
dichloromethane. PCC (13.53 g, 6.28 x 102 mol) was added slowly in portions with
vigorous stirring. After the addition was complete, the mixture was allowed to stir at
room temperature for an additional 6 hours. The darkened reaction mixture (which
demonstrated complete consumption of starting material by TLC analysis) was filtered
through a pad of silica gel, which was rinsed with an additional three 10 mL portions of
CH2CI2. Rotary evaporation of the solvent followed by fractional distillation at reduced
pressure afforded 4.46 g (53.1%) 134 as a colorless liquid, bp 100-102° C / 50 mm Hg.
5-Bromo-5.5-difluoropentanal (134): 1H NMR: 5 1.97 (2H, m), 2.41 (2H, m), 2.59
(2H, t, 3Jhh = 7.2 Hz), 9.80 (1H, s); 13C NMR: 5 16.6 (t, 3JCF = 3.5 Hz), 42.1, 43.2 (t, 2JCF =
22.0 Hz); 122.5 (t, 1JCF = 303.4 Hz), 200.5; 19F NMR: <|> -44.4 (t, 3JFH = 14.4 Hz); HRMS
for C5H7F2BrO: calc. 199.9648, found (M+H) 200.9726.
Preparation of 6-Bromo-6,6-difluorohex-1-ene (135)
A 100 mL three-necked round-bottomed flask equipped with self-equalizing
addition funnel, argon inlet and magnetic stir bar was charged with 9.12 g (2.55 x 10'2
mol) of methyltriphenylphosphonium bromide and 20 mL anhydrous THF. The flask was
cooled to 0° C and 9.4 mL of a 2.5 M solution of butyllithium in hexanes (2.35 x 10'2 mol)
was added dropwise. After addition was complete, the mixture was stirred for an
additional 30 minutes at 0° C. 5-Bromo-5,5-difluoropentanal (134) (4.28 g, 2.13 x 10‘2
mol) was dissolved in 20 mL anhydrous THF and added dropwise to the reaction
mixture. After addition the mixture was allowed to warm to room temperature and stir for
an additional 6 hours. The contents were poured into 50 mL of water and extracted with
five 20-mL portions of ether. The combined ether fractions were dried over MgS04,
filtered, and the solution concentrated by distillation through a 15 cm vigreaux column.

124
Upon removal of most of the ether and residual THF the product was distilled at ambient
pressure, yielding 2.04 g (48.1%) 135, bp 120-123° C. An analytically pure sample was
obtained by preparative GC for spectroscopic analysis and kinetic experiments.
6-Bromo-6.6-difluorohex-1 -ene (135): 1H NMR: 8 1.74 (2H, m), 2.15 (2H,
overlapping dt, J = 7 Hz), 2.35 (2H, m), 5.05 (2H, m), 5.77 (1H, m); 13C NMR: 8 23.1,
32.3, 43.6 (t, 2Jcf = 21.6 Hz), 115.9, 123.1 (t, 1JCF = 303.5 Hz), 137.0; 19F NMR: <(> -43.9
(t, 3Jfh = 14.7 Hz); HRMS for C6H9F2Br: calc. 197.9855, calc (M-Br) 119.0672, found
119.0667; CHN for C5H9F2Br: calc. 36.21% C, 4.56% H, found 36.28% C, 4.56% H.
Preparation of 6,6-Difluorohex-1-ene (136)
6-Bromo-6,6-difluorohex-1-ene (135) (1.0 g, 5.02 x 10'3 mol) was dissolved in
mesitylene (1 mL) in a 10 ml_ round-bottomed flask equipped with a septum-capped
sidearm inlet and small stir bar. This was attached to an ice-water-cooled micro
distillation apparatus. Tributyltin hydride (1.6 g, 5.50 x 10'3 mol) was slowly injected into
the flask through the septum. When the addition was complete, heating was begun with
an oil bath. After 15 minutes at 50° C, the temperature was quickly raised and all volatile
material was flash distilled until the bath temperature reached 150° C. The distillate was
subjected to preparative GC separation, affording pure 136
6.6-Difluorohex-1-ene (136V 1H NMR: 8 1.57 (2H, m), 1.83 (2H, m), 2.12 (2H,
overlapping dt, J = 7 Hz), 5.00 (2H, m), 5.79 (1H, m), 5.81 (1H, tt, 3JHH = 4 Hz, 2JHF = 57
Hz); 13C NMR: 8 21.3, 32.9, 33.8 (t, 2JCF = 20.5 Hz), 115.4, 117.3 (t, 1JCF = 237.6 Hz),
120.5; 19F NMR: <|> -116.4 (dt, 3JFH = 14.6 Hz, 2JFH = 59.8 Hz); HRMS for C6H10F2: calc.
120.0751, found 120.0756; CHN for C6H10F2: calc. 59.98% C, 8.39% H, found 59.96% C,
8.47% H.
Preparation of 1.1-Difluoro-2-methvlcvclopentane (138)
Into a 50 mL three-necked round-bottomed flask equipped with magnetic stirrer
and septum was placed 0.5 g (5.09 x 103 mol) 2-methylcyclopentanone dissolved in 10

125
mL of anhydrous CH2CI2. DAST (0.9 g, 5.58 x 10‘3 mol) was then injected and the
mixture stirred at room temperature overnight. The reaction was dispensed onto ca. 2 g
of ice, the layers separated, and the organic layer washed with 1 mL of saturated
aqueous NaHC03. After drying, all volatile material was flash distilled and subjected to
preparative GC, affording pure 138
1,1-Difluoro-2-methvlcyclopentane (138): 1H NMR: 6 1.04 (3H, d, 3JHh = 7 Hz),
1.40 (1H, m), 1.73 (2H, m), 2.04 (4H, overlapping m); 13C NMR: 5 12.1, 19.9, 30.9, 34.4
(t, 2Jcf = 25.1 Hz), 40.7 (t, 2JCf = 23.5 Hz), 132.4 (t, 1JCF = 249.4 Hz); 19F NMR: (|) -100.2
(1F, d of overlapping dt, 3JFh = 12.2 Hz, 2JFf = 225.8 Hz), -107.6 (1F, d of overlapping dt,
3Jfh = 17.1 Hz, 2Jff = 224.6 Hz); HRMS forC6H10F2: calc. 120.0751, found 120.0748.
Preparation of 6-Bromohex-1-en-5-ol (142)
To a 500 mL three-necked round-bottomed flask equipped with magnetic stirrer
was added 100 mL acetic acid, 50 mL of saturated aqueous potassium bromide, and 50
mL THF. The flask was cooled to 0° C and 5.0 g (5.09 x 10'2 mol) of 1,2-epoxy-5-
hexene dissolved in 10 mL of THF was added dropwise with stirring. The
heterogeneous mixture was stirred at 0° C for an additional two hours, then allowed to
warm to room temperature and stir overnight. Most of the THF was removed by rotary
evaporation, 100 mL of ether and 50 mL of water was added, and the aqueous layer
washed with saturated aqueous NaHC03 until the acetic acid was removed. Drying over
MgS04 followed by rotary evaporation of the solvent afforded 8.03 g (88.1%) 142 which
was used in the next step without further purification.
6-Bromohex-1-en-5-ol (142): 1H NMR: 8 1.61 (2H, overlapping dt, J = 8 Hz), 2.15
(2H, m), 2.67 (1H, s), 3.35 (1H, m), 3.48 (1H, m), 3.76 (1H, m), 4.99 (2H, m), 5.77 (1H,
m); 13C NMR: 8 29.6, 34.0, 40.0, 70.2, 115.2, 137.5; HRMS for CeH^BrO: calc.
177.9993, calc (M+H) 178.9993, found 179.0058.

126
Preparation of 1-Bromo-5-hexen-2-one (143)
6-Bromohex-1-en-5-ol (142) (7.25 g, 4.05 x 10'2 mol) dissolved in 10 ml_ ether
was added dropwise to a mixture of 60 ml. of Jones’ reagent and 25 mL ether at room
temperature with stirring. After 4 hours the dark green reaction mixture was diluted with
50 mL of water. The layers were separated and the aqueous layer extracted with three
20 mL portions of ether. The combined organic extracts were washed twice with 20 mL
of saturated aqueous NaHC03 and once with 20 mL of water. Drying and rotary
evaporation of the solvent afforded 5.92 g (82.6%) 143 which was used without further
purification.
1-Bromo-5-hexen-2-one (143): 1H NMR: 6 2.35 (2H, overlapping dt, J = 6 Hz),
2.74 (2H, t, 3Jhh = 7 Hz), 3.88 (2H, s), 5.01 (2H, m), 5.78 (1H, m); 13C NMR: 5 27.7, 34.2,
38.8, 115.7, 136.3, 201.2; HRMS for C6H9BrO: calc. 175.9836, found 175.9850.
Preparation of 6-Bromo-5.5-difluorohex-1-ene (144)
A 100 mL three-necked round-bottomed flask equipped with an argon inlet,
rubber septum, and magnetic stirrer was charged with 2.1 g (1.19 x 102 mol) 1-bromo-5-
hexen-2-one (143) in 20 mL of anhydrous CH2CI2. The flask was cooled to 0° C and 1.9
mL (2.32 g, 1.44 x 102 mol) of diethylaminosulfurtrifluoride was slowly injected into the
reaction mixture with stirring. After two hours at 0° C, the flask was allowed to warm to
room temperature and stirring continued for an additional 48 hours. The contents were
carefully dispensed onto 20 g of ice, the layers separated, and the aqueous layer
extracted twice with 5 mL CH2CI2. The combined organic extracts were washed once
with 10 mL of saturated aqueous NaHC03 and once with 10 mL of water. After drying
over MgS04 the solution was carefully concentrated via gentle ambient pressure
distillation. A total of 1.33 g (56.2%) 144 was obtained as a colorless liquid, bp
117-119° C, which was further purified by preparative GC for spectroscopic analysis and
kinetic experiments.

127
6-Bromo-5,5-difluorohex-1-ene (144): 1H NMR: 8 2.06 - 2.31 (4H, m), 3.53 (2H, t,
3Jhf = 13 Hz), 5.07 (2H, m), 5.82 (1H, m); 13C NMR: 8 26.2 (t, 3JCF = 4.5 Hz), 31.3 (t, 2JCF
= 33.6 Hz), 33.8 (t, 2JCF = 24.1 Hz), 115.8, 121.1 (t, 1JCF = 241.4 Hz), 136.2; 19F NMR: <|>
-99.3 (m); HRMS forC6H9F2Br: calc. 197.9855, found 197.9850; CHN for C6H9F2Br: calc.
36.21% C, 4.56% H, found 36.16% C, 4.57% H.
Preparation of 5,5-Difluorohex-1-ene (144)
6-Bromo-5,5-difluoro-1-hexene (144) (1.0 g, 5.02 x 10'3 mol) was treated with
1.6 g (5.50 x 10'3 mol) of tributyltin hydride in a manner identical to the independent
preparation of 136. Flash distillation followed by preparative GC separation afforded
pure 144.
5,5-Difluorohex-1-ene (144): 1H NMR: 8 1.60 (3H, t, 3JHF = 18 Hz), 1.94 (2H, m),
2.24 (2H, m), 5.03 (2H, m), 5.83 (2H, m); 13C NMR: 8 23.3 (t, 2JCF = 28.1 Hz), 26.9 (t,
3Jcf = 5.0 Hz), 37.2 (t, 2Jcf = 25.1 Hz), 115.2, 123.9 (t, 1JCF = 236.4 Hz), 136.9; 19F NMR:
<|> -91.3 (m); HRMS for C6H10F2: calc. 120.0751, found 120.0743.
Preparation of 1,1-Difluoro-3-methylcvclopentane (147)
3-Methylcyclopentanone (0.5 g, 5.09 x 103 mol) and 0.9 g (5.58 x 10‘3 mol)
DAST were reacted in 10 mL of anhydrous CH2CI2 in a manner identical to the
preparation of 138 Flash distillation and preparative GC separation afforded pure 147.
1,1-Difluoro-3-methylcvclopentane (147): 1H NMR: 8 1.05 (3H, d, 3JHh = 6Hz),
1.36 (1H, m), 1.61 (1H, m), 1.87 - 2.31 (5H, m); 13C NMR: 8 20.0, 31.6, 32.0 (t, 3JCF = 4.3
Hz), 36.0 (t, 2Jcf = 25.0 Hz), 44.0 (t, 2JCF = 23.6 Hz), 133.0 (t, 1JCF = 246.9 Hz); 19F NMR:
<(> -88.9 (1F, dm, 2JFF = 217.1 Hz), -90.2 (1F, dm, 2JFF = 227.1 Hz); HRMS for C6H10F2:
calc. 120.0751, found 120.0759.

128
Preparation of 1,1-Difluorocvclohexane (149)
Cyclohexanone (0.5 g, 5.09 x 10'3 mol) and 0.9 g (5.58 x 10‘3 mol) DAST were
reacted in 10 mL of anhydrous CH2CI2 in a manner identical to the preparation of 138.
Preparation and flash distillation afforded pure 149.
1,1 -Difluorocvclohexane (149): 1H NMR: 8 0.97 (2H, m), 1.29 (4H, q, 3JHH = 6
Hz), 1.58 (4H, m); 13C NMR: 5 22.8, 24.4, 34.1 (t, 2JCF = 23.5 Hz), 123.6 (t, 1JCF = 239.9
Hz); 19F NMR: <|> -95.7 (2F, br s).
Preparation of 5,5,6,6-Tetrafluorohex-1-ene (153), 1,1,2,2-Tetrafluoro-3-methylcvclo-
pentane (155), and 1,1,2,2-Tetrafluorocvclohexane (157)
6-Bromo-5,5,6,6-tetrafluorohex-1-ene (101) (5.0 g, 2.13 x 10‘2 mol) dissolved in 5
mL of mesitylene was added to a 50 mL three-necked round-bottomed flask equipped
with ice water condenser, argon inlet, magnetic stir bar and rubber septum. Tributyltin
hydride (7.5 g, 2.58 x 10'2 mol) and 0.05 g (3.04 x 104 mol) 2,2'-azobisisobutyronitrile
(AIBN) in 5 mL of mesitylene were taken up into a syringe. The flask was heated at
50° C and irradiated with a 150 W flood lamp placed at a distance of ca. 1 m while the
nBu3SnH solution was delivered to the reaction mixture, via syringe pump, over a 24
hour period. After the addition was complete, all volatile material was flash distilled from
the reaction mixture until the bath temperature reached 150° C. Purification by
preparative GC afforded pure samples of 153, 155, and 157.
5.5,6.6-Tetrafluorohex-1-ene (1531: 1H NMR. 5 2.06 (2H, m), 2.33 (2H, m), 5.08
(2H, m), 5.72 (1H, tt, 3JHF = 3 Hz, 2JHF = 54 Hz), 5.84 (1H, m); 13C NMR: 8 24.6 (t, 3JCF =
4.0 Hz), 29.2 (t, 2Jcf = 22.1 Hz), 110.3 (tt, 2JCF = 41.1 Hz, 1JCF = 247.7 Hz), 115.9, 117.8
(tt, 2Jcf = 29.0 Hz, 1JCF = 244.8 Hz), 136.1; 19F NMR: -116.7 (2F, t, 3JFH = 17.1 Hz),
-136.0 (2F, d, 2Jfh = 56.2 Hz); HRMS for C6H8F4: calc. 156.0562, found 156.0562.
1,1,2,2-Tetrafluoro-3-methylcvclopentane (155): 1H NMR 8 1.12 (3H, d, 3JHH = 7
Hz), 1.47 (1H, m), 2.00 (1H, m), 2.07 - 2.49 (3H, m); 13C NMR: 8 11.4, 23.7 (m), 29.8 (t,
2JCF = 22.8 Hz), 36.2 (t, 2JCF = 21.0 Hz), 117.6 - 125.8 (2C, m); 19F NMR: <\> -110.1 (1F,

129
dm, 2Jff = 234.4 Hz), -120.7 (1F, dm, 2JFF = 239.3 Hz), -126.0 (1F, dt, 3JFH = 12.2 Hz,
2Jff = 236.8 Hz), -132.9 (1F, dm, 2JFF = 235.6); HRMS for C6H8F4: calc. 156.0562, found
156.0563; CHN for C6H8F4: calc. 46.16% C, 5.16% H, found 46.17% C, 5.35% H.
1,1,2,2-Tetrafluorocvclohexane (157): 1H NMR: 6 1.69 (4H, br s), 2.06 (4H, br s);
13C NMR: 8 21.0, 31.7 (t, 2JCF = 22.1 Hz), 117.0 (tt, 2JCF = 28.1 Hz, 1JCF = 250.4 Hz); 19F
NMR: -119.7 (4F, br s); HRMS for C6H8F4: calc. 156.0562, found 156.0571.
Competitive Kinetic Procedures
Competition Kinetics: Hydrogen Atom Abstraction (kH) versus Addition (kaM) and
Hydrogen Atom Abstraction versus Cyclization (km)- General Procedure.
Into each of a set of six Pyrex NMR tubes were added a known amount of C6D6,
varying, known amounts of trapping agent or trapping agent and styrene, and a known
amount of trifluorotoluene as an internal 19F NMR standard. Each tube was sealed with
rubber septa secured with PTFE tape, frozen in a dry ice-isopropanol slush, and
subjected to three successive freeze-pump-thaw cycles followed by pressurization with
argon. Into each frozen tube was then injected a known amount of the radical precursor
(in the case of 119, a stock solution in degassed C6D6) followed by warming to room
temperature (in the case of 119, the tubes were flame-sealed before warming) with
vigorous shaking. The tubes were then subjected to UV photolysis in a Rayonet reactor
(254 nm lamps) until complete consumption of starting material was demonstrated by19F
NMR analysis. Product ratios for varied concentrations of trapping agent or ratios of
trapping agent to styrene allow determination of the ratios kH / kCn or kH / kaM. Yields are
determined by integration of product resonances versus that of internal standard (
-63.24) in the 19F NMR.
1,1-Difluorohex-1-vl Radical (77)
Ratios of [81] / [83] were determined by integration of the -CF2H and -CF2-
resonances (at <|> -116.0 and -99.1, respectively) in the 19F NMR.

130
Table 5-1. Data for Competitive Determination of kH (nBu3SnH) / kaM (Styrene) for 1,1-
Difluorohex-1-yl Radical (77).
L80J
Í CrH,CH=CH, 1
r nBu,SnH 1 / fCfiH.CH=CH, 1
f 81 1/i 83 1
% Yield
0.094
2.01
0.719
2.30
95
0.094
1.81
0.847
2.73
95
0.094
1.61
1.01
3.26
96
0.094
1.41
1.21
3.93
97
0.094
1.21
1.49
4.88
96
0.094
1.01
1.87
6.21
95
Slope = kH / kaM = 3.39 ± 0.02
kH (nBu3SnH) = 9.1 (± 1.7) x 106 M'1 s'1
2,2-Difluorohex-1-vl Radical (88)
Ratios of [93] / [98] were obtained by integration of the respective -CF2-
resonances at Table 5-2. Data for Competitive Determination of kH (nBu3SnH) / kaM (Styrene) for 2,2-
Difluorohex-1-yl Radical (88).
I92J
Í CfiH.CH=CH, 1
[ nBihSnH 1 / ÍCRI-UCH=CH, 1
Í 931 / f 98 1
% Yield
0.087
4.08
0.223
5.96
96
0.087
3.97
0.243
6.62
98
0.087
3.85
0.263
7.16
97
0.087
3.73
0.285
7.79
98
0.087
3.62
0.308
8.33
97
0.087
3.50
0.333
9.01
100
Slope = kH / kaM = 27.3 + 0.6
kH (nBu3SnH) = 1.4 (±0.5) x 107 M'1 s'1

131
1.1.2.2-Tetrafluorohex-1-vl Radical (100)
Ratios of [108] / [110] were obtained by the sum of integrals of the -CF2- and
-CF2H resonances of 108 (at § -117.1 and -136.1, respectively) versus that of the two
-CF2- resonances of 110 (at <|) -116.0 and -116.4, respectively).
Table 5-3. Data for Competitive Determination of kH (nBu3SnH) / kadd (Styrene) for 1,1-
2.2-Tetrafluorohex-1-yl Radical (100).
L102J
[ CrH,CH=CH, 1
f nBu.SnH 1 / rCfiH.CH=CH, 1
r 1081 /r 1101
% Yie
0.085
1.86
0.506
2.33
98
0.085
1.68
0.611
2.90
97
0.085
1.49
0.741
3.56
97
0.085
1.30
0.909
4.30
97
0.085
1.12
1.13
5.34
97
0.085
0.931
1.45
6.65
99
Slope = kH / kaM = 4.56 ±0.10
(nBu3SnH) = 9.2 (± 0.8) x 107 M"1
s'1
Table 5-4. Data for Competitive Determination of kH ((TMS)3SiH) / /cadd (Styrene) for 1,1-
2,2-Tetrafluorohex-1-yl Radical (100).
U02j
Í CfiH.CH=CH, 1
i (TMS)-íSiH 1 / fCfiFUCH=CHo 1
r 1081/rnoi
% Yield
0.074
1.29
0.988
1.10
90
0.074
1.19
1.11
1.20
92
0.074
1.08
1.26
1.32
93
0.074
0.970
1.44
1.52
94
0.074
0.863
1.67
1.71
92
0.074
0.755
1.96
1.98
95
Slope = kH / kaM = 0.913 ± 0.017.
kH ((TMS)3SiH) = 1.8 (±0.1) x 107 M'1 s1

132
2-fPerfluorohexvneth-1-vl Radical (111)
Ratios of [113] / [117] were determined by integration of the respective -CF2-
resonances at <)> -117.1 and -115.0.
Table 5-5. Data for Competitive Determination of kH (nBu3SnH) / kadd (Styrene) for 2-
[Perfluorohexyl]eth-1-yl Radical (111).
L112J
F CRFUCH=CH, 1
F nBu.SnH 1 / FCRHRCH=CH, 1
r 1131 / r 117 l
% Yield
0.064
2.96
0.217
3.11
100
0.064
2.83
0.248
3.58
100
0.064
2.69
0.283
4.13
100
0.064
2.56
0.320
4.72
100
0.064
2.42
0.362
5.44
100
0.064
2.29
0.408
6.14
100
Slope = kH / kaM = 16.0 ± 0.1
kH (nBu3SnH) = 2.1 (± 0.3) x 106 M 1 s1
Pentafluoroethvl Radical (118)
Ratios of 120 / 122 were determined the sum of integrals of the CF3- and -CF2H
resonances of 120 (at -86.6 and -138.6, respectively) versus that of the CF3- and -CF2-
resonances of 122 (at -86.1 and -119.2, respectively).
Table 5-6. Data for Competitive Determination of kH (nBu3SnH) / kaM (Styrene) for
Pentafluoroethyl Radical (118).
f 119]
r CRFUCH=CH, l
r nBu.SnH 1 / rCRHRCH=CH, 1
F1201 / r 122 1
% Yield
0.082
2.20
0.434
1.15
89
0.082
2.08
0.485
1.31
85
0.082
1.96
0.542
1.46
85
0.082
1.84
0.608
1.62
82
0.082
1.71
0.682
1.79
83

133
Table 5-6-- continued
f 1191 f CfiHsCH=CH? 1 f nBu.SnH 1 / rCfiH,CH=CH, 1 f 1201/ f 1221 % Yield
0.082 1.59 0.768 2.05 83
Slope = kH / /(add = 2.62 ± 0.06
kH (nBu3SnH) = 3.2 (± 0.3) x 108 M'1 s'1
1,1-Difluorohex-5-en-1-yl Radical (128)
Ratios of 136 to 138 were determined by integration of the -CF2H resonance of
136 (at <)) -116.2) versus the sum of integrals for the diastereotopic -CF2- resonances of
138 (at <)) -100.3 and -107.8) in the 19F NMR.
Table 5-7. Data for Competitive Determination of /cH (nBu3SnH) / kC5 for 1,1-Difluorohex-
5-en-1-yl Radical (128).
[1351
Í nBu3SnH 1
r1361/r138 1
% Yield
0.054
0.673
1.53
88
0.054
0.807
1.91
100
0.054
0.942
2.28
89
0.054
1.08
2.57
94
0.054
1.21
2.93
95
0.054
1.35
3.29
92
Slope = kH / kC5 = 2.57 ± 0.05.
/(es = 3.5 (± 0.6) x 106 s1
2,2-Difluorohex-5-en-1-yl Radical (140)
Ratios of 145 to 147 and 145 to 149 were determined by integration of the
-CH2CF2CH3 resonance of 145 (at <)> -91.5) versus the sum of integrals for the
diastereotopic -CF2- resonances of 147 (at <(> -88.9 and -90.4) and the -CF2- resonance
of 149 (at
134
Table 5-8. Data for Competitive Determination of kH (nBu3SnH) / kc5 and /cH (nBu3SnH) /
/(c6for2,2-Difluorohex-5-en-1-yl Radical (140).
L144J
í nBu3SnH l
Í1451/Í1471
r1451/r1491
% Yield
0.053
0.286
2.64
24.3
98
0.053
0.343
3.40
30.9
100
0.053
0.400
4.03
37.8
100
0.053
0.457
4.71
48.2
99
0.053
0.514
5.44
54.4
98
0.053
0.572
6.26
61.1
98
For C5:
Slope = kH / kcs =
12.510.2.
kC5 = 1.1 (± 0.38) x 106 s'1
For Cf
3: Slope = kH / kce =
= 13215.
kc6= 11 (± 0.34) x 10s s1
1,1,2,2-Tetrafluorohex-5-en-1-vl Radical (152)
Ratios of 153 to 155 and 153 to 157 were determined by the sum of integrals of
the -CF2- and -CF2H resonances of 153 (at <|> -116.7 and -135.9, respectively) versus the
sum of integrals for the diastereotopic -CF2- resonances of 155 (at <)> -109.9, -120.8,
-126.0, and -132.9) and the -CF2-
resonance of 157 (at <|> -119.6, for kC6) in the 19F NMR.
Table 5-9. Data for Competitive Determination of kH ((TMS)3SiH) / kC5 and kH
((TMS)3SiH) / /rC6for 1,1,2,2-Tetrafluorohex-5-en-1-yl Radical (152).
[101 i
f (TMS),SiH 1
f1531/Í1551
f1531/Í1571
% Yield
0.052
0.581
0.988
4.35
96
0.052
0.686
1.19
5.28
95
0.052
0.792
1.44
6.39
94
0.052
0.897
1.67
7.36
95
0.052
1.00
1.84
8.21
98
0.052
1.11
2.11
9.37
99

135
For C5: Slope = kH / kC5 = 2.11 ± 0.05.
/fC5 = 8.7(± 0.41) x106 s'1
For C6: Slope = kH / kc& = 9.44 ±0.14.
kce= 1.9 (± 0.11) x 106 s1

APPENDIX A
SELECTED 19F NMR SPECTRA
The 19F NMR spectra of radical precursors, a representative sample from each
kinetic run, and isolated products of addition, hydrogen abstraction, and cyclization are
graphically illustrated in this appendix. Full spectral characterization data are presented
numerically in their respective areas of Chapter 5.
136

Figure A-1.
1
9 -S -19 -IS -2® -25
1-Bromo-1,1-difIuorohexane (80).
to
-o

34.6]
Figure A-2. Data Point 1 of kH / kaM Determination for 1,1-Difluorohex-1-yl Radical (77).

Figure A-3. 1
,1-Difluorohexane (81).
SI'828?C â– 
co
CD
-100
-II#
PP«

-98.7 -98.8 -98.9 -99.9 -99-1
-!•
» -39 <9 -59
Figure A-4. 3,3-Difluoro-1-phenyloctane (83)
A

B
Figure A-5. 1-lodo-2,2-difluorohexane (92)

I
~91
-93
-94
I
-97
ppm
-9U
1
-95
-96
~~l
-98
T'~
-99
Figure A-6. Data Point 1 of kH / kaM Determination for 2,2-Difluorohex-1-yl Radical (88).

Figure A-7. 2,2-Difluorohexane (93).

I
-10
T-
T 1 1 1 1 1—
-100 ppm
Figure A-8. 4,4-Difluoro-1-pheny!octane (98).
144

t—'—■—'—»—i—'—1—■—'—i—*—'—'—'—r~
-M -3* -«• -S«
Figure A-9. 1-Bromo-1,1,2,2-tetrafluorohexane (102)
4^
cn
a.u
8.M

Figure A-10. Data Point 1 of kH I kaM Determination for 1,1,2,2-Tetrafluorohex-1-yl Radical (100)
146

Figure A-11.
V S-*
12.48
12.65
1,1,2,2-Tetrafluorohexane (108).
147

i
-10
T
-30
—I i » ' r | . . , . 1 1 . > 1 | 1 r
-40 -50 -60 -70
-80
i
-90
■*—I—1
-100
i—I—«
-110
Figure A-12. 3,3,4,4-Tetrafluoro-1-phenyloctane (110).

Figure A-13. 1,1,2,2-Tetrafluoro-1-iodobutane (107).
149

JL
• 7"
-80
- r -i
-85
-90
-95
-100
Figure A-14. l-lodo-2-[perfluorohexyl]ethane (112).
cn
o
9.Si
9.46 9.SO
9.S2
9.39

*—T—' 1 | 1 1 »'*■» * »■* W*T-J
5.94 7.8b 7.73
1.91 7.01 7.92
Figure A-15. Data Point 1 of kHl kaaa Determination for 2-[Perfluorohexyl]eth-1-yl Radical (111)

-80
-8b
11 .4411 45
11.46
en
N)
Figure A-16. 1-[Perfluorohexyl]ethane (113).

8.90 6.89
9.01
en
co
Figure A-17. 1-[Perfluorohexyl]-4-phenylbutane (117)

Figure A-18 lodopentafluoroethane (119).

-¿5 -¿O
—I—■-
-100
-105
—I—
-110
—I—■“
-115
-120
V
S .03
I r
-125
l—'—'—'—'—I—*"
-130 -135
ppm
Figure A-19. Data Point 1 of kH / kaM Determination for Pentafluoroethyl Radical (118).
155

0.47
s.ss
Figure A-20. Pentafluoroethane (120).

Figure A-21. 1,1,1,2,2-Pentafluoro-4-phenylbutane (122).
en
-vi
—i—
-105
-110
-115
1
-120
ppm
4.95

-4365 -43.?• -43.7S
Figure A-22. 6-Bromo-6,6-difluorohex-1-ene (135)

I».M ^ UK
II.M
Figure A-23. Data Point 1 of kHl kc Determination for 1,1-Difluorohex-5-en-1-yl Radical (128).
159

Figure A-24. 6,6-Difluorohex-1-ene (136).

Figure A-25. 1
*
l!
ti
/
J
23.23
, 1 -D¡fluoro-2-methylcyclopentane (138).
2 2
ii
\¡
i
U
en
23.4!

-9Í.(
-99'. 1
-98.9
-99~¡
-99.3
J
1 1
-10
-30
-40 '
- SO
Figure A-26. 6-Bromo-5,5-difluorohex-1-ene (144)
-28028.20
28050.17

en
co
Figure A-27. Data Point 1 of kHl kc Determination for 2,2-Difluorohex-5-en-1-yl Radical (140).

Figure A-28 5,5-Difluorohex-1-ene (145).
164

Figure A-29. 1,1-Difluoro-3-methylcyclopentane (147).
--24959.26
--24973.91
--24988.56
13
-25540.33
/—-25550.10
/ -25557.42
'—-25562.30
^--25572.07
W-25579.39
-25586.72
-25594.04
S9U

Figure A-30. 1,1-Difluorocyclohexane (149).
166

i
Figure A-31. 6-Bromo-5,5,6,6-tetrafluorohex-1-ene (101)
411
4.«4

Jj
n pnr,.,..
-110 -112 -114
4.4$
T-m-rr
-116
4.08
"r"M ’'
—118
XT-» v. p in f t
-120 -122
4.11
4.44
JU
i . r mi , . rrrp. r. .f i .
-124 -126 -128 -130
4.S4
I I | I I • I , I I I • | IT
-132 -134
4.S8
-136 ppm
8.81
0.17
Figure A-32. Data Point 1 of kHl kc Determination for 1,1,2,2-Tetrafluorohex-5-en-1-yl Radical (152).

IET
-40
-200
-80
-100
'tt r t
-120
-140
-160
-180
ppm
*
0.33
V
13.90
Figure A-33. 5,5,6,6-Tetrafluorohex-1-ene (153).

TrT~l ' T
T-n . . . . | n , . , , ,
-110 -112
-114
~n rT
-116
rrr7
-118
1 1 I ' ' ' 1 1 ; ' ' ' I 1 ' 1 1 » » 1 1 1 I 1 » 1 1 I ■ ' » '» | 1 « ' ' T « » ' . » | » >
-120 -122 -124 -126 -128
Tf-j-r-r
-130
17.19
17.15
17.14
Figure A-34. 1,1,2,2-Tetrafluoro-3-methylcyclopentane (155)

Figure A-35 1,1,2,2-Tetrafluorocyclohexane (157).

APPENDIX B
B3LYP/6-31G(d) TOTAL AND ZERO-POINT ENERGIES
FOR DATA IN TABLES 3-3 AND 3-4
The total electronic and unsealed zero-point energies for hydrofluorocarbon
radicals and closed shell species used in the determination of the bond dissociation
energies reported in Tables 3-3 and 3-4 are provided in this appendix. In instances
where more than one conformer is provided, that of lowest energy was used in the C-C
and C-H BDE computation.
Table B-1. Total and (Unsealed) Zero-Point Energies for Fluorinated Alkanes.
Species
E(B3LYP/6-31G(d)), a.u.
ZPE. a.u.
CH3CH3
-79.8304 167
0.075231
CH3CF3
-377.5549 235
0.052872
CH3CF2H
-278.3015 940
0.061076
cf3cf2h
-576.0077 790
0.037905
ch3ch2ch3
-119.1442 464
0.104110
ch3cf2ch3
-317.6263 447
0.088479
cf3ch2ch3
-416.8699 039
0.081649
cf3cf2ch3
-615.3356 199
0.065266
ch3ch2cf2h 3
-317.6163 135
0.089851
ch3ch2cf2h 6
-317.6162 650
0.090006
ch3cf2cf2h b
-516.0884 655
0.073687
ch3cf2cf2h 3
-516.0853 571
0.073705
CH3CH2CH2CH3c
-158.4580 400
0.132860
172

173
Table B-1-- continued
Species
EÍB3LYP/6-31 G(d)1. a.u.
ZPE. a.u.
ch3ch2ch2ciV
-158.4567 065
0.132957
CH3CH2CF2CH3c
-356.9408 361
0.117249
ch3ch2cf2ciV
-356.9401 189
0.117349
CH3CF2CF2CH3c
-555.4151 664
0.101158
CH3CF2CF2CH3d
-555.4110 960
0.101141
CF3CH2CH2CH3c
-456.1836 478
0.110224
CF3CH2CH2CH3d
-456.1826 866
0.110445
3 Methyl and Hydrogen Gauche. b Methyl and Hydrogen Anti. c Methyl (or Methyl and
Trifluoromethyl) Groups Anti. d Methyl (or Methyl and Trifluoromethyl) Groups Gauche
Table B-2. Total and (Unsealed) Zero-Point Energies for Fluorinated Radicals.
Species
E(B3LYP/6-31G(d)1. a.u.
ZPE. a.u.
H
-0.5002 728
0
ch3
-39.8382 909
0.029849
cf3
-337.5509 879
0.012153
ch3ch2
-79.1578 663
0.059647
ch3cf2
-277.6352 216
0.047481
cf3ch2
-376.8760 847
0.037830
cf3cf2
-575.3382 007
0.024522
ch3ch2ch2 3
-118.4713 699
0.088731
ch3ch2ch2 b
-118.4711 179
0.088945
ch3ch2cf2 b
-316.9496 940
0.076574
ch3ch2cf2 3
-316.9496 712
0.076368
CH3CF2CH2 c
-316.9493 447
0.073435
ch3cf2ch2 b
-316.9489 182
0.073225

174
Table B-2-- continued
Species
E(B3LYP/6-31 GidIV a.u.
ZPE. a.u.
CH3CF2CF2
b
-515.4178 170
0.060271
CH3CF2CF2
c
-515.4173 201
0.060276
CF3CH2CH2
a
-416.1952 103
0.066154
CF3CH2CH2
b
-416.1951 213
0.066242
a Radical p Orbital Aligned with p C-H Bond. B p C-C Bond Alignment. c p C-F bond
alignment.

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BIOGRAPHICAL SKETCH
Michael David Bartberger was born September 6, 1970, in Ft. Lauderdale,
Florida, the first of two children, and spent his childhood years in the Margate / North
Lauderdale area.
Michael received a B.S. degree in chemistry from the University of Central
Florida in May of 1992, during which time he was introduced to organic fluorine
chemistry by University of Florida alumnus and former Dolbier group member, Professor
Seth Elsheimer.
In August of 1992, Michael began graduate study at the University of Florida,
having practically joined the research group of Professor William R. Dolbier, Jr. even
before moving to Gainesville. In addition to experimental physical organic studies,
Michael developed an interest in theoretical methodology, having been exposed to the
results of AM1 calculations attractively displayed on a Tektronix CAChe workstation.
With a few days of instruction from Dr. Max Muir (a postdoc from another group working
on some semiempirical calculations as a favor to the Dolbier group, leaving UF shortly
thereafter for a position at MSI) Michael set out to hone his own computational chemistry
skills. Spending many late nights at the workstation, he developed his proficiency with
the MOPAC, ZINDO, Gaussian, ACES II, and GAMESS program systems, leading to a
number of productive collaborations with other members of the department, both within
and outside the Dolbier group. The highlight of his graduate career came in August of
1996, as an invited participant in the Elucidation of Reaction Mechanisms by Ab Initio
Methods Symposium held at the 212th National Meeting of the American Chemical
Society in Orlando, Florida.
185

186
Michael is a recipient of the Department of Chemistry Excellence in Teaching
Award, the Shell Foundation Fellowship, and has twice received the Outstanding
Presentation Award administered by the Florida Section of the American Chemical
Society, once as an undergraduate.
At the time of this writing, Michael had accepted a position as a postdoctoral
research associate in the Department of Chemistry and Biochemistry at the University of
California, Los Angeles, under the direction of Professor K. N. Houk.

I certify that I have read this study and that in my opinion it conforms to
acceptable standards of scholarly presentation and is fully adequate, in scope and
quality, as a dissertation for the degree of Doctor of Philosophy.
William R. Dolbier, Jr., Chair
Professor of Chemistry
I certify that I have read this study and that in my opinion it conforms to
acceptable standards of scholarly presentation and is fully adequate, in scope and
quality, as a dissertation for the degree of Doctor of Philosophy.
O- 1
Merle A. Battiste
Professor of Chemistry
I certify that I have read this study and that in my opinion it conforms to
acceptable standards of scholarly presentation and is fully adequate, in scope and
quality, as a dissertation for the degree of Doctor of Philosophy.
^ S.
Kirk S. Schanze
Professor of Chemistry
I certify that I have read this study and that in my opinion it conforms to
acceptable standards of scholarly presentation and is fully adequate, in scope and
quality, as a dissertation for the degree of Doctor of Philosophy.
Graduate Research Professor of
Chemistry
I certify that I have read this study and that in my opinion it conforms to
acceptable standards of scholarly presentation and is fully adequate, in scope and
quality, as a dissertation for the degree of Doctor of Philosophy.
J^Mri R. Sabin
Professor of Physics

This dissertation was submitted to the Graduate Faculty of the Department of
Chemistry in the College of Liberal Arts and Sciences and to the Graduate School and
was accepted as partial fulfillment of the requirements for the degree of Doctor of
Philosophy.
May, 1998
Dean, Graduate School

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