Distance dependence of boron on bond dissociation energies and electron affinities of simple alkylboranes

Material Information

Distance dependence of boron on bond dissociation energies and electron affinities of simple alkylboranes a computational approach
Custer, Thomas Corey
Publication Date:
Physical Description:
viii, 45 leaves : ; 28 cm

Thesis/Dissertation Information

Master's ( Master of science)
Degree Grantor:
University of Colorado Denver
Degree Divisions:
Department of Chemistry, CU Denver
Degree Disciplines:


Subjects / Keywords:
Boron ( lcsh )
Chemical affinity ( lcsh )
Scission (Chemistry) ( lcsh )
Boron ( fast )
Chemical affinity ( fast )
Scission (Chemistry) ( fast )
bibliography ( marcgt )
theses ( marcgt )
non-fiction ( marcgt )


Includes bibliographical references (leaves 41-45).
Department of Chemistry
Statement of Responsibility:
by Thomas Corey Custer.

Record Information

Source Institution:
|University of Colorado Denver
Holding Location:
Auraria Library
Rights Management:
All applicable rights reserved by the source institution and holding location.
Resource Identifier:
74818053 ( OCLC )
LD1193.L46 2006m C87 ( lcc )

Full Text
Distance Dependence of Boron on Bond Dissociation Energies and Electron
Affinities of Simple Alkylboranes: A Computational Approach
Thomas Corey Custer
B.A., Monmouth College, 2002
A thesis submitted to the
University of Colorado at Denver
in partial fulfillment of
the requirements for the degree of
Master of Science
Chemistry Department

Thesis for Master of Science
degree by
Thomas Corey Custer
has been approved
Hai Lin
S i ijwok

Custer, Thomas Corey (M.S. Chemistry)
Distance Dependence of Boron on Bond Dissociation Energies and Electron
Affinities of Simple Alkylboranes: A Computational Approach
Thesis directed by Professor Robert Damrauer
The distance dependence of boron on the electron affinity (EA) of alkylborane
radicals (CH2(CH2)nBH2 as n = 0 to 4) and the bond dissociation energy
(BDE) of a terminal C-H group on simple alkylboranes (H-CH2(CH2)nBH2 as
n = 0 to 4) were calculated using computational methods. The computed EAs
of BH3 and CH3 are compared to their experimentally determined values to
ensure correct computational methods were chosen. In addition, the computed
C-H and B-H BDEs of (CH3)3B, BH3, and CH4 are compared to the
experimentally determined values. Anions of CH2(CH2)nBH2 where n = 1 to 4
are ring-closed structures at their equilibrium geometries, where boron
becomes tetracoordinate. The large geometry change is accompanied by a
large EA; the greatest EA value is observed for n = 4. The presence of strong
hyperconjugative effects and delocalization of the radical are responsible for
surprisingly low BDEs of all alkylboranes; the smallest of which is observed
for n: 0.
This abstract accurately represents the content of the candidate
recommend its publication.
Robert Damrauer

Never have I felt the touch of God like I had through my most recent events.
Each persons experience with God is unique; however it is our responsibility
to never let it go unnoticed. Therefore, I dedicate this thesis to God, for
without Him none of this would have been possible. In addition, I must
acknowledge the tremendous people that God has placed in my life that
pushed and prayed for me in the difficult times and commended me in the
good times. A special thanks to my parents Tom and Becky Custer for their
undying love and support in every one of my endeavors. For a long time, you
were the only ones who believed in me.

Thanks to my advisor Robert Damrauer for his contributions and support on
my research. Thanks to my committee members for taking time from their
busy schedules to participate in the examination and give input on my

Tables..................................................... viii
1. Introduction................................................1
2. Computational Methods......................................11
3. Results and Discussion.....................................13
3.1 BH2CH2 and BH2CH3.........................................15
3.2 BH2CH2CH2 and BH2CH2CH3...................................22
3.3 BH2CH2CH2CH3 and BH2CH2CH2CH2.............................28
3.4 BH2CH2CH2CH2CH3 and BH2CH2CH2CH2CH2..................... 34
4. Conclusion.................................................40

FIGURE 1: BH2CH3 NEUTRAL...........................20
FIGURE 2: BH2CH2 RADICAL...........................20
FIGURE 3: BH2CH2 ANION.............................20
FIGURE 5: BH2CH2CH2 ANION..........................23
FIGURE 6: BH2CH2CH3 NEUTRAL....................... 24
FIGURE 7: BH2CH2CH2 RADICAL........................27
FIGURE 8: BH2CH2CH2CH3 NEUTRAL.....................30
FIGURE 9: BH2CH2CH2CH2 RADICAL.....................31
FIGURE 10: BH2CH2CH2CH2 ANION......................32
FIGURE 11: BH2CH2CH2CH2CH2 ANION...................35
FIGURE 12: BH2CH2CH2CH2 RADICAL....................37
FIGURE 13: BH2CH2CH2CH2CH3 NEUTRAL.................39

Table I: Experimental Bond Dissociation Energies, Gas Phase Acidities, and
Electron Affinities of Simple Alkanes........................4
Table II: Experimental Electron Affinities............................4
Table III: Computed BDE of Boron and Carbon Species (kcal/mol)........8
Table IVa: Experimental Bond Dissociation Energies of Boranes.........9
Table V: Computed Adiabatic Electron Affinities and Vertical Attachment
Table VI: Computed Bond Dissociation Energies.........................18

1. Introduction:
Since the ease of synthesizing1-4 and purchasing organoboron reagents has
increased, organoboron reagents have become practical choices for both large and
small scale synthetic procedures.1 The dichotomy of borons electropositivity
(relative to other 1st row elements) and yet wanting to complete its octet is what
contributes to organoborons schizophrenic tendencies and yet versatile uses.
Organoboron reagents are used in hydroborations of sp2 and sp hybridized carbons,
haloborations of terminal alkynes, and reductions of ketones and aldehydes, to
name a few.1'5
The ease of such synthetic procedures affords opportunities to make
organoborons with a wide range of functionality. Therefore, organoborane reagents
for asymmetric reductions1, Suzuki-Miyaura couplings,6"9 sp3 sp3 carbon
couplings,10 introduction of functional groups,11,12 asymmetric radical
additions,13'17intramolecular cyclization,18 elimination,19,20 radical initiators,14 and
many other reactions are now commonplace. Many of said synthetic techniques
have supplanted other chemical processes due to their functionality,5,10
stereospecificity,22 mild reaction conditions,23,24 and high yields.25 Understanding

these organoboranes as synthetic intermediates, in order to better select reagents
and reaction conditions to attain desired products, is an ongoing research area.
Gas phase ion studies are an excellent tool to measure thermodynamic
properties of molecules in the absence of solvent.26'30 Thermodynamic parameters
can then be utilized to better understand transition states and thus products observed
n i in
in various chemical reactions. Gas phase acidity (Equation 5) is an example,
where it may be used to elucidate anion stabilities, thus giving insight in predicting
reaction propensities of a molecule within a given class.31,33 Gas phase acidities
can be difficult to measure directly and quantitatively since the proton of interest
might not be acidic enough to extract exclusively with even the most basic of
species.33,34 Gas phase acidities can also be calculated indirectly by utilizing
thermodynamic cycles (equation 26-31>33>35
AHacid(RH) = D[R-H] EA(R) + IP(H) (equation 1)
Here D[R-H] is the homolytic bond dissociation energy (BDE) (equation 2) of the
acidic group of interest, EA(R) is the electron affinity (equation 3) of the radical
R to its anionic state R', and IP(H) is the ionization potential for hydrogen
(equation 4), a value that is well known.26'30,35

R-H R + H BDE (equation 2)
R + e" R" -EA (equation 3)
H -^ff + e- IP (equation 4)
R-H-^R' + H* AHacid (equation 5)
Experimental techniques of measuring EAs and BDEs will not be discussed here,
having been reviewed elsewhere. As a word of caution, temperature conditions
at which these parameters are typically measured can vary and potentially affect the
calculated gas-phase acidity value. Therefore, correction factors need to be added
to equation 1 to ensure accurate values.
DePuy et al have developed a novel method for measuring gas phase
acidities of alkanes,40 the results of which can be found in Tables I & II. Using the
catalytic cycle in equations 2-5 with measured bond dissociation energies41-43 and
gas phase acidities,40 EA values could be calculated for various radicals (Table I).
Measuring and therefore interpreting EA and BDE trends for various radicals can
be very useful in understanding factors that stabilize / destabilize and therefore
increase the likelihood of carbanion formation.

Table I: Experimental Bond Dissociation Energies, Gas Phase Acidities, and
Electron Affinities of Simple Alkanes_________________________________
Neutral Species BDE (kcal/mol) at 298 Ka DHacid (kcal/mol) at 298 Kb EA (kcal/mol) at 298 Kc
ch3-h 104.8 416.7 0.7 1.8
ch3ch2-h 100.1 420.1 2.0 -6.4
(CH3)2CH-H 96.3 419.4 2.0 -9.5
Cyclobutyl-H 96.3 417.4 2.0 -7.5
Cyclopentyl-H 95.5 416.1 2.0 -7
Sec-butyl-H 96.3 415.7 2.0 -5.8
CH3CH2CH2-H 100.1 415.6 2.0 -1.9
Tert-butyl-H 93.6 413.1.2.0 -5.9
Isobutyl-H 100.1 412.9 2.0 0.8
Cyclopropyl-H 106.3 411.5 2.0 8.4
Neopentyl-H 100.1 408.9 2.0 4.8
a All data was taken from ref. 41,42, and 43. b Acidities were measured using silane cleavage branching
ratios taken from ref. 40. c Data was calculated using BDE and DH^ data in table by using thermo-
chemical cycles. All data were taken from ref. 40.
Table II: Experimental Electron Affinities
Radical Species_______________EAa_________
C 29.06 0.005
CH 28.60 0.23
ch2 14.99 0.23
ch3 1.85 0.69
CH2Br 18.22 3.22
CH2C1 17.07 3.69
Benzyl radical 15.75 0.30
(CH3)2BCH2 42.44 7.84
B 6.46 0.001
bh3 0.88 0.46
BO 57.89 0.23
BF, 49.12 3.0
a all data have kcal mol'1 units, and were taken from ref. 27.

The electron affinities (EA) of a number of simple radical groups have been
determined (Equation 3 and Tables I & II).2732,35,37,44 The EA of methyl and
related alkyl-substituted radicals range from slightly stable (EA of CH3 = 1.8 + 0.7
kcal mol'1) to unstable (EA of CH3CHCH3 = -9.5 2 kcal mol'1). The general trend
observed is a destabilizing one with increased alkyl substitution alpha to the anionic
center (Table I). Although carbon has an incomplete octet, the addition of an
electron is hindered by the inductive electron donation from the methyl-substituents
onto the radical center. Interestingly, there is a slight increase in electron affinity
from CH3CH2 radical (-6.4 kcal mol') to (Ct^C radical (-5.9 kcal mol'1).
Although the EA for tert-butyl radical is still negative, the addition of a third methyl
substituent seems to offset some of the anion destabilizing inductive effects
observed upon mono and dimethyl-substitution.35 Another trend worthy of notice is
anion stabilization upon increased alkyl-substitution located beta to the radical
center (EA of isobutyl radical = 0.8 kcal mol'1, propyl radical = -1.9 kcal mol'1, and
ethyl radical = -6.4 kcal mol'1). The additional methyl group seems to have offset
any destabilizing inductive effects that are typical of a-substituted methyl
substituents 45
Neighboring group effects on the EAs of non-carbon substituted methyl
radicals relative to methyl radical have been investigated both experimentally and
computationally (Table II).2746-48 The general trend is anion stabilization for

substituents that are electron withdrawing through both inductive and / or
delocalization effects.23,46-48 Boron is particularly good at delocalizing electron
density from an adjacent anion, particularly a methyl anion (EA for (CH3)2BCH2 =
42.4 + 7.8 kcal mol-1 )27 relative to its methyl-substituted analogue (EA for
(CH3)2CHCH2 = 1.2 + 2.8 kcal mol-1)40 Ion cyclotron resonance (ICR)
experiments show that the proton affinity (PA) of (CH^BCEb'1 (PA = 365 5 kcal
mol-1)49 is less than what is observed for its alkyl-substituted analogue
(CH3)2CHCH2-1 (PA = 162.0 kcal mol-1).27 The small proton affinity indicates that
proton addition requires the disruption of the u-delocalization between the unpaired
electrons of carbon and the empty p-orbital on boron.45 Borons ability to
delocalize electron density from an anion is also consistent with borons role in
various Lewis acid reactions.45,50-52 Since borons p-orbital is responsible for
delocalizing the electron density from methyl anion, any additional substituents on
boron that can donate, electron density into its vacant 2 p-orbital will destabilize this
methyl anion interaction.53 Therefore, one can expect that the greatest interaction
between methyl anion and boron will be found when borons substituents are
exclusively electron withdrawing.
Although boron can stabilize anionic charge indirectly by accepting electron
density from the neighboring anion, it is less capable of stabilizing the addition of
an electron directly. Elemental boron has an electron affinity of 6.5 + 0.001 kcal

mol'1, which is much less than that of carbon (29.1 0.005 kcal mol'1).54 The EA
of BH2 relative increases when it is substituted with more electronegative elements
(Table II).56,57 Additional experimental information is limited due to the difficulty
of making gas phase ions of molecules containing boron, which stems from borons
inherent reactive propensities with oxygen and its role as a strong Lewis acid.51,52,
Bond dissociation energies (BDE) have been measured for a number of
simple alkanes both experimentally and computationally (Tables I & III). Such
experiments show that inductive and electron delocalization effects of neighboring
substituents can drastically affect the stability of a carbon centered radical.58'63
Substituents that are inductively electron donating and electron withdrawing by
delocalization will stabilize radical centers. Methyl substituents are radical
stabilizing since they inductively donate electron density. This effect is more
pronounced as a-methyl substitution increases (BDE of (CLL^C-H = 93.6 kcal
mol'1) relative to methane (BDE of CH3-H =104.8 kcal mol'1). Boron substituents
are radical stabilizing (BDE of BH2CH2-H = 94.5 kcal mol'1) relative to its methyl
substituted analogue (BDE of CH3CH2-H =100.1 kcal mol'1) since the unpaired
electron is delocalized into the empty p-orbital of boron.58,59 Borons role as a
radical stabilizing substituent has also been shown in free radical bromination
and hydrogen abstraction experiments50 for alkylboranes, borinates, and

boronates.50 Delocalization of the methyl radical onto boron results in a partial
positive charge on carbon and a partial negative charge on boron.66 It would then
seem intuitive that electron withdrawing substituents on boron would further
stabilize a partial negative charge and therefore further decrease the BDE of the C-
H bond of interest. When an oxygen substituent replaces a hydrogen atom on
boron, the BDE of an adjacent C-H bond increases (BDE of MeO(Me)BCH2-H =
96.6 kcal mol'1) relative to the hydrogen-substituted analogue (BDE of Me2BCH2-H
= 94.4 kcal/mol). The increase in BDE occurs from boron sharing its vacant p-
orbital with both the non-bonding electrons on oxygen and the unpaired electron on
the carbon radical (Table III).63, M> 67
Table HI: Computed BDE of Boron and Carbon Species (kcal/mol)
Neutral Species G-2a CBS-4 b B3LYP/6-31G*c
BHj-H 105.2 104 -
BFH-H 104.8 103.4 - '
bf2-h 108.8 107.1 -
(CH^B-H 104 102.4 -
BHrCH, 104.2 103.5 -
BtCH^-CH, - 102.4 -
BHzCHz-H 94.5 94 -
MejBCHj-H - - 94.4
MeO(Me)BCH2-H - - 96.6
(MeO^BCHj-H - - 98.1
a Calculated bond dissociation energies are carried out using G-2 theory. All data were taken from ref. 67.
b Calculated bond dissociation energies are carried out using CBS-4 level of theory. All data were taken
from ref. 67. c All calculated bond dissociation energies are carried out using B3LYP /6-31G* level of
theory. All data were taken from ref 63.

Experimental BDEs of B-H groups are listed in Table IV.69 Notice the
similarities in the BDE boron hydride (BDE of BH2-H = 104.9 kcal mol'1) relative
to that of methane (BDE of CH3-H = 104.8). Rablen et al have used G-2 and CBS-
4 calculations to show good agreement between experimentally measured values
(Table I & IV) and those calculated (Table III).68 Their studies also revealed
similarities in the BDEs of B-C and B-H groups68 regardless of the differences in
the electronegativities of H and C. However, since the difference in
electronegativity between carbon and boron is greater than the difference between
boron and hydrogen, a slight increase in bond polarity occurs in the B-G bond. This
then forces a rehybridization in the carbon orbitals towards more s-character, thus
explaining similarities in B-H and B-C BDEs.
Table IVa: Experimental Bond
Dissociation Energies of Boranes
Neutral Species BDE (kcal/mol) at 298 K
B-H 84.30
BH-H 78.59
bh2-h 104.92
All data came from ref. 69.
Clearly, borons role in anionic and radical stabilization results from its
vacant p-orbitals ability to accept electron density through delocalization effects.
These anion and radical stabilizing features are what give boron its desirable

properties in some of the reactions previously mentioned. Boron, much like silicon,
has the capacity to delocalize electron density into one of its unoccupied valence
orbitals (its functionality). Recently, some computational work done by Damrauer
et al probed the distance dependence of silicon on carbanion stability;68 the results
from which suggest that extending this research to boron may prove fruitful.
Therefore, it is our objective to utilize proven computational methods to determine
the distance dependence of borons functionality and how it affects the BDEs and
EAs of various alkylboranes.

2. Computational Methods:
All input geometries were optimized using Moller-Plesset second-order
perturbation theory (MP2) with a 6-31++G(d,p) basis set; no geometry constraints
were used. Second derivative analytical energy gradient calculations (analytical
Hessians) are performed in order to determine the relative coordinates of the nuclei
and where they lie along that molecules potential energy surface (minima =
positive definite Hessian; nth-order saddle point = n negative eigenvalues). All
reported geometries are global minima for the indicated molecular configuration.
Single point energies were computed using density functional theory (DFT) with a
6-311+G(3df,2p) basis set and the B3LYP exchange-correlation functional.32'71-73
Curtiss et al demonstrated the efficacy of said computational methods for
computing ionization potentials and EAs for a diverse test group where results
differed by a range of 2.3 to 4.6 kcal mol'1 when compared to experimental
values. The energies obtained here were corrected for zero point energy
contributions (ZPC) using the vibrational analysis of the optimized structures. All
data are reported as energies at 0 K and have not been corrected for temperature

The BDE and EA computations were carried out by optimization of input
geometric parameters for the neutral, radical, and anionic species, followed by their
frequency and single point energy calculations. For comparative purposes,
additional calculations were performed, where the single point energies were
computed using density functional theory (DFT) with the B3LYP exchange-
correlation functional and the Dunning aug-cc-pVTP basis set.74'77 Vertical
attachment energy (VAE) computations were performed using single point energy
calculations using density functional theory (DFT) with aug-cc-pVTP basis set.
VAE calculations for both the anion and radical states are carried out using the
optimized radical geometry calculated using MP2/6-31++G(d,p) basis set. The
resulting VAEs are not zero point energy corrected. All computations were
carried out using the GAMESS suite of programs. MacMol Plot has been used to
visualize the molecular structures79

3. Results and Discussion:
Reliable methods for calculating thermodynamic parameters of an atom or
molecule computationally are available and have been reviewed elsewhere.32,80-83
Specifically, there is literature precedence of attaining accurate calculations of
electron affinities (EA uncertainties of 2.3 to 4.6 kcal mol'1).32 Adiabatic electron
affinity (EA) is defined as the difference between the total energies of the neutral
and anion at their respective equilibrium nuclear configurations32 The addition of
an electron usually results in a equilibrium geometry change. However, in the
instances where electron addition is stable only for a brief amount of time,
geometry relaxation into a more electron stabilizing configuration cannot occur. In
such instances, anionic molecules are better represented by the vertical attachment
energy (VAE). VAE is determined from the energy difference between the radical
and anion, where both species are described using only the equilibrium geometry of
the ground state neutral radical.32
Bond dissociation energies have been reviewed elsewhere; typical
computations are accurate to + 3.43 kcal mol-1.80-84,88 A bond dissociation energy
(BDE) is the minimum energy required to homolytically dissociate a bond in a
molecule into its respective radicals. Herein, BDE will be calculated as the

difference in energy between the equilibrium geometries of the neutral molecule
and that of the sum of the energies for both newly formed radicals (Equation 6).39
AEt0tai = (Ex + Eh) EX-h AEtotai = BDE (equation 6)
To properly represent radical and anionic orbitals in a molecule, diffuse,
functions and electron correlation on exchange functionals are necessary for
geometry optimizations, vibrational analyses, and single point energy
calculations.32 Basis functions 6-31++G(d,p) with Moller-Plesset second order
perturbation theory (MP2) for geometry optimizations and zero point correction
energies are sufficient to accurately calculate various types of alkyl anions, radicals,
and neutral species and therefore are utilized herein. An accurate and
computationally cost efficient method for calculating single point energies is the
density functional theory (DFT) B3LYP coupled with a 6-311+G(3df,2p) basis
set.33,65 Nevertheless, we have compared calculations of BH2, BH3", BH3, CH3",
CH3 and CH4 with their known experimental values. Taking into account that
temperature correction factors have not been added to the calculated parameters,
reasonably close BDEs have been calculated for CH4 and BH3 (101.09 and 102.46
kcal/mol) relative to their measured values (104.8 and 104.9 kcal mol'1). In
addition, reasonably accurate (+ 2 kcal mol'1), non-temperature corrected EA values
for CH3 and BH3 have been calculated (-0.2 and 1.6 kcal mol'1) relative to their

measured values (1.8 and 0.9 kcal mol'1).27 To test the basis description for this
class of molecules, single point energy calculations using the aug-cc-pVTP
Dunning basis set coupled with the B3LYP functional are also reported herein.
Dunnings augmented basis set showed acceptable results for CH4 and BH3 BDEs
(102.1 and 102.5 kcal mol'1) and EAs (0.7 and 3.4 kcal mol'1), indicating that the
values calculated using DFT are sufficiently accurate. VAEs for CH3 and BH3 are
calculated to be -2.12 and 1.15 kcal mol'1, which indicate that the geometry changes
upon electron addition do not largely affect the anion stability.
Having previously established that boron substitution causes large methyl
radical and anion stabilization effects relative to its methyl substituted analogue, a
study of the distance dependence of the boron moiety on EA and BDE of
alkylboranes seemed to be an interesting relationship to probe. The systems to be
explored are of the type BH2(CH2)nCH3 where n = 0 to 4.
3.1 BH2CH2 and BH2CH3. The BDE of BH2CH3 and EA of BH2CH2 have
been computed from their optimized equilibrium geometries (Tables V and VI).
The BDE for the C-H group in BH2CH3 is 90.5 kcal mol'1 and the EA for BH2CH2
radical is 42.0 kcal mol'1. Although there is no experimental evidence for the EA
and BDE values for these species, the EA of the methyl-substituted structural
analogue (CH3)2BCH2 is 42.4 kcal mol'1.45 Since it was shown earlier in this paper

that (CH3)3B and BH3 have similar C-H BDEs and EAs,35 it should be no surprise
that there are strong similarities between the calculated hydrogen-substituted and
experimentally measured methyl-substituted EAs. In addition, previous
computational studies have reported a BDE for (CH3)2BCH2-H to be
94.4 kcal mol', which is 4 kcal/mol larger than our results for BH2CH2-H
(90.5 kcal mol"1). As it is observed with tert-butyl radical, methyl substituents can
stabilize adjacent electron deficient atoms by donating electron density via
hyperconjugative effects. Since boron is electron deficient (incomplete octet),
methyl substituents will donate electron density onto boron. Boron then becomes
less electrophilic and thus less efficient at stabilizing an adjacent carbon radical
than its hydrogen substituted analogue H2BCH2-H. Thus, the BDE and EA reported
for BH2CH3 and its associated radical are consistent with what is observed

Table V: Computed Adiabatic Electron Affinities and
Vertical Attachment Energies______________________________
Radical Species EA (kcal/mol) a VAE (kcal/mol)c
ch3 -0.18 (0.67)b -2.12
bh2 23.64 (2.57)b -0.12
bh3 1.58 (3.40)b 1.15
bh2ch2 42.03 (42.08)b 38.69
bh2ch2ch2 45.70 (47.97)b -6.76
bh2ch2ch2ch2* 46.67 (46.54)b 0.21
bh2ch2ch2ch2ch2 61.81 (61.79)b 27.55
a Optimizations and vibrational analyses were carried out at the MP2/6-
31++G(d,p). Each geometry was subsequently followed by a single point energy
calculation at the B3LYP/6-311+G(3df,2p) level. The energies were corrected
for zero point energy contributions by an analysis of the vibrations. All EAs
reported herein are the corrected energies of the difference between the radical
and anion molecular energies. b These values represent single point energies ca-
lculated at BLY3P/aug-cc-pVTP level. All energies were corrected for zero poi-
nt energy contributions through analysis of vibrations. EA values are calculated
as the difference between the radical and anion molecular energies. 0 Vertical a-
ttachment energies were determined as the difference between the radical and an-
ionic energies calculated using the optimized radical geometries for both species.
All geometries were optimized using MP2/6-311++G(d,p) but were not corrected
for zero point energy contributions. All single point energy calculations were
calculated using BLY3P/aug-cc-pVTP. The lowest energy geometry is not
reported here since it contains two imaginary frequencies. Thus, the next lowest
energy species, differing only by 0.6 kcal/mol and very similar in geometry, but no
imaginary frequency, is used to calculate BDE.

Table M: Cmputed Bold Qssodaticn Engles
Saturated Species BH? (kcal/rrol)
H£-H 101.09 (10Z05)b
hb-h 102.46 (102.50)b
90.45 (90.54)b
ipa^a^H 95.43 (95.48)b
ipa^a^a^a^H 95.97 (96.10)b
a QtfimizatiQns andvibrational analyses vrere carried out at the MP2/6-
31+-Kj(4p)- EadigeonxtoyvrassubsequentiyMovsedtyasinglepoint energy
calculation at the B3LYP/6311-Kj(3d£2p) leveL The enogjesvsere corrected
for zero point energy contributions by an analysis of the vibrations. AllBCEs
reported herein are the corrected energies ofthe difference between theneutral
saturated energies and the summation ofthe molecular energies for both the ra-
dicals ^nerated after homofytic cleavage, bThesevalues represent single point
energies calculated at BLYP/augcc-pVIP. All oieigies-v^ee corrected fa-z-
ero point enei^caitributic^lhrexi^ana^isofvibratiais. BEE\aluesare
calculated as the diffeenee between the neutral saturated molecule and the sum
of the oiagjes fa- the radicals post hemolytic cleavage. *Thelowest0iergyg-
eometryisnofreportedheresinreitcontainstwoinBgjnaryfiequendes. Thus,
the next lowest energy species, diffeing only by 0.6 kcal/mol andvery similar in
geometry, but no imaginary frequency, is used to calculate BEE

The calculated BDE and EA for BH2CH3 and its associated radical can be
best explained by a thorough interpretation of the associated geometries of the
equilibrium structures for all three states: radical, anionic, and neutral (by neutral
we mean the species with an additional hydrogen, in this case BH2CH3). Anion and
radical stabilizing features are the result of electron delocalization, where the
electron density from the radical and anionic center is delocalized into the vacant p-
orbital of boron. This is confirmed by the shortening of the B-C bond distances.
The B-C bond shortens as the molecule moves from its neutral (1.56 A), to the
radical (1.53 A), and then to its anionic state (1.47 A) (Figures 1-3). Bond
shortening is typically observed upon the formation of a second bond and results
from rehybridization to more s-character to optimize orbital overlap.48 In this
instance, the second bond is formed as a result of the delocalization of electrons
from the carbon atom into the vacant p-orbital of boron. Both the radical and
anionic states are planar, which indicates a change from sp3 to sp2 hybridization
state on the carbon. Therefore, the radical and unshared electrons on the anion now
occupy a carbon p-orbital in order to optimize the degree of overlap with the vacant
p-orbital on boron. Rotating the BH2 moiety perpendicular from planarity results in
an increase in energy (H2BCH2 barrier height for the radical is 9.6 kcal mol'1)
indicating the loss of orbital overlap.64,84

1.10 | '56
H,*""? 1.09 i f 1.09 h3

Angles Degrees
H1-C-H2 106.36
H1-C-H3 106.36
H2-C-H3 110.16
H1-C-B 104.19
H2-C-B 114.46
H3-C-B 114.47
H5-B-H4 118.07
H5-B-C 120.92
H4-B-C 120.91
1.53 9 rv#*
oo & \ 1.19%

Angles Degrees
H1-C-H2 114.14
H1-C-B 122.95
H2-C-B 122.92
H3-B-H4 119.41
H3-B-C 120.29
H4-B-C 120.30
Angles Degrees
H1-C-H2 112.98
H1-C-B 123.51
H2-C-B 123.51
H3-B-H4 115.92
H3-B-C 122.06
H4-B-C 122.02

In ethyl cation, hyperconjugation from the adjacent methyl group occurs,
whereby electron density from an appropriately aligned sigma bond will delocalize
into the vacant p-orbital on the carbocation. Since methylborane is similar
electronically to ethyl cation, one would expect to observe similar hyperconjugative
effects from a methyl sigma orbital when appropriately aligned with borons vacant
p-orbital. Computational studies have shown that any C-H group that is aligned
parallel with borons vacant p-orbital (see figure 1) results in a decrease in its H-C-
B bond angle and an increase in C-H bond length (B-C-Hi = 104.19 and C-Hi =
1.10 A) relative to the other C-H groups on that atom (B-C-H2 = 114.46 and C-H2
1.09 A).83 Pross et al85 and Flood et al86 explain this phenomenon in terms of two-
electron interactions and bond-bond repulsion models; however, Mo et al
demonstrated that when the methyl group is rotated from the staggered to the
eclipsed conformation (Figure 4) of CH3BH2 the C-B bond shortens (staggered =
1.571 A and eclipsed = 1.569 A) followed by an decrease of the aligned C-H bond
length (staggered = 1.096 A and eclipsed = 1.084 A).83 Therefore, an interaction
between boron and the C-Hi bond when aligned parallel with borons vacant p-
orbital is broken upon rotation from the staggered to the eclipsed conformation.
Again, these findings are in good agreement with the optimized geometries
calculated in this study (Figure 1). The C-H sigma bond involved in
hyperconjugation with the vacant p-orbital on boron is destabilized since electron

density is being drawn away from it, thereby causing the reduction in BDE (90.45
kcal mol'1) relative to ethane (100.1 kcal mol'1).
This is a demonstration of staggered vs. eclipsed conformation of the adjacent
methylene moiety. In the staggered position, the H1 group is aligned parallel
with boron's vacant p-orbital. When the is on the right side indicates interaction
through hyperconjugativity. When placed on the left side, indicates when in a
hypercojugativity is not in effect.
3.2 BH2CH2CH2 and BH2CH2CH3. Geometry optimizations of the
BH2CH2CH2 anion resulted in the formation of a cyclic structure (Figure 5), where
boron becomes tetracoordinate. Here, the carbanion center interacts with the boron
through dative bond formation between the non-bonding electron pair on carbon
and that of the vacant p-orbital on boron. The tetracoordinate boron observed here
is best described as a distorted tetrahedron, which is indicated through the external
bond angles about the boron atom (H-B-H bond angle = 116.2, H-B-C bond angle
= 117.8). The ring portion of the molecule has endocyclic bond angles (C-C-B =
61.7 and C-B-C = 56.6) similar to that of cyclopropane (C-C-C = 60.0),87
however, with a much stressed C-B-C bond angle.

Angles Degrees
H1-C1-H2 111.87
H1-C1-B 120.77
H1-C1-C2 116.35
H2-C1-B 120.73
H2-C1-C2 116.35
H4-C2-H3 111.86
H4-C2-C1 116.32
H4-C2-B 120.75
H3-C2-C1 116.35
H3-C2-B 120.78
H5-B-H6 116.16
H5-B-C1 117.75
H5-B-C2 117.79
H6-B-C1 117.72
H6-B-C2 117.71
C1-B-C2 56.60
C2-C1-B 61.71
C1-C2-B 61.69
In addition to stressed endocyclic bond angles, increases in bond lengths are
found, but only for atoms bound directly to boron. For example, elongation of the
B-C bond (1.62 A), but without C-C bond lengthening (1.53 A), is observed relative
to the neutral specie (Figure 6) (B-C =1.56 A and C-C = 1.53 A). In addition, the
C-H bond lengths (1.09 A) are unchanged relative to the neutral molecule (C-H =

1.09 A), whereas a slight elongation in the B-H bond length (1.22 A) occurs relative
to its neutral (Figure 6) (B-H =1.19 A). Therefore, cyclization can be thought of in
rehybridization terms, where the boron orbitals go from sp to sp increasing p-
character, and causing bond elongation. All attempts to locate a lower energy
acyclic configuration at this level of computational theory have failed. A similar
phenomenon was reported previously by Damrauer et al using similar methodology,
for silicon substituted homologues.

As mentioned earlier, the adiabatic electron affinities (EA) are defined as
the total energy differences between the equilibrium geometries of both the radical
and the anionic species. The electron affinity of the ethylborane radical is large
(45.7 kcal mol'1) relative to CH3CH2CH2 (-1.9 kcal mol"1). The stability of the
anion relative to the radical is the result of cyclization, which permits direct boron-
carbon electron delocalization. Dunnings aug-cc-pVTP basis set coupled with
DFT was used to recalculate the energies of the optimized structures (for more on
methods see the Computational Methods section) to ensure correct basis selection,
but with little change in the EA (EA = 48.0 kcal/mol).
Large increases in molecular energy upon electron addition can be followed
by large geometry changes in order to better stabilize extra charge. Geometry
optimizations will find a lowest energy conformation for a molecules state, but the
timescale required to reach this lowest energy conformation is not always permitted
relative to electron loss. In such cases, vertical attachment energies (VAE) are
more representative of the system than EAs. VAE is calculated as the difference in
energy between the anion and radical, where the equilibrium geometry of the
radical is used to describe the geometry of both states. The VAE for ethylborane is
-6.7 kcal mol'1, which is significantly smaller than the EA = 48.0 kcal mol"1,
indicating that the ring-closure is necessary for anion stability. Therefore,

experimentation is necessary to determine which thermodynamic parameter (VAE
vs. EA) best represents this system.
Although the anion ring-closed, the radical remained acyclic, meaning
boron stayed in its tricoordinate state (Figures 6 and 7). The equilibrium geometry
of the radical does not permit the direct interaction between boron and carbon. Yet,
the BDE of a terminal C-H group on ethylborane (95.4 kcal mol'1) is significantly
lower than propane (100.1 kcal mol'1). As is observed for the neutral and radical
states of the ethylborane molecule (Figures 6 and 7), the staggered position is
preferred for the adjacent C-H moiety (Figure 4), suggesting that the adjacent C-H
moiety is involved in hyperconjugation onto boron. In addition, the C-H bond
lengths and H-C-B bond angles on each of the CH2 groups adjacent to boron are
noticeably different. As is the case for the radical of ethylborane, a difference of
0.02 A is observed between the two C-H bond lengths on the adjacent carbon and
13.7 between the two H-C-B bond angles (Figure 7). This effect is also observed
for the neutral molecule (Figure 6), where adjacent C-H bond lengths differ by 0.01
A and H-C-B bond angles differ by 8.3. This phenomenon has also been observed
elsewhere for neutral ethylborane, where results showed that the adjacent CH2
group to boron had a difference in C-H bond lengths by 0.01 A and H-C-B bond
angles differ by 7.2 Therefore, the preferred staggered formation, differences in
C-H bond lengths, and H-C-B bond angles indicate that a hyperconjugative effect

between the appropriately aligned C-H moiety with borons 2 p-orbital occurs. In
addition, the interaction is largest for the radical state.
Angles Degrees
H 1-C1-H2 118.05
H 1-C1-C2 121.49
H2-C1-C2 120.30
C1-C2-B 117.29
H3-C2-B 112.49
H3-C2-C1 111.91
H 3-C2-H4 104.03
H4-C2-B 98.83
H4-C2-C1 1 10.69
H5-B-C2 120.67
H6-B-C2 120.25
H5-B-H6 118.91
Previous studies have also shown that an interaction between the adjacent
C-H group and boron can be shown by the B-C bond elongation of neutral
ethylborane (from 1.576 to 1.595 A) upon rotating the BH2 moiety into a
hyperconjugative deactivating position (Figure 4).83 Although the radical (Figure
7) B-C and the aligned C-H bond lengths (1.56 A and 1.11 A) are slightly larger

than for the neutral (Figure 6) (1.57 A and 1.10 A), the aligned C-H bond has a
smaller H4-C2-B bond angle (98.8) relative to the neutral equivalent H5-C2-B
bond angle (102.4) (Figure 7); in addition to a smaller C1-C2 bond length for the
radical (1.49 A) (Figure 6) than the neutral (1.53 A) (Figure 7). The shorter C-C
bond for the radical in addition to a more shallow H4-C2-B bond angle for the
radical indicates that the hyperconjugation effect is not exclusive to atoms directly
bound to boron; rather this effect occurs throughout the molecule. In other words,
the electron density from C2-H4 sigma bond for the radical specie is delocalized
onto both the boron and Cl p-orbitals. This hyperconjugative effect is what
contributes to the decrease in BDE for ethylborane (95.4 kcal mol1) relative to its
methyl-substituted analogue (100.1 kcal mol"1). The BDE calculated using
Dunnings aug-cc-pVTP basis set with DFT (BDE = 95.5 kcal mol"1) is in good
3.3 BH2CH2CH2CH3 and BH2CH2CH2CH2. Much like the case with
ethylborane, propylborane has acyclic neutral and radical geometries (Figures 8 and
9) and ring-closed anions (Figure 10), where the nonbonding electron pair on the
carbanion are shared by the vacant p-orbital on boron. Upon ring-closure boron
becomes tetracoordinate, resulting in a rehybridization of its orbitals: sp2 to sp3.
This is further confirmed by the elongation of bonds on boron (Figure 10), as seen
with B-H (1.23 A) and B-C bond lengths (1.67 A) relative to the neutral (1.19 A)

and B-C (1.56 A). The tetrahedral nature of boron is also observed from the bond
angles surrounding boron (H-B-H = 110.4, H-B-C =119.4, H-B-C = 110.4, and
C-B-C = 83.8). Although there are similarities between the ethylborane and
propylborane anionic species, the propylborane ring is much less stressed. This is
confirmed by the inner bond angles of the cyclic species (83.8, 86.0, 92.7, and

H1 Hy Jj6
Angles Degrees
H1-C1-H3 107.77
H1-C1-H2 108.00
H1-C1-C2 110.64
H2-C1-C2 111.65
H2-C1-H3 107.95
H3-C1-C2 110.69
C1-C2-C3 113.04
H4-C2-H5 106.25
H4-C2-C1 109.62
H4-C2-C3 109.15
H5-C2-C1 109.36
H5-C2-C3 109.19
C2-C3-B 117.94
H6-C3-H7 103.43
H6-C3-C2 109.73
H6-C3-B 104.11
H7-C3-C2 110.68
H7-C3-B 109.78
H8-B-H9 118.22
H8-B-C3 119.99
H9-B-C3 121.76

Angles Degrees
H1-C1-H2 117.98
H1-C1-C2 121.08
H2-C1-C2 120:94
C1-C2-C3 113.91
H3-C2-H4 121.10
H3-C2-C1 109.71
H3-C2-C3 109.16
H4-C2-C 1 109.71
H4-C2-C3 109.16
C2-C3-B 117.49
H5-C3-H6 103.51
H 5-C3-C2 110.24
H5-C3-B 107.18
H6-C3-C2 110.25
H6-C3-B 107.21
H8-B-H7 118.28
H8-B-C3 119.76
H7-B-C3 121.96

h2 h8
Hl 1.10^ Cl -E ; *^^1.23
1.54 1 .67
1 1.54
1.10 . .u2 -c '3
Hf" i,10 ,oj 1 "'"He
h4 H 5
Angles___________D eg rees
H1-C1-H2 107.96
H1-C 1-B 1 1 1 .75
H 1-C 1-C2 109.55
H2-C 1-B 123.28
H2-C 1-C2 1 16.52
H3-C2-H4 1 07.39
H3-C2-C1 109.35
H3-C2-C3 109.38
H4-C2-C 1 118.60
H4-C2-C3 1 18.59
C 1-C2-C3 92.72
H5-C3-H6 107.93
H5-C3-B 11 1 .75
H5-C3-C2 109.52
H6-C3-B 123.30
H6-C3-C2 1 16.57
C2-C3-B 85.99
H8-B-H7 110.75
H8-B-C1 110.43
H8-B-C3 110.42
H7-B-C 1 1 19.42
H7-B-C3 1 19.39
C3-B-C 1 83.83
C2-C 1-B 85.99

The EA for the BH2CH2CH2CH2 radical is 46.7 kcal mol'1 (Table V).
Similarities between the EA of propylborane (46.7 kcal mol'1) and ethylborane
(45.7 kcal mol"1) are observed. Stabilization of the additional electron, relative to
its radical, is heavily favored in the ring-closed geometry. However, VAE
calculations indicate that without ring-closure of the anionic species, the anion is
slightly stable (0.2 kcal mol'1). Therefore, since there are large differences between
calculated values for EA and VAE, experimentation is necessary to discern whether
ring-closure occurs. It should be mentioned here that the lowest energy radical that
was calculated contained two imaginary frequencies. Since we failed to discover a
lower energy species without imaginary frequencies, we chose to report the second
lowest energy species without imaginary frequencies. Both species had similar
geometries and only differed by 0.6 kcal mol'1 in molecular energy.
The BDE for propylborane (96.6 kcal mol'1) is very similar to what is
observed for ethylborane (95.4 kcal mol"1). Clearly, the addition of a methylene
group between the terminal carbon and boron does not play a significant role in the
long-range hyperconjugative effects of boron. However, because the geometry
reported for the radical is not the lowest in energy, a detailed discussion is not

3.4 BH2CH2CH2CH2CH3 and BH2CH2CH2CH2CH2. As seen with the
other alkylboranes anions, the BH2CH2CH2CH2CH2 anion (Figure 11) will ring-
close, which is driven by borons ability to accept electrons into its vacant p-orbital.
The difference between butylborane and the previously investigated anions is that
the inner-angles of the five-member ring (Figure 11) are less stressed. As proof, the
geometry surrounding boron is much closer to tetrahedral geometry (C-B-C =
101.5, H-B-H = 108.8,H-B-C= 111.9 and 111.5) than the previous species. The
B-C bond length increases (1.67 A) relative to the neutral species (1.56 A)
indicating a rehybridization from sp2 to sp3.

Angles______________D egrees
H1-C1-H2 1 06.28
H 1 -C 1-B 114.83
H1-C1-C2 111.33
H2-C 1-B 111.48
H2-C1-C2 108.07
B-C1-C2 104.76
H3-C2-H4 106.93
H3-C2-C1 109.79
H3-C2-C3 1 08.26
H4-C2-C1 1 14.73
H4-C2-C3 1 12.43
C1-C2-C3 1 04.57
H5-C3-H6 1 06.93
H5-C3-C2 108.18
H5-C3-C4 109.79
H6-C3-C2 1 12.46
H6-C3-C4 114.72
C2-C3-C4 1 04.61
C3-C4-B 1 04.79
H7-C4-H8 1 06.26
H7-C4-C3 111.18
H7-C4-B 1 14.89
H8-C4-C3 108.08
H8-C4-B 1 1 1 .45
C4-B-C1 1 01 .46
H9-B-H 1 0 108.8
H9-B-C4 1 1 1 .39
H9-B-C 1 1 1 1 .77
H10-B-C4 1 1 1.86
H10-B-C1 1 1 1 .46

The EA of butylborane radical is 61.8 kcal mol'1 (Table V), which is the
largest observed of all the molecules investigated in this study. The stability is
attributed to the cyclization and therefore delocalization of electron density away
from the carbanion center. Since the inner angles of this ring structure are more
relaxed than the previously mentioned species in this discussion, a greater increase
in the EA is observed (Table V). Correct basis selection was confirmed by the
similarities in the calculated EA using Dunnings aug-cc-pVTP basis set
(61.8 kcal mol'1).
For most of the molecules studied herein, the difference in energy between
VAE and EA has been substantial. However, for butylborane, the VAE was found
to be surprisingly high (27.6 kcal mol"1). A closer examination of the radical
equilibrium geometry (Figure 12) shows a preferential alignment of the vacant 2-p
orbital on boron with that of the orbital housing the unpaired electron (Figure 12).
Since VAE is calculated using the radical geometry for both the anion and the
radical, both states appear to be in orbital overlapping distance relative to the empty
p-orbital on boron (terminal carbon and boron distance is 2.42 A).


As described above, the distance between the carbon radical and the boron
vacant p-orbital is close enough to allow some overlap. Therefore, electron
delocalization of the unpaired electron into the vacant p-orbital of boron would
further stabilize the formation of the radical. Interestingly, the BDE of butylborane
(96.0 kcal mol"1) is close in energy to what was calculated for ethyl and
propylborane (Table VI). A closer investigation into the structure of the radical
(Figure 12) indicates the presence of hyperconjugative effects, where the H-C-B
bond angles vary (107.9 and 112.73) and a reduction in the C1-C2 bond length on
the radical specie (1.49 A) and the neutral (Figure 13) (1.53 A). In addition, no
significant change in BDE relative ethyl and propyl BDE (Table IV) indicates that
increasing the carbon chain size has little to no effect on BDEs. Given that the
values for ethyl, propyl, and butylboranes are very similar, it seems more likely that
hyperconjugativity is the major contributor to the BDE observed. However,
conclusions cannot be drawn without experimentation.

Angles Degrees
H 1 -C 1 -H 2 1 0 7 .9 0
H 1 -C 1 -H 3 1 0 7 .9 3
H 1 -C 1 -C 2 1 1 1 .5 6
H 2 -C 1 -H 3 1 0 7 .8 3
H 2 -C 1 -C 2 1 1 0 .7 6
H 3 -C 1 -C 2 1 1 0 .7 2
C 1 -C 2 -C 3 1 1 2 .8 7
H 4 -C 2 -H 5 1 0 6 .4 8
H 4 -C 2 -C 1 1 0 9 .7 1
H 4 -C 2 -C 3 1 0 8 .8 8
H 5 -C 2 -C 1 1 0 9 .7 1
H 5 -C 2 -C 3 1 0 8 .9 9
C 2 -C 3 -C 4 1 1 3 .6 6
H 6 -C 3 -H 7 1 0 6 .2 4
H 6 -C 3 -C 2 1 0 9 .1 8
H 6 -C 3 -C 4 1 0 9 .3 2
H 7 -C 3 -C 2 1 0 8 .8 4
H 7 -C 3 -C 4 1 0 9 .3 3
C 3 -C 4 -B 1 1 7 .7 2
H 8 -C 4 -H 9 1 0 3 .7 5
H 8 -C 4 -C 3 1 0 9 .4 6
H' 8 -C 4 -B 1 0 2 .8 8
H 9 -C 4 -C 3 1 1 0 .8 2
H 9 -c 4 -B 1 1 0 .9 8
H 1 0 - B H 1 1 1 1 8 .2 5
H 1 0 - B C 4 1 2 1 .3 4
H 1 1 - B C 4 1 2 0 .3 5

4. Conclusions:
Computations have been performed on a series of alkylboron species, where
the distance dependence of boron on radical and anion stabilities have been
investigated through calculations of bond dissociation energies (BDE) and electron
affinities (EA). Results have shown that experimentation is necessary to determine
the relevancy of vertical attachment energy (VAE) to the adiabatic electron affinity
(EA) reported herein. In addition, experimentation might also show that zero point
energy corrections of calculated VAEs are necessary. Although the effect that
boron has on anion stabilities cannot be determined exactly without further
experimentation, it can be concluded that when the vacant p-orbital on boron
becomes directly involved in the delocalization of the nonbonding electrons on the
carbanion, a stabilizing effect will be observed. In addition, BDE is greatly reduced
by the presence of boron relative to its carbon analogue. Although the greatest
decrease in BDE occurs when boron is in direct interaction with the C-H bond of
interest, methyl hyperconjugative effects between boron and properly aligned
adjacent C-H groups can also cause a molecular effect which results in decreased
BDE of a terminal C-H group relative to its methyl-substituted analogue.

1. Burkhardt, E. R.; Matos, K. Chem. Rev. 2006, 10.1021/cr0406918.
2. Suzuki, A. Acc. Chem. Res. 1982, 15,178-184.
3. Matteson, D. S. Tetrahedron 1989,45,1859-1885.
4. Miyaura, N.; Suzuki, A. Chem. Rev. 1995, 95,2457-2483.
5. Knochel, P. J Am. Chem. Soc. 1990, 112, 7431-7433.
6. Iverson, C. N.; Smith, M. R., III. J. Am. Chem. Soc. 1999, 121, 7696-7697.
7. Mandal, A. K.; Schneekloth, J. S., Jr.; Crews, C. M. Org. Lett. 2005, 7, 3645-3648.
8. Ridgway, B. H.; Woerpel, K. A. J. Org. Chem. 1998, 63,458-460.
9. Matos, K.; Soderquist, J. A. J. Org. Chem. 1998, 63,461-470.
10. Cardenas, D. J. Angew, Chem. lnt. Ed. 2003,42, 384-387.
11. La ne, C. F.; Brown, H. C. J. Am. Chem. Soc. 1970, 92, 7212-7213.
12. B illaud, C.; Goddard, J.; Gall, T. L.; Mioskowski, C. Tettrahedron Lett. 2003,44,4451-
13. Gue nnouni, N.; Lhermitte, F.; Cochard, S.; Carboni, B. Tetrahedron 1995,6999-7018.
14. Ollivier, C.; Renaud, P. Chem. Rev. 2001, 101, 3415-3434.
15. Mat teson, D. S. J. Org. Chem. 1960, 82,4228-4233.
16. Batey, R. A.; Pedram, B.; Yong, K.; Baquer, G. Tetrahedron Lett. 1996,37,6847-6850.
17. Risberg, E.; Fi scher, A.; Somfai, P. Tetrahedron 2005, 61, 8443-8450.
18. Batey, R. A.; Smil, D. V. Angew. Chem. lnt. Ed. 1999, 38, 1798-1800.
19. Pasto, D. J .; Snyder, R., Sr. O.S.F. J. Org. Chem. 1966, 31,2777-2783.
20. Brow n, H. C.; Cope, O. J. J. Am. Chem. Soc. 1964, 86, 1801-1807.
21. He wavitharanage, P.; Danilov, E. O.; Neckers, D. C. J. Org. Chem. 2005, 70, 10653-

22. Are ntsen, K.; Caddick, S.; Cloke, F. G. N.; Herring, A. P.; Hitchcock, P. B. Tetrahedron
Lett. 2004,45,3511-3515.
23. Taft, R. W.; Raks hys, J. W., Jr. J. Am. Chem. Soc. 1965, 87,4387-4388
24. Ollivier, C.; Renaud, P. Angew, Chem. lnt. Ed. 2000, 39, 925-928.
25. Kabalka, G. W.; Malladi, R. R.; Te jedor, D.; Kelley, S. Tetrahedron Lett. 2000, 999-1001.
26. NIST has an excelle nt Gas-Phase Ion tutorial at http://webbook ,
27. The NIST Standard Reference Database is at
28. Bower s, M. T., Ed. Gas-Phase Ion Chemistry, Academic Press: New York, 1979; Vol. 1.
29. Bower s, M. T., Ed. Gas-Phase Ion Chemistry, Academic Press: New York, 1979; Vol. 2.
30. Bower s, M. T., Ed. Gas-Phase Ion Chemistry, Academic Press: New York, 1984; Vol. 3.
31. Fattahi, A.; Kass, S. R. J. Org. Chem. 2004, 69, 9176-9183.
32. Rien stra-Kiracofe, J. C.; Tschumper, G. S.; Schafer, H. F., III.; Nandi, S.; Ellison, G. B.
Chem. Rev. 2002, 102,231-282.
33. Graul, S. T.; Squires, R. R. J. Am. Chem. Soc. 1990, 112,2517-2529.
34. Graul, S. T.; Squire s, R. R. J. Am. Chem. Soc. 1989, 111, 892-899.
35. Graul, S. T.; Squires, R. R. J. Am. Chem. Soc. 1990, 112,2506-2516.
36. Graul, S. T.; Squires, R. R. J. Am. Chem. Soc. 1988,110,607-608.
37. DePu y, C. H.; Bierbaum, V. M.; Damrauer, R. J. Am. Chem. Soc. 1984, 106,4051-4053.
38. Sablier, M.; Fujii, T. Chem. Rev. 2002, 102,2855-2924.
39. Blank sby, S. J.; Ellison, G. B. Acc. Chem. Res. 2003, 36,255-263.
40. DePu y, C. H.; Gronert, S.; Barlow, S. E.; Bierbaum, V. M.; Damrauer, R. J. Am. Chem.
Soc. 1989,111,1968-1973.
41. Griller, D.; Kanabus -Kaminska, J. M.; Maccoll, A. J. Mol. Struct. 1988, 163,125-131.
42. Ru ssell, J. J.; Seetula, J. A.; Gutman, D. J. Am. Chem. Soc. 1988, 110, 3092-3099.
43. Ru ssell, J. J.; Seetula, J. A.; Timonen, R. S.; Gutman, D.; Nava, D. F. J. Am. Chem. Soc.
44. Ellison, B.G.; Engelking, P. C.; Lineberger, W. C. J. Am. Chem. Soc. 1978, 100,2556-

45. Murph y, M. K.; Beauchamp, J. L. J. Am. Chem. Soc. 1976, 98, 1433-1440.
46. Here, W. J.; Radom, L.; Schleyer P. v. R.; Pople, J. A. Ab Initio Molecular Orbital Theory,
John Wiley & Sons, Inc.: New York, 1986; pp 346 379.
47. El -Nahas, A. M.; Schleyer, P. von R. J. Comput. Chem. 1994, 15, 596-626.
48. Pross, A.; DeFrees, D. J.; Lev i, B. A.; Pollack, S. K.; Radom, L.; Hehre, W. J. J. Org.
Chem. 1981,46, 1693-1699.
49. Murph y, M. K.; Beauchamp, J. L. Inorg. Chem. 1977, 16, 2437-2443
50. Laidi g, K. E.; Streitwieser, A. J. Comput. Chem. 1996, 17, 1771-1781.
51. Haartz, J. C.; McDaniel, D. H. J. Am. Chem. Soc. 1973, 95, 8562-8565.
52. Larson, J. W.; McMahon, T. B. J. Am. Chem. Soc. 1985, 107, 766-773.
53. Veke n, B. J., van der; Sluyts, E. J. J. Am. Chem. Soc. 1997, 119, 11516-11522.
54. Wickha m-Jones, C. T.; Moran, S.; Ellison, G. B. J. Chem. Phys. 1989, 90, 795-805.
55. Kaldor, A.; Porter, R. F. J. Am. Chem. Soc. 1971, 93,2140-2145.
56. Pople, J. A.; Schleyer, P. von R.; Kaneti, J.; Spitznagel, G. W. Chem. Phys. Lett. 1988,
145, 359-364.
57. Vianello, R.; Ma ksic, Z. B. Inorg. Chem. 2005,44, 1095-1102.
58. Coolidge, M. B.; Bo rden, W. T. J. Am. Chem. Soc. 1988, 110, 2298-2299.
59. Song, K.; Liu, L.; Guo, Q. Tetrahedron 2004, 60, 9909-9923.
60. Bordwell, F. G.; Harrelson, J. A., Jr.; Zha ng, X. J. Org. Chem. 1991, 56, 4448-4450.
61. Ochters ki, J. W.; Petersson, G. A.; Wiberg, K. B. J. Am. Chem. Soc. 1995, 117, 11299-
62. Pasto, D. J .; Krasnansky, R.; Zercher, C. J. Org. Chem. 1987, 52, 3062-3072.
63. Walton, J. C.; McCarro 11, A. J., Chen, Q.; Carboni, B.; Nziengui, R. J. Am. Chem. Soc.
2000, 122, 5455-5463.
64. Hoff mann, R.; Radom, L.; Pople, J. A.; Schleyer, P. von R; Hehre, W. J.; Salem, L. J. Am.
Chem. Soc. 1972,94,6221-6223.
65. Grote wold, J.; Lissi, E. A.; Scaiano, J. C. J. Organometal. Chem. 1969,431-434.
66. Mattenso n, D. S. Prog. Boron Chem. 1970, 117 176.

67. Rablen, P. R.; Hart wig, J. F. J. Am. Chem. Soc. 1996,118,4648-4653.
68. Da mrauer, R.; Crowell, A. J.; Craig, C. F. J. Am. Chem. Soc. 2003, 125,10759-10766.
69. Bauer, S. H. Chem. Rev. 1996,96,1907-1916.
70. Moller, C.; Plesset, M. S. Phys. Rev. 1934,46,618-622.
71. Becke, A. D. J. Chem. Phys. 1993,98,5648-5652.
72. Lee, C.; Ya ng, W.; Parr, R. G. Phys. Rev. B 1988,37, 785-789.
73. Curti ss, L. A.; Redfem, P. C.; Raghavachari, K.; Pople, J. A. J. Chem. Phys. 1988,109,
74. Basis sets were obtained from the Extensible Computational Chemistry Environmental
Basis Set Database, V., as developed and distributed by the Molecular Science Computing
Facility, Environmental and Molecular Sciences Laboratory which is part of the Pacific
Northwest Laboratory, P.O. Box 999, Richland, WA 99352, USA, and funded by the U.S.
Department of Energy. The Pacific Northwest Laboratory is a multiprogram laboratory
operated by Battelle Memorial Institute for the U.S. Department of Energy under contract
DE-AC06-76RLO 1830. Contact David Feller or Karen Schuchardt for further
, ......v------------
76. Kendall, R. A.; Dunning, T. H., Jr.; Harrison, R. J. J. Chem Phys. 1992,96,6796-68C
77. Woon, D. E.; Dunni ng, T. H., Jr. J. Chem. Phys. 1993,98,1358-1371.
78. Sch midt, M. W.; Baldridge, K. K.; Boatz, J. A.; Elbert, S. T.; Gordon, M. S.; Jensen, J.
Koseki, S.; Matsunaga, N.; Nguyen, K. A.; Su, S.; Windus, T. L.; Dupuis, M.;
Montgomery, J. A., Jr. J. Comput. Chem. 1993,14,1347-1363.
79. Bode, B. M.; Gordon, M. S. J. Mol. Graphics Model. 1998, 16,133-138.
80' SD'Paridnso"'C' J-; MayerP- M'; Rad0'.LJ Phys. Chem. A 2001,105,675
81. Parkinson, C. J.; Mayer, P. M.; Radom, L. PieorChem. Acc. J599 lm o, 0.
S3.Mo,y,i/ao, g.p vR 2im,2305-23B.
85. Press, A.; m ^ J'M Soc. ft.
*t,ood-*. nun. o.- mN;V'J-An-'9S'


87. Ru ggiero, G. D.; Williams, I. H. J. Chem. Soc., Perkin Trans. 2001, 2, 733-737.
88. Petersson, G. A.; Malick, D. K.; Wilson, W. G.; Ochterski, J. W.; Montgomery, J. A., Jr.;
Frisch, M. J. J. Chem. Phys. 1998, 109, 10570-10579.