Acidity, basicity, and the stability of hydrogen bonds: Complexes of RO- + HCF3

Article abstract of DOI:10.1021/ja9817592

Ion-molecule complexes of RO^- (R = Me, Et, i-Pr) and HCF₃ have been studied with Fourier transform ion cyclotron resonance spectrometry. The RO^- complexation energies with HCF₃ were measured relative to RO^-·H₂O. These complexes, [ROHCF₃]^-, have complexation energies on the order of -20 kcal/tool and have low deuterium fractionation factors and are, therefore, hydrogen bonded. The structure of the complexes was studied by isotopic equilibrium experiments and ab initio calculations. All of the complexes studied have the structure RO^-·HCF₃ even when HCF₃ is a stronger acid than ROH. The structure of the complexes can be understood through electrostatic arguments rather than the difference in acidity between the ion and neutral.

Full text of DOI:10.1021/ja9817592

J. Am. Chem. Soc. 1998, 120, 10863-10870
10863
Acidity, Basicity, and the Stability of Hydrogen Bonds: Complexes
of RO- + HCF3
Michael L. Chabinyc and John I. Brauman*
Contribution from the Department of Chemistry, Stanford UniVersity, Stanford, California 94305-5080
ReceiVed May 20, 1998
Abstract: Ion-molecule complexes of RO- (R ) Me, Et, i-Pr) and HCF3 have been studied with Fourier
transform ion cyclotron resonance spectrometry. The RO- complexation energies with HCF3 were measured
relative to RO-‚H2O. These complexes, [ROHCF3]-, have complexation energies on the order of -20 kcal/
mol and have low deuterium fractionation factors and are, therefore, hydrogen bonded. The structure of the
complexes was studied by isotopic equilibrium experiments and ab initio calculations. All of the complexes
studied have the structure RO-‚HCF3 even when HCF3 is a stronger acid than ROH. The structure of the
complexes can be understood through electrostatic arguments rather than the difference in acidity between the
ion and neutral.
Introduction
structure and stability of hydrogen-bonded complexes is related
to the structure and stability of their ionic and neutral compo-
nents.
Hydrogen bonds are one of the most important noncovalent
interactions in chemistry. A hydrogen bond is defined most
generally as an intermolecular (or intramolecular) interaction
specifically involving a proton donor A-H and a proton
acceptor B.1 Its existence is often characterized by spectroscopic
features such as intense IR bands, unusual NMR shifts, and
structural features such as short contact distances (smaller than
van der Waals radii).2-4 Hydrogen bonds are observed in both
neutral and ionic systems. Understanding the ability of
hydrogen bonds to stabilize ions is especially important because
they are essential to a wide variety of phenomena, including
solvation of ions,5,6 stabilization of intermediates in proton-
transfer reactions,7,8 enzymatic stabilization of complexes and
transition states,9,10 and molecular recognition in biological and
nonbiological systems.11 In addition, knowledge of the structure
of gas-phase proton-bound ions is critical to the interpretation
of results from kinetic method determinations of thermodynamic
quantities.12 In this paper we address the question of how the
Several key issues remain unresolved about hydrogen bonding
in ionic systems. Hydrogen bonds to ions are thought to be
stronger than those to neutral molecules, but the difference in
magnitude is still debated.13,14 Empirically, a linear free energy
relationship (LFER) is seen between the complexation energy
and the difference in acidity and basicity of the molecules
involved for systems in both the gas phase and solution.15-19
Our knowledge of the dependence of hydrogen bond strength
on structure is limited, however, as most data are known for
hydrogen bonds between structurally similar molecules (e.g.,
proton bound amine dimers, alcohol-alkoxides).19-22 Solvation
is known to play a large role in both the stability and the
structure of hydrogen-bonded complexes.23 By uncovering the
intrinsic character of hydrogen bonds we can begin to understand
the basis by which these factors affect the interaction.
Gas-phase studies can provide information about the intrinsic
stability of hydrogen-bonded ionic intermediates by eliminating
solvation effects. Hydrogen-bonded complexes of both positive
and negative ions have been widely studied. Most of these
complexes are considered to have strong hydrogen bonds and
generally involve nitrogen and oxygen for both positive and
negative ions.15,17,19,22 For example, alcohol-alkoxide dimers
(13) Guthrie, J. P. Chem. Biol. 1996, 3, 163-170.
(14) Gerlt, J. A.; Kreevoy, M. M.; Cleland, W. W.; Frey, P. A. Chem.
Biol. 1997, 4, 259-267.
(15) Kebarle, P. Annu. ReV. Phys. Chem. 1977, 28, 445-476.
(16) Shan, S.-O.; Loh, S.; Herschlag, D. Science 1996, 272, 97-101.
(17) Caldwell, G.; Rozeboom, M. D.; Kiplinger, J. P.; Bartmess, J. E. J.
Am. Chem. Soc. 1984, 106, 4660-4667.
(18) Meot-Ner, M.; Sieck, L. W. J. Am. Chem. Soc. 1986, 108, 7525-
7529.
(19) Larson, J. W.; McMahon, T. B. J. Am. Chem. Soc. 1987, 109, 6230-
6236.
(20) Kreevoy, M. M.; Liang, T. M. J. Am. Chem. Soc. 1980, 102, 3315-
3322.
(21) Stahl, N.; Jencks, W. P. J. Am. Chem. Soc. 1986, 108, 4196-4205.
(22) Meot-Ner, M. J. Am. Chem. Soc. 1984, 106, 1257-1264.
(23) Perrin, C. L.; Nielson, J. B. J. Am. Chem. Soc. 1997, 119, 12734-
12741.
(1) Pimentel, G. C.; McClellan, A. L. The Hydrogen Bond; Freeman:
San Francisco, 1960.
(2) Buckingham, A. D. In Theoretical Treatments of Hydrogen Bonding;
Hadzi, D., Ed.; John Wiley & Sons: New York, 1997.
(3) Perrin, C. L.; Nielson, J. B. Annu. ReV. Phys. Chem. 1997, 48, 511-
544.
(4) Emsley, J. Chem. Soc. ReV. 1980, 9, 91-124.
(5) Jorgensen, W. AdV. Chem. Phys. 1988, 70, 469-488.
(6) Ions and Ion Pairs in Organic Reactions; Szwarc, M., Ed.; Wiley-
Interscience: New York, 1972.
(7) Eigen, M. Angew. Chem. Int. Ed. Engl. 1964, 3, 1-19.
(8) Bell, R. P. The Proton in Chemistry; Cornell University Press: Ithaca,
NY, 1973.
(9) Jencks, W. P. Catalysis in Chemistry and Enzymology; Dover
Publications: New York, 1987.
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Francisco, 1979.
(11) ComprehensiVe Supramolecular Chemistry; Lehn, J.-M.; Atwood,
J. L.; Davies, J. E. D.; MacNicol, D. D.; Vo¨gtle, F., Ed.; Pergamon: New
York, 1996.
(12) Cooks, R. G.; Patrick, J. S.; Kotaiho, T.; McLuckey, S. A. Mass
Spectrom. ReV. 1994, 13, 287-339.
10.1021/ja9817592 CCC: $15.00 © 1998 American Chemical Society
Published on Web 10/09/1998

10864 J. Am. Chem. Soc., Vol. 120, No. 42, 1998
Chabinyc and Brauman
have complexation energies in the range from 20 to 28 kcal/
mol,17,18 and proton-bound amine dimers have energies of 18
to 25 kcal/mol.22 Linear free energy relationships are often
observed between the complexation energy and the difference
in acidity and basicity of the neutral and ion. Because these
relationships are known only for a limited range of structural
types, e.g., alcohols, amines, carboxylic acids, etc., it is not
possible to fully understand their origin.22
Experimental Section
Chemicals. Fluoroform-d1, DCF3, was synthesized by a literature
procedure.31 The synthesized gas was purified by a trap-to-trap vacuum
distillation. The product obtained contained 5% HCF3 as an impurity
as determined by mass spectrometry. Dimethyl peroxide was synthe-
sized by a standard literature procedure.32 The product was purified
by a trap-to-trap vacuum distillation and characterized by mass
spectrometry. All other chemicals were obtained commercially and
used without further purification. Alkyl ethers were obtained from
Aldrich Chemical. N2O (99% pure grade) was obtained from Matheson.
All samples used were subjected to multiple freeze-pump-thaw cycles
before introduction into the ICR spectrometer.
Many anionic hydrogen-bonded complexes appear to be the
most stable when they consist of an ion and neutral which form
a nearly conjugate acid-base pair. That is, the better the match
of the gas-phase acidities of the neutral and the conjugate acid
of the ion, the more stable the complex. There are, however,
examples in which matched acidities are not sufficient to form
strong hydrogen bonds. For example, carbon acids are as acidic
as many alcohols in the gas phase; toluene and methanol have
comparable acidities. Nevertheless, the hydrogen-bonded com-
plex of methoxide and toluene is not known, and our efforts to
generate this complex have been unsuccessful.24 Caldwell and
Bartmess have examined complexes of phenylacetylide with a
series of alcohols whose acidity is comparable to that of
phenylacetylene.17 These ion-molecule complexes have com-
plexation energies ranging from 21.4 kcal/mol for methanol to
26.6 kcal/mol for benzyl alcohol.17 Meot-Ner has studied
complexes of cyclopentadienide with alcohols which have
strengths near 20 kcal/mol, but whether these complexes are
best termed “hydrogen bonded” is unclear.25 Thus, the acidity
and basicity of the neutral and ion are not the only factors
governing the stability of hydrogen-bonded intermediates.19,26
Clearly polarity has an effect on the stability of the complex,
but its role is uncertain due to the lack of experimental studies.
Instrumentation. All experiments were performed on an IonSpec
Fourier transform ion cyclotron resonance (FT-ICR) spectrometer.
Details of the spectrometer have been given previously.33 The magnetic
field strength was 0.6 T. The temperature in the cell is estimated to
be 350 K.30 Background pressures were on the order of 2.0-5.0 ×
10-9 Torr, and operating pressures ranged from 0.7 to 1.6 × 10-6 Torr.
Pressure measurements were made with an ion gauge (Granville Phillips
330), which was calibrated against a capacitance manometer (MKS
170 Baratron with a 315BH-1 sensor). We estimate the absolute
pressure measurements to have an error of (20%.
Ion-Molecule Chemistry. Reactions of CF3- have previously been
studied in the flowing afterglow environment.34 We have reinvestigated
its reaction with alkyl formates, the Riveros reaction,35,36 to see if similar
results would be obtained in the ICR (Scheme 1). CF3- was produced
from electron impact on CF4 or the proton-transfer reaction of HCF3
with CH3O-, which was generated by electron impact on dimethyl
-
peroxide. CF3 was isolated and allowed to react with various alkyl
formates, HCO2R (R ) Me, Et). Although ion-molecule complexes
of HCF3, [ROHCF3]-, can be isolated from the reaction of CF3 with
-
the formates, they further undergo a solvated Riveros reaction with
the alkyl formate precursor to produce [ROHCF3ROH]-, a cluster of
three molecules (eq 3, Scheme 1).37 These results are consistent with
the flowing afterglow results.34 Attempts to obtain [ROHCF3]- from
the exchange reaction of HCF3 with RO-‚HOR were not fruitful.
Alcohol‚alkoxide complexes are sufficiently stable that they do not
undergo a solvent switch with HCF3 (eq 4, Scheme 1).
To better understand the factors affecting hydrogen bonding,
we have chosen to study ion-molecule complexes of fluoro-
form, HCF3. Fluoroform is known to form hydrogen bonds
with other neutrals such as ammonia27 and has a similar acidity28
and polarity29 as the simple aliphatic alcohols. Furthermore,
proton-transfer reactions between fluoroform and alkoxides are
rapid, indicating that there are no large barriers on the potential
surface as observed for other carbon acids such as substituted
toluenes.30 Fluoroform should therefore provide a good case
for comparison with other hydrogen-bonded complexes.
Scheme 1
In this paper we report studies of the hydrogen-bonded
intermediates in the reaction of HCF3 with several alkoxides.
We have characterized these complexes through equilibrium
binding studies as well as isotopic equilibrium fractionation
experiments. Our studies show that the structure of the complex
is not solely determined by the overall thermochemistry of the
proton-transfer reaction. Our results suggest that electrostatics
provide an important key to understanding the structure and
strength of the hydrogen-bonded complexes.
Alkoxide-water complexes, RO-‚H2O, were synthesized from the
proton-transfer/elimination reaction of hydroxide with dialkyl ethers
(Scheme 2).38 O-• was generated by electron impact upon N2O; OH-
was then produced by the H atom abstraction reaction of O-• with either
HCF3 or the alkyl ethers present. The elimination reaction of OH-
with the alkyl ethers was competitive with proton transfer from HCF3
at pressure ratios of ether:HCF3 of ∼4:1, so reasonable quantities of
RO-‚H2O could be obtained (eq 6, Scheme 2). Alkoxide‚water
complexes of MeO-, EtO-, and i-PrO- were generated from tert-butyl
(24) Gatev, G. G.; Zhong, M.; Brauman, J. I. J. Phys. Org. Chem. 1997,
10, 531-536.
(31) Andreades, S. J. Am. Chem. Soc. 1964, 86, 2003-2010.
(32) Hanst, P. L.; Calvert, J. G. J. Phys. Chem. 1959, 63, 104.
(33) Wilbur, J. L. Ph.D. Dissertation, Stanford University, 1993.
(34) McDonald, R. N.; Chowdhury, A. K. J. Am. Chem. Soc. 1983, 105,
7267-7271.
(25) Meot-Ner, M. J. Am. Chem. Soc. 1988, 110, 3858-3862.
(26) French, M. A.; Ikuta, S.; Kebarle, P. Can. J. Chem. 1982, 60, 1907-
1918.
(27) Fraser, G. T.; Lovas, F. J.; Suenram, R. D.; Nelson, D. D., Jr.;
Kemperer, W. J. Chem. Phys. 1986, 84, 5983-5988.
(28) Lias, S. G.; Bartmess, J. E.; Liebman, J. F.; Holmes, J. L.; Levin,
R. D.; Mallard, W. G. 1988, 17, Suppl. 1.
(29) CRC Handbook of Chemistry and Physics, 75th ed.; Lide, D. R.,
Ed.; CRC Press: Boca Raton, 1994.
(35) Blair, L. K.; Isolani, P. C.; Riveros, J. M. J. Am. Chem. Soc. 1973,
95, 1057-1060.
(36) DePuy, C.; Grabowski, J. J.; Bierbaum, V. M.; Ingemann, S.;
Nibbering, N. M. M. J. Am. Chem. Soc. 1985, 107, 1093-1098.
(37) RO-‚HOR is also observed, presumably from decomposition of the
[ROHCF3ROH]- product.
(30) Han, C.-C.; Brauman, J. I. J. Am. Chem. Soc. 1989, 111, 6491-
6496.
(38) DePuy, C. H.; Bierbaum, V. M. J. Am. Chem. Soc. 1981, 103, 5034-
5038.

Complexes of RO- + HCF3
J. Am. Chem. Soc., Vol. 120, No. 42, 1998 10865
Table 1. Equilibrium Constants and Thermochemical Values for
the Exchange Reactiona
RO-‚H2O + HCF3 h [ROHCF3]- + H2O
∆G°
(kcal/mol)
∆H°
(kcal/mol)
∆H°ab initio
(kcal/mol)
RO-
Keq
MeO-
EtO-
0.29 ( 0.06
0.34 ( 0.07
0.39 ( 0.08
0.86 ( 0.1
0.75 ( 0.1
0.65 ( 0.1
0.38
0.27
0.17
0.70
0.60
0.60
i-PrO-
a All values measured at 350 K. ∆H°ab initio was calculated at the
MP2/6-311++G**//HF/6-311++G** level and corrected to 350 K
with vibrational frequencies scaled by 0.89.
Table 2. Gas-Phase Acidities
AH h A- + H+
Figure 1. Fractional ion intensities as a function of time until
equilibrium for the reaction. MeO-‚H2O + HCF3 h [MeOHCF3]-
+
AH
∆H° (kcal/mol)
∆H°ab initiod (kcal/mol)
H2O: 2, [CH3OHCF3]-; b, [CH3OH2O]-. PHCF ) 2.2 × 10-7 Torr,
3
PH 0 ) 1.9 × 10-7 Torr.
MeOH
EtOH
i-PrOH
t-BuOH
HCF3
381.5 ( 0.1a
378.6 ( 0.8b
376.7 ( 0.8b
375.9 ( 0.8b
377.8 ( 0.5c
383.6
379.5
377.1
375.8
380.5
2
methyl ether, ethyl ether, and isopropyl ether, respectively. No side
reactions with the alkyl ethers were observed. No other ion-molecule
complexes such as [HOHCF3]- were observed. Equilibrium was found
to be measurable between alkoxide complexes of H2O and HCF3 (eq
7, Scheme 2).
a From: Meot-Ner, M.; Siek, L. W. J. Phys. Chem. 1986, 90, 6687-
6690. b From Ervin, K. M.; Gronert, S.; Barlow, S. E.; Gilles, M. K.;
Harrison, A. G.; Bierbaum, V. M.; DePuy, C. H.; Lineberger, W. C.;
Ellison, G. B. J. Am. Chem. Soc. 1990, 112, 5750-5759. c From ref
28 re-anchored to values from footnote b. d ∆H°ab initio was calculated
at the MP2/6-311++G**//HF/6-311++G** level and corrected to 298
K with vibrational frequencies scaled by 0.89.
Scheme 2
fractionation factor, Φ, was determined by the method outline above.
All of the RO-‚H2O was consumed before the equilibrium constant
was measured. The deuterium fractionation factor, Φ (the equilibrium
constant for eq 8), is defined in eq 9. The ratios, [M + 1]/[M], were
corrected for the contribution from the natural abundance of 13C in the
[M] ion to the [M + 1] peak. We have assigned the errors in Φ based
upon the statistical error in the measured values.
[RODCF3]-/[ROHCF3]-
Deuterium isotope fractionation experiments were performed in a
similar manner. The [ROHCF3]- complexes were synthesized as shown
in Scheme 2. The equilibrium constant for exchange of DCF3 and HCF3
into the alkoxide complex was then measured (eq 8).
Φ )
(9)
[DCF3]/[HCF3]
Theory. All ab initio calculations were performed with Gaussian9439
on an IBM RS6000-590. Geometries were optimized at the Hartree-
Fock level with the 6-311++G** basis set. All stationary points were
characterized by vibrational frequency analyses. MP2 single-point
calculations were performed at the HF/6-311++G** geometry to
account for electron correlation. Frequencies were scaled by 0.89 for
use in thermochemical calculations.40 This level of theory has been
shown to reproduce experimental gas-phase acidities with an error of
approximately (1-2 kcal/mol.41 The experimental and ab initio gas-
phase acidities are given in Table 2. Visualization of electrostatic
potential surfaces was performed with Spartan and MacSpartan.42
Proton-transfer potential energy surfaces were calculated by optimiz-
ing the geometry of the complexes at fixed values of the C-H or O-H
distance. MP2 single-point calculations were performed at each of these
points along the reaction path.
The calculation of the eigenvalues in model potentials was ac-
complished with Truhlar’s FDBVM program.43 The program uses the
finite difference boundary value method for obtaining the eigenvalues
and eigenfunctions of an arbitrary potential function. Model double
well potentials were constructed to represent the calculated reaction
barrier and energy difference between wells for the proton-transfer
reactions.20 Both fourth- and sixth-order polynomials were used, and
[ROHCF3]- + DCF3 h [RODCF3]- + HCF3
(8)
Equilibrium Measurements. All equilibrium measurements were
obtained as an average of at least five trials at several pressure ratios
on different days. At the pressures used, the system reached equilibrium
in ∼1 s. A typical plot for the equilibrium between [ROHCF3]- and
RO-‚H2O is shown in Figure 1. Several methods were used to test
whether a true equilibrium had been achieved. After a constant ratio
of ion intensities was obtained, one species was ejected and the reaction
was followed in time until a constant ratio was reached again. The
final ratio of ion intensities was found to be independent of ejected
ion. The equilibrium constant obtained at several pressure ratios was
found to be in good agreement. The major source of error in these
experiments is the accuracy of the measurement of the absolute
pressures. We have assigned the error in the equilibrium constants
based on our estimate for the error in the absolute pressure readings
(20%). The relative values should be more accurate because the relative
pressure errors should be smaller (∼10%). The relative values are more
important for our discussion than the absolute values.
The isotopic exchange measurements were performed in a similar
manner. To obviate the difficulty of absolute pressure measurements
we determined the ratio of DCF3/HCF3 in situ by measuring the ratio
of the DCF2+/HCF2+ peaks in the positive ion spectra. The ratio was
found to be nearly independent of electron-impact energy and delay
time of detection. The ratio of ion intensities was also very similar to
the ratio of pressures as measured by the ion gauge. The equilibrium
(39) Frisch, M. J.; et al. Gaussian94ReVC.3; Gaussian, Inc.: Pittsburgh,
PA, 1995.
(40) Scott, A. P.; Radom, L. J. Phys. Chem. 1996, 100, 16502-16513.
(41) Smith, B. J.; Radom, L. Chem. Phys. Lett. 1995, 245, 123-128.
(42) MacSpartan; Wavefunction, Inc.: Irvine, 1996.
(43) Truhlar, D. G. J. Comput. Phys. 1972, 10, 123-132.

10866 J. Am. Chem. Soc., Vol. 120, No. 42, 1998
Chabinyc and Brauman
Table 3. Complexation Energies of Alkoxides with Watera
RO- + H2O h RO-‚H2O
each gave similar results.44 The reduced mass of the vibration was
taken to be that of hydrogen or deuterium.45
Results
RO-
∆H° (kcal/mol)
∆H°ab initio (kcal/mol)
RO-‚H2O h RO-‚HCF3. The equilibrium constants for
exchange of HCF3 with alkoxide-water complexes, RO-‚H2O
(R)Me, Et, i-Pr), were measured at 350 K (eq 10, Scheme 3).
Experimental values for Keq and ∆G° are shown in Table 1.
The value of ∆H° can be derived from ∆G° if ∆S° for the
exchange reaction is known. However, obtaining accurate
estimates of the entropies for ion-molecule complexes is
difficult.17,46 The vibrational modes with the greatest contribu-
tion to the entropy are the six low-frequency modes created
upon complexation. These modes are very anharmonic and are
generally not reproduced accurately by standard ab initio
harmonic frequency calculations.46,47 Due to this difficulty we
have chosen to assume the intrinsic ∆S° for eq 10 is zero and
to correct for symmetry only.48 We expect that this will
introduce an error of no more than (1.0 kcal/mol in the derived
value of ∆H°. The relative error between values should be
smaller because the ∆S° of complexation for these structurally
similar complexes should be nearly equivalent. Values for ∆H°
obtained from ab initio calculations agree well with the
experimental values (Table 1).
MeO-
EtO-
-23.9
[-22.3]
[-21.2]
-20.9
-23.0
-21.5
-21.0
-20.1
i-PrO-
t-BuO-
a Experimental values are from ref 18 and have errors of (1 kcal/
mol; values in brackets are interpolated from experimental values.
∆H°ab initio was calculated at the MP2/6-311++G**//HF/6-311++G**
level and corrected to 350 K with vibrational frequencies scaled by
0.89.
Table 4. Complexation Energies of Alkoxides with Fluoroforma
RO- + HCF3 h [ROHCF3]-
RO-
∆H° (kcal/mol)
∆H°ab initio (kcal/mol)
MeO-
EtO-
-23.5
[-22.0]
[-21.1]
-22.3
-20.9
-20.4
i-PrO-
a Values were derived as described in the text; values in brackets
were derived from interpolated values. ∆H°ab initio was calculated at the
MP2/6-311++G**//HF/6-311++G** level and corrected to 350 K
with vibrational frequencies scaled by 0.89.
Scheme 3
values for the water complexation energies of EtO- and i-PrO-
have been obtained by interpolation from this plot and are listed
in Table 3.
We have also performed ab initio calculations to determine
the binding energies of the RO-‚H2O (R ) Me, Et, i-Pr, t-Bu)
complexes (Table 3). The calculated values of ∆H° for
MeO-‚H2O and t-BuO-‚H2O are systematically lower than the
experimental values (Table 2). One possible source of this
discrepancy is the inability to accurately calculate the O-H
stretch of water in the ion-molecule complex that is expected
to be highly anharmonic.50,51 This frequency is expected to be
red-shifted in the complex, which would lead to an increase in
the binding energy because of the change in zero-point energy.
However, these discrepancies are small (less than 1 kcal/mol)
and the relative values agree well with the experimental values.
A plot of the ab initio values versus the ab initio acidities of
ROH gives a slope of 0.4 that is in reasonable agreement with
the slope of 0.5 from the experimental values. We therefore
believe that we can safely use interpolated values from the
experimental LFER plot for the water complexation energies
of EtO- and i-PrO-.
RO- + HCF3 h RO-‚HCF3. The values of the fluoro-
form-alkoxide complexation energies can be derived from the
thermochemical cycle in Scheme 3. We have used the
experimental value of the complexation energy of MeO-‚H2O
to derive the HCF3 binding energy of MeO-. Interpolated
values of the water complexation energies were used to obtain
the HCF3 binding energies of EtO- and i-PrO-. The values
show the expected decrease in binding energy with acidity as
seen in other systems (Table 4). The derived values from the
thermochemical cycle also compare well with the [ROHCF3]-
complexation energies calculated purely by ab initio methods
(Table 4). Again the calculated values are systematically too
low compared to the derived values, but the relative values are
nearly identical. If the complexation energies are plotted versus
RO- + H2O h RO-‚H2O. We cannot measure the relative
fluoroform‚alkoxide, [ROHCF3]-, binding energies directly due
to preferential formation of the alcohol‚alkoxide, RO-‚HOR′
dimer (eq 13). However, they can be derived via the thermo-
chemical cycle in Scheme 3.
The values of the water complexation energies of the alkoxides
are needed to derive the fluoroform complexation energies. The
ion-molecule complexation energy of MeO- and t-BuO- to
H2O has been measured.18,49 Values for complexes of EtO-
and i-PrO- are not available in the literature, and we are unable
to determine them experimentally using our instrumentation.
However, these values can be estimated or calculated. The
binding energies of alcohols to alkoxides are known to follow
a linear free energy relationship with the acidity of the alcohol.
We expect water binding energies to exhibit similar behavior.
Indeed, a plot of the experimentally known RO-‚H2O com-
plexation energies versus the acidity of ROH has a slope of 0.5
similar to the value of 0.4 for alcohol-alkoxide complexes.17
The gas-phase acidity values used are listed in Table 2. The
(44) Laane, J. Appl. Spectrosc. 1970, 24, 73-80.
(45) Examination of the normal coordinates shows that this is a reasonable
approximation for the systems considered here.
(46) Paul, G. J. C.; Kebarle, P. J. Phys. Chem. 1990, 94, 5184-5189.
(47) East, A. L. L.; Radom, L. J. Chem. Phys. 1997, 106, 6655-6674.
(48) Benson, S. W. Thermochemical Kinetics, 2nd ed.; Wiley: New York,
1976.
(49) We believe the literature value for the complexation of energy of
tert-butoxide to water in ref 18 is too large due to the estimated entropy
value. We have assumed here that the difference in ∆G° of complexation
between MeO- and t-BuO- is equal to the difference in ∆H°.
(50) Yates, B. F.; Schaefer, H. F., III; Lee, T. J.; Rice, J. E. J. Am. Chem.
Soc. 1988, 110, 6327-6332.
(51) Del Bene, J. E.; Jordan, M. J. T. J. Chem. Phys. 1998, 108, 3205-
3212.

Complexes of RO- + HCF3
J. Am. Chem. Soc., Vol. 120, No. 42, 1998 10867
large downfield shifts in NMR spectra.3,4 In the gas phase these
techniques are difficult to perform on ionic systems, so
hydrogen-bonded complexes have been characterized by their
complexation energies (∆H° for eq 18) relative to typical ion-
molecule complexes, e.g., 20-30 kcal/mol vs 10-15 kcal/mol.53
A- + HB h A-‚HB
(18)
The magnitude of the complexation energy that an ion-
molecule complex must have to be considered as hydrogen
bonded is not defined. Most systems that have been studied
have large complexation energies.
In the gas phase a linear free energy relationship between
the complexation energy and the difference in acidity and
basicity of the hydrogen-bonded molecule and ion is often
observed.17,19 The slope from these plots for negative ion
complexes such as [ROHF]- and [ROHOR′]- is typically near
0.5.17,54 McMahon has also derived a predictive model for the
complexation energy based on the electronegativities of the
heteroatoms, A and B, and acid-base acidities.19 Because there
is a lack of data for complexes that contain dissimilar hydrogen
bond donors and acceptors, it is not well understood how acidity
and basicity contribute to the hydrogen bond strength. For
example, phenylacetylide, whose basicity is comparable to that
of alkoxides, forms strong hydrogen-bonded complexes with
alcohols, but its complex with phenylacetylene has not been
observed.17 By studying hydrogen-bonded complexes of dis-
similar ions and neutrals we can begin to understand the basis
of these relationships.
To understand the factors which contribute to the stability of
the complex, the identity of the hydrogen bond donor and
acceptor within the complex must be known. The structures
of hydrogen-bonded intermediates of proton transfer reactions
are generally assumed to reflect the acidity difference of the
endpoints of the proton-transfer reaction. That is, the structure
is A-‚HB if AH is the stronger acid (compared to HB), and
AH‚B- if BH is the stronger acid. This assumption has been
tested in several ways including photodetachment spectro-
scopy,55-57 isotopic equilibrium effects,58 and computational
analysis.59-61 Proton-transfer reactions of various alcohols with
F- have been widely studied. Deuterium isotopic fractionation
experiments and computational studies suggest that the potential
energy surface (PES) near the intermediates is nearly flat,58,61
indicating that the two possible complexes RO-‚HF and
ROH‚F- have similar energies. Photodetachment spectroscopy
has shown that despite the flatness of the PES, the structure of
the intermediate complex is influenced by the acidity difference
of the two reactants.56,62 The isolated intermediate resembles
Figure 2. Alkoxide-fluoroform complexation energies as a function
of acidity of ROH. The slope of the least squares line is 0.5.
Table 5. Deuterium Fractionation Factors, Φ, for
Alkoxide-Fluoroform Complexesa
[ROHCF3]- + DCF3 h [RODCF3]- + HCF3
RO-
Φ (HCF3/DCF3)
Φ′(ROH/ROD)
MeO-
EtO-
0.57 ( 0.03
0.68 ( 0.02
0.74 ( 0.02
[0.37]
[0.44]
[0.48]
i-PrO-
a All values measured at 350 K. Values in brackets were derived as
indicated in the text.
the acidity of the alcohol, the slopes of the derived and ab initio
values are 0.5 and 0.3, respectively (See Figure 2).
Fractionation Factors. The deuterium isotopic fractionation
factors, Φ, of the [ROHCF3]- complexes were measured by
the equilibrium exchange reaction with DCF3 and HCF3 (eq
8). The measured values are shown in Table 5. Fractionation
factors, by definition, are measured relative to one of the
compounds in the hydrogen-bonded dimer, in our case HCF3
and DCF3. Although we cannot experimentally measure the
fractionation factor, Φ′ (eq 17), relative to ROH and ROD, we
can easily derive it from a thermochemical cycle (Scheme 4).
We have experimentally measured the equilibrium constant, Φ,
for eq 14 and can obtain the equilibrium constant for eq 15
from the difference in zero-point energies of HCF3/DCF3 and
ROH/ROD.52 The resulting values of Φ′ are shown in Table
5.
Scheme 4
(53) Note that the difference between these values is ∼10-15 kcal/mol,
which is similar to the estimated strength of the hydrogen bonds in solution,
refs 13 and 16.
(54) Larson, J. W.; McMahon, T. B. J. Am. Chem. Soc. 1983, 105, 2944-
[RODCF3]-/[ROHCF3]-
2950.
(55) Moylan, C. R.; Dodd, J. A.; Han, C.-C.; Brauman, J. I. J. Chem.
Phys. 1987, 86, 5350-5357.
Φ′ )
(17)
[ROD]/[ROH]
(56) Mihalick, J. E.; Gatev, G. G.; Brauman, J. I. J. Am. Chem. Soc.
1996, 118, 12424-12431.
Discussion
(57) de Beer, E.; Kim, E. H.; Neumark, D. M.; Gunion, R. F.; Lineberger,
W. C. J. Chem. Phys. 1995, 99, 13627-13636.
Hydrogen bonds are defined by their unique properties
relative to other intermolecular interactions, that is, their strength
and structure. We focus here on hydrogen bonds to ions. In
solution, ionic hydrogen bonds are characterized spectroscopi-
cally by energetic shifts in IR and UV absorption spectra and
(58) Wilkinson, F. E.; Szulejko, J. E.; Allison, C. E.; McMahon, T. B.
Int. J. Mass. Spectrom. Ion Proc. 1992, 117, 487-505.
(59) Cao, H. Z.; Allavena, M.; Tapia, O.; Evleth, E. M. J. Phys. Chem.
1985, 89, 1581-1592.
(60) Gronert, S. J. Am. Chem. Soc. 1993, 115, 10258-10266.
(61) Wladkowski, B. D.; East, A. L. L.; Mihalick, J. E.; Allen, W. D.;
Brauman, J. I. J. Chem. Phys. 1993, 100, 2058-2088.
(62) Bradforth, S. E.; Arnold, D. W.; Metz, R. B.; Weaver, A.; Neumark,
D. M. J. Phys. Chem. 1991, 95, 8066-8078.
(52) Shimanouchi, T. Tables of Molecular Vibrational Frequencies
Consolidated; National Standard Reference Data Service: Washington,
1972; Vol. I.

10868 J. Am. Chem. Soc., Vol. 120, No. 42, 1998
Chabinyc and Brauman
RO-‚HF when ROH is a stronger acid than HF and resembles
ROH‚F- when HF is the stronger acid. This relationship
between the structure of the intermediates and the thermody-
namic endpoints is not always the case, however. In the proton-
transfer reaction of H- with H2O, the most stable intermediate
has been shown to be H-‚H2O although H2O is a stronger acid
than H2.57,63 The H2‚OH- complex is observed to be in
equilibrium, but only comprises a small fraction of the popula-
tion.57 This suggests the structure of the intermediate is not
solely determined by the acidity difference.64 To understand
these effects we have attempted to fully characterize the
hydrogen-bonded complexes of HCF3 and several alkoxides.
Strength. The complexation energies between RO- and
HCF3 range from -21.1 to -23.5 kcal/mol (Table 4). These
energies are larger than those for typical ion-dipole complexes
(10-15 kcal/mol). For example, the binding energy of Cl- to
HCF3 is 16 kcal/mol.28 The fluoroform complexation energies
to alkoxides are almost the same as those for H2O to alkoxides.
However, the HCF3 complex energies are not as large as the
complexation energies of alcohol‚alkoxides or alcohol‚
phenylacetylide complexes which are near 25-28 kcal/mol.
Therefore the complexes, [ROHCF3]-, can be said to exhibit a
reasonable, but not exceptionally strong, hydrogen bond.
The complexation energies follow the basicity of the alkox-
ides similarly to the LFER seen for other hydrogen-bonded
complexes (Figure 2). The complexation energy of [EtOHCF3]-,
where the acidities of the two molecules are nearly matched
(∆∆H°acid ) 0.8 kcal/mol), does not show any enhanced
stabilization.13,14 The slope of the least-squares fit line is 0.5,
which is equivalent to that for the water‚alkoxide complexes
(0.5) and also similar to the slope (0.4) for alcohol‚alkoxide
complexes.17 The values for the fluoroform‚alkoxide complexes
cover a limited range, so we do not wish to make a quantitative
comparison of these slopes. We only note here that there is a
striking similarity despite the substantial structural differences
between fluoroform and an aliphatic alcohol.
Figure 3. Zero-point energies of H and D atoms in double well
potentials with varying barrier heights.
this case the fractionation factor would decrease and reach a
lower limit as the barrier becomes negligible compared to the
zero-point energy.20,65 The contributions due to other normal
modes, such as bending, are expected to be small compared to
that of the C-H stretch.66
The value of the fractionation factor is also affected by
symmetry aspects of the potential surface. For systems with a
single well potential, the fractionation factor should be nearly
constant for a series of complexes. McMahon and co-workers
have shown that the fractionation factors for [ROHF]- are nearly
constant over the range of alcohol acidities.58 This is consistent
with quantum calculations which predict that these systems have
a very flat potential surface near the hydrogen-bonded com-
plexes.61 For a double well potential, the fractionation factor
is expected to increase as the difference in the energy between
the wells increases. As the difference in energy between the
wells increases, a concomitant increase in barrier height is also
expected.67 This increase in barrier should cause the potential
of the H stretch to become more harmonic, and thus the
fractionation factor should approach the limiting value of 1. To
substantiate this qualitative prediction, model calculations were
performed to calculate the anharmonic H/D stretching frequency
in a double well (see Figure 4). These calculations indeed show
a trend to larger values of Φ with a larger difference in energy
between wells. Similar results have been obtained by Kreevoy
and Liang,20 and also Huskey.65
To analyze the possible structure of the [ROHCF3]- com-
plexes, we first examine the magnitudes of the measured
fractionation factors. If the PES has a high barrier and the
structure of the complex is RO-‚HCF3, we would expect the
fractionation factor relative to HCF3/DCF3, Φ, to have a value
near 1.0. If the structure of the complex is ROH‚CF3- and the
barrier is high we would expect the fractionation factor relative
to ROD/ROH, Φ′, to have a value near 1.0. For a low barrier
complex, Φ and Φ′ should both be less than 1.0 as the difference
Structure. To probe the structure of the [ROHCF3]-
complexes we have carried out deuterium fractionation experi-
ments. The value of the fractionation factor, Φ, arises from
the differences of the zero-point energies (ZPE) in the separated
reactants and in the complex.20 The enthalpy change of the
exchange reaction (eq 8) is approximately equal to the change
in ZPE (eq 19).
∆H° = (ZPE[RODCF ] - ZPE[ROHCF - ) - (ZPE
-
-
]
3
[DCF3]
3
ZPE[HCF ]) (19)
3
The difference in zero-point energy of the C-H bond in the
free molecule and in the complex can be understood from a
simple one-dimensional model corresponding to the proton-
transfer reaction coordinate.20 If the barrier to proton transfer
is high, the potential of the stretch is nearly harmonic and
therefore the vibrational frequency is not strongly perturbed in
the complex. The fractionation factor in this case would be
1.0. As the barrier between the two complexes decreases, the
stretching frequency becomes more anharmonic and the differ-
ence in ZPE between H and D decreases51,65,20 (Figure 3). In
(63) Miller, T. M.; Viggiano, A. A.; Stevens-Miller, A. E.; Morris, R.
A.; Henchman, M.; Paulson, J. F.; Doren, J. M. V. J. Chem. Phys. 1994,
100, 5706-5714.
(64) There are similar interesting observations involving proton-bound
cluster ions where the structure is determined by interactions between
multiple protonation sites of differing basicity. See: Bo¨hringer, H.; Arnold,
F. Nature 1981, 290, 321-322. Graul, S. T.; Squires, R. R. Int. J. Mass.
Spectrom. Ion Proc. 1989, 94, 41-61.
(65) Huskey, W. P. J. Am. Chem. Soc. 1996, 118, 1663-1668.
(66) Westheimer, F. H. Chem. ReV. 1961, 61, 265-273.
(67) Marcus, R. A. Annu. ReV. Phys. Chem. 1964, 15, 155-196.

Complexes of RO- + HCF3
J. Am. Chem. Soc., Vol. 120, No. 42, 1998 10869
Figure 5. Calculated potential energy surface for the reaction of EtO-
with HCF3 at the MP2/6-311++G**//HF/6-311++G** level of theory.
Relative energies include corrections from zero-point energies.
would expect the structure of the i-PrO- complex to resemble
i-PrO-‚HCF3, as isopropyl alcohol is more acidic than fluoro-
form by 1.1 kcal/mol.28 Since ethanol and fluoroform have
nearly matched acidities (HCF3 is 0.8 kcal/mol more acidic,
based on ∆∆H°acid), we might assume that the two structures,
EtO-‚HCF3 and EtOH‚CF3-, have similar energies. However,
Φ is smaller for [MeOHCF3]- than for [EtOHCF3]-. Our model
predicts that the fractionation factor should reach a minimum
when the wells have equal energies and then begin to increase
as the difference in energy increases. This suggests that the
two structures, EtO-‚HCF3 and EtOH‚CF3-, do not have equal
energies and that EtO-‚HCF3 is more stable. The observation
that the fractionation factor is lowest for [MeOHCF3]-, where
the acidities are not matched, suggests that the two structures,
MeO-‚HCF3 and MeOH‚CF3-, have similar energies (or at least
are closer in energy than for the EtO- complex). We believe
that the continual decrease in fractionation factor with difference
in acidity indicates that all of the complexes have the RO-‚HCF3
structure.
Figure 4. Fractionation factors calculated for H and D in asymmetric
double-well potentials relative to HCF3/DCF3 with varying differences
in energy between the two minima (∆Ewell).
in zero-point energies of the protiated and deuterated complex
should be smaller than the difference in zero-point energies
between HCF3 and DCF3 or ROD and ROH.
The barriers on the PES’s for the proton-transfer reactions
between RO- and HCF3 are small, as indicated by the high
efficiencies of these reactions.68 The fractionation factors should
therefore reflect the limiting case of the low barrier. The values
of Φ, the fractionation factor relative to HCF3/DCF3, for the
[ROHCF3]- complexes are given in Table 5. All values are
less than 1.0 as expected for a low-barrier PES. The values of
Φ′, the fractionation factor relative to ROH/ROD, are also given
in Table 5. The values of Φ′ are lower than those for Φ because
the difference in zero-point energy between ROH and ROD is
larger than that for HCF3 and DCF3. These values can be
compared to those for MeO-‚HOMe69,70 and EtO-‚HOEt71
which are both near 0.40. The absolute magnitude of the
fractionation factor does not indicate the structure of the ion-
molecule complex directly in our case.
The trend in the fractionation factors suggests that the
structure of these three complexes is similar. The experimental
values of Φ for the [ROHCF3]- complexes decrease with
decreasing acidity of the alcohol. This behavior is consistent
with a model of the proton-transfer potential surface as a double
well; as the difference in energy between the wells decreases,
the difference in zero-point energies for H and D in the complex
decreases causing the fractionation factor to decrease (see Figure
4). The lowest fractionation factor should occur when the two
wells have the same energy, that is, the PES where the vibration
is the most anharmonic. The fractionation factor should then
begin to increase as the difference in energy between wells
increases (i.e., the plot should be nearly symmetric around
matched energies).65 The agreement with our model calcula-
tions is not expected to be quantitative due to the simplicity of
our model.72 From simple considerations of the acidities, we
To further investigate the structure of these complexes we
performed ab initio calculations to map the proton-transfer
potential energy surfaces. The calculated ab initio potential
energy surface for the proton-transfer reaction between EtO-
and HCF3 is shown in Figure 5. Although the proton transfer
is nearly thermoneutral, the complexes have different stabilities.
The complex with the structure EtO-‚HCF3 is calculated to be
more stable than EtOH‚CF3- by 3.0 kcal/mol. For the proton-
transfer reaction of MeO- with HCF3, the MeO-‚HCF3 complex
is more stable by 1.4 kcal/mol despite MeOH being 3.1 kcal/
mol less acidic than HCF3 (see Figure 6). With this energy
difference reasonable populations of both structures will be
present, but MeO-‚HCF3 should be dominant. The i-PrO-‚
HCF3 complex is calculated to be more stable than i-PrOH‚
-
CF3 by 4.8 kcal/mol. We believe these calculations are
consistent with all the experimental data.73 The calculated well
depths for the RO-‚HCF3 complexes agree with our experi-
-
mentally determined values better than those for ROH‚CF3
.
The calculated barriers are low, consistent with the magnitude
of the fractionation factors and kinetic data. The trend in the
fractionation factors suggests a changing double well potential
(68) We have measured the rate constant for the EtO- + HCF3 proton-
transfer reaction. Its value is 4.11 × 10-10 cm3 molecule-1 s-1 and the
collision capture rate from the Su-Chesnavich model is 1.65 × 10-9 cm3
molecule-1 s-1. S. L. Craig and M. L. Chabinyc, unpublished results.
(69) Weil, D. A.; Dixon, D. A. J. Am. Chem. Soc. 1985, 107, 6859-
6865.
(72) The one-dimensional model may not produce quantitatively accurate
vibrational frequencies, refs 50 and 51, and also neglects the other modes,
such as bending, that are affected by isotopic substitution. However, these
effects are expected to be small compared to the shift of the H/D stretching
frequency, ref 66.
(70) Barlow, S. E.; Dang, T. T.; Bierbaum, V. M. J. Am. Chem. Soc.
1990, 112, 6832-6838.
(71) Ellenberger, M. R.; Farneth, W. E.; Dixon, D. A. J. Phys. Chem.
1981, 85, 4-7.
(73) Fluoroform is calculated to be less acidic than ethanol by 1.0 kcal/
mol at this level of theory compared to the experimental difference of -0.8
kcal/mol. This deficiency should not affect the conclusions from the
calculated PES’s.

10870 J. Am. Chem. Soc., Vol. 120, No. 42, 1998
Chabinyc and Brauman
away from the carbon into the fluorine atoms. As the hydrogen
bond donor approaches the ion, it sees a lower effective charge
-
in the case of CF3 than for the alkoxides. This can also be
understood based on the larger electron affinity (∼2-3 kcal/
mol) of CF3 relative to the alkoxy radicals.28 Because CF3- is
able to internally stabilize its charge effectively by delocaliza-
tion, it is not stabilized as much as an alkoxide by a hydrogen
bond. Thus, the hydrogen bond accepting ability of the anions
is different.76 To examine the donating abilities of HCF3 and
ROH, we can consider the stability of ion-molecule complexes
where the molecules are expected to be the hydrogen bond
donor. EtOH is a better hydrogen bond donor than HCF3 by
4.1 kcal/mol, based on their F- affinities,54 and the EtO-‚HCF3
complex readily reacts with EtOH to form EtO-‚HOEt. On
the basis of EtOH we can characterize the alcohols as generally
Figure 6. Calculated potential energy surface for the reaction of MeO-
with HCF3 at the MP2/6-311++G**//HF/6-311++G** level of theory.
Relative energies include corrections from zero-point energies.
-
being better hydrogen bond donors than HCF3. Even so, CF3
is such a poor hydrogen bond acceptor that the structure
EtO-‚HCF3 is calculated to be 3 kcal/mol more stable than
EtOH‚CF3-. In short, in contrast to alcohols which are good
hydrogen bond donors and whose conjugate bases are good
acceptors, fluoroform is a reasonably good donor, but its
conjugate base is a poor acceptor.
Our results indicate that care must be taken in analyzing data
from studies of hydrogen-bonded complexes. The ability to
donate and accept hydrogen bonds may not be a simple function
of the acid-base properties of the molecule. If generalizations
are to be made from ion-molecule complexation energies about
the hydrogen bond accepting and donating abilities of ions and
neutrals, the structure of the complex must be known.
energy surface as seen in the calculations. RO-‚HCF3 is the
only structure consistent with all the data.
The calculated barrier in these proton-transfer reactions is
small in agreement with the observation of rapid proton-transfer
kinetics in these systems.68 The barrier is high enough for these
reactions, however, so that the ground-state wave function of
the minimum energy complex would be expected to be localized
in the RO-‚HCF3 well, based on our estimates of the C-H
stretching frequency. The PES for the reaction of H- with H2O
has been widely studied,57,74 and is calculated to have a small
barrier, about 3.5 kcal/mol higher than the minimum energy
complex. In this case, 2-D discrete variable representation
calculations suggest the wave function is localized in the
H-‚H2O well.57 The complexes studied here should behave
Summary
We have studied the hydrogen-bonded complexes on the
proton-transfer surface of several alkoxides with fluoroform.
The ion-molecule complexes are strongly bound and are best
described as hydrogen bonded. The structures of the complexes
have been studied by isotopic fractionation equilibria and by
ab initio calculations. Our data show that the structure of these
complexes does not reflect the difference in acidity of the two
acids. This is in contrast to the usual assumption that acidity
determines structure. The difference can be understood in terms
of electrostatics in the isolated acid and base and their influence
on the structure of the complexes.
-
similarly. In contrast, while the ROH‚CF3 structure is
calculated to be a stationary point on the potential surface, the
zero-point energy is high enough that hydrogen should be able
to freely access the RO-‚HCF3 structure.
Conclusions. We believe our data require the structure of
all three (MeO-, EtO-, i-PrO-) proton-transfer complexes to
be RO-‚HCF3. For the case of nearly matched acidities (EtO-)
the two intermediates have substantially different energies. The
MeO-‚HCF3 complex is especially striking, because this is a
case where the acidity difference between the reactants does
not determine the structure of the hydrogen-bonded complexes
in the reaction. In other words, although HCF3 is a stronger
acid than MeOH, this acidity difference is not reflected in the
structure of the hydrogen-bonded complex.
Acknowledgment. We are grateful for support from the
National Science Foundation. M.L.C. thanks the NSF and the
ACS Division of Organic Chemistry for fellowship support.
Our results can be rationalized by examining the ability of
the ions and neutrals to respectively accept and donate hydrogen
bonds. The electrostatic surface of the alkoxides shows that
the charge is localized mainly on the oxygen.75 Conversely,
JA9817592
(75) Wiberg, K. B. J. Am. Chem. Soc. 1990, 112, 3379-3385.
(76) Electrostatics do not provide a quantitative means to calculate the
hydrogen bond strength since the charge distributions are certainly perturbed
at the close range of interaction of the ion and neutral. They do, however,
provide a useful model to interpret the energy difference between the two
isomeric hydrogen bonded complexes.
-
CF3 gains its stability as an anion by polarizing the charge
(74) Chalasinski, G.; Kendall, R. A.; Simons, J. J. Chem. Phys. 1987,
87, 2965-2975.