Intrinsic acidity and basicity of 2,2,2-trifluoroethanethiol. The first experimental and theoretical study

Article abstract of DOI:10.1021/jo960524a

The gas phase acidity and basicity of 2,2,2-trifluoroethanethiol (TFET), i.e., the standard Gibbs energy changes for the following two reactions have been determined by means of Fourier transform ion cyclotron resonance spectroscopy: CF₃CH₂SH(g) → CF₃CH₂S^-(g) + H⁺(g) and CF₃CH₂SH₂⁺ (g) → CF₃CH₂SH(g) + H⁺(g). Also determined were the equilibrium constants for the 1:1 associations in dilute solution between TFET and pyridine N-oxide, 3,4-dinitrophenol (both in cyclohexane), and molecular iodine (in tetrachloromethane). Quantum-mechanical treatments at the G2(MP2) level were carried out on TFET, 2,2,2-trifluoroethanol, ethanethiol, and ethanol as neutral, protonated, and deprotonated species. Topological analyses of the charge densities and the Laplacians thereof were performed on all of them. This combination of experimental and theoretical information leads to a vastly enlarged view of structural effects on the reactivity of alcohols and thiols as well as to a satisfactory rationalization of the reactivity of TFET.

Full text of DOI:10.1021/jo960524a

J . Org. Chem. 1996, 61, 5485-5491
5485
In tr in sic Acid ity a n d Ba sicity of 2,2,2-Tr iflu or oeth a n eth iol. Th e
F ir st Exp er im en ta l a n d Th eor etica l Stu d y†
M. T. Molina
Instituto de Qu ı´ mica M e´ dica, C.S.I.C. c/ J uan de la Cierva 3, E-28006 Madrid, Spain
W. Bouab and M. Esseffar
D e´ partement de Chimie, Facult e´ des Sciences (Semlalia), Marrakesh, Morocco
M. Herreros, R. Notario, and J .-L. M. Abboud*
Instituto de Qu ı´ mica F ı´ sica “Rocasolano”, C.S.I.C. c/ Serrano 119, E-28006 Madrid, Spain
O. M o´ and M. Y a´ n˜ ez*
Departamento de Qu ı´ mica, C-IX, Universidad Aut o´ noma de Madrid,
E-28049, Cantoblanco (Madrid), Spain
Received March 19, 1996X
The gas phase acidity and basicity of 2,2,2-trifluoroethanethiol (TFET), i.e., the standard Gibbs
energy changes for the following two reactions have been determined by means of Fourier transform
-
+
+
ion cyclotron resonance spectroscopy: CF
3
CH
2
SH(g) f CF
3
CH
2
S (g) + H (g) and CF
3 2 2
CH SH (g)
+
f CF CH SH(g) + H (g). Also determined were the equilibrium constants for the 1:1 associations
3 2
in dilute solution between TFET and pyridine N-oxide, 3,4-dinitrophenol (both in cyclohexane),
and molecular iodine (in tetrachloromethane). Quantum-mechanical treatments at the G2(MP2)
level were carried out on TFET, 2,2,2-trifluoroethanol, ethanethiol, and ethanol as neutral,
protonated, and deprotonated species. Topological analyses of the charge densities and the
Laplacians thereof were performed on all of them. This combination of experimental and theoretical
information leads to a vastly enlarged view of structural effects on the reactivity of alcohols and
thiols as well as to a satisfactory rationalization of the reactivity of TFET.
-
+
In tr od u ction
AH(g) a A (g) + H (g)
(2)
We have been involved for several years in the quan-
titative study of the basicities and acidities of homologous
series of oxygen and sulfur compounds, using both
theoretical and experimental methods. Following are the
basicity and acidity criteria most frequently investi-
gated:
alcohols and a few alkanethiols are presently available.6,7
2) The equilibrium constants, K , for the formation of
(
c
the 1:1 hydrogen-bonded complexes between selected
8
hydrogen bond (HB) donors, AH, and oxygen (e.g., ethers
1
9
,10
and alcohols ) and sulfur n-bases (e.g., thiols and
thioethers11 and thiocarbonyl compounds
11,12
) in highly
(1) The proton affinities (PAs) and gas phase basicities
dilute solutions in cyclohexane or tetrachloromethane
solvents S (reaction 3).
Also established was a preliminary quantitative rank-
ing of HB acidities of monomeric alcohols and a few
alkanethiols through the determination of the equilibri-
(GBs) of these compounds. For a given base B in the gas
phase, PA(B) and GB(B) respectively stand for the
standard enthalpy and Gibbs energy changes for reaction
2
Extensive sets of PA and GB values for thiocarbonyl3
1
.
+
+
BH (g) f B(g) + H (g)
∆H°H
+
(g), ∆G°H
+
(g) (1)
(4) Lias, S. G.; Bartmess, J . E.; Liebman, J . F.; Holmes, J . L.; Levin,
R. D.; Mallard, W. G. J . Phys. Chem. Ref. Data, Suppl. 1 1988, 17,
1
-861.
and carbonyl2,4,5 compounds have been determined. Val-
ues are also available for several alkanols and alkanethi-
ols.2
(5) Homan, H.; Herreros, M.; Notario, R.; Abboud, J .-L. M.; Essefar,
M.; M o´ , O.; Y a´ n˜ ez, M.; Foces-Foces, C.; Ramos-Gallardo, A.; Mart ´ı nez-
Ripoll, M.; Vegas, A.; Molina, M. T.; Casanovas, J .; J im e´ nez, P.; Roux,
M. V.; Turri o´ n, C. Submitted for publication.
,4
(
6) NIST Standard Reference Database 19B, NIST Negative Ion
In a similar vein, the gas phase acidity of neutral
species AH is defined as the standard Gibbs energy
change for reaction 2. Values of gas phase acidities for
Energetics Database, Computerized Version 3.01; National Institute
of Standards and Technology, Gaithersburg, MD 20899; 1993.
(7) (a) Taft, R. W.; Taagepera, M.; Abboud, J .-L. M.; Wolf, J . F.; De
Frees, D. J .; Hehre, W. J .; Bartmess, J . E.; McIver, R. T., J r. J . Am.
Chem. Soc. 1978, 100, 7765-7767. (b) Abboud, J .-L. M.; Elguero, J .;
Liotard, D.; Essefar, M.; El Mouhtadi, M.; Taft, R. W. J . Chem. Soc.,
Perkin Trans. 2 1990, 565-571; 1279.
(8) Abboud, J .-L. M.; Bellon, L. Ann. Chim. (Paris) 1970, 5, 63-74.
(9) Abboud, J .-L. M.; Sra ¨ı di, K.; Guih e´ neuf, G.; Negro, A.; Kamlet,
M. J .; Taft, R. W. J . Org. Chem. 1985, 50, 2870-2873.
(10) Laurence, C.; Berthelot, M.; Helbert, M.; Sra ı¨ di, K. J . Phys.
Chem. 1989, 93, 3799-3802.
(11) Abboud, J .-L. M.; Roussel, C.; Gentric, E.; Sra ¨ı di, K.; Lauransan,
J .; Guih e´ neuf, G.; Kamlet, M J .; Taft, R. W. J . Org. Chem. 1988, 53,
1545-1550.
(12) Laurence, C.; Berthelot, M.; Le Questel, J .-Y.; El Ghomari, M.
J . J . Chem. Soc., Perkin Trans. 2 1995, 2075-2079.
†
This work is dedicated in memoriam to Professor Robert W. Taft.
Abstract published in Advance ACS Abstracts, J uly 15, 1996.
X
(
1) (a) Gal, J .-F.; Maria, P.-C. Prog. Phys. Org. Chem. 1990, 17, 159-
38. (b) Maria, P.-C.; Gal, J .-F.; de Franceschi, J .; Fargin, E. J . Am.
Chem. Soc. 1987, 109, 483-492.
2) Lias, S. G.; Liebman, J . F.; Levin, R. D. J . Phys. Chem. Ref. Data
984, 13, 695-808.
3) (a) Abboud, J .-L. M.; M o´ , O.; de Paz, J . L. G.; Y a´ n˜ ez, M.; Bouab,
2
(
1
(
W.; Mokhlisse, R.; Ballesteros, E.; Herreros, M.; Homan, H.; L o´ pez-
Mardomingo, C.; Notario, R. J . Am. Chem. Soc. 1993, 115, 12668-
1
2676. (b) Decouzon, M.; Exner, O.; Gal, J .-F.; Maria, P.-C. J . Phys.
Org. Chem. 1994, 71, 511-517.
S0022-3263(96)00524-5 CCC: $12.00 © 1996 American Chemical Society

5
486 J . Org. Chem., Vol. 61, No. 16, 1996
Molina et al.
(7)
S
\
δ∆GH (g) ) -RT ln Kp
+
AH + B
um constants, K
Kc
(3)
In every case, the reversibility of reactions 5 and 6 was
confirmed by means of double resonance experiments.
c
, for the formation of the 1:1 complexes
between these compounds in pyridine N-oxide (PyO) in
In the case of equilibrium 6, AHref and CF
3
CH
2
SH were
dilute solution in cyclohexane at 23.3 °C, (reaction 4).1
3,14
-
initially deprotonated by iso-C H11O , generated in situ by
5
electron ionization of isoamylnitrite.
AH + PyO a AH···OPy
AH ) ROH, RSH)
Kc
(4)
The pressure readings for the various neutral reagents, as
determined by the Bayard-Alpert gauge of the FTICR spec-
trometer, were corrected by means of the appropriate calibra-
tion coefficients obtained for each reagent by plotting the
readings of the ion gauge against the absolute pressures
provided by a capacitance manometer (Baratron, MKS).
Routinely, three calibration runs were performed on each gas.
The reproducibility of the calibration can be estimated at ∼3%.
The formal relative sensitivity of the ion gauge is ∼10%.
(
Our interest in the title compound originates in the
fact that this is one of the simplest thiols in which the
sulfhydryl group is close to a substituent exerting a large
positive field effect. This opens up the possibility of
vastly enlarging the experimentally accessible range of
intrinsic acidities and basicities for thiols.
c
K values for the 1:1 associations between TFET and 3,4-
dinitrophenol, PyO, and I
2
in solution were determined by
A more comprehensive appraisal of structural effects
on the acidities and basicities of sp oxygen and sulfur
acids and n-bases is thus possible. CF CH SH has
3 2
recently become commercially available. Hence, infor-
mation on its HB acidity and basicity is likely to be quite
interesting for its potential use as a solvent.
means of UV-visible spectroscopy. Experiments were per-
formed on a Cary 219 spectrophotometer using matched 1-cm
silica window cells in the case of the charge-transfer complexes
with iodine. For the associations with 3,4-dinitrophenol and
PyO, 10-cm matched silica window cells were used. The
experimental methods have already been described in detail:
for reactions 3 and 4, see refs 9 and 13. For charge-transfer
3
In this work, the intrinsic (gas phase) acidity and
20,21
complexation, Drago’s method
was used. The equilibrium
basicity of CF
3
CH
2
SH were determined by means of
constants we report are the averages of four independent
measurements. Relative uncertainties are reasonably large,
Fourier transform ion cyclotron resonance (FTICR) spec-
troscopy. The HB acidity and basicity of this compound
were probed through the experimental determination of
∼
15%. This is not surprising, on account of the very small
values of the various constants.
The purifications of 3,4-dinitrophenol and PyO have already
been described in refs 9 and 13, respectively. Solvents of
spectrograde quality were refluxed over and distilled from a
the equilibrium constants, K
c
, for its associations with
3
,4-dinitrophenol (HB donor) and PyO (HB acceptor).
Also determined was the equilibrium constant pertaining
to the formation of the 1:1 charge-transfer complex1
with molecular iodine in tetrachloromethane solvent at
4 4
sodium-potassium alloy (cyclohexane) or P O10 (CCl ). TFET
5-18
(Aldrich) was carefully distilled immediately prior to use.
To our knowledge, no precise information is available so far
on the physiological effects of 4. Thus, appropriate care should
be exerciced when handling this compound.
2
5.0 °C. This is intended to provide an estimate of its
electron-donating ability.
Using the gas phase data as a vantage point, an ab-
initio study of ethanol (1), 2,2,2-trifluoroethanol (2),
ethanethiol (3), and 2,2,2-trifluoroethanethiol (4) as
neutral, protonated, and deprotonated species was per-
formed at the G2(MP2) level.
Com p u ta tion a l Deta ils
To obtain the most reliable energetics possible, the
proton affinities and the deprotonation energies of the
neutrals under investigation were obtained in the frame-
work of the G2 theory.22 Since in its original formulation
this theoretical scheme would be prohibitively expensive
for systems of this size, we shall use the more economic
G2(MP2) formalism, which has proved to yield proton
Exp er im en ta l Section
The gas phase basicity and acidity were determined from
equilibrium proton-transfer reactions conducted in a modified
Bruker CMS-47 FT ICR mass spectrometer.19 Working condi-
affinities not significantly different from those obtained
3
a
tions were similar to those already described. The average
23
at the G2 level.
Although in the G2 formalism the
cell temperature is ∼333 K. These FTICR measurements
geometries are obtained at the MP2/6-31G* level, in the
present study we shall employ MP2/6-311+G(d,p) opti-
mized geometries, since it is well established that, to
+
(g), for the
H
provide the standard Gibbs energy change, δ∆G
proton-exchange reactions 5 and 6:
properly reproduce anionic systems, the inclusion of
+
+
CF CH SH (g) + Bref(g) a CF CH SH(g) + B H (g) (5)
3
2
2
3
2
ref
24a
diffuse functions in the basis set is crucial.
Although
this requirement does not apply to protonated species,
for the sake of consistency we have used the same scheme
for neutrals, cations, and anions. The calculations were
performed using the Gaussian 94 package of programs.
The geometries were initially optimized at the HF/6-
-
-
CF CH SH(g) + A (g) a CF CH S (g) + AHref(g) (6)
3
2
ref
3
2
AHref and Bref are reference acids and bases, respectively. For
each of these equilibria,
24b
3
11+G(d,p) level. The corresponding harmonic vibra-
(
13) Frange, B.; Abboud, J .-L. M.; Benamou, C.; Bellon, L. J . Org.
Chem. 1982, 47, 4553-4557.
14) Abboud, J .-L. M.; Sra ¨ı di, K.; Abraham, M. H.; Taft, R. W. J .
tional frequencies were evaluated at the same level of
accuracy by means of analytical second derivatives
techniques and used to characterize the stationary points
(
Org. Chem. 1990, 55, 2230-2232.
(
15) Mulliken, R. S. J . Am. Chem. Soc. 1950, 72, 600-608.
(16) Mulliken, R. S.; Person, W. B. Molecular ComplexessA Lecture
and Reprint Volume; Wiley-Interscience: New York, 1969.
17) Briegleb, G. Elektronen-Donator-Acceptor Komplexe; Springer:
Berlin, 1961.
(20) Rose, N. J .; Drago, R. S. J . Am. Chem. Soc. 1959, 81, 6138-
6141.
(21) Guih e´ neuf, G.; Abboud, J .-L. M.; Bouab, W. Can. J . Chem. 1987,
65, 2106-2108.
(22) (a) Curtiss, L. A.; Raghavachari, K.; Trucks, G. W.; Pople, J .
A. J . Chem. Phys. 1991, 94, 7221-7226. (b) Curtiss, L. A.; Raghava-
chari, K.; Pople, J . A. J . Chem. Phys. 1993, 98, 1293-1298.
(23) Smith, B. J .; Radom, L. Chem. Phys. Lett. 1994, 231, 345-351.
(
(
18) Foster, R. Organic Charge-Transfer Complexes; Academic
Press: London, 1969.
(19) Fourier Transform Mass Spectrometry, Evolution, Innovation
and Applications; Buchanan, M. V., Ed.; ACS Symposium Series 359;
American Chemical Society: Washington, DC, 1987.

Acidity/Basicity of 2,2,2-Trifluoroethanethiol
J . Org. Chem., Vol. 61, No. 16, 1996 5487
of the potential energy as local minima and to evaluate
the zero-point energies, which were scaled by the empiri-
Ta ble 1. Exp er im en ta l Deter m in a tion of th e Ga s P h a se
Ba sicity of CF 3CH2SH
2
4c
∆∆GH+(std)a δ∆GH+(g)a ∆∆GH+ a
∆GBa
cal factor 0.893 298.
These geometries were then
reference
refined at the MP2/6-311+G(d,p) level.
27.5b
Cl3CCN
Cl3CCH2OH
CF3CO2CH3
-0.82
1.11
3.22
26.6(8)
27.1(1) 26.9
27.0(2) (SD = 0.2)
c
To investigate the charge redistributions undergone by
the neutral systems upon protonation and deprotonation,
we have carried out a topological analysis of the electronic
26.0
23.8b
a
All values in kcal/mol (1 cal ) 4.184 J ). b See text. c Deter-
mined in this work. Cl3CCH2OH was found to be 1.50 kcal/mol
more basic than Cl3CCN and 2.20 kcal/mol less basic than
CF3CO2CH3.
2
25
charge density, F, and its Laplacian, ∇ F. In general,
bond reinforcements and bond activations imply signifi-
cant changes of the charge density and its Laplacian at
the corresponding bond critical points, defined as points
where the charge density is minimum along the bond
path and maximum in the other two directions. Bader
and co-workers have shown that the existence of these
critical points is associated with the existence of a bond.
Exp er im en ta l Resu lts
(1) Ga s P h a se Ba sicities. Table 1 presents the
results of proton-transfer equilibria (5) between 4 and a
series of standard reference bases. The values of δ∆GH (g)
given in this table are defined by means of eq 7.
All gas phase basicities, ∆GB, are referred to ammonia.
+
2
More importantly, the values of F and ∇ F at the bond
critical points (bcps) permit classification of the nature
of the interaction as covalent, ionic, etc. and may be used
as a quantitative measure of the strength of the bond,
since negative values of the Laplacian indicate that
charge density increases in that region, while positive
values are generally associated with charge depletion.
Hence, an increase (decrease in the absolute value) of
the (negative) Laplacian indicates2 that the bond is
somewhat activated, while an increase is associated with
a bond reinforcement process. Another index of interest
which may be defined in this topological analysis is the
Thus, with respect to this reference, ∆GB(4) ) -∆∆GH (g)
+
for reaction 8:
CF CH SH(g) + NH +(g) a
3
2
4
+
CF CH SH (g) + NH (g)
∆∆GH
+
(g) (8)
3
2
2
3
6a
∆
∆GH+(g) is the average of the ∆∆G values obtained
through eq 9:
∆
∆G ) δ∆GH
+
(g) + ∆∆GH
+
(std)
(9)
elipticity of the bond, defined as ꢀ ) 1 - λ
1
/λ
2
, where λ
1
and λ are the two negative values of the Hessian of F,
2
where ∆∆GH+(std) pertains to reaction 10:
evaluated at the bcp. Obviously, these two values will
be identical only for those bonds which, as the single and
the triple bonds, have cylindrical symmetry, for which
the elipticity will be zero.
+
+
B (g) + NH (g) a B H (g) + NH (g) (10)
ref
4
ref
3
H
The values of ∆∆G (std) used in this work have been
determined in Prof. Taft’s laboratory and are given in
refs 2 and 4.
+
This topological analysis is carried out on the wave
function correct to first order, to take explicitly into
account the electron correlation effects. For this purpose,
we have used the AIMPAC27 series of programs.
Since predicted structural variations can be qualita-
tively rationalized in terms of the bonding indexes, such
as hybridization parameters, we have also carried out a
natural bond order analysis28 at both the HF and MP2
levels. Values derived in this way are compatible with
the classical notion of hybridization, as introduced by
Pauling.
Inasmuch as the entropy change for reaction 1 cannot
be obtained directly from the FTICR experiments, we
have used the value computed at the 6-31G**//6-31G**
-
1
-1
level, 24.2 cal mol
K . This value combines the
+
contribution from S(H ) as determined from the Sackur-
Tetrode equation with those from ∆S(rot) and ∆S(vib)
for the molecules and ions as determined by means of
the partition functions calculated at the 6-31G**//6-31G**
level. The standard state is 298 K and 1 atm (0.1 MPa).
Combining these values with the most recent values
(24) (a) Hehre, W. J .; Radom, L.; Schleyer, P. v. R.; Pople, J . A. Ab
-
1
initio Molecular Orbital Theory; J ohn Wiley & Sons: New York, 1986;
pp 86-88. (b) Frisch, M. J .; Trucks, C. W.; Schlegel, H. B.; Cill, D. M.
W.; J ohnson, B. G.; Robb, M. A.; Cheeseman, J . R.; Keith, T.; Petersson,
G. A.; Montgomery, J . A.; Raghavachari, K.; Al- Laham, M. A.;
Zakrzewski, V. G.; Ortiz, J . V.; Foresman, J . B.; Peng C. Y.; Ayala, P.
Y.; Chen, W.; Wong, M. W.; Andr e´ s, J . L.; Replogle, E. S.; Gomperts,
R.; Martin, R. L.; Fox, D. J .; Binley, J . S.; Defrees, D. J .; Baker, J .;
Stewart, J . P.; Head-Gordon, M.; Gonz a´ lez, C.; Pople, J . A. Gaussian
of GB(NH
3 3
) and PA(NH ), 195.3 and 203.5 kcal mol ,
29
respectively, we obtain GB(4) ) 168.4 ( 0.2 and PA(4)
-
1
)
175.6 ( 0.2 kcal mol
.
(2) Ga s P h a se Acid ity. The Gibbs energy change for
reaction 3, ∆Gacid(g)(av), is the average of the ∆Gacid(g)
values obtained through eq 11:
9
4, Revision B.3; Gaussian Inc.: Pittsburgh, PA, 1995. (c) Pople, J .
A.; Schlegel, H. B.; Krishnan, R.; Defrees, D. J .; Binkley, J . S.; Frisch,
M. J .; Whitesode, R. A.; Hout, R. F.; Hehre, W. J . Int. J . Quantum
Chem. Symp. 1981, 15, 269.
∆
Gacid(g) ) δ∆Gacid(g) + ∆Gacid(std)
(11)
(12)
(
25) (a) Bader, R. W. F.; Ess e´ n, H. J . Chem. Phys. 1984, 80, 1943-
wherein ∆Gacid(std) pertains to reaction 12:
1
960. (b) Bader, R. W. F.; MacDougall, P. J .; Lau, C. D. H. J . Am.
Chem. Soc. 1984, 106, 1594-1605. (c) Bader, R. W. F. Atoms in
Molecules. A Quantum Theory; Oxford University Press: New York,
AHstd(g) a A- (g) + H (g)
+
std
1
990.
26) (a) Alcam ´ı , M.; M o´ , O.; Y a´ n˜ ez, M.; Abboud, J .-L. M.; Elguero,
(
J . Chem. Phys. Lett. 1990, 172, 471-477. (b) Tipping, A. E.; J im e´ nez,
P.; Ballesteros, E.; Abboud, J .-L. M.; Y a´ n˜ ez, M.; Essefar, M.; Elguero,
J . J . Org. Chem. 1994, 59, 1039-1046. (c) Esseffar, M.; El Mouhtadi,
M.; L o´ pez, V.; Y a´ n˜ ez, M. J . Mol. Struct. (Theochem) 1992, 255, 393.
Experimental results are summarized in Table 2. From
these results, we obtain ∆Gacid(4) ) 335.6 ( 0.3 kcal
mol . The corresponding standard enthalpy change is
∆Hacid(4) ) 342.4 ( 0.3 kcal mol using the computed
see above) entropy change ( 22.8 cal mol
-
1
-
1
(
27) The AIMPAC program was kindly provided by J . R. Cheseman
and R. W. F. Bader.
28) (a) Carpenter, J . E.; Weinhold, F. J . Mol. Struct. (Theochem)
-1
-1
(
K ).
(
1
7
1
988, 169, 41-62. (b) Reed, A. E.; Weinhold, F. J . Chem. Phys. 1983,
8, 4066-4073. Reed, A. E.; Curtiss, L. A.; Weinhold, F. Chem. Rev.
988, 88, 899-926.
(29) Szulejko, J . E.; McMahon, T. B. J . Am. Chem. Soc. 1993, 115,
7839-7847.

5
488 J . Org. Chem., Vol. 61, No. 16, 1996
Molina et al.
Ta ble 2. Exp er im en ta l Deter m in a tion of th e Ga s P h a se
Acid ity of CF 3CH2SH
latter, the SH hydrogen is gauche with respect to the
substituted carbon atom, while in the former it is anti.
This seems to point to the existence of an intramolecular
hydrogen bond in 4 that does not appear in 2. However,
no bcp was found between the SH hydrogen and the
closest fluorine atom. Hence, very likely the preference
for a gauche conformation in 4 is simply the result of the
attractive electrostatic interaction between the positively
charged H(S) hydrogen and the closest negatively charged
fluorine atom. This will be particularly favored in 4 due
to the small value of the CSH bond angle and the large
S-H bond length.
∆
Gacid-
(std)
δ∆Gacid-
∆Gacid-
∆Gacid(g)
a,b
a
a
(av)a
reference
(g)
(g)
â-naphthol
336.5
336.2
333.1
-0.54
-0.68
2.35
335.9(6)
335.5(2)
335.4(4)
4
-nitroaniline
335.6
(SD ) 0.3)
C6H5CO2H
a
All values in kcal/mol. b From ref 6.
Notice that the uncertainties reported for acidities and
basicities are rather small, as far as the overlaps with
the reference compounds are concerned. The “absolute”
values are obviously less precise, perhaps by as much as
Due to the charge transfer that takes place from the
neutral to the incoming proton, protonation of ethanol
1
or 2 kcal mol-1. The quality is most certainly not
significantly worse than this, because of the multiple
overlaps involved in the construction of the acidity and
basicity scales.
implies a significant lengthening of the C-O bond and a
26b
shortening of the C-C linkage.
However, when the
methyl group is substituted by a trifluoromethyl group,
the C-C bond slightly lengthens. This simply reflects
the greater ability of the methyl group to accommodate
a positive charge.
Since sulfur is less electronegative and more polariz-
able than oxygen, the aforementioned geometrical changes
are smaller for the corresponding sulfur-containing sys-
tems.
(
3) Hyd r ogen -Bon d in g Acid ity a n d Ba sicity. The
equilibrium constants K , for the formation of the 1:1
complexes in c-C 12 solution at 23.3 °C between 4 and
,4-dinitrophenol and PyO are respectively equal to 1.11
c
6
H
3
(
0.17 and 1.52 ( 0.23 dm3 mol . These are rather
-1
small values and, thus, are somewhat uncertain. The
experimental study of reaction 4 also provides the value
of Kdim, the dimerization constant of 4 in c-C
6
H
12 solution
It is worth noting that, although the F‚‚‚H distances
at 23.3 °C:
+
+
in 2H and 4H are about 2.5 Å (1 Å ) 100 pm), no bcps
were found in the corresponding interatomic regions.
Hence, as for the neutrals, we must conclude that, strictly
speaking, no intramolecular hydrogen bonds are present.
Upon deprotonation, the changes are also important,
in particular for the oxygen-containing species. Depro-
tonation implies that the σ orbital involved in the O-H
bond becomes a lone pair, with a concomitant decrease
of the electronegativity of the oxygen atom. Accordingly,
the s character of the carbon hybrid orbital involved in
the C-O bond increases, and the bond becomes shorter.
By orthogonality, the other three hybrids increase their
p character, and the C-C and C-H bonds become longer.
Consistent with this, the CCO bond angle increases,
while the CCH bond angles decrease. These hybridiza-
tion changes are well reproduced by the corresponding
NBO analysis (see Table 3). A similar behavior is found
for the sulfur-containing derivatives.
CF CH SH + HS-CH CF a
3
2
2
3
CF CH SH‚‚‚S(H)CH CF
3
Kdim (13)
3
2
2
This value is quite small, 0.10 ( 0.05 dm mol-1.
4) Ch a r ge Tr a n sfer Ba sicity. The equilibrium
constant K for reaction 14 in tetrachloromethane solu-
tion at 25.0 °C, as determined in this work, equals 1.20
3
(
c
3
-1
(
0.18 dm mol
CF CH SH + I a CF CH (H)S‚‚‚I
2
:
Kc (14)
3
2
2
3
2
Discu ssion
Str u ctu r es a n d Bon d in g Ch a r a cter istics. The
MP2/6-311+G(d,p) geometries of the four species under
investigation and their protonated and deprotonated
+
+
+
+
-
-
-
-
forms 1H , 2H , 3H , 4H , 1 , 2 , 3 , and 4 have been
schematized in Figure 1 and are also given as supporting
information. Their bonding characteristics are sum-
marized in Table 3. Since a detailed discussion of
geometries is not the aim of this paper, we shall concen-
trate our attention on the structural differences between
species 2 and 4 and on the effects associated with their
protonation and deprotonation processes. The most
important difference between 2 and 4 is that, in the
CF
3
substitution has almost negligible effects on the
O-H and S-H bonds. This seems to be consistent with
previous findings,2 which showed that CF
6b,c
substitution
3
in pyrazole has small effects on the azolic system.
It is also interesting to realize that, upon deprotona-
tion, the charge density at the most electronegative atoms
becomes larger than that in the neutral, since the system
becomes electron-excessive. This favors the π back-
Ta ble 3. Bon d in g Ch a r a cter istics of th e System s (Y3C-CH-XH; X ) O, S; Y ) H, F ) In clu d ed in Th is Stu d ya
bond
1
1-
1H+
2
2-
2H+
3
3-
3H+
4
4-
4H+
C-C
F
∇
ꢀ
0.253
0.231
-0.510
0.045
0.357
-0.650
0.010
0.273
-0.752
0.060
0.253
-0.642
0.056
0.278
-0.773
0.004
0.243
-0.566
0.004
0.240
-0.548
0.011
0.244
-0.576
0.004
0.268
-0.710
0.018
0.267
-0.707
0.006
0.268
-0.705
0.021
2F -0.625
0.037
C-X
C-Y
X-H
F
∇
ꢀ
0.249
0.315
-0.714
0.010
0.261
-0.837
0.005
0.167
0.003
0.118
0.271
-0.915
0.007
0.341
0.261
-0.459
0.039
0.267
-0.151
0.146
0.363
-0.253
0.027
0.324
-0.770
0.013
0.246
-0.137
0.256
0.189
-0.412
0.028
0.274
-0.276
0.128
0.336
-0.219
0.022
0.178
-0.285
0.102
0.272
-0.910
0.009
0.217
-0.666
0.153
0.172
-0.258
0.002
0.263
-0.851
0.005
0.172
-0.257
0.003
0.274
-0.933
0.011
0.225
-0.709
0.038
0.181
-0.303
0.122
0.274
-0.140
0.132
0.219
-0.674
0.146
0.170
-0.254
0.013
0.266
-0.121
0.163
0.179
-0.292
0.014
0.290
-0.233
0.113
0.229
-0.718
0.229
2F -0.409
0.024
F
∇
ꢀ
0.271
2F -0.903
0.006
F
∇
ꢀ
0.365
2F -2.500
0.027
-2.525
0.021
a
Charge densities (F in e‚au-3), Laplacians of the charge density (∇ F in e‚au-5), elipticities (ꢀ), and percentage of s character (s/c).
2
The natural atomic charges (qN) of the atoms involved in the bonds are also given.

Acidity/Basicity of 2,2,2-Trifluoroethanethiol
J . Org. Chem., Vol. 61, No. 16, 1996 5489
F igu r e 1. Optimized structures for 1-4 and their corresponding protonated and deprotonated forms. Bond lengths are in
angstroms, and bond angles are in degrees.
donation from the fluorine atoms to the carbon atom,
which is reflected in an increase of the elipticity of the
C-F bonds (see Table 3).
an alkyl group by a trifluoroethyl group brings about an
important increase in acidity and a substantial decrease
in basicity. Both acidity and basicity are seen to increase
with the size of the alkyl groups, largely as a consequence
of polarizability effects.7
Theoretically calculated proton affinities and deproto-
nation enthalpies for compounds 1-4 are presented in
Table 5. It can be seen that the quantitative agreement
with our results is excellent. Furthermore, these calcula-
tions reproduce well the fact that 3 is slightly more basic
than 1, while it is significantly more acidic.
Figures 2 and 3 present the relationships between the
proton affinities and deprotonation enthalpies of alcohols
and thiols. Some relevant features of these plots are as
follows:
(1) Structural effects on the proton affinities and
deprotonation enthalpies of thiols and alcohols are lin-
As for protonation, the effects of the deprotonation
processes on sulfur-containing systems are smaller than
those described above for oxygen-containing compounds.
As we will discuss later, this dampening of the charge
redistributions undergone by sulfur derivatives may be
also related to the fact that trifluoromethyl substitution
effects on the intrinsic basicities and acidities are also
slightly smaller for sulfur- than for oxygen-containing
compounds.
Acid ities a n d Ba sicities. For comparison purposes,
we have gathered in Table 4 the proton affinities and
deprotonation enthalpies in the gas phase of sets of
representative alkanols and alkanethiols. From these
results, it is clear that, in all cases, the substitution of

5
490 J . Org. Chem., Vol. 61, No. 16, 1996
Molina et al.
Ta ble 4. Exp er im en ta l Ion iza tion En th a lp ies, ∆Ha cid ,
a n d P r oton Affin ities, P A, of Selected Alcoh ols (R-OH)
a n d Th iols (R-SH)
alcohols
thiols
R
∆Hacida,b
PAa,c
∆Hacida,b
PAa,c
Me
Et
n-Pr
i-Pr
n-Bu
i-Bu
t-Bu
t-BuCH2
CF3CH2
380.6
377.4
376.0
375.4
375.4
374.7
374.6
372.6
361.8
181.7d
188.3
190.8
191.1
191.1
192.4
193.7
193.6
169.0
356.9
355.2
354.2
353.4
353.7
353.1
352.5
351.7
342.4
187.4
190.8
191.6
194.1
196.9
175.6
a
All values in kcal/mol. b From ref 6. c From ref 4. d From ref
2
9.
Ta ble 5. G2-Ca lcu la ted Tota l En er gies (E), P r oton
Affin ties (P A), Acid ities (∆Ha cid ), a n d Dip ole Mom en ts (µ)
PA
(kcal/mol)
AA
(kcal/mol)
a
a
compd
E (Hartrees)
µ (D)
1
1
-154.753 3284
-154.152 1926
-155.048 2019
-452.251 6636
-451.679 3357
-452.519 6033
-477.359 4227
-476.794 2003
-477.659 4959
-774.842 0471
-774.295 7616
-775.117 3293
186.4
169.5
189.6
174.1
378.4
360.2
355.7
343.8
1.79
4.27
2.46
3.59
5.83
7.46
1.66
5.07
-
1
2
2
2
3
3
3
4
4
4
H+
-
F igu r e 3. Experimental deprotonation enthalpies of thiols
versus experimental deprotonation enthalpies of homologous
alcohols.
H+
-
H+
have portrayed as a dotted line the correlation obtained
upon inclusion of the 2/4 data). These results show for
the first time the pattern of similarity between alcohols
and thiols to be quite broad.
1.87
6.50
7.16
-
H+
a
These values include the corresponding thermal corrections
(2) In both cases, the introduction of the trifluoro-
evaluated at 298.15 K.
methyl group more than doubles the range of structural
effects.
(3) In every case, thiols are both more acidic and more
basic than the corresponding alcohols. This reflects the
fact that the polarizability of sulfur is larger than that
of oxygen. On the other hand, the slopes are significantly
less than 1. Thus, substituent effects are appreciably
attenuated in thiols with respect to alcohols, possibly as
a consequence of the S-C bonds being longer than the
O-C ones. It is noteworthy that this is exactly the same
pattern of sensitivity to substituent effects on the intrin-
3
sic basicities of carbonyl and thiocarbonyl compounds.
The much greater acidity of 3 relative to 1 is in clear
3
0
contrast with their behavior in aqueous solution. This
may be rationalized in terms of the differences between
the electronic structures of both neutrals. On the one
hand, due to the larger size of sulfur, the S-H linkage
is weaker than the O-H bond, as reflected by a much
smaller charge density at the corresponding bcp (see
Table 3). On the other hand, since oxygen is more
-
+
electronegative than sulfur, the O -H polarity of the
hydroxylic bond is much higher than that of the S-H
linkage of ethanethiol, which also contributes to enhance
the stability of the former. As a consequence, the O-H
bond fission for the alcohol is energetically less favorable
than the S-H bond fission in the thiol. In aqueous
solution, the situation is different. The great polarity of
the O-H linkage of ethanol strongly favors the interac-
tions with the solvent and the heterolytic fission of the
bond to yield an ethoxide anion, which is more stabilized
by the interactions with the solvent than the thioethoxide
anion. Hence, the gap between the acidities of both
F igu r e 2. Experimental proton affinities of thiols versus
experimental proton affinities of homologous alcohols.
early related to a high degree of precision. The correla-
tion involving PAs is of slightly lower quality. This likely
reflects the fact that the experimental determination of
the PAs of secondary and tertiary alcohols is somewhat
less precise because of the competing dehydration of the
protonated species. In these plots, the data for the 2/4
couple were not included in the correlations represented
as solid lines. We see that the corresponding point is
either on the line or very close to it (in plots 2 and 3, we
(30) (a) Fehrst, A. R. J . Am. Chem. Soc. 1971, 93, 3504-3515. (b)
Pohl, E. R.; Hupe, D. J . J . Am. Chem. Soc. 1978, 100, 8130-8133.

Acidity/Basicity of 2,2,2-Trifluoroethanethiol
J . Org. Chem., Vol. 61, No. 16, 1996 5491
systems decreases considerably, and 3 is only about 7
kcal/mol more acidic than 1.
Hyd r ogen -Bon d in g Acid ity a n d Ba sicity. The
Ch a r ge-Tr a n sfer Ba sicity. We showed a few years
ago35 that there is a good linear relationship between the
proton affinities in the gas phase and the Gibbs energy
equilibrium constant for the complexation between 4 and
changes pertaining to the 1:1 charge-transfer complex-
PyO14 can be linked to the R
H
scale of hydrogen-bonding
ation between S(sp ) n-bases and molecular iodine in CH
3
-
2
2
acidity of monomeric species.3
1,32
The computed value,
Cl
solution at 298 K.
The correlation equation was eq 16,
2
0
.08, is very small and ranks 4 significantly below
H
dichloromethane (R
2
) 0.13).
∆G° (CH Cl ) )
I 2 2
2
Regarding the hydrogen-bonding basicity, the equilib-
rium constant for the association between 4 and 3,4-
(-3.67 ( 0.09) - (0.113 ( 0.007)∆PA (16)
wherein ∆G° 2(CH Cl ) is the Gibbs energy change for
H
dinitrophenol can be linked to the â
2
scale of hydrogen-
bonding basicity of monomeric species.3
librium constant is about one-half of the value for C
1,33
The equi-
I
2
2
reaction 17 and ∆PA is the proton affinity relative to
2
H
H
5
-
-
3
ammonia of the base B, an S(sp ) n-donor base.
1
1
9
SH and some 240 times smaller than that for C
2
5
OH (âH
) 0.44). The â
H
value for 4 equals 0.12 and is
B + I a B‚‚‚I
(17)
4
solution were linked to
2
2
2 2
0,33
substantially smaller than that for CF
3
CH
2
OH (0.18).1
∆
G° 2
I
values determined in CCl
Cl solution through eq 18,
It is clear that 4 is an extremely weak hydrogen-
bonding base and also a very weak hydrogen-bonding
acid. Consideration of the self-association process (eq 14)
is appropiate at this point.
As indicated earlier, the equilibrium constant Kdim for
this reaction is of the order of 0.1 dm3 mol-1. Now, it
has been shown that, for thousands of 1:1 HB complexes
values in CH
2
2
∆G° (CH Cl ) )
I
2
2
2
(
-0.46 ( 0.14) + (1.04 ( 0.05)∆G° (CCl ) (18)
I2 4
If we apply eq 18 to our data for 4, we get ∆G°I2(CH
2
Cl
) value one would
predict on the basis of eq 16 and using the experimental
2
)
-
1
)
I 2 2
-0.57 kcal mol . The ∆G° 2(CH Cl
in solution in CCl
very satisfactory degree of precision.
4 6
or c-C H12 solvents, eq 15 holds to a
3
1,34
Using the values
(15)
(4) ) 0.12, this equation predicts a
-
1
∆
PA (-27.9 kcal/mol) is -0.52 kcal mol . On account
H
H
of the combined uncertainties of these correlations, the
agreement is remarkably good and provides strong sup-
port for the contention that, in selected cases, charge-
transfer complexation with molecular iodine in solution
displays a pattern of structural effects similar to that of
gas phase basicity.35
log K ) 7.354R 2â - 1.094
c
2
RH
(4) ) 0.08 and âH
2
2
3
-1
value for the dimerization constant of 0.095 dm mol
,
in excellent (and perhaps somewhat fortuitous) agree-
ment with the experimental value. It is remarkable that
equilibrium constants of this order of magnitude are
considered to define the onset34 of “true” HB interactions.
Ack n ow led gm en t. This work was supported by
Grants PB 93-0289-C02 and PB 93-0142-C03 of the
Spanish D.G.I.C.Y.T.
4
seems to be precisely at this onset and is, therefore,
very feebly self-associated.
Su p p or tin g In for m a tion Ava ila ble: MP2/6-311+G(d,p)
+
+
+
+
(
31) Abraham, M. H.; Grellier, P. L.; Prior, D. V.; Morris, J . J .; Taft,
geometries and G2 total energies of 1-4, 1H , 2H , 3H , 4H ,
-
-
-
-
R. W.; Morris, J . J .; Taylor, P. J .; Laurence, C.; Berthelot, C.; Doherty,
R. M.; Kamlet, M. J .; Abboud, J .-L. M.; Sra ¨ı di, K.; Guih e´ neuf, G. J .
Am. Chem. Soc. 1988, 110, 8534-8535.
1 , 2 , 3 , and 4 (3 pages). This material is contained in
libraries on microfiche, immediately follows this article in the
microfilm version on the journal, and can be ordered from the
ACS; see any current masthead page for ordering information.
(32) Abraham, M. H.; Grellier, P. L.; Prior, D. V.; Duce, P.; Morris,
J . J .; Taylor, P. J . J . Chem. Soc., Perkin Trans. 2 1989, 699-712.
33) Abraham, M. H.; Grellier, P. L.; Prior, D. V.; Morris, J . J .;
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(
J O960524A
(
Treatments of Solvent/ Solute Interactions, Theoretical and Computa-
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Amsterdam, 1994.
(35) Abboud, J .-L. M.; Notario, R.; Berthelot, M.; Claramunt, R. M.;
Cabildo, P.; Elguero, J .; El Ghomari, M. J .; Bouab, W.; Mokhlisse, R.;
Guih e´ neuf, G. J . Am. Chem. Soc. 1991, 113, 7489-7493.