Kinetics of proton transfer from cationic carbon acids in water and aqueous DMSO. Effect of activating groups and solvent on intrinsic rate constants

Article abstract of DOI:10.1021/jo051166r

Acidity constants and rates of reversible deprotonation of acetonyltriphenylphosphonium ion (1H⁺), phenacyltriphenylphosphonium ion (2H⁺), N-methyl-4-phenacylpyridinium ion (3H⁺), and N-methyl-4-(phenylsulfonylmethyl)pyrid

Full text of DOI:10.1021/jo051166r

Kinetics of Proton Transfer from Cationic Carbon Acids in Water
and Aqueous DMSO. Effect of Activating Groups and Solvent on
Intrinsic Rate Constants
Claude F. Bernasconi,* Douglas E. Fairchild, Robert L. Montan˜ez, Perdram Aleshi,
Huaiben Zheng, and Edward Lorance
Department of Chemistry and Biochemistry, University of California, Santa Cruz, California 95064
bernasconi@chemistry.ucsc.edu
Received June 9, 2005
Acidity constants and rates of reversible deprotonation of acetonyltriphenylphosphonium ion (1H+),
phenacyltriphenylphosphonium ion (2H+), N-methyl-4-phenacylpyridinium ion (3H+), and N-methyl-
4-(phenylsulfonylmethyl)pyridinium ion (4H+) by amines in water, 50% DMSO-50% water (v/v),
and 90% DMSO-10% water (v/v) have been determined. From the respective Brønsted plots, log
ko values for the intrinsic rate constants of the various proton transfers were obtained. Solvent
transfer activity coefficients of the carbon acids and their respective conjugate bases were also
determined which helped in understanding how the pKa values and intrinsic rate constants depend
on the solvent. Some of the main conclusions are as follows: (1) The pKa values of 1H+, 2H+, and
3H+ are significantly higher than that of 4H+ because of a stronger resonance stabilization of the
+
corresponding conjugate bases 1, 2 and 3, respectively. (2) The electronic effects of the PPh3 and
the N-methyl-4-pyridylium group are similar but the mix between inductive and resonance effect
is different. (3) All four acids become more acidic upon addition of DMSO to the solvent. In all
cases, the main factor is the stronger solvation of H3O+ in DMSO; for 1H+, 2H+, and 3H+ but not
4H+ this factor is significantly attenuated by stronger solvation of the carbon acid in DMSO. (4)
The intrinsic rate constants for proton transfer are relatively high for all four carbon acids and
show little solvent dependence; this contrasts with nitroalkanes which have much lower intrinsic
rate constants and show a strong solvent dependence. These results can be understood by a detailed
analysis of the interplay between inductive, resonance, and solvation effects.
Introduction
resonance or solvation have a disproportionately small
effect on the transition state because their development
in a product lags behind bond changes or their loss from
a reactant is ahead of bond changes (“transition state
imbalance”). As a result, the intrinsic barriers are higher
(or the intrinsic rate constants are lower) for reactions
that lead to more delocalized or more solvated products.
On the other hand, inductive effects by electron-with-
drawing substituents that are close to the reaction site
The study of proton transfers from carbon acids has
provided numerous insights about the factors that affect
intrinsic barriers (∆Goq)1 or intrinsic rate constants (ko)1
of chemical reactions.4 An important conclusion is that
product-stabilizing factors such as charge delocalization/
(1) The intrinsic barrier (intrinsic rate constant) of a reaction with
a forward rate constant k1 and a reverse rate constant k-1 is defined
tend to decrease ∆Gq or increase ko due to the imbalance
as ∆Goq ) ∆G1q ) ∆Gq-1 when ∆G° ) 0 (as ko ) k1 ) k-1 when K1 ) 1).2,3
For proton transfers statistical factors are usually included.3
(2) Marcus, R. A. J. Phys. Chem. 1968, 72, 891.
o
because their stabilizing effect on the transition state is
disproportionately large.
Most studies of proton transfers at carbon have dealt
with neutral carbon acids that are activated by one or
two π-acceptor groups.4c Cationic carbon acids have
received less attention, although kinetic data on several
(3) Keeffe, J. R.; Kresge, A. J. In Investigation of Rates and
Mechanisms of Reactions; Bernasconi, C. F., Ed.; Wiley-Interscience:
New York, 1986; Part I, p 747.
(4) (a) Bernasconi, C. F. Acc. Chem. Res. 1987, 20, 301. (b) Ber-
nasconi, C. F. Acc. Chem. Res. 1992, 25, 9. (c) Bernasconi, C. F. Adv.
Phys. Org. Chem. 1992, 27, 119.
10.1021/jo051166r CCC: $30.25 © 2005 American Chemical Society
Published on Web 08/17/2005
J. Org. Chem. 2005, 70, 7721-7730
7721

Bernasconi et al.
1-methyl-3-(phenylacetyl)pyridinium,5 1-methyl-4-(phen-
ylacetyl)pyridinium,5 4-phenacylpyridinium,6 4-(phenyl-
sulfonyl)methylpyridinium cations,7 dimethyl-9-fluoren-
ylsulfonium ions,8 triphenyl-9-fluorenylphosphonium,9
2-acetyl-1-methylpyridinium ion,10 and several cationic
rhenium Fischer carbene complexes11,12 have been re-
ported.
In this paper, we present rate data on the reversible
deprotonation of two triphenylphosphonium ion deriva-
tives (1H+, 2H+) and two N-methylpyridinium ion de-
rivatives (3H+, 4H+) by amines in water, 50% DMSO-
50% water (v/v), and/or 90% DMSO-10% water (v/v).
constants for equilibrium approach are given by eq 3 for
all carbon acids of this study.
H2O
kobsd ) k1H O + k-1HaH+ + k OH[OH-] + k-1
+
2
1
k1 [B] + k-1BH[BH+] (3)
B
CH
The pKa values of the various carbon acids were
determined spectrophotometrically via eq 4 where AC and
ACH refer to the absorbance of C (pH . pKaCH) and CH+
(pH , pKaCH), respectively, and A to the absorbance of
CH
6-9 solutions of intermediate pH close to pKa
.
AC - A
pKaCH ) pH + log
A - ACH
(4)
Kinetic Results for Acetonyltriphenylphospho-
nium Ion (1H+) and Phenacyltriphenylphospho-
nium Ion (2H+). Kinetic measurements were performed
in water, 50% DMSO-50% water (v/v), and 90% DMSO-
10% water (v/v). The choice of amines was limited by the
time resolution of the stopped-flow apparatus. For the
purpose of comparison with other carbon acids (see the
Discussion) where most of the relevant literature data
refer to reactions with alicyclic secondary amines, mea-
surements with such secondary amines would have been
desirable. However, the reactions with these amines were
too fast for the stopped-flow method. The same was true
for n-butylamine in water and 90% DMSO, but measure-
ments with n-butylamine and morpholine were feasible
in 50% DMSO. In 90% DMSO even the reaction with
2-methoxyethylamine was too fast.
The main focus is on the determination of intrinsic rate
constants. We are particularly interested in the effect of
the interplay between cationic substituent and π-acceptor
and also on the solvent effect on the intrinsic rate
constants. The determination of solvent activity coef-
ficients of the carbon acids and their respective conjugate
bases will be used to analyze the various contributions
to the solvent dependence of the intrinsic rate constants
and the pKa values of the four carbon acids.
Results
Under our reaction conditions, the contribution of the
H2O
k1H O, k-1HaH , k1OH[OH-], and k-1
terms to kobsd was
+
2
General Features. The reactions of 1H+-4H+ with
amine buffers can be described by eq 1 where CH+ refers
to the cationic carbon acid, C to its conjugate base, B to
the amine, and BH+ to its respective conjugate acid. For
either small or negligible; since the amine dependence
of kobsd was determined at constant pH, eq 3 simplifies
to eq 5 with C being a constant.
kobsd ) C + k1 [B] + k-1BH[BH+]
(5)
B
Plots of kobsd versus [B] yielded slopes given by eq 6 from
BH
which k1B could be calculated since KaCH is known; k-1
was calculated as k1BKaBH/KaCH. The various rate con-
1H+-3H+, enol formation (EH+) by rapid protonation of
the enolate oxygen, e.g., eq 2, might, in principle, play a
role in the kinetics at low pH. However, under our
CH
stants and pKa values are summarized in Table 1.
aH+
aH+
BH
B
slope ) k1B + k-1
BH ) k1 1 +
(6)
CH
(
)
Ka
Ka
N-Methyl-4-phenacylpyridinium Ion (3H+) and
N-Methyl-4-(phenylsulfonylmethyl)pyridinium Ion
(4H+).13 Kinetic determinations were made in water, 50%
DMSO-50% water (v/v), and 90% DMSO-10% water (v/
v) with piperidine, piperazine, 1-(2-hydroxyethyl)pipera-
zine (PZ-CH2CH2OH), and morpholine. For 4H+ rates in
water and 50% DMSO-50% water were also measured
with n-butylamine, 2-methoxyethylamine, 2-chloroeth-
ylamine, and glycinamide. Data analysis was performed
as described for 1H+ and 2H+. The results are sum-
marized in Table 2.
reaction conditions no evidence for the involvement of the
EH
enol could be found, indicating that pH . pKa under
all conditions. Hence, the observed pseudo-first-order rate
(5) Bunting, J. W.; Stefanidis, D. J. Am. Chem. Soc. 1988, 110, 4008.
(6) Stefanidis, D.; Bunting, J. W. J. Am. Chem. Soc. 1990, 112, 3163.
(7) Wodzinski, S.; Bunting, J. W. J. Am. Chem. Soc. 1994, 116, 6910.
(8) Murray, C. J.; Jencks, W. P. J. Am. Chem. Soc. 1990, 112, 1880.
(9) Bernasconi, C. F.; Fairchild, D. E. J. Phys. Org. Chem. 1992, 5,
409.
(10) (a) Tobin, J. B.; Frey, P. A. J. Am. Chem. Soc. 1996, 118, 12253.
(b) Bernasconi, C. F.; Moreira, J. A.; Huang, L. L.; Kittredge, K. W. J.
Am. Chem. Soc. 1999, 121, 1674.
Solvent Transfer Activity Coefficients. To assess
how changes in the solvation of CH+, C and H+ may
(11) Bernasconi, C. F.; Ragains, M. L. J. Am. Chem. Soc. 2001, 123,
11890.
(12) Bernasconi, C. F.; Ragains, M. L.; Bhattacharya, S. J. Am.
Chem. Soc. 2003, 125, 12328.
(13) Rates of deprotonation of 3H+ and 4H+ by OH- and pKa values
in water have been reported by Bunting et al.6,7
7722 J. Org. Chem., Vol. 70, No. 19, 2005

Kinetics of Proton Transfer from Cationic Carbon Acids
CH
TABLE 1. Rate Constants and pKa Values for the Reversible Deprotonation of Acetonyltriphenylphosphonium Ion
(1H+) and Phenacyltriphenylphosphonium Ion (2H+) in Water, 50% DMSO-50% Water (v/v), and 90% DMSO-10% Water
(v/v) at 20 °Ca
a Error limits in k1 and k-1 ( 6% or better. b µ ) 0.5 M (KCl). c µ ) 0.06 M (KCl).
B
BH
CH
affect the solvent effects on pKa and log ko, we deter-
mined the solvent transfer activity coefficients, IγCIIH, for
the transfer of C from solvent I to solvent II, e.g., I )
H2O, II ) 90% DMSO-10% water (v/v). These determi-
nations were made by carrying out distribution experi-
ments between n-heptane and the various solvents of
interest, as described by Watarai14 and by Bernasconi
and Bunnell.15 These experiments yielded partition coef-
from Well’s work.17
IδII∆G° (CH+) ) 2.303RT log IγII
(8)
(9)
tr
CH+
IδII∆G° (C) ) 2.303RT log IγICI
tr
IδII∆G° (H+) ) 2.303RT log IγIHI+
(10)
tr
II
I
ficients from which log γC was calculated as described
in the Experimental Section. From these IγCII values, the
solvent activity coefficient, IγCIIH+, for the transfer of CH+
from solvent I to solvent II, could then be obtained by
applying the Parker16 formalism of eq 7.
Discussion
pKaCH Values. The pKaCH values and relevant solvent
transfer activity coefficients for 1H+-4H+ in water, 50%
DMSO-50% water (v/v), and 90% DMSO-10% water
(v/v) are summarized in Table 3 along with the previously
reported pKaCH values and activity coefficients 5H18 and
6H15 in the same solvents. The following points are
noteworthy.
I∆IIpKaCH ) log IγCII + log IγHII+ - log IγICIH+ (7)
CH
In this equation I∆IIpKa
represents the difference
CH
between the pKa
values in the two solvents, i.e.,
II
II
I
I∆IIpKa ) pKaCH(II) - pKaCH(I), while log γCH+, log γC
CH
I
II
I
and log γH+ are related to the standard free energies of
transfer by eqs 8-10, respectively; log γH+ is known
II
I
(1) The pKaCH values in water of the three cationic acids
1H+, 2H+, and 3H+ are quite similar but significantly
lower than the pKaCH of 4H+.19 This reflects the stronger
π-acceptor effect of the acyl groups compared to the
(14) (a) Watarai, H.; Suzuki, N. Bull. Chem. Soc. Jpn. 1980, 53,
1848. (b) Watarai, H. Bull. Chem. Soc. Jpn. 1980, 53, 3019.
(15) Bernasconi, C. F.; Bunnell, R. D. J. Am. Chem. Soc. 1988, 110,
2900.
(16) Parker, A. J. Chem. Rev. 1969, 69, 1.
J. Org. Chem, Vol. 70, No. 19, 2005 7723

Bernasconi et al.
CH
TABLE 2. Rate Constants and pKa Values for the Reversible Deprotonation of N-Methyl-4-phenacylpyridinium Ion
(3H+) and N-Methyl-4-(phenylsulfonylmethyl)pyridinium Ion (4H+) in Water, 50% DMSO-50% Water (v/v), and 90%
DMSO-10% Water (v/v) at 20 °Ca
B
BH
a Error limits in k1 and k-1 ( 6% or better. b µ ) 0.5 M (KCl). c µ ) 0.06 M (KCl). d 1-(2-Hydroxyethyl)piperazine.
CH
phenylsulfonyl group (σR(CH3CO) ) 0.16,21 σR(PhCO) )
0.16,21 σR(PhSO2) ) 0.1221); the stronger inductive effect
of the PhSO2 group (σF(CH3CO) ) 0.25,21 σF(PhCO) )
0.29,21 σF(PhSO2) ) 0.5921) is not able to offset its weaker
π-acceptor effect.22 The somewhat lower pKa of 2H+
compared to that of 1H+ is apparently the result of the
somewhat stronger inductive effect of the PhCO group
relative to that of the CH3CO group.
(2) The acidifying effects of the PPh3+ and the N-meth-
CH
(17) Wells, C. F. In Thermodynamic Behavior of Electrolytes in
Mixed Solvents; Furter, W. F., Ed.; Advances in Chemistry 177;
American Chemical Society: Washington, DC, 1979; p 53.
(18) Bernasconi, C. F.; Montan˜ez, R. L. J. Org. Chem. 1997, 62,
8162.
yl-4-pyridylium groups lead to decreases in pKa
by
(22) A reviewer has suggested that the differences between σR
)
0.16 for PhCO and 0.12 for PhSO2 does not seem sufficient to account
for the large difference in the pKa values of the respective compounds
(3H+ and 4H+). In view of the close proximity of the PhCO and PhSO2
groups to the reaction center one actually should expect a rather large
difference in the pKa values.
(19) At 25 °C, µ ) 0.1 M, pKaCH(3H+) ) 7.66 and pKaCH(4H+) )
11.34.20
(20) Wodzinski, S.; Bunting, J. W. Can. J. Chem. 1992, 70, 2635.
(21) Hansch, C.; Leo, A.; Taft, R. W. Chem. Rev. 1991, 91, 165.
7724 J. Org. Chem., Vol. 70, No. 19, 2005

Kinetics of Proton Transfer from Cationic Carbon Acids
CH
TABLE 3. Solvent Transfer Activity Coefficients and pKa Values at 20 °Ca
a log 0γ5H0+ ) -1.93; log 0γH90+ ) -3.05, Reference 17. b For 5H and 6H the acid form is neutral (log IγCIIH) and the conjugate base is anionic
II
I
(log γC-). c Reference 18. d Reference 15.
10.5-11.5 pKa units relative to the respective reference
ketones (pKaCH(CH3COCH3) ) 19.3,23 pKaCH(PhCOCH3)
) 18.324). There is very little difference in the overall
electronic effect of the two groups, although the mix
between inductive and resonance effects is different: for
1H+ and 2H+ the inductive effect exerted by the posi-
tively charged PPh3 group is probably dominant, i.e.,
besides the large weight of 1a (2a), the contribution of
1b (2b) to the structure of the conjugate base is more
important than that of 1c (2c).25,26 In contrast, for 3H+
the resonance effect symbolized by 3c is dominant while
3b is probably less important.
CH
However, as the still lower pKa
of 5H indicates,
neither the +PPh3 nor the N-methylpyridyl group is as
J. Org. Chem, Vol. 70, No. 19, 2005 7725

Bernasconi et al.
effective as the nitro group in stabilizing the respective
conjugate base. This is undoubtedly the result of the
strong resonance effect of the nitro group where 5c- is
the dominant resonance structure.18
Solvent Effects on pKaCH. The four cationic acids,
1H+-4H+, become more acidic in the presence of DMSO
while the two neutral acids, 5H and 6H, become less
CH
acidic. These solvents effects on the pKa values are
most easily understood by considering the solvent trans-
fer activity coefficients (Table 3) which measure the
solvational stabilization/destabilization of the various
species that results from changing the solvent; note a
FIGURE 1. Brønsted plots for the reaction of 1H+ with
primary aliphatic amines in water (O, b), 50% DMSO-50%
water (0, 9), and 90% DMSO-10% water (4, 2).
II
I
positive value of log γX means that species X becomes
less solvated and hence destabilized upon transfer from
solvent I to solvent II while a negative value means the
opposite. The relationship between the solvent effect on
solvated in 90% DMSO. Hence, these two effects partially
offset each other but, since the stabilization of CH+ in
90% DMSO is stronger than that of C, the net result is
to decrease the acidity in 90% DMSO and thus to
attenuate the effect of the stronger solvation of the
CH
the pKa value and the various activity coefficients is
given by eq 7. Applied, e.g., to the change from water to
90% DMSO, eq 7 becomes eq 11. For the six acids, 1H+-
4H+, 5H and 6H, the numerical results of an analysis
according to eq 11 are summarized in eqs 12-17.
0
hydronium ion on ∆90pKaCH. The enhanced solvation of
CH+ in 90% DMSO is consistent with the well-known
properties of DMSO as a cation-solvating medium16,28
while the enhanced solvation of C is simply a manifesta-
tion of the better solubility of neutral organic molecules
in DMSO than in water. It is noteworthy that for the
N-methylpyridinium derivative 3H+ the solvational sta-
bilization of both the acid and the conjugate base is much
weaker than for the phosphonium derivatives 1H+ and
2H+; however, because the effects on CH+ and C again
partially offset each other, 0∆90pKaCH(3H+) is not dra-
0∆90pKaCH ) pKaCH(90) - pKaCH(0) ) log 0γ9C0
+
log 0γ9H0+ - log 0γ9C0H+ (11)
0∆90pKaCH(1H+) ) -0.86 ) -1.43 - 3.05 + 3.62
(12)
0∆90pKaCH(2H+) ) -1.46 ) -2.14 - 3.05 + 3.72
(13)
matically different from, say, ∆90pKaCH(2H+).
0
0∆90pKaCH(3H+) ) -2.11 ) -0.95 - 3.05 + 1.89
(3) For 4H+, neither CH+ nor C is strongly affected by
the change in solvent; interestingly CH+ is slightly less
solvated in 90% DMSO which contributes to the acidify-
ing effect in this solvent. The small log 0γ9C0H+ and log 0γ9C0
values in this case may be traced to the sulfonyl group:
the oxygen atoms of this group carry a significant partial
negative charge which is better solvated by water than
by DMSO. Hence the features that enhance solvation in
DMSO for 1H+-3H+ and their respective conjugate bases
are offset by the counter effect of the sulfonyl group.
(4) For 5H and 6H, the conjugate base which is an
anion, is destabilized in 90% DMSO which leads to a
decrease in acidity. This destabilization is particularly
strong for the conjugate base of 6H (6-) which has a
highly concentrated negative charge on the oxygens and
is much more strongly solvated by hydroxylic solvents
(14)
0∆90pKaCH(4H+) ) -3.60 ) -0.28 - 3.05 - 0.27
(15)
0∆90pKaCH(5H)27 ) 1.72 ) 1.67 - 3.05 + 3.10 (16)
0∆90pKaCH(6H)27 ) 4.52 ) 6.70 - 3.05 + 0.87 (17)
The following points are of interest.
(1) The stronger solvation of the hydronium ion in 90%
DMSO than in water (log 0γH90+ ) -3.05) leads to an
acidifying effect of 3.05 pKa units for all acids.
(2) For 1H+, 2H+, and 3H+, both the acid (CH+) and
their respective conjugate base (C) are significantly better
-90
0
than by DMSO (log γC ) 6.70); the more dispersed
negative charge in 5- reduces this effect (log 0γ-C90
1.67).
)
(23) Chiang, Y.; Kresge, A. J.; Schepp, N. P. J. Am. Chem. Soc. 1989,
111, 3977.
(24) Chiang, Y.; Kresge, A. J.; Wirz, J. J. Am. Chem. Soc. 1984, 106,
6392.
(25) Johnson, A. W. Ylid Chemistry; Academic Press: New York,
1966; p 292.
(26) Schlosser, M.; Titus, J.; Schaub, B. Heteroatom Chem. 1990, 1,
151.
7726 J. Org. Chem., Vol. 70, No. 19, 2005

Kinetics of Proton Transfer from Cationic Carbon Acids
FIGURE 4. Brønsted plots for the reaction of 4H+ with
primary aliphatic amines in water (O, b) and 50% DMSO-
50% water (0, 9).
FIGURE 2. Brønsted plots for the reaction of 2H+ with
primary aliphatic amines in water (O, b), 50% DMSO-50%
water (0, 9), and 90% DMSO-10% water (4, 2).
FIGURE 5. Brønsted plots for the reaction of 4H+ with
secondary alicyclic amines in water (O, b), 50% DMSO-50%
water (0, 9), and 90% DMSO-10% water (4, 2).
FIGURE 3. Brønsted plots for the reaction of 3H+ with
secondary alicyclic amines in water (O, b), 50% DMSO-50%
water (0, 9), and 90% DMSO-10% water (4, 2).
1H+, 3H+, 4H+, 5H, and 6H but not for 2H+. We do not
attach much significance to these trends which are
difficult to interpret.30
As is typical for proton transfers involving carbon acids
that have resonance stabilized conjugate bases, the log
ko values are much lower4 than for the proton transfers
involving “normal” acids.33 However, there are sig-
Intrinsic Rate Constants. The proton-transfer rate
constants for 1H+ and 2H+ are reported in Table 1, those
for 3H+ and 4H+ in Table 2. Figures 1-5 show the
respective Brønsted plots where q and p represent
statistical factors.3,29 The Brønsted coefficients and log
ko values for the intrinsic rate constants are summarized
in Table 4; the latter values were obtained as log ko )
BH
CH
log(k1B/q) ) log(k-1/p) at pKa - pKa + log(p/q) ) 0.
Most Brønsted coefficients are in the range of 0.4-0.6
which is typical for proton transfer at carbon. There is a
small trend toward increasing R and decreasing â values
as DMSO is added to the solvent; this trend is seen for
(29) q is the number of equivalent basic sites on the amine, while p
is the number of equivalent protons on the protonated amine.
(30) For example, for the cationic acids 1H+, 3H+, and 4H+ the
increased acidity in the presence of DMSO means that the deproto-
nation becomes thermodynamically more favorable; according to the
Hammond31-Leffler32 postulate, this should lead to a transition state
with less proton transfer which is consistent with a smaller â value.
However, for the two neutral acids 5H and 6H DMSO decreases the
acidity, and hence, the observed trend in â is inconsistent with the
Hammond-Leffler postulate.
(27) For 5H and 6H, the acid is neutral (CH) and its conjugate base
(C-) is anionic.
(28) Buncel, E.; Wilson, H. Adv. Phys. Org. Chem. 1977, 14, 133.
(31) Hammond, G. S. J. Am. Chem. Soc. 1955, 77, 334.
J. Org. Chem, Vol. 70, No. 19, 2005 7727

Bernasconi et al.
TABLE 4. Brønsted Coefficients and Intrinsic Rate Constants at 20 °C
a Reference 18. b Bernasconi, C. F.; Kliner, D. A. V.; Mullin, A. S.; Ni, J.-X. J. Org. Chem. 1988, 53, 3342.
nificant differences in the log ko values as well as their
solvent dependence.
primary amines.4c,35 These results show that for a given
family of amines, the intrinsic rate constants for 1H+,
2H+, and 4H+ are significantly higher than for 3H+ and,
in the more aqueous solvents, also higher than for 5H
and 6H.
(1) For 1H+, 2H+, and 4H+ the log ko values are
relatively high and, for the reactions with primary
amines, range from about 3.1-3.8. For 4H+, data with
secondary alicyclic amines are also available. As is typical
for such proton transfers, the logko values for the second-
ary amines are 0.7-0.9 log units higher than for the
(34) (a) Eigen, M. Angew. Chem., Int. Ed. Engl. 1964, 1, 3. (b)
Ahrens, M.-L.; Maass, G. Angew. Chem., Int. Ed. Engl. 1968, 7,
818.
(35) The higher ko values are the result of differences in the solvation
energies of the respective protonated amines and the fact that at the
transition state solvation of the incipient protonated amines lags
behind proton transfer.4c,36,37
(32) Leffler, J. E.; Grunwald, E. Rates and Equilibria of Organic
Reactions; Wiley: New York, 1963; p 156.
(33) For normal acids, log ko is typically in the range of 8.5-9.5.34
7728 J. Org. Chem., Vol. 70, No. 19, 2005

Kinetics of Proton Transfer from Cationic Carbon Acids
For the reactions of 1H+ and 2H+ the transition state
may be described by 7. Since, due to the lag in delocal-
ization,4 the incipient negative charge on 7 is mainly
localized on the carbon, the positively charged PPh3 group
stabilizes this charge by its inductive effect. This results
stabilization of the conjugate base (4a-) and hence there
is diminished overall resonance stabilization of the
conjugate base. This reduces the ko-lowering resonance
effect and contributes to the relatively high ko values.
(4) The low ko values for the reactions of 5H in water
and 50% DMSO-50% water is mainly the result of the
strong resonance stabilization of the corresponding car-
banion (5a- and 5c-) which is reflected in the high acidity
of 5H discussed above. Hydrogen bonding solvation by
water which, like resonance/delocalization, lags behind
proton transfer at the transition state,4 contributes to the
lowering of ko. This latter factor becomes even more
prominent for 6- and explains why ko for nitromethane
is so low in the two most aqueous solvents (more on this
below).
Solvation Effect on ko. As discussed in much detail
elsewhere,4c solvation of products and desolvation of
reactants which are not synchronous with proton transfer
at the transition state can have a significant effect on
intrinsic rate constants and hence on their solvent
dependence. More specifically, solvation of incipient
products invariably lags behind proton transfer (develops
“late”) at the transition state while desolvation of reac-
tants is ahead of proton transfer (lost “early”). Even
though both of these effects lower the intrinsic rate
constants, a change in solvent can lead either to an
increase or decrease in ko. Following the formalism
developed earlier for the analysis of the solvent effect on
the intrinsic rate constant4c we can write eq 18 for the
reaction of a cationic carbon acid and eq 19 for a neutral
carbon acid with a neutral base.
in an increase in ko because in the conjugate base, where
the resonance structure 1a (2a) is dominant, the induc-
CH
tive effect is diminished and thus the effect on the pKa
is rather small. This increase in ko by the inductive effect
thus counteracts the typical ko-lowering effect of the lag
in the resonance development behind the proton transfer.
(2) For the reaction of 3H+ the transition state can be
represented as 8. In this case, there is at best a minimal
transition state stabilization by the inductive effect of the
N-methylpyridyl group because the positive charge is
farther away from the negative carbon than in 7, and
hence, there is no significant counter effect to the
ko-lowering resonance effect of the conjugate base. This
IδIIlogo ) δCH+ + δ + δ + δBH+ + δTS
(18)
(19)
B
C
feature, in combination with the greater resonance
stabilization of the conjugate base because of two major
resonance structures (3a and 3c) explains why ko is lower
than for 1H+ and 2H+.
IδIIlogo ) δCH + δB + δC- + δ + + δTS
BH
IδIIlog ko ) log ko(II) - log ko(I) represents the difference
For the reactions of 4H+ the transition state and
conjugate base are shown as 9 and 4-, respectively, with
4c- being the dominant resonance structure. In this case,
+
between the logko values in the two solvents, δCH (δCH),
-
+
δB, δC (δC ) and δBH are the contribution to the solvent
effect from the early desolvation of the carbon acid, the
early desolvation of the amine base, the late solvation of
the deprotonated carbon acid and the late solvation of
the protonated amine, respectively, while δTS includes
possible effects other than those arising from nonsyn-
chronous solvation/desolvation.38
+
The contributions of δB and δBH have been shown to
be small and quite independent of the specific reaction;
+
for H2O f 50% DMSO: δB ≈ -0.03, δBH ≈ -0.16; for
H2O f 90% DMSO: δB ≈ -0.01, δBH ≈ -0.25.4c For the
+
-
reaction of nitromethane (6), δCH, δC , and δTS for the
change from water to 90% DMSO were estimated as
-0.08, 3.15, and 0.84, respectively.4c This means that the
solvent effect on logko (0δ90 log ko ) 3.65; Table 4) is
-
dominated by the δC term. This is consistent with the
-90
0
very large log γC value of 6.70 (Table 3) which indi-
cates very much weaker solvation of the nitronate ion in
it is mainly the electron-withdrawing inductive effect of
the PhSO2 group which stabilizes the incipient negative
charge at the transition state and offsets the ko-lowering
effect of the lag in resonance development involving the
N-methylpyridinium system. Furthermore, since the
PhSO2 group is a weaker π-acceptor than the RCO
groups, it participates only minimally in the resonance
(36) Bell, R. P. The Proton in Chemistry, 2nd ed.; Cornell University
Press: Ithaca, 1973; Chapter 10.
(37) Jencks, W. P. Catalysis in Chemistry and Enzymology; McGraw-
Hill: New York, 1968; p 178.
+
-
+
(38) A further breakdown of δCH (δC), δB, δC (δC ), δBH , and δTS
into components that include the respective solvent activity coefficient
and a measure of the imbalance between solvation/desolvation and
proton transfer discussed in ref 4c is not necessary for the present
discussion and will not be attempted here.
J. Org. Chem, Vol. 70, No. 19, 2005 7729

Bernasconi et al.
bromomethane. 1H NMR δ (250 MHz, DMSO-d6): 4.36 (s, 3H,
CH3), 4.94 (s, 2H, CH2), 7.5-7.8 (m, H, Ph), 8.0-8.1 (m, 4H,
pyr).
The iodide salt of 4H+ was synthesized as described by
Wodzinski and Bunting20 except that the methylation step was
carried out with iodomethane instead of bromomethane. 1H
NMR δ (250 MHz, DMSO-d6): 4.34 (s, 3H, CH3), 5.22 (s, 2H,
CH2), 7.65-7.93 (m, 7H, ph and pyr), 8.95 (d, 2H, pyr).
Buffers and Other Reagents. Amines were obtained from
Aldrich as analytical grade and purified as follows. Piperidine,
morpholine, n-butylamine, 1-(2-hydroxyethyl)piperazine, and
2-methylethylamine were refluxed over CaH2 for 1 h and
distilled under nitrogen. Aminoacetonitrile hydrochloride,
methoxyamine hydrochloride, glycinamide hydrochloride, and
glycinethylester hydrochloride were recrystallized twice from
1:1 2-proponol/ethanol; piperazine was used without further
purification. Reagent grade DMSO was obtained from Fisher
Scientific. KOH and HCl solutions originated from “DILUT-
IT” analytical concentrates (J. T. Baker). Ultrapure water was
obtained form a Millipore MILLI-Q Plus water system.
90% DMSO than in water, i.e., the ko depressing effect
of the late solvation is strongly attenuated in the less
solvating solvent.
In light of the results for nitromethane the following
conclusions may be drawn for the other systems. (1) The
only other case where log ko increases significantly in 90%
DMSO (0δ90 log ko ) 1.18) is 5H which is the only other
system where log 0γC90 is positive (1.67). This implies
that in this case, too, the late solvation of the conjugate
base is the dominant factor while the other terms in eq
19 play a minor role. (2) The solvent effect on ko for all
the four cationic carbon acids is very small. This could
be due to a near cancellation of the various terms in eq
18. However, since there are large variations in the log
II
II
I
IγC and log γCH+ values for the various systems, it
seems more likely that none of the terms in eq 18 are
large. Hence, strong solvent effects only seem to occur
when the conjugate base is an anion that is prone to strong
hydrogen bond solvation in water.
General Procedures. The kinetic methodology, solution
CH
preparation, pH measurements, pKa determinations, and
measurements of solvent transfer activity coefficients were
essentially the same as described in ref 18. In most cases, the
pKa of the carbon acid was low enough that the rates could be
measured after mixing the carbon acid with the appropriate
amine buffer. For 4H+ in water the pKa was too high for such
methodology; in this case, 4 was generated in a dilute KOH
solution of 4H+ and this solution was subsequently mixed with
the amine buffer in a double-mixing stopped-flow apparatus.
Experimental Section
Substrates. 1H+ was purchased from Aldrich as the neutral
ylide 1. A hot methanolic solution of this material was vacuum
filtered to remove some insoluble gray solids. The filtrate was
diluted with water and chilled for 2 h. A crystalline solid was
collected and recrystallized from 2:1 water/methanol (v/v) to
give white crystals, mp 203-204 °C (lit.38 mp 205-206 °C).
The bromide salt of 2H+ was synthesized according to the
procedure of Ramirez and Dershowitz39 using bromo-
acteophenone40 as precursor.
Acknowledgment. This research was supported by
Grant Nos. CHE-9307659, CHE-0098553, and CHE-
00446622 from the National Science Foundation. E.L.
was supported by REU Grant No. CHE-9322464 (SURF
Program at UCSC) from the National Science Founda-
tion.
The iodide salt of 3H+ was obtained by following the
procedure of Stefanidis and Bunting5,6 except that the meth-
ylation step was carried out with iodomethane instead of
(39) Ramirez, F.; Dershowitz, S. J. Org. Chem. 1957, 22, 41.
(40) Cooper, R. M.; Davidson, L. H. In Organic Synthesis; Blatt, E.
H., Ed.; Wiley: New York, 1943; Collect. Vol. II, p 480.
JO051166R
7730 J. Org. Chem., Vol. 70, No. 19, 2005