The alkaline hydrolysis of sulfonate esters: Challenges in interpreting experimental and theoretical data

Article abstract of DOI:10.1021/jo402420t

Sulfonate ester hydrolysis has been the subject of recent debate, with experimental evidence interpreted in terms of both stepwise and concerted mechanisms. In particular, a recent study of the alkaline hydrolysis of a series of benzene arylsulfonates (Babtie et al., Org. Biomol. Chem. 10, 2012, 8095) presented a nonlinear Br?nsted plot, which was explained in terms of a change from a stepwise mechanism involving a pentavalent intermediate for poorer leaving groups to a fully concerted mechanism for good leaving groups and supported by a theoretical study. In the present work, we have performed a detailed computational study of the hydrolysis of these compounds and find no computational evidence for a thermodynamically stable intermediate for any of these compounds. Additionally, we have extended the experimental data to include pyridine-3-yl benzene sulfonate and its N-oxide and N-methylpyridinium derivatives. Inclusion of these compounds converts the Br?nsted plot to a moderately scattered but linear correlation and gives a very good Hammett correlation. These data suggest a concerted pathway for this reaction that proceeds via an early transition state with little bond cleavage to the leaving group, highlighting the care that needs to be taken with the interpretation of experimental and especially theoretical data.

Full text of DOI:10.1021/jo402420t

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pubs.acs.org/joc
The Alkaline Hydrolysis of Sulfonate Esters: Challenges in
Interpreting Experimental and Theoretical Data
Fernanda Duarte,† Ting Geng,† Gael Marloie,† Adel O. Al Hussain,‡ Nicholas H. Williams,*,‡
̈
and Shina Caroline Lynn Kamerlin*,†
†Department of Cell and Molecular Biology, Uppsala University, Uppsala 751 05, Sweden
‡Department of Chemistry, The University of Sheffield, Sheffield S3 7HF, United Kingdom
S
* Supporting Information
ABSTRACT: Sulfonate ester hydrolysis has been the subject
of recent debate, with experimental evidence interpreted in
terms of both stepwise and concerted mechanisms. In
particular, a recent study of the alkaline hydrolysis of a series
of benzene arylsulfonates (Babtie et al., Org. Biomol. Chem. 10,
2012, 8095) presented a nonlinear Brønsted plot, which was
explained in terms of a change from a stepwise mechanism
involving a pentavalent intermediate for poorer leaving groups
to a fully concerted mechanism for good leaving groups and
supported by a theoretical study. In the present work, we have
performed a detailed computational study of the hydrolysis of these compounds and find no computational evidence for a
thermodynamically stable intermediate for any of these compounds. Additionally, we have extended the experimental data to
include pyridine-3-yl benzene sulfonate and its N-oxide and N-methylpyridinium derivatives. Inclusion of these compounds
converts the Brønsted plot to a moderately scattered but linear correlation and gives a very good Hammett correlation. These
data suggest a concerted pathway for this reaction that proceeds via an early transition state with little bond cleavage to the
leaving group, highlighting the care that needs to be taken with the interpretation of experimental and especially theoretical data.
into the chemical similarities and differences between the
different substrates at the reactant level.
1. INTRODUCTION
Compared to the large body of work on phosphoryl transfer
reactions,1,2 sulfuryl transfer has received rather less attention.
However, recent years have seen a revival of interest in sulfate3,4
and other related group-transfer reactions.5,6 Sulfate hydrolysis
has directly important biological roles in, for instance, cellular
signaling7 and detoxification,8 and sulfonate esters can be used
by bacteria as sulfur sources in environments with low sulfate
concentrations.9,10 Furthermore, there has been an increased
awareness of the prevalence of catalytic promiscuity in enzymes
Although less studied than their phosphate counterparts,
there has been substantial research effort invested into
deciphering the mechanism and nature of the transition states
for the hydrolyses of sulfate3,4,7,16−19 and neutral sulfo-
nate6,20−29 monoesters. The hydrolysis of a sulfonate
monoester can, in principle, proceed through multiple
mechanisms, and arguments have been made in favor of both
stepwise addition−elimination mechanisms6,26−28 (involving a
pentacoordinate intermediate), as well as concerted mecha-
nisms24,29 (involving a single transition state) (Figure 1).
catalyzing phosphoryl and sulfuryl transfer reactions,11,12
phenomenon that has been suggested to play an important role
both in natural protein evolution13 and artificial enzyme
design.14
a
Recently, a phosphonate monoester hydrolase (PMH) from
Burkholderia caryophilli was found to be one of the most
promiscuous hydrolases characterized to date,15 catalyzing at
least five classes of hydrolytic reaction in addition to its native
phosphonate monoesterase activity. One of these activities is
the hydrolysis of xenobiotic sulfonate monoesters, and this
PMH is the only known enzyme capable of catalyzing the
hydrolysis of these compounds through direct S-OR cleavage
(as opposed to C−O cleavage).15 Elucidating the precise
molecular basis for this promiscuity will enhance our
understanding of the evolution of protein function, and an
important first step in this direction is obtaining detailed insight
Figure 1. Potential stepwise (A) and concerted (B) pathways for the
alkaline hydrolysis of aryl benzenesulfonates.
Received: October 31, 2013
© XXXX American Chemical Society
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Figure 2. (A) Brønsted correlation for the alkaline hydrolysis of aryl benzenesulfonates. Linear least squares fits to the low (blue) and high (green)
pKa data are shown as dotted lines, as reported in ref 6. The solid line is the least squares fit for a stepwise mechanism involving a change in rate-
limiting step. (B) Hammett correlation for the same data. The solid is the linear least squares fit, giving ρ = 1.83 0.10 and R2 = 0.981.
Although not inconceivable, a dissociative DN + AN
mechanism has been generally ruled out.23 In particular,
Williams and co-workers24 examined the reaction of a range of
oxyanions with basicities spanning a range of over 8 pKb units
with 4-nitrophenyl 4-nitrobenzenesulfonate. They also exam-
ined the reaction of phenoxide with substituted phenyl
benzenesulfonic acids. In both cases, they found good linear
correlations spanning the pKa of the incoming nucleophile,
causing them to argue that a single transition state provides the
best explanation for the process. Experimental studies on the
alkaline hydrolysis of diaryl sulfate diesters have supported a
similar mechanistic pathway.30
to become concerted, with bond cleavage accompanying the
proton transfer. The difference with the sulfonate ester
hydrolysis reactions discussed here is that the initial attack is
always assumed to be the rate-limiting step, even when an
intermediate is proposed.32
The status of the potential intermediate was investigated
using hybrid effective fragment potential (EFP)33−35 and
conductor-like polarizable continuum model36 (C-PCM)
calculations (QM(HF)/EFP/PCM approach) using a number
of explicit water molecules surrounding the sulfonate. These
showed that a pentavalent intermediate was only stable with the
poorer leaving groups, but not for the better leaving groups,
supporting the proposed mechanistic change. One concern
with these calculations is that it was necessary to introduce at
least 8 discrete water molecules into the calculations to obtain
the intermediate shown in Figure 2 of ref 6. We have discussed
the problems such a mixed solvation model introduces in terms
of quantitative accuracy of the calculations elsewhere37 and
guide interested readers to this work for further information.
Additionally, these calculations were performed at a relatively
low level of theory (HF), raising the possibility that this
intermediate is actually a simulation artifact (note also that only
a single stationary point with no corresponding transition
state(s) is presented in ref 6, and that all the water molecules
are placed where they selectively stabilize the nucleophile).
Thus it is possible that the calculations do not provide
unambiguous support for the proposed mechanism.
Overall, there is sufficient doubt about the interpretation and
analysis of these compounds to test whether an intermediate is
theoretically viable,6,24,26−29 so we have revisited the LFER
presented in ref 6, generating detailed free energy surfaces for
the hydrolysis of each compound, following our previous work
on the hydrolysis of phosphate and sulfate esters.3,38−41 We
also discuss the broader challenges of obtaining reliable
quantitative results for reactions involving both anionic
nucleophiles and polarizable atoms such as sulfur using
conventional density functional (DFT) approaches.
Very recent work6 has suggested a more complex scenario,
presenting a Brønsted plot for the alkaline hydrolysis of aryl
benzenesulfonates that breaks at pKa,lg ∼ 8.5, with βlg values of
−0.27 and −0.97 for leaving groups with pKas lower and higher
than 8.5, respectively (Figure 2A; dashed lines). A curve based
on a stepwise reaction involving a change in the rate-
determining step31 also fits the data very convincingly (Figure
2A, solid line; see also discussion in ref 6). However, this
interpretation would suggest that at the breakpoint, the loss of
either hydroxide or phenolate occur at equal rates, which is
implausible given the very different nucleofugalities of these
two leaving groups. For a stepwise process, the rate-limiting
step is expected to be the attack of the hydroxide nucleophile
for all the compounds studied. Thus, it was proposed that the
break in the Brønsted plot corresponds to a switch from a two-
step mechanism involving a pentavalent intermediate for poorer
leaving groups (Figure 1, pathway A), to a concerted
mechanism with good leaving groups (Figure 1, pathway B).
It was suggested that for substrates with a sufficiently good
leaving group (i.e., with a pKa lower than the breakpoint in the
plot), the attack of hydroxide leads to an intermediate that is
too unstable to have a significant lifetime, and so the reaction
has to become concerted.6 Although a change in mechanism
would usually show curvature that is concave upward in a
Brønsted plot, a similar concave-downward Brønsted plot is
observed for the reactions of phosphate monoester mono-
anions and has been rationalized in related terms. However, in
this analysis, the rate-limiting step is the breakdown of a
tautomer, which is assumed to form in a fast pre-equilibrium
step. When this tautomer is so reactive that it is predicted to
have a lifetime shorter than a bond vibration, the reaction has
Finally, we note that while the experimental data appear to
generate a smooth, nonlinear Brønsted plot (Figure 2A), a
Hammett plot of the same data yields a simple linear
correlation (Figure 2B), raising the question of whether
correlating the reactivity with the leaving group pKa is
appropriate. This has been noted in previous work concerning
B
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sulfonate esters: Buncel et al. concluded that σ (rather than σ−,
analogous to pKa) was a more appropriate parameter to use for
the ethoxide-promoted cleavage of sulfonate esters in
anhydrous ethanol,26 and Um et al. reported a similar result.28
These analyses led to the suggestion that the reaction proceeds
through a stepwise reaction, where the rate-limiting step is the
formation of a pentavalent intermediate. However, in
subsequent work, Um et al. concluded that a significant
contribution from both σ and σ− (derived from a Yukawa−
Tsuno analysis) provided the best fit under largely aqueous
conditions.29 These data include substituents in the 2 position,
which might be expected to generate scatter due to steric
interactions, thus confusing an interpretation based solely on
electronic variation. Most recently, Um and Buncel suggest that
this is the best analysis to use for the reaction of ethoxide with
aryl benzenesulfonate esters, and that the best description of
the reaction is concerted, with significant bond formation to the
nucleophile but with relatively small bond cleavage to the
leaving group in the transition state.42 Williams et al. also
considered these factors in the formation of a cyclic sulfonate
with aryloxy leaving groups, and showed that σ− (effectively a
Brønsted plot) rather than σ provided the most appropriate
analysis.23 Thus, there is ambiguity about the best way to
represent the kinetic data to reveal a reliable correlation.
Hence, to complement these theoretical studies, we also
studied the hydroxide-promoted hydrolysis of pyridin-3-yl
benzenesulfonate (1), along with its N-oxide (2) and N-
methylpyridinium (3) derivatives (Figure 3). The lowest pKa
Introduction. In our previous computational studies of the LFER for
related compounds,39,41 we have generated 2-D potential energy
surfaces for only one compound in the series, and we used this as a
starting point to obtain transition states with perturbed leaving groups
for the remainder of the series. However, in the present work, as there
is potential for two different mechanisms depending on leaving group,
we have generated individual potential energy surfaces for all
compounds of interest. These were obtained in the space defined by
S−O distances to the departing leaving group (x-axis) and the
incoming nucleophile (y-axis), spanning a range from 1.5 to 2.8 Å on
the x-axis, and 1.5 to 3.3 Å on the y-axis.
Figure 4. Aryl benzenesulfonates examined in this work.
We originally started with surfaces corresponding to three
compounds at either extremes of the series and at the break point
(i.e., the 3-F-4-NO2, 4-Cl and 3,4-dimethyl compounds), for which we
mapped the surface using a finer grid of 0.1 Å increments in the bond
distances on each axis. For the remaining compounds, we used a
slightly coarser grid of 0.15 Å increments. At each point on the plot,
the two degrees of freedom corresponding to the reaction coordinate
were kept frozen, while all other degrees of freedom were allowed to
freely optimize, and the surface was obtained by carefully pushing the
reaction coordinate in all relevant directions until the complete surface
was obtained. All points on the surface were obtained using Truhlar’s
M06-2X functional,43 which is a dispersion corrected hybrid
metaexchange-correlation functional that has been rigorously para-
metrized for organic compounds. Solvation was simulated implicitly,
using Cramer and Truhlar’s SMD solvation model.44 In order to save
computational cost, initial geometry optimizations were performed
using the smaller 6-31+G* basis set, followed by a single point
energetic correction to the obtained geometries using the larger 6-
311+G** basis set. In each case, the final surface was used to locate an
approximate transition state geometry, which was then optimized to a
saddle point using an unconstrained transition state optimization. The
resulting structures were characterized by frequency calculations, and
the minimum energy path connecting reactants to products through
this transition state was evaluated by calculating the intrinsic reaction
coordinate45,46 (IRC = ξ) in both the forward and reverse directions in
order to determine whether an intermediate could be found.
Figure S1A (Supporting Information) shows an overlay of the
calculated IRC from each transition state for each compound
(normalized to reactants), using the M06-2X functional, and Figure
S1B (Supporting Information) shows the change in S−Onuc and S−Olg
distances along the reaction coordinate for a representative compound
(4h). From these figures, it can be seen that, apart from the energetics
of the product state, the energy profiles for the hydrolysis are very
similar, with early reactant-like transition states. The final structure
from following the IRC in the reactant direction was then subjected to
an unconstrained geometry optimization in order to obtain the
geometry of the reactant complex, which was then used to evaluate the
activation barrier for the reaction. For comparison, the transition states
obtained using M06-2X were then reoptimized using a number of
other DFT functionals: the popular B3LYP47−49 functional, the
dispersion corrected ω-B97X-D,50 and finally CAM-B3LYP51 (which
was designed to correct delocalization error). New IRC calculations
were run from these TS in order to verify the identity of the
reoptimized transition states. We have also explored the effect of
including an increasing number of discrete water molecules (2, 8, and
10) in the calculations on the nature of the transition state using the
Hartree−Fock (HF) method, for direct comparison with previous
Figure 3. Structures of (1) pyridine-3-yl benzene sulfonate, as well as
(2) its N-oxide and (3) N-methylpyridinium derivatives.
leaving groups in Figure 2A all have substituents that provide
strong resonance stabilization of the oxyanion product; if this
effect is not significant in the transition state, this may account
for why they are less reactive than predicted by the line defined
by the poorer (nonresonance stabilized) leaving groups.
Therefore, these additional compounds were selected as they
also have very good leaving groups, but due to inductive and
field effects rather than direct conjugation of the anion with the
π-system of a substituent. If the Brønsted plot really does reveal
limiting behavior dependent on leaving group ability, these
compounds will fall close to previously reported data; if they do
not, and the data can be described by a smooth Hammett
correlation, then the experimental evidence for discontinuous
behavior will no longer be apparent.
On the basis of the results of our combined theoretical and
experimental studies presented in this report, we conclude that
a concerted reaction mechanism involving a single transition
state provides a better interpretation of the combined data
presented here and in ref 6. Importantly, we believe that this
system provides an excellent demonstration of the pitfalls in the
interpretation of experimental and especially computational
data.
2. COMPUTATIONAL METHODS
Figure 4 shows the substituted aryl benzenesulfonates studied in this
work and originally presented in ref 6. This set of compounds was then
supplemented with compounds 1−3 (Figure 3), as discussed in the
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work.6 Finally, free energies in the absence of explicit water molecules
were obtained by evaluating the zero point energies and entropies
from calculated vibrational frequencies using the 6-311+G** basis set,
SMD solvation model and relevant DFT functional. Note that we did
not include the free energy corrections when examining systems with
extra discrete water molecules, as the frequency calculations become
increasingly unreliable as the total number of degrees of freedom is
increased. All calculations presented in this work have been performed
using the Gaussian 09 quantum chemical suite of programs.52
Potential energy surfaces (PES) for hydroxide attack on 3-
fluoro-4-nitrophenyl (pKa 6.11), 3-nitrophenyl (pKa 8.35), 4-
chlorophenyl (pKa 8.61), and 3,4-dimethyl benzene sulfonate
(pKa 10.36) are shown in Figure 6A−D, with the PES for the
remaining compounds being presented in Figure S2A−D
(Supporting Information). The corresponding transition state
geometries, obtained using optimization with M06-2X/6-
311+G**, are shown in Figure 7A−D and Figure S3A−D
(Supporting Information), respectively. The compounds high-
lighted in Figures 6 and 7 correspond to the two pKa extremes
of the LFER presented in Figure 1 in ref 6, as well as the two
compounds that lie on the possible break-point of this plot. As
can be seen from these figures, in all cases the potential energy
surface indicates the presence of a single, concerted transition
state, with no intermediate. The approximate transition state
from each surface was then subjected to a full geometry
optimization and characterized as described in the Computa-
tional Methods, and the resulting energetics (including a
breakdown of different contributions to the calculated
activation barrier) as well as S−O distances to the incoming
nucleophile and departing leaving group at the transition state
are highlighted in Tables 1 and 2, respectively. The fact that
these are concerted transition states is further verified by IRC
calculations in both forward and reverse directions, which all
lead to either reactant or product complexes. Note that all
energetics presented in Table 1 are relative to the
corresponding reactant complex obtained by following the
IRC as far as possible in the reactant direction, and concluding
with a final geometry optimization to obtain the stationary
point. In terms of geometric parameters (Table 2), while there
appears to be negligible change in the S−Olg distance moving
across the series, there is a gradual increase in S−Onuc distance,
from 2.35 Å for the compound with the poorest leaving group
(3,4-dimethyl benzene sulfonate) to 2.47 Å for the compound
with the best leaving group (3-fluoro-4-nitro benzene
sulfonate), suggesting that the better the leaving group, the
earlier there is commitment to the reaction, resulting in a
gradually more expansive transition state. The shifting
transition states observed across these series of compounds
follow a trend we previously observed for both methyl
arylphosphate diesters41 and fluorophosphates39 (although
these transition states are very slightly more expansive than
the ones obtained here).
3. RESULTS AND DISCUSSION
3.1. Free Energy Surfaces and Initial Characterization
of Transition States. To probe the potential mechanism(s)
for the hydrolysis of the compounds listed in Figure 4, our
starting point involved the generation of 2-D potential energy
surfaces for hydroxide attack on each of these compounds, as
outlined in the Computational Methods. Here, the first
question involves identifying the relevant conformation of the
two aromatic rings with respect to each other for our
optimization, as they could take either of two conformations
as illustrated in Figure 5. In the first of these, the two aromatic
rings have a weak π-stacking interaction, whereas in the second,
the rings avoid each other. The latter conformation is the one
used in the calculations of ref 6; however, we were only able to
obtain optimized transition states in implicit solvent using the
conformation in which the two rings stack together. We have
performed an energy scan of the O−S−Olg−C dihedral angle in
the parent compound (Figure 5) to characterize the effect of
adjusting the position of the two aryl rings with respect to each
other.
As an aside, an important issue to take into account in our
calculations is the challenges of obtaining quantiative accuracy
in calculations involving hydroxide as a nucleophile (see also
discussion and other examples in the literature39,54−57),
particularly when using an implicit solvent model. That is,
including a charged nucleophile creates a major underestimate
in the calculated energetics; similar problems have been seen
before, both in our simulations of hydroxide attack on the
phenyl phosphate dianion,39 and in independent studies of the
hydrolysis of p-nitrophenyl phosphate55 and acetate56 (to name
a few examples). We believe that this underestimate arises from
the combination of a number of factors. The first issue is simply
due to the known problems with correctly solvating anionic
species and the hydroxide ion in particular using dielectric
continuum models (see discussion in the literature57). That is,
while the precise value of the solvation free energy of the
hydroxide ion remains controversial58,59 due in part to the
challenges evaluating ΔsG*(H+) and misunderstandings about
corrections for gas/solution phase standard states,60 dielectric
continuum models tend to significantly understimate hydroxide
Figure 5. Scan of the O−S−Olg−C dihedral angle in the ground state
of the parent compound.
From this figure, it can be seen that there are two minima
along the reaction coordinate, corresponding to each of the
conformations outlined above, and that the conformation with
the weak π-stacking interaction appears to be the favored
conformation in the ground state. Therefore, this is the
conformation we have used for all our subsequent calculations
in implicit solvent. It is important to bear in mind that this
could be an artifact of the implicit solvent model. However, as
discussed in Section 3.2, we would like to emphasize that
performing the calculations using the alternate conformation
(corresponding to that used in ref 6) in the presence of a
number of explicit water molecules does not change the
qualitative nature of our results, and either conformation would
in principle form a viable starting point to similar effect.
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Figure 6. Free-energy surfaces for the alkaline hydrolysis of aryl benzenesulfonate. Shown here are 3-fluoro-4-nitrophenyl, 3-nitrophenyl, 4-
chlorophenyl, and 3,4-dimethylphenyl benzenesulfonates. The approximate positions of the relevant transition states are indicated by ‡, and the
actual optimized structures are shown in Figure 7
much. Tying in with this is the static correlation error.61 These
issues can be magnified in large systems, because the
delocalization error increases with system size.62 Additionally,
one would expect them to be compounded when more
challenging atoms such as sulfur,67,68 which is polarizable and
has low-lying vacant d-orbitals, are introduced into the
calculation, and where the quantitative reliability of DFT
approaches becomes extremely unpredictable and functional/
basis set dependent.68−70 However, it should be noted that
much more reasonable energetics have been obtained for
sulfate hydrolysis using a neutral nucleophile,3 making it likely
that the presence of the charged nucleophile is a significant part
of the problem. Finally, implicit solvent models are known to
underestimate activation barriers due to the fact that they
neglect nonequilibrium solvation effects.71
Clearly, the majority of these issues are extremely
challenging, and resolving them is out of the scope of the
present work. However, one trivial source of error that can be
addressed is the solvation free energy of the hydroxide ion.
That is, as there is no bond formation to the nucleophile in the
reactant complex, one can simply adjust the calculated value of
the solvation free energy of the reactant complex by the error in
the calculated solvation free energy of hydroxide ion compared
to the experimental value, thus introducing a constant
correction to all calculated solvation free energies (see also
related discussion72). The only caveat with this approach is the
continued uncertainty as to what this solvation free energy
actually is. Values in the range of −90.6 to −110.0 kcal/
mol58−60,73−75 have been reported for the solvation free energy
of the hydroxide ion in the literature, although at present most
sources appear to converge on −104.5 kcal/mol58,60 as being
the most reliable estimate. The value obtained using 6-
311+G**/M06-2X/SMD is −97.3 kcal/mol, which is clearly
far too low; therefore, we have decided to follow the advice of
the literature58,60 and use −104.5 kcal/mol as the “exper-
imental” solvation free energy for hydroxide, based on the
Figure 7. Geometries of optimized transition states for alkaline
hydrolysis of (A) 3-fluoro-4-nitrophenyl, (B) 3-nitrophenyl, (C) 4-
chlorophenyl, and (D) 3,4-dimehtylphenyl benzenesulfonates. These
structures were obtained by optimization of the approximate
geometries highlighted on the surfaces shown in Figure 6. The
corresponding transition states for the remaining compounds shown in
Figure 3 are presented in Figure S3 (Supporting Information).
solvation, which would in turn lead to an underestimate of the
calculated activation barrier. In combination with problems
with incorrect solvation of the hydroxide ion, DFT approaches
tend to underestimate barrier heights,61−64 although this
problem is somewhat mitigated by the M06-series of
functionals,65 which was part of our rationale for using this
functional in this work. One source of this problem is the
delocalization error,62,66 which refers to the tendency of the
approach to artificially spread the electron density out too
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a
Table 1. Energy Decomposition for ΔG‡ at the M06-2X/6311+G** Level of Theory
calc
b
b
c
substituted phenol
pKa
ΔE‡
ΔΔGsolv
ΔZPE
−TΔS‡
ΔG‡
ΔG‡
gas
calc
exp
3F-4-NO2
4-NO2
4-CN
6.11
7.14
7.95
8.35
8.61
9.38
9.95
10.36
−21.9
−22.9
−22.2
−21.8
−21.8
−20.9
−19.5
−19.9
33.4
34.8
34.6
34.1
34.4
34.0
33.1
33.9
−0.1
−0.1
−0.2
0.1
3.2
3.1
3.0
3.2
3.3
3.8
3.6
2.2
14.5
17.3
18.0
18.1
18.2
18.3
19.3
20.2
20.9
15.0
15.2
15.5
15.9
16.8
17.3
15.8
3-NO2
3-CN
0.1
4-Cl
0.0
H
0.2
3,4-dimethyl
−0.5
a
b
All energies are in kcal/mol, relative to the reactant complex. Zero point energies and entropies were obtained from frequency calculations at
c
323.15 K. All experimental values have been corrected for the entropic cost of bringing the reacting fragments into the reacting cage (K = 0.017
M−1), following ref 53. This is important for consistency, as our reference state in the calculations is a reactant complex (obtained by following the
minimum energy path from the relevant transition state) and not the individual fragments at infinite separation from each other.
Table 2. Comparison of S−O Distances to the Nucleophile
(S−Onuc) and Leaving Group (S−Olg) at the Reactant and
Transition States for the Alkaline Hydrolysis of the Benzene
accuracy with experiment, with calculated activation barriers
that lie within ∼4 kcal/mol of the experimental value. It should
be pointed out that the inclusion of this correction is based on
the assumption that the error in the solvation free energy of
hydroxide is absent (or reduced by a constant value) at the
transition state along this series of compounds. As the distance
of the hydroxide ion from the reacting center increases with
increasing acidity of the leaving group (see below), there is a
risk that this is not completely true. Nevertheless, while
obtaining absolute quantitative accuracy is challenging in these
cases, the relative trends should still be informative.
To validate our results, we repeated our transition state
optimizations and IRC calculations using a number of different
functionals with various corrections implemented into them to
see how much these would effect our results. Our starting point
was just to compare our approach to the standard and popular
B3LYP functional,47−49 which does not include a complete
dispersion treatment (in contrast to M06-2X). We then also
included CAM-B3LYP,51 which has been developed to correct
the delocalization error, and another dispersion corrected
functional, ω-B97X-D,50 which also includes a correction for
long-range effects, as discussed in the Computational Methods.
A comparison of calculated total activation energies using these
different functionals, as well as the resulting rate constants
(obtained using transition state theory), are shown in Table 3
and Figure 8, respectively. From Figure 8, it can be seen that
while we have fairly consistent deviations between calculated
a
Sulfonates Shown in Figure 4
b
b
b
RS
TS
ΔΔRS → TS
substrate
pKa
S−Onuc S−Olg S−Onuc S−Olg S−Onuc S−Olg
3F-4NO2
4-NO2
4-CN
3-NO2
3-CN
4-Cl
6.11
7.14
7.95
8.35
8.61
9.38
9.95
10.36
3.62
3.63
3.59
3.63
3.61
3.67
3.63
3.64
1.65
1.64
1.64
1.64
1.64
1.63
1.63
1.63
2.47
2.41
2.41
2.42
2.40
2.38
2.38
2.35
1.72
1.72
1.72
1.71
1.71
1.71
1.71
1.71
−1.15
−1.21
−1.19
−1.21
−1.21
−1.29
−1.25
−1.29
0.07
0.08
0.08
0.07
0.08
0.08
0.08
0.08
H
3,4-CH3
a
All geometry optimizations were performed using the M06-2X
b
functional, and all distances are shown in Å. RS and TS denote
reactant and transition states, respectively. ΔΔRS → TS denotes the
change in distance upon moving from RS to TS respectively.
related discussion about the accuracy of values for the free
energy of hydration of H+. This results in a correction of 7.2
kcal/mol for the solvation free energy of OH−, which we have
added to the ΔΔGsolv values shown in Table 2 (see Table S1
(Supporting Information) for the error and corresponding
correction we obtain using other functionals). As can be seen
from Table 1, once the calculated solvation free energies are
adjusted for this correction, we obtain better quantitative
Table 3. Calculated Activation Barriers (ΔG‡calc, kcal/mol) and Rate Constants (log kcalc) for the Different Functionals Tested
a
in This Work
substituted phenol
functional
3-F-4-NO2
6.11
4-NO2
7.14
4-CN
7.95
3-NO2
8.35
3-CN
8.61
4-Cl
9.38
H
3,4-dimethyl
10.36
pKa
ΔG‡
9.95
M06-2X
14.5
3.00
15.0
2.68
15.2
2.52
15.5
2.32
15.9
2.06
16.8
1.44
17.3
1.11
15.8
2.16
calc
log kcalc
B3LYP
ΔG‡
19.7
21.4
22.4
21.3
24.3
23.0
23.1
24.6
calc
log kcalc
−0.52
20.2
−1.64
21.2
−2.34
21.2
−1.55
20.8
−3.63
21.0
−2.71
22.1
−2.81
22.3
−3.83
23.0
ω-B97X-D
CAM-B3LYP
experiment
ΔG‡
calc
log kcalc
−0.81
19.3
−1.50
21.6
−1.53
22.9
−1.22
21.7
−1.36
22.3
−2.15
23.2
−2.25
23.5
−2.73
21.7
ΔG‡
calc
log kcalc
−0.25
17.3
−1.75
18.0
−2.65
18.1
−1.82
18.2
−2.22
18.3
−2.84
19.3
−3.10
20.2
−1.82
20.9
ΔG‡
exp
log kexp
1.12
0.68
0.56
0.51
0.42
−0.25
−0.86
−1.28
a
The corresponding experimental values (ΔG‡exp and log kexp) are also presented here for comparison. All energies are presented in kcal/mol. Note
that the experimental energetics have been corrected for the cost of bringing the reacting fragments from infinite separation into the encounter
complex, as discussed in the caption to Table 1.
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Figure 8. A comparison of calculated (pink circles) and experimental (blue circles) log k obtained when taking into account just ΔE‡ (A) and also
when including zero point energy and entropy corrections (B) using the M06-2X43 (1), B3LYP47−49 (2), ω-B97X-D50 (3), and CAM-B3LYP51 (4)
functionals, respectively.
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Figure 9. (A) S−Onuc and (B) S−Olg distances at the transition state for different leaving groups with different DFT functionals, respectively.
and experimental values across the series, we are not able to
qualitatively reproduce the experimental values and instead
obtain very scattered plots. The problem here is that the
observed changes in activation barrier across the series are so
small (4 kcal/mol between the most reactive and least reactive
compounds) and are overshadowed by the error margin in the
entropies obtained from the vibrational frequencies. This is
supported by the fact that much better qualitative agreement is
obtained between theory and experiment if only the trends
obtained from ΔE‡ are used, without the addition of any free
energy corrections.
using implicit solvation, we have also explored the effect of
introducing discrete water molecules into our calculations.
Although increasingly popular, there are a number of potential
pitfalls associated with such an approach. Most critically, as
discussed in ref 37 and references cited therein, it is unclear
whether using a mixed implicit/explicit solvation model with a
limited number of discrete water molecules reproduces the
correct polarization boundary conditions between the explicit
solvent and the bulk continuum. If it does not, this would
potentially result in overpolarization of the water molecules in
the first solvation shell, or, to put it more simply, discrete water
molecules embedded into a continuum do not necessarily
behave like water molecules embedded into a large, explicit
water sphere. In addition to this, including explicit water
molecules in the geometry optimization makes it necessary to
take into account the entropic cost associated with releasing
these water molecules from the artificial QM-optimized
positions (which can be quite large37). Nevertheless, if treated
carefully, as shown by the literature,55,56 the inclusion of a
cluster of explicit water molecules can provide more reasonable
energetics when dealing with alkaline nucleophiles, where
workers other than us have also observed the continuum model
alone to drastically underestimate the calculated activation
barrier.
As a starting point, we have performed a full QM(HF)
(rather than QM(HF)/MM with the water molecules treated
classically) optimization of the phenyl benzene sulfonate
intermediate presented in ref 6 using the Cartesian coordinates
provided in the Supporting Information of ref 6 to determine
whether this is a true stationary point or a Hartree−Fock
artifact.76−78 While the structure presented in ref 6 is not a true
minimum using a full QM description in which the water
molecules are treated as part of the quantum system (based on
frequency calculations which give an imaginary frequency of
−142.7633 cm−1), it was nevertheless possible to use this as a
starting point to obtain both an intermediate structure as well
as transition states for both addition and elimination steps of
the reaction. These were then connected to each other by
means of IRC calculations to obtain a consistent profile and
relevant stationary points (Figure 10). From Figure 10, it can
be seen that while this reaction is tending toward a mechanism
that is stepwise in terms of bonding pattern when including
explicit water molecules, the apparent “intermediate” from the
optimization is in fact not thermodynamically stable. Note that
The corresponding energy breakdowns for each functional
are shown in Table S2 (Supporting Information), and a
comparison of the changes in S−Onuc/lg distances for each
functional are shown in Table S3 (Supporting Information). In
all cases, the calculated solvation free energies have been
corrected for the experimental solvation free energy of the
hydroxide ion, as was done in Table 1 for M06-2X and outlined
above. Finally, Figures 9A and B show changes in S−Onuc and
S−Olg distances at the transition state for different leaving
groups with different DFT functionals, respectively. The first
thing that can be observed is that, while there are clearly
changes in absolute values upon moving to different func-
tionals, overall trends are not affected by changing the
functional. In all cases, we obtain very early transition states
with very little bond cleavage to the leaving group (S−Olg
distances in a range of 1.70 to 1.80 Å), and some bond
formation beginning to occur to the nucleophile (S−O
distances in the range of 2.28 to 2.53 Å). Additionally, in all
cases, S−Onuc distance increases with increasing acidity of the
leaving group, whereas the S−Olg distance remains largely
unchanged. This is also reflected in the calculated charges
(Figure S5, Supporting Information), which range from −0.5 to
−0.8 on the leaving group oxygen atom and −1.08 to −1.32 on
the nucleophile oxygen, depending on level of theory and
approach used to calculate the charges. Despite the challenges
with quantitatively reproducing the experimental activation
barriers, the qualitative data these calculations provide all lead
to a convergent picture that is in agreement with experimental
observables. The crucial factor here is that all functionals tested
provide the same trend, making it less likely that this is an
accidental observation.
3.2. Exploring the Effect of Explicit Water Molecules
on the Transition States. Following on from our calculations
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its simplicity, implicit solvation is sufficient to obtain a reliable
description of the reacting system in this case (see also
discussion in the Supporting Information of ref 41).
3.3. Discussion of Kinetic Data. The rates of reaction of
compounds 1−3 under the same conditions as the data shown
in Figure 2 are added to the previously reported data in Figure
11, and the raw data is shown in Figure S6 (Supporting
Information). Inspection of Figure 11A shows that the apparent
smooth curvature of the Brønsted plot has disappeared to
become a linear, somewhat scattered correlation. Figure 11B
shows that the new data contributes to a Hammett plot that
retains a much better linear correlation. This correlation shows
a slight concave-downward curvature, consistent with Ham-
mond movement of the transition state. The much better
correlation with σ rather than pKa suggests that the leaving
group has not broken its bond to the sulfur to any significant
extent in the transition state, and so mesomeric interactions do
not have a large impact. These data, which suggest little bond
cleavage to the leaving group in the transition state and a small
amount of structural variation with leaving group ability, is also
consistent with the theoretical picture. According to the
calculations, the main changes that occur lie in the interaction
of the nucleophile with the sulfur atom. This will be affected by
the electronic character of the leaving group oxygen, which will
perturb the electrophilicity of the sulfur atom and so affect the
reaction rate. The minimal variation in leaving group bond
lengths suggests that bond cleavage is not extensive; a slight
lengthening is to be expected as the nature of the sulfur atom in
the bond is altered. The better leaving groups have a longer
bond to the nucleophile in the transition state, suggesting that
the nature of the electrophilic sulfur has changed less, and a
correspondingly smaller variation with σ is observed in this
region of the plot; i.e., there is slightly concave downward
curvature of the Hammett plot.
Figure 10. Energy profile for the hydrolyses of (A) phenyl benzene
diol and (B) the corresponding 4-cyano substituted compound, at the
HF/6-31+G** level of theory, including eight explicit water molecules
and the CPCM continuum solvent model. Here, the two aromatic
rings were placed in the same conformation as in ref 6, on the basis of
having used the coordinates provided in the Supporting Information of
this work as a starting conformation.
although the profiles in Figure 10 do not include a correction
for the configurational entropy, this would be expected to be
the same for both transition states and the intermediate, and
will therefore be unlikely to change the relative energetics of
these species. Additionally, as would be expected, as one moves
to a better leaving group (e.g., X = 4-CN, Figure 10), the
reaction becomes unquestionably concerted and the shoulder
observed with the poorer leaving group vanishes. However, it is
important to emphasize that even the “intermediate” obtained
with the parent compound is not thermodynamically stable and
appears to be just an inflection point along the intrinsic reaction
coordinate. Additionally (and more critically), it is striking that
all the water molecules in the calculation in ref 6 are placed so
as to stabilize the nucleophile, with none stabilizing the leaving
group. This could create an artifactual intermediate by
artificially stabilizing only one side of the reacting species. To
test this, we added 2 extra water molecules to stabilize the
leaving group. Upon doing this, we were no longer able to
obtain a stable intermediate at the HF level of theory,
highlighting the dangers of selective microsolvation using
explicit water molecules. Overall, it would appear that despite
The ρ value is 1.61; this can be compared with the ρ value for
the overall equilibrium process to provide some insight into the
reaction progress as described by the effective charge on the
leaving group. Williams has measured the β for the equilibrium
as 1.8 (which also leads to an assignment of an effective charge
of 0.879 on the leaving oxygen atom in the starting substrate).
This β value can be converted to an equilibrium value for ρ by
using the ρ value for the ionization of phenols (2.1), and so the
equilibrium ρ = 2.1 × 1.8 = 3.8. This in turn leads to an
Figure 11. (A) Brønsted correlation for the alkaline hydrolysis of aryl benzenesulfonates (black points reported in ref 6; red points correspond to
compounds 1−3). The red dotted line is the linear least-squares fit to all the data, giving β = −0.67 0.07 and R2 = 0.912. (B) Hammett correlation
for the same data. The red dotted line is the linear least-squares fit to all the data, giving ρ = 1.61 0.07 and R2 = 0.983.
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estimate of the Leffler index for this reaction as 1.61/3.8 = 0.42,
which further reinforces the early nature of the transition state.
This corresponds to a change of effective charge on the leaving
oxygen atom from +0.8 to about 0 in the transition state (0.8 −
0.42 × 1.8 = 0.04). Williams has thoroughly described this
method of analysis,80 showing that the principal assumptions
linking the analysis of type II LFER apply to both Hammett
and Brønsted correlations. In molecular terms, as the hydroxide
attacks the sulfur, the leaving group is converted from a
strongly electron-withdrawing group to an essentially neutral
substituent. The ρ value reflects this change in character, rather
than any significant cleavage of the sulfur−oxygen bond.
The conclusions from this analysis are broadly consistent
with the earlier analysis of Buncel, concerning the reactions of
sulfonate esters with ethoxide in ethanol.26,27 These authors
show a clearly improved correlation with σ rather than σ− and
concluded that this could be intepreted in either the rate-
limiting formation of an intermediate or a transition state that
closely resembles this. The large ρ values measured for these
reactions led Buncel to suggest that the reaction closely
resembles the formation of a pentavalent species in the rate-
limiting step, or a transition state that closely resembles a
putative pentavalent intermediate. Likewise, the recent work of
Um and Buncel suggests a small resonance demand from the
leaving group in the transition state in anhydrous ethanol is
consistent with this description of the reaction character.42 The
experimental and theoretical analysis here suggests that the
reaction in water is rather earlier in character.
We note that it is difficult to find aryloxy leaving groups that
have pKas lower than 8 that do not have 2-substituents and/or
strong resonance interactions with the substituent. These
substituted pyridin-3-yl types of compounds are the only
readily available aryloxy leaving groups that we have been able
to find that are not substituted in the 2 position and have low
pKas due to inductive and field rather than resonance effects
(and so do not have dissimilar σ and σ− parameters, in contrast
to most phenols with low pKas). For many slow reactions, good
leaving groups are required for practical purposes: however, it is
possible that these data are skewed by the extensive use of
substituents that rely on resonance interactions to perturb the
pKa significantly. Herschlag and Zalatan81 have noted that these
types of compounds may be outliers in the plot of methyl aryl
phosphate diesters and analyzed a series of reactions to justify
treating 4-NO2 and 4-CN substitutents as outliers in Brønsted
plots; in ref 6, it was noted that perhaps the explanation
advanced in terms of intermediate instability might provide an
alternative explanation for that data too. However, the data
presented here suggests that this explanation is not required. A
reasonable approach to correlating the phosphate diester series
is to apply a Yukawa−Tsuno (or Young−Jencks) analysis to the
data, which reveals a significant (∼33%) resonance demand to
be included alongside the inductive and field effects (Figure S7,
Supporting Information). Thus, in these reactions, the degree
of bond cleavage at the transition state is greater, and
delocalization plays a more significant role in its stabilization.
This is supported by theoretical calculations on this series using
the B3LYP functional,41 which show shifting transition states
with P−Onuc distances similar to those presented in Table S3
(Supporting Information), but with slightly longer P−Olg
distances compared to the arylsulfonates presented in the
present work. The greater degree of bond cleavage presumably
allows for better synchronization of the charge delocalization
with the charge development on the leaving group, although
again there is limited variation in P−Olg distances across the
series. One can reasonably expect a continuum: as the
correlation becomes better with σ−, the Hammett ρ or
Brønsted β values will increase. At low dependence, one
might expect delocalization to be less important (as a result of
either an early transition state or imperfect synchronization of
charges between nucleophile and leaving group) and Hammett
to provide the more appropriate correlation. Combining these
parameters through a Yukawa−Tsuno (or Young−Jencks82)
analysis probably provides a better approach to correlating the
effect of varying phenolic leaving groups on rates of reaction.
Recent work83 has proposed the introduction of an additional
saturation term (for electron-releasing substituents) in an
extended Yukawa−Tsuno equation to analyze phenolates, but
there are seldom enough data in these reactivity-based LFER to
justify the inclusion of a third term. Furthermore, because for
slow reactions the phenols chosen tend to be more rather than
less acidic than the parent and so these leaving groups feature
less heavily than those with strongly electron withdrawing
groups, in the data presented here, leaving groups with strongly
electron-releasing groups do not feature. Lastly, in considering
how to correlate reactivity data with the electronic properties of
a leaving group, we can suggest that the substituted pyridyl
leaving groups provide a useful series that extend the series of
leaving groups only affected by inductive and field effects down
to low pKa, which is likely to be useful in analyzing leaving
group ability, perhaps better than, and certainly a complement
to, the widely utilized 4-nitrophenyl substituent.
3.4. A Revised Mechanistic Picture Based on
Computational and Experimental Evidence. On the
basis of the computational and experimental evidence
presented in this work, the data are most simply analyzed in
terms of a concerted pathway, where bond cleavage to the
leaving group is not greatly advanced in the transition state. It is
possible that there is a transient intermediate, but it is not
required and there is no evidence for it.
The previous explanation6 suggested that there might be a
transition from a marginally stable intermediate to one that has
no lifetime. This may be true but would not be expected to
make a significant difference to the structure of the rate-limiting
transtion state, and hence would not lead to markedly different
sensitivity to the properties of the leaving group due to this
qualitative transition. The scenario here is somewat different to
the case where the rate-limiting step leads to an intermediate
that becomes too unstable to exist, at which point the reaction
is forced to become concerted. This can lead to downward
curvature but is similar to a change in rate-limiting step, rather
than a change in mechanism.
4. CONCLUSIONS
We find that in all cases the free energy surface provides a
single, concerted reaction pathway, with compact transition
states that become gradually more dissociative with increasing
acidity of the leaving group. In addition, we have explored the
effect of including 2, 6, and 8 explicit water molecules in the
calculation and demonstrate that the intermediate obtained
using HF/6-31++G* has no appreciable lifetime. This
intermediate appears to be a simulation artifact, because it
also disappears as one adds extra water molecules to stabilize
the leaving group as well as the nucleophile. Overall, these
results are consistent with our previous studies on the alkaline
hydrolysis of methyl arylphosphate diesters,41 as well as the
corresponding fluorophosphates,39 and experimental work
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8.42.84 3-Hydroxypyridine-N-oxide: 6.45.85 3-Hydroxypyridine-N-
methyl iodide: 4.96.84 The Hammett sigma values used for these
substituents were as follows: 3-hydroxypyridyl, 0.67;86 3-hydroxypyr-
idyl-N-oxide, 1.59;87 3-hydroxy-N-methylpyridinium, 2.31.86 All other
sigma values were taken from the literature.88 See the Supporting
Information for a discussion of these parameters.
arguing in favor of a rate-determining transition state with little
bond cleavage to the leaving group.26−29 Additionally, we
demonstrate, in alignment with the observations of other
workers,68−70 that performing reliable quantitative studies of
hydrolysis of sulfur-containing compounds is extremely
challenging, particularly if alkaline nucleophiles are involved.
However, through validation using several different theoretical
approaches, we are reasonably confident with the qualitative
information obtained from our calculations. Specifically, on the
basis of the free energy surfaces presented in this work, as well
as transition states obtained at several levels of theory both with
and without the inclusion of discrete water molecules, we
demonstrate that the alkaline hydrolysis of all compounds in
this series proceeds via a single-step mechanism with a
concerted transition state, in line with previous experimental
interpretations,24 arguing against a change in mechanism or
rate-determining step as one moves across the series. Finally,
exploring the effect of adding explicit water molecules shows
that while seemingly a stationary point, the structure shown in
ref 6 does not correspond to a thermodynamically stable
intermediate even at the HF level of theory, and particularly if
additional water molecules are included to stabilize the leaving
group in addition to the nucleophile, any hint even of an
intermediate vanishes.
The results presented here are qualitatively very similar to
those obtained for similar calculations on fluorophosphates39
and methyl arylphosphate diesters41 and show a shifting
transition state that becomes gradually more reactant-like as the
leaving group becomes more acidic. The key difference between
these compounds is in the solvation effects, with the neutral
sulfonate esters being substantially solvent destabilized at the
TS (because of better solvation of the hydroxide ion in the
ground state), whereas the transition states of the monoanionic
phosphate diesters and fluorophosphates appear to be stabilized
by solvent. Therefore, while the transition states for these
compounds are geometrically quite similar, as in the related
comparison of phosphate and sulfate monoesters,3 there are
some fundamental chemical differences between these com-
pounds, putting significant challenges on unusual enzymes like
BcPMH, which can nevertheless catalyze both reactions within
the same active site.15 The precise molecular basis for why this
happens is still unresolved, but the present work is an
important step toward understanding such catalytic promiscuity
at the atomic level.
Pyridin-3-yl benzenesulfonate89 (1). Benzenesulfonyl chloride
(2.56 mL, 20 mmol) and pyridin-3-ol (1.90 g, 20 mmol) were stirred
in THF (50 mL) at room temperature, and triethylamine (3.35 mL, 24
mmol) and pyridine (1.93 mL, 24 mmol) were added dropwise. The
reaction was stirred at room temperature for 16 h, the precipitate was
filtered off, and the solvent was removed in vacuo. The residue was
dissolved in dichloromethane (20 mL) and washed with water (10
mL) and then sodium bicarbonate solution (10 mL). The solution was
dried over sodium sulfate before removing the solvent and purifying
the crude product by column chromatography on silica (67% 40−60
petroleum ether:33% ethyl acetate) to yield 3.54 g of 1 (74%) as a
colorless solid: mp 46−47 °C; δH (400 MHz, CDCl3) 7.32 (1 H, dd,
4.6, 8.3), 7.48 (1 H, ddd, 1.3, 2.7, 8.3), 7.58 (1 H, t, 7.5), 7.73 (1 H, t,
7.5), 7.86 (1 H, d, 7.5), 8.18 (1 H, d, 2.7), 8.53 (1 H, dd, 1.3, 4.6); δC
(100 MHz, CDCl3) 124.2, 128.5, 129.4, 130.1, 134.69, 134.74 (C−S),
144.0, 146.4 (C−O), 148.3; ESI-MS positive ion mode m/z 236
(MH+, 100%); HRMS (TOF mode) calculated for C11H10NO3S,
236.0381, found 236.0370.
1-Oxidopyridin-3-yl benzenesulfonate (2). 3-Pyridyl benzene-
sulfonate (0.50 g, 2.1 mmol) and 3-chloroperbenzoic acid (0.55 g, 3.2
mmol) in chloroform (50 mL) were stirred at room tempertaure for
24 h. The solvent was removed in vacuo, and the crude product was
purified by column chromatography on silica (95% ethyl acetate:5%
methanol) to yield 0.24 g of 2 (45%) as a white solid: mp 97−98 °C;
δH (400 MHz, CDCl3), 7.14 (1 H, d, 8.5), 7.26 (1 H, dd, 6.5, 8.5),
7.62 (1 H, t, 7.5), 7.77 (1 H, t, 7.5), 7.89 (1 H, s), 7.90 (1 H, d, 7.5),
8.13 (1 H, d, 6.5); δC (100 MHz, CD3OD) 122.9, 126.6, 128.4, 129.6,
134.0 (C−S), 134.6, 135.3, 138.1, 148.2 (C−O); ESI-MS positive ion
mode m/z 252 (MH+, 100%); HRMS (TOF mode) calculated for
C11H10NO4S, 252.0331, found 252.0331.
1-Methyl-3-[(phenylsulfonyl)oxy]pyridinium iodide89 (3). 3-
Pyridyl benzenesulfonate (0.50 g, 2.1 mmol) and methyl iodide (1.32
mL, 21 mmol) in acetone (50 mL) were refluxed for 16 h. The solvent
was removed in vacuo, and the residue was dissolved in the minimum
amount of methanol. The product was precipitated by the dropwise
addition of diethyl ether, filtered, and washed with diethyl ether to
yield 0.2 g of 3 (25%) as a yellow solid: mp 144.8−145 °C; δH (400
MHz, D2O) 4.24 (3 H, s), 7.57 (1 H, t, 7.5), 7.76 (1 H, t, J = 7.5), 7.81
(1 H, d, 7.5), 7.90 (1 H, dd, 6.1, 8.7), 8.09 (1 H, d, 8.7), 8.68 (1 H, d,
6.1); δC (100 MHz, CD3OD) 48.4 (CH3), 128.6, 128.8, 130.0, 133.5
(C−S), 135.7, 139.0, 141.0, 144.7, 148.3 (C−O); ESI-MS positive ion
mode m/z 250 (M+ − I−, 100%); HRMS (TOF mode) calculated for
C12H12NO3S, 250.0538, found 250.0527.
5. EXPERIMENTAL METHODS
ASSOCIATED CONTENT
■
Compound 1 was synthesized from phenyl sulfonyl chloride and 3-
hydroxypyridine in the presence of pyridine and triethylamine. 1 was
converted to 2 using 3-chloroperbenzoic acid in chloroform, and to 3
by reaction with excess methyl iodide in refluxing acetone.
S
* Supporting Information
Complementary plots for 4-nitrophenyl, 4-cyanophenyl, 3-
nitrophenyl, and phenyl benzenesulfonates, as well as structures
of relevant transition states, atomic charges at the transition
state, absolute energies, and Cartesian coordinates of key
stationary points. Discussion of sigma values used in Hammett
plots. This material is available free of charge via the Internet at
http://pubs.acs.org.
The rates of reaction of compounds 1−3 were monitored by UV
spectroscopy at 50 °C in solutions of KOH with the ionic strength
maintained at 0.5 M with KCl. 1: 300 nm, [KOH] 0.01−0.1 M. 2: 320
nm, [KOH] 0.004−0.07 M. 3: 320 nm, [KOH] 0.002−0.008 M.
Reactions were initated by adding 5−15 μL of a stock solution in
DMSO or dioxan (0.05 M) to 3 mL of KOH solution that had
equilibrated at 50 °C. All reactions showed excellent first order
behavior, and observed rate constants were obtained by fitting the
absorbance change to the integrated first order rate equation. Each
compound showed good first order dependence on hydroxide when
the observed rate constants were plotted against [KOH] (see the
Supporting Information). The rate constant calculated for 0.1 M KOH
was used in the Brønsted and Hammett plots shown, to allow direct
comparison with the data as reported in ref 6. The pKa values of the
leaving groups were obtained from the literature. 3-Hydroxypyridine:
AUTHOR INFORMATION
■
Corresponding Authors
*E-mail: n.h.williams@sheffield.ac.uk.
*E-mail: kamerlin@icm.uu.se.
Notes
The authors declare no competing financial interest.
K
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ACKNOWLEDGMENTS
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The European Research Council has provided financial support
under the European Community’s Seventh Framework
Programme (FP7/2007-2013)/ERC Grant Agreement No.
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