Enhanced Chelate Cooperativity in Polar Solvents

Article abstract of DOI:10.1021/jacs.7b01765

High-throughput UV-vis titrations in combination with chemical double-mutant cycles (DMCs) have been used to study the competition of a polar solvent for formation of intramolecular H-bonds. Twenty-four different zinc porphyrin-pyridine complexes were investigated in mixtures of toluene and phenol. DMCs were used to determine effective molarities (EM) for the formation of intramolecular phenol-amide H-bonds as a function of solvent composition. The values of EM increase by an order of magnitude with increasing concentrations of the more polar solvent, phenol. Phenol solvates the amide groups on the ligands strongly, increasing the steric bulk and destabilizing the complexes. These adverse steric interactions are removed when intramolecular H-bonds are formed and therefore provide an increased driving force for formation of cooperative interactions. The result is that the effects of competitive interactions with polar solvents that reduce binding affinity are attenuated to a significant extent by a corresponding increase in EM in multivalent complexes.

Full text of DOI:10.1021/jacs.7b01765

Article
pubs.acs.org/JACS
Enhanced Chelate Cooperativity in Polar Solvents
Stefan Henkel, Maria Cristina Misuraca, Yudi Ding, Maxime Guitet, and Christopher A. Hunter*
Department of Chemistry, University of Cambridge, Lensfield Road, Cambridge CB2 1EW, U.K.
S
* Supporting Information
ABSTRACT: High-throughput UV−vis titrations in combination with chemical
double-mutant cycles (DMCs) have been used to study the competition of a polar
solvent for formation of intramolecular H-bonds. Twenty-four different zinc
porphyrin−pyridine complexes were investigated in mixtures of toluene and phenol.
DMCs were used to determine effective molarities (EM) for the formation of
intramolecular phenol−amide H-bonds as a function of solvent composition. The
values of EM increase by an order of magnitude with increasing concentrations of
the more polar solvent, phenol. Phenol solvates the amide groups on the ligands
strongly, increasing the steric bulk and destabilizing the complexes. These adverse
steric interactions are removed when intramolecular H-bonds are formed and
therefore provide an increased driving force for formation of cooperative
interactions. The result is that the effects of competitive interactions with polar
solvents that reduce binding affinity are attenuated to a significant extent by a
corresponding increase in EM in multivalent complexes.
molecule.34,35 This approach allows the interplay between
INTRODUCTION
■
solvent−solvent and solvent−solute interactions to be
Systematic studies of the mechanisms governing molecular
quantified. When solvent mixtures are involved, molecular
recognition events are fundamental to understanding and
recognition probes are particularly valuable in focusing
controlling biological processes and the advancement of
measurements on solvation equilibria at specific sites on the
material science and nanotechnology. Almost all systems of
solute. The association constant (K) for formation of a 1:1
interest in these fields involve multiple cooperative interactions
complex between a H-bond donor (D) and acceptor (A) in a
between macromolecular surfaces, and many factors can
mixture of a nonpolar solvent (S1) and a polar solvent (S2) was
contribute to the thermodynamic properties of these interfaces:
found to obey eq 1 for solvents that are not strongly self-
electrostatic interactions, H-bonds, aromatic interactions, and
associated.36
desolvation.1−15
KS1
We have been investigating chelate cooperativity in terms of
effective molarity (EM)16−20 and the relationship of this
parameter with the number and nature of the intermolecular
interactions and the geometric complementarity and conforma-
tional flexibility of the binding partners.21−33 These experi-
ments show that the thermodynamic properties of multivalent
systems can be described as the sum of contributions from
individual interaction sites. The value of EM for an intra-
molecular interaction does not depend on the intrinsic strength
of the noncovalent interaction but is a function of the
supramolecular architecture, specifically geometric complemen-
tarity.27 Although preorganization does affect EM values, the
influence of conformational flexibility on the overall stability of
supramolecular complexes is modest. Flexible recognition
interfaces can adapt, leading to surprisingly stable complexes
even when complementarity appears poor. The next step to a
better understanding of cooperativity in multivalent complexes
is to investigate how chelate cooperativity is affected by the
solvent environment, specifically the impact of preferential
solvation in solvent mixtures.
K =
(1 + KD[S2])(1 + KA[S2])
(1)
Here, KS1 is the D·A association constant in S1 and KD and KA
are the D·S2 and A·S2 association constants in S1.
Figure 1 illustrates eq 1 graphically: the binding constant for
the D·A complex is independent of the concentration of S2
when the concentration of the polar solvent is low. However,
when the concentration of S2 is increased, the polar solvent
starts competing with S1 for solvation of the solutes. At this
point, preferential solvation by S2 competes with the
interaction of A with D leading to a less stable complex at a
higher concentration of S2. A linear dependence of log K on
log[S2] (slope ≈ −1) was observed for solvent mixtures where
S2 contained a good H-bond acceptor and no strong H-bond
donor sites (e.g., di-n-octyl ether) or S2 contained a H-bond
donor site and no strong acceptors (e.g., 1,1,2,2-tetrachloro-
ethane).36 When S2 contained both H-bond donor and
acceptor sites, a greater dependence on log[S2] was observed.
For example, the slope was −2 when S2 was an alcohol.37 The
We have previously shown that molecular recognition-based
probes provide insight into the thermodynamic properties of
the solvation shell at individual binding sites on the surface of a
Received: February 20, 2017
© XXXX American Chemical Society
A
DOI: 10.1021/jacs.7b01765
J. Am. Chem. Soc. XXXX, XXX, XXX−XXX

Journal of the American Chemical Society
Article
The chemical double-mutant cycle (DMC) shown in Figure
3 can be used to determine the value of EM in such systems.
Figure 1. Association constant (log K) for formation of an
intermolecular H-bond between a donor (D) and an acceptor (A)
in mixtures of a nonpolar solvent (S1) and a more polar solvent (S2)
as a function of S2 concentration (log[S2]). The horizontal dashed
line corresponds to the association constant of the D·A complex in S1,
and the vertical dashed line corresponds to the association constant of
S2 with one of the solutes in S1.
slope of the plot of log K versus log[S2] effectively quantifies
the number of S2−solute interactions that compete with
formation of the D·A complex. In multivalent association, we
might therefore expect that the slope of this plot could be used
to measure the number of intermolecular interactions involved
in binding, in effect a H-bond inventory.
Figure 3. Chemical double mutant cycle used to determine the EM for
formation of an intramolecular amide−phenol H-bond.
The DMC was introduced as an analytical tool to quantify
individual contributions to the overall stability of supra-
molecular complexes.28,38,39 The free energy contribution due
to the intramolecular H-bond in complex A can be determined
using eq 2. An alternative expression (eq 3) can be used to
describe this free energy contribution in terms of an apparent
association constant for formation of intramolecular H-bond in
complex A, Kintra. For porphyrin−ligand complexes that can
form only one intramolecular H-bond, Kintra can be directly
related to the value of EM (eq 4). For porphyrin−ligand
complexes that can form two or more H-bonds, Kintra represents
the overall equilibrium constant for formation of singly and
multiply H-bonded states
APPROACH
■
The porphyrin−ligand system depicted in Figure 2 has been
exploited for the quantitative analysis of chelate coopera-
°
°
°
°
°
ΔΔG = ΔGA − ΔGB − ΔGC + ΔGD
(2)
Kintra = e−ΔΔG°/RT − 1
Kintra = σKref EM
(3)
(4)
where σ is a statistical factor accounting for the degeneracies of
the complexes.
Here, we describe binding studies of porphyrin−ligand
complexes carried out in mixtures of two solvents, toluene (S1)
and phenol (S2). The impact of the polar solvent competing
for the intramolecular amide−phenol H-bonds was evaluated
through application of DMCs to determine the dependence of
the intramolecular equilibrium constant Kintra on the concen-
tration of S2.
The porphyrins and pyridine ligands used are shown in
Figure 4. Synthesis of these compounds has been reported in
previous publications.24,28 Porphyrins P1a−P4a are equipped
with phenol moieties, which are able to interact with the H-
bond acceptors on the ligand arms. Porphyrins P1b−P4b have
methoxy groups instead and cannot form H-bonds with the
ligands. Ligands L2e and L3e have either one or two arms
bearing amide H-bond acceptor sites that can interact with the
porphyrin phenol groups. It was shown previously that
porphyrins P1a, P2a, and P3a all form H-bonds with the
amide groups of L2e and L3e, but in P4a the donor and
acceptor sites are too far apart and no H-bonds are observed.24
Figure 2. Stepwise equilibria in the formation of a complex between a
zinc porphyrin and a pyridine ligand equipped with H-bonding sites.
K0 is the association constant for the zinc−pyridine interaction, Kref is
the association constant for formation of an intermolecular phenol−
amide H-bond, EM is the effective molarity associated with the
formation of the intramolecular phenol-amide H-bond illustrated, and
σ is a statistical factor accounting for the degeneracies of the
complexes.
tivity.18,22−33 The zinc−nitrogen interaction between the
porphyrin and the pyridine ligand provides a strong
coordination bond that allows investigation of formation of
weaker peripheral H-bonds. The ability of the system to form
intramolecular H-bonding interactions can be quantified in
terms of effective molarity (EM). Strictly speaking, the H-bond
in Figure 2 is an intermolecular interaction, but we will use the
term intramolecular because if we consider stepwise equilibria
where the zinc−nitrogen interaction is formed first the
formation of the H-bond is intramolecular.
B
DOI: 10.1021/jacs.7b01765
J. Am. Chem. Soc. XXXX, XXX, XXX−XXX

Journal of the American Chemical Society
Article
low concentrations of S2, the stabilities of the complexes are
not perturbed by the presence of the second solvent. For P4,
complexes A−D all have similar stabilities, which confirms that
there are no H-bonding interactions in this system at any
concentration of S2. For porphyrins P1−P3, complex A is
significantly more stable than complexes B, C, and D. This
observation is consistent with the presence of strong intra-
molecular amide−phenol H-bonds in these systems.
At higher concentrations of S2, the polar solvent starts to
compete with the intramolecular phenol−amide H-bond,
leading to a drop in the stability of complex A when [S2] is
greater than 10 mM for P1−P3. The concentration at which S2
begins to preferentially solvate the amide groups of the free
ligand is given by the association constant for formation of an
intermolecular phenol−amide H-bond, Kref = 86 M−1, which
was measured previously using the p-cresol·N,N-diethylaceta-
mide complex in toluene (Figure 6a).24 When Kref[S2] < 1, the
amide groups of the free ligand are mainly solvated by S1, but
when Kref[S2] > 1, they are mainly solvated by S2.
At even higher concentrations of S2, the association
constants for complexes B, C, and D also start to decrease
(Figure 5), which suggests that phenol competes for the zinc−
nitrogen coordination bond. The decrease in log K observed for
complex A therefore comprises two different contributions
because phenol competes for both the intramolecular H-bond
and the zinc−nitrogen interaction (Figure 6b). However, the
DMC is specifically designed so that any effects due to changes
in the strength of the zinc−nitrogen interaction will cancel out.
The data in Figure 5 were used in eqs 2 and 3 to determine
Kintra, and the results are shown in Figure 7. The values
obtained for the ligand with two H-bonding sites L3e are
roughly twice the values found for the one-armed ligand L2e,
confirming the additivity of the free energy contributions from
multiple interactions in these systems.
Figure 4. Porphyrins and ligands used in binding studies.
The ester carbonyl groups of L2e and L3e do not form H-
bonds in any of the complexes because they are too far from the
peripheral phenol H-bond donors. Similarly, ligand L3b does
not form H-bonds with any of the porphyrins, so this ligand
was used as the non-H-bonding control ligand in complexes B
and D of the DMC.
RESULTS AND DISCUSSION
■
Binding of the pyridine ligands to each of the eight porphyrins
was studied by UV−vis titration experiments. A shift of about
10 nm in the absorption maximum of the Soret band (425 nm)
is indicative of binding of the pyridine nitrogen to the zinc of
the porphyrin, and this change can be used to determine the
association constant for the complex. Due to the large number
of association constants measured for each porphyrin−ligand
system, automated protocols were developed using a UV−vis
plate reader equipped with internal syringes that were used to
vary the solvent composition. The experimental data were fit to
1:1 binding isotherms using purpose-written software, which
allowed analysis of 18 titration data sets in a single experimental
run.
For the P4a complexes where no H-bonds are formed, Kintra
is close to zero and constant over the entire range of phenol
concentrations (not shown in Figure 7). For the P1a−P3a
complexes, the log Kintra values are constant for phenol
Figure 5 shows the association constants for formation of the
porphyrin−ligand complexes plotted as a function of S2
concentration (see the SI for tabulated data). As expected, at
Figure 5. Association constants (log K) for formation of 1:1 porphyrin−ligand complexes in toluene−phenol mixtures plotted as a function of the
concentration of phenol (log[S2]). For each porphyrin−ligand combination, data for the four complexes used to construct the DMC in Figure 3 are
plotted together; i.e., A corresponds to the Pa·Le complex, B to Pa·Lb, C to Pb·Le, and D to Pb·Lb. The vertical dashed line corresponds to −log
Kref, the association constant for formation of a phenol−amide H-bond in toluene.
C
DOI: 10.1021/jacs.7b01765
J. Am. Chem. Soc. XXXX, XXX, XXX−XXX

Journal of the American Chemical Society
Article
Within the L2e complexes the slopes vary by a factor of 2
between different porphyrins. Similarly the slopes for the L3e
complexes span a range between −0.5 and −1.0. The slope of
the logKintra versus log[S2] plot clearly does not provide the H-
bond inventory suggested in the Introduction.
Phenol can self-associate at high concentrations, and
although the self-association constant in aromatic solvents is
low,40 it is possible that aggregation could affect the results
shown in Figure 7 and Table 1. In order to check the behavior
of phenol at high concentrations in mixtures with toluene, the
intermolecular H-bond formed between 4-phenylazophenol
and N,N-dihexylacetamide was studied in the same solvent
mixtures. The results are shown in Figure 8. The behavior
Figure 6. (a) Intermolecular phenol−amide H-bond formed between
p-cresol and N,N-diethylacetamide used to estimate Kref.24 (b) S2
(PhOH) competes for both the intramolecular phenol−amide H-bond
in complex A and the zinc−nitrogen coordination bond.
Figure 8. Values of log K for formation of a complex between 4-
phenylazophenol and N,N-dihexylacetamide as a function of phenol
concentration in toluene (log[S2]). The gradient of the linear part of
the curve is −1.
illustrated in Figure 1 is observed. When [S2] < 10 mM, log K
is constant, and when [S2] > 10 mM, log K decreases in
proportion to log[S2] with a slope of −1. Since phenol is a
good H-bond donor and a poor H-bond acceptor, the
competition for formation of the intermolecular H-bond is
dominated by interactions of S2 with the amide H-bond
acceptor. In other words, KD in eq 1 is small, and KA is large,
which leads to the expected slope of −1 in the limit of large
values of [S2] (eq 5).
Figure 7. Values of log Kintra for formation of intramolecular amide−
phenol H-bonds plotted as a function of phenol concentration in
toluene (log[S2]).
KS1
K ≈
concentrations less than 10 mM, but above this concentration,
log Kintra decreases as the concentration of phenol increases.
The slopes of the log Kintra versus log[S2] plots in this region
are shown in Table 1. The slopes for the L3e complexes are
(1 + KA[S2])
(5)
where KS1 = 260 M−1 and KA = Kref = 86 M−1 in toluene.
It should be possible to use eqs 1−5 in the same way to
rationalize the relationship between log Kintra and log[S2]
observed for the intramolecular H-bonds in the porphyrin−
ligand complexes. Neglecting the KD term in eq 1 gives eq 5,
which successfully describes the data shown in Figure 8 for the
intermolecular H-bond. Combining eqs 4 and 5 gives eq 6 for
the L2e complexes that can make one intramolecular H-bond.
Equation 7 is the equivalent relationship for the L3e complex
that can make two intramolecular H-bonds.
Table 1. Effective Molarities (EM) and Slopes (mexp) of
logKintra versus log[S2] Plots24
L2e
L3e
EM (mM)
mexp
EM (mM)
mexp
P1a
P2a
P3a
P4a
140
15
−0.6
−0.3
−0.4
0
130
14
−1.0
−0.5
−0.6
0
36
27
KS1EM
1 + KA[S2]
K
intra(L2e) = 4
(6)
(7)
2
⎛
⎞
approximately double the values found for the L2e complexes,
as might be expected for systems that make twice as many H-
bonds. In the L2e complexes, the polar solvent competes with
one intramolecular H-bond, and in the L3e complexes S2
competes with two intramolecular H-bonds. However, the
slopes are all substantially less negative than might be expected
based on the number of H-bonds present in the complexes.
KS1EM
1 + KA[S2]
KS1EM
K
intra(L3e) = 8
+ 12
⎜
⎟
⎠
1 + K [S2]
⎝
A
Here, KS1 = KA = Kref = 86 M−1 in toluene.
These equations show that when KA[S2] < 1, the presence of
the second polar solvent has no impact on the stability of
complex A. When KA[S2] > 1, the interaction of S2 with the
D
DOI: 10.1021/jacs.7b01765
J. Am. Chem. Soc. XXXX, XXX, XXX−XXX

Journal of the American Chemical Society
Article
amide groups of the free ligand competes for formation of the
intramolecular H-bond in complex A. For the systems
described here, KS1 ≈ KA, so the relative populations of the
two zinc porphyrin−pyridine complexes shown in Figure 6b are
determined by the relationship between EM and [S2]. When
[S2] > EM, the intramolecular H-bond will not be significantly
populated because the amide groups of the ligand will
preferentially interact with the solvent. When [S2] < EM, the
fully bound state of complex A with the intramolecular H-bond
will dominate. However, even in systems where the intra-
molecular H-bond is highly populated due to a high value of
EM, the presence of a polar solvent will still reduce the
observed association constant for formation of complex A when
KA[S2] > 1, because the free ligand will be solvated by S2.
Thus, the concentration at which S2 starts to affect the stability
of a cooperatively assembled complex is determined by KA and
not by EM.
10 shows the difference in the free energy of complexation of
complexes C and D as a function of [S2] (ΔGC° − ΔGD°). At
Figure 10. Free energy difference between complexes C and D of the
DMC (ΔGC° − ΔGD°) as a function of phenol concentration in
toluene (log[S2]).
Equation 6 is similar to eq 5, which suggests that in the limit
of large values of [S2], the slope of the log Kintra versus log[S2]
plot for an intramolecular H-bond should be −1, as observed
for the corresponding intermolecular H-bond in Figure 8.
However, the experimental slopes observed for the intra-
molecular H-bonds formed by the one-armed ligand L2e are
significantly less negative (Table 1). Moreover, eq 6 suggests
that the limiting slope should be independent of the value of
EM, which is not the case. For the two-armed ligand L3e, the
relationship between the logKintra and log[S2] is more
complicated (eq 7). In the limit of large values of [S2], the
slope is expected to be −2 when EM > [S2] and −1 when EM
< [S2]. The experimental results in Table 1 are substantially
less negative than these limiting values.
low concentrations of S2, the differences are small, but when
[S2] > 10 mM, significant concentration-dependent differences
are observed. The positive value of ΔGC° − ΔGD° indicates
that there is a destabilization of complex C associated with the
presence of the amide groups in the Le ligands. The shape of
the plots in Figure 10 is similar to that of the plots in Figure 7,
which suggests that the destabilization measured by the changes
in ΔGC° − ΔGD° are due to phenol binding to the amide
groups of the ligands (KA = 86 M−1). There are no H-bonding
interactions in complexes C and D, so the differences must be
related to steric effects associated with phenol binding to the
amide groups. The steric effects are similar for all of the
porphyrins indicating that it is interactions with the porphyrin
core that are responsible rather than clashes with the
substituents on the m-phenyl substituents.
Equations 6 and 7 can be solved to determine values of EM
as a function of [S2], and the results are shown in Figure 9. In
Similar steric effects must be present when the amide ligands
(Le) bind to the phenol-substituted porphyrins (Pa), and these
interactions could explain the solvent-dependence of EM
illustrated in Figure 9. The consequences of steric interactions
with S2 are illustrated in Figure 11. When the amide acceptors
Figure 9. Relationship between EM for formation of intramolecular
amide−phenol H-bonds as a function of phenol concentration in
toluene (log[S2]).
Figure 11. Schematic representation of steric effects associated with
solvation by a polar solvent (S2) on intramolecular interactions
between A and D. When A is solvated by S2, the increased steric
demand destabilizes the partially bound state, and this strain is relieved
on formation of the A·D interaction in the fully bound state.
all cases, EM is not a constant but increases significantly with
[S2] when the concentration of phenol exceeds 10 mM. The
values of EM obtained for the one-armed and two-armed
ligands are practically identical and vary in the same way with
[S2]. Over the concentration range studied, EM increases by an
order of magnitude, and the behavior is similar for all six of the
porphyin-ligand complexes, which suggests that neither the
overall supramolecular architecture nor the geometric com-
plementarity play a significant role in determining the solvent
dependence of EM.
on the ligands (A) are solvated by phenol (S2), steric clashes
between S2 and the porphyrin destabilize the partially bound
state in which the zinc−nitrogen coordination bond is formed
but the A·D H-bond is not. Formation of the A·D H-bond
removes S2 and the steric clashes, leading to a higher
equilibrium constant for the intramolecular process. The result
is that, as the concentration of S2 increases and the ligands are
increasingly solvated by S2, the intermediate partially bound
Closer examination of the experimental data displayed in
Figure 5 reveals that complex C is less stable than complexes B
and D at high concentrations of S2 in all eight systems. Figure
E
DOI: 10.1021/jacs.7b01765
J. Am. Chem. Soc. XXXX, XXX, XXX−XXX

Journal of the American Chemical Society
Article
state becomes less stable (cf. complex C), and EM for
formation of the intramolecular interaction in complex A
increases. These results are consistent with a previous
observation that EMs measured for these systems in
tetrachloroethane are somewhat larger than the values in
other solvents.30
thanks the German Research Foundation (DFG) for a
Postdoctoral Research Fellowship.
REFERENCES
■
(1) Persch, E.; Dumele, O.; Diederich, F. Angew. Chem., Int. Ed. 2015,
54 (11), 3290−3327.
(2) Ariga, K.; Ito, H.; Hill, J. P.; Tsukube, H. Chem. Soc. Rev. 2012, 41
(17), 5800−5835.
CONCLUSIONS
■
(3) Kay, E. R.; Leigh, D. A.; Zerbetto, F. Angew. Chem., Int. Ed. 2007,
46 (1−2), 72−191.
The effects of preferential solvation on chelate cooperativity
have been quantified using a family of zinc porphyrin−pyridine
ligand complexes in mixtures of toluene and phenol. Chemical
double mutant cycles have been used to measure the effective
molarities for the formation of intramolecular phenol−amide
H-bonds in these systems. In all cases, EM increases
dramatically with increasing concentrations of the more polar
solvent, phenol. The result is that the effects of competitive
interactions with polar solvents that reduce binding affinity are
attenuated to a significant extent by a corresponding increase in
EM in multivalent complexes.
(4) Fasting, C.; Schalley, C. A.; Weber, M.; Seitz, O.; Hecht, S.;
Koksch, B.; Dernedde, J.; Graf, C.; Knapp, E. W.; Haag, R. Angew.
Chem., Int. Ed. 2012, 51 (42), 10472−10498.
(5) Badjic, J. D.; Nelson, A.; Cantrill, S. J.; Turnbull, W. B.; Stoddart,
J. F. Acc. Chem. Res. 2005, 38 (9), 723−732.
(6) Philp, D.; Stoddart, J. F. Angew. Chem., Int. Ed. Engl. 1996, 35
(11), 1154−1196.
(7) Mulder, A.; Huskens, J.; Reinhoudt, D. N. Org. Biomol. Chem.
2004, 2 (23), 3409−3424.
(8) Mammen, M.; Choi, S. K.; Whitesides, G. M. Angew. Chem., Int.
Ed. 1998, 37 (20), 2755−2794.
The origin of the effect can be identified as strong solvation
of the amide H-bond acceptor groups by the polar solvent,
which increases the effective steric bulk of the ligands. These
steric interactions have a destabilizing effect on the zinc
porphyrin−pyridine complexes, but if an intramolecular H-
bond is made, the amide group must be desolvated, removing
the adverse steric interactions. The result is an increase in EM
by an order of magnitude compared with the corresponding
intramolecular interactions in toluene.
The behavior observed in these systems is expected to be
general: polar solvents will enhance cooperativity between
multiple interaction sites that are in close proximity. In polar
solvents, polar functional groups are strongly solvated. If
multiple polar groups are in close proximity on the surface of a
molecule, then formation of an intermolecular interaction at
any one site will lead to steric interactions with the strongly
bound solvation shell at neighboring sites. These adverse steric
interactions will be removed, if all of the functional groups are
desolvated to form cooperative intermolecular interactions.
Thus, enhanced cooperativity is expected for strongly solvated
systems.
(9) O’Sullivan, M. C.; Sprafke, J. K.; Kondratuk, D. V.; Rinfray, C.;
Claridge, T. D. W.; Saywell, A.; Blunt, M. O.; O’Shea, J. N.; Beton, P.
H.; Malfois, M.; Anderson, H. L. Nature 2011, 469 (7328), 72−75.
(10) Gargano, J. M.; Ngo, T.; Kim, J. Y.; Acheson, D. W. K.; Lees, W.
J. J. Am. Chem. Soc. 2001, 123 (51), 12909−12910.
(11) Dam, T. K.; Roy, R.; Das, S. K.; Oscarson, S.; Brewer, C. F. J.
Biol. Chem. 2000, 275 (19), 14223−14230.
(12) Bundle, D. R.; Kitov, P. I.; Sadowska, J. M.; Mulvey, G.;
Armstrong, G. D.; Ling, H.; Pannu, N. S.; Read, R. J. Nature 2000, 403
(6770), 669−672.
(13) Anderson, H. L. Inorg. Chem. 1994, 33 (5), 972−981.
(14) De Greef, T. F. A.; Smulders, M. M. J.; Wolffs, M.; Schenning,
A.; Sijbesma, R. P.; Meijer, E. W. Chem. Rev. 2009, 109 (11), 5687−
5754.
(15) Mahadevi, A. S.; Sastry, G. N. Chem. Rev. 2016, 116 (5), 2775−
2825.
(16) Motloch, P.; Hunter, C. A., Thermodynamic Effective Molarities
for Supramolecular Complexes. In Advances in Physical Organic
Chemistry; Williams, I. H., Williams, N. H., Eds.; Academic Press,
2016; Vol. 50, pp 77−118.
(17) Hunter, C. A.; Anderson, H. L. Angew. Chem., Int. Ed. 2009, 48
(41), 7488−7499.
(18) Hogben, H. J.; Sprafke, J. K.; Hoffmann, M.; Pawlicki, M.;
Anderson, H. L. J. Am. Chem. Soc. 2011, 133 (51), 20962−20969.
(19) Ercolani, G. J. Phys. Chem. B 2003, 107 (21), 5052−5057.
(20) Ercolani, G.; Schiaffino, L. Angew. Chem., Int. Ed. 2011, 50 (8),
1762−1768.
(21) Hunter, C. A.; Ihekwaba, N.; Misuraca, M. C.; Segarra-Maset,
M. D.; Turega, S. M. Chem. Commun. 2009, 26, 3964−3966.
(22) Sun, H. M.; Hunter, C. A.; Llamas, E. M. Chem. Sci. 2015, 6 (2),
1444−1453.
(23) Sun, H. M.; Guo, K.; Gan, H. F.; Li, X.; Hunter, C. A. Org.
Biomol. Chem. 2015, 13 (29), 8053−8066.
(24) Adams, H.; Chekmeneva, E.; Hunter, C. A.; Misuraca, M. C.;
Navarro, C.; Turega, S. M. J. Am. Chem. Soc. 2013, 135 (5), 1853−
1863.
ASSOCIATED CONTENT
■
S
* Supporting Information
The Supporting Information is available free of charge on the
ACS Publications website at DOI: 10.1021/jacs.7b01765.
Detailed experimental procedures and tables (PDF)
AUTHOR INFORMATION
■
Corresponding Author
*herchelsmith.orgchem@ch.cam.ac.uk
ORCID
Stefan Henkel: 0000-0003-3362-8825
Maria Cristina Misuraca: 0000-0001-7024-4535
Christopher A. Hunter: 0000-0002-5182-1859
(25) Sun, H. M.; Navarro, C.; Hunter, C. A. Org. Biomol. Chem. 2015,
13 (17), 4981−4992.
(26) Misuraca, M. C.; Grecu, T.; Freixa, Z.; Garavini, V.; Hunter, C.
A.; van Leeuwen, P.; Segarra-Maset, M. D.; Turega, S. M. J. Org. Chem.
2011, 76 (8), 2723−2732.
Notes
(27) Hunter, C. A.; Misuraca, M. C.; Turega, S. M. J. Am. Chem. Soc.
2011, 133 (3), 582−594.
The authors declare no competing financial interest.
(28) Hunter, C. A.; Misuraca, M. C.; Turega, S. M. J. Am. Chem. Soc.
2011, 133 (50), 20416−20425.
ACKNOWLEDGMENTS
■
We acknowledge financial support from the Engineering and
Physical Sciences Research Council (EP/K025627/2). S.H.
(29) Hunter, C. A.; Misuraca, M. C.; Turega, S. M. Chem. Sci. 2012, 3
(2), 589−601.
F
DOI: 10.1021/jacs.7b01765
J. Am. Chem. Soc. XXXX, XXX, XXX−XXX

Journal of the American Chemical Society
Article
(30) Chekmeneva, E.; Hunter, C. A.; Misuraca, M. C.; Turega, S. M.
Org. Biomol. Chem. 2012, 10 (30), 6022−6031.
(31) Hunter, C. A.; Misuraca, M. C.; Turega, S. M. Chem. Sci. 2012, 3
(8), 2462−2469.
(32) Sun, H.; Hunter, C. A.; Navarro, C.; Turega, S. J. Am. Chem. Soc.
2013, 135 (35), 13129−13141.
(33) Jinks, M. A.; Sun, H. M.; Hunter, C. A. Org. Biomol. Chem. 2014,
12 (9), 1440−1447.
(34) Cabot, R.; Hunter, C. A. Chem. Soc. Rev. 2012, 41 (9), 3485−
3492.
(35) Amenta, V.; Cook, J. L.; Hunter, C. A.; Low, C. M. R.; Vinter, J.
G. Org. Biomol. Chem. 2011, 9 (21), 7571−7578.
(36) Amenta, V.; Cook, J. L.; Hunter, C. A.; Low, C. M. R.; Vinter, J.
G. J. Phys. Chem. B 2012, 116 (49), 14433−14440.
(37) Amenta, V.; Cook, J. L.; Hunter, C. A.; Low, C. M. R.; Sun, H.;
Vinter, J. G. J. Am. Chem. Soc. 2013, 135 (32), 12091−12100.
(38) Cockroft, S. L.; Hunter, C. A. Chem. Soc. Rev. 2007, 36 (2),
172−188.
(39) Camara-Campos, A.; Musumeci, D.; Hunter, C. A.; Turega, S. J.
Am. Chem. Soc. 2009, 131 (51), 18518−18524.
(40) Magee, M. D.; Walker, S. Can. J. Chem. 1971, 49 (7), 1106−
1114.
G
DOI: 10.1021/jacs.7b01765
J. Am. Chem. Soc. XXXX, XXX, XXX−XXX