Isothermal calorimetric titrations on charge-assisted halogen bonds: Role of entropy, counterions, solvent, and temperature

Article abstract of DOI:10.1021/ja2119207

We have conducted isothermal calorimetric titrations to investigate the halogen-bond strength of cationic bidentate halogen-bond donors toward halides, using bis(iodoimidazolium) compounds as probes. These data are intended to aid the rational design of h

Full text of DOI:10.1021/ja2119207

Article
pubs.acs.org/JACS
Isothermal Calorimetric Titrations on Charge-Assisted Halogen
Bonds: Role of Entropy, Counterions, Solvent, and Temperature
Sebastian M. Walter, Florian Kniep, Laxmidhar Rout, Franz P. Schmidtchen, Eberhardt Herdtweck,
and Stefan M. Huber*
Department Chemie, Technische Universitat Munchen, Lichtenbergstraße 4, 85747 Garching, Germany
̈
̈
S
* Supporting Information
ABSTRACT: We have conducted isothermal calorimetric
titrations to investigate the halogen-bond strength of cationic
bidentate halogen-bond donors toward halides, using bis-
(iodoimidazolium) compounds as probes. These data are
intended to aid the rational design of halogen-bond donors as
well as the development of halogen-bond-based applications in
solution. In all cases examined, the entropic contribution to
the overall free energy of binding was found to be very
important. The binding affinities showed little dependency on
the weakly coordinating counteranions of the halogen-bond donors but became slightly stronger with higher temperatures. We
also found a marked influence of different solvents on the interaction strength. The highest binding constant detected in this
study was 3.3 × 106 M−1.
assisted XBs between cationic XB donors and neutral or anionic
Lewis bases, to the best of our knowledge only one set of
binding constants is presently known:10b,c the interaction of
halides with a bidentate bromo-imidazoliophane receptor
reported by Beer et al. The binding of this receptor to bromide
corresponds to a binding constant of K = 889 M−1 in CD3OD/
INTRODUCTION
■
Weak non-covalent interactions called “halogen bonds” 1,2
(XBs) occur between Lewis bases and compounds which
feature an electrophilic halogen substituent. The latter is usually
achieved by using strongly electron-withdrawing perfluorinated
backbones.2 Alternatively, cationic core structures have also
been employed occasionally3 in order to design strong XB
donors (i.e. halogen-based Lewis acids). Halogen bonds find
their main use in crystal engineering, and fittingly investigations
regarding XBs have mostly been performed in the solid state.4
In these cases, the strength of the bond was usually inferred by
the amount to which the interaction distance falls below the
sum of the van der Waals radii of the involved atoms.4
Although the occurrence of XBs in solution5 (and in the gas
phase)6 is also well established, relatively few applications based
on XBs in solution have been developed so far.7
D2O (9:1) at 295 K.7e
We have recently introduced bis(haloimidazolium)-based
activators for the solvolysis of benzhydryl bromide, which
served as a test substrate for C−X bond activation.7f These
activators presumably act through the formation of an XB
complex with the substrate and/or the liberated bromide,
providing additional driving force for the overall reaction. Very
little is known, however, about the thermodynamics of XB
complex formation involving cationic XB donors in solution, as
already indicated above. Open questions include the relative
contributions of enthalpy and entropy, as well as the
dependency of the interaction strength on the counteranions,
solvent, and temperature. In general, means to maximize the
interaction strength12 would provide a solid basis for the
rational application of cationic XB donors for various
applications, including the C−X bond activation of diverse
systems (which represents one of our long-term goals). In the
following, we aim to address these open questions by
isothermal titration calorimetry (ITC),13,14 using our previous
XB donors7f as probes (see Scheme 1).
Determinations of the XB interaction strength in solution
have mostly been limited to isolated, scattered cases,8 with only
very few systematic studies.9 The overwhelming majority of
these measurements were performed on neutral XB donors,
especially elemental bromine or iodine. To be precise, more
than 600 interaction free energies have been determined for
XBs with iodine in alkanes at 298 K,9c with typical binding
constants in the range of K ≈ 0.5−1000 M−1. The strongest
such interaction was found for the complex of iodine with
quinuclidine (K = 1.6 × 105 M−1).10a In contrast, only about a
dozen such data are known for XB donors with a carbon
backbone (all of them being neutral compounds).11 Typical
binding constants for such interactions are in the range K ≈ 5−
100 M−1,11 and the maximal value of K = 1.9 × 104 M−1 was
obtained for the binding of chloride to a tridentate XB-based
receptor introduced by Taylor and co-workers.7b For charge-
Received: January 9, 2012
Published: May 11, 2012
© 2012 American Chemical Society
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Scheme 1. Synthesis and Derivatization of
Bis(haloimidazolium)-Based Halogen-Bond Donors
a
a
Reagents and conditions: (i) 2 equiv of ROTf, CH2Cl2, or 2 equiv of
Me3OBF4, CH2Cl2; (ii) 2 equiv of NaZ, MeOH. For yields of
previously unkown compounds, see Supporting Information. X1, X2 =
−
−
Br, I, H; R = methyl, octyl; Z = PF6 , BPh4 .
RESULTS AND DISCUSSION
■
Figure 1. Exemplary isothermal titration calorimetry data for the
titration of m-II-1Me/OTf with NBu4Br in CH3CN at 30 °C.
Derivatization of XB Donors. An important prerequisite
for these investigations was the further derivatization of our
imidazolium-based XB donors 1 so that the solubility of the
compounds could be increased and the available range of
counteranions could be broadened (Scheme 1). Thus, while
our previous N-bis(methylated) XB donors were reasonably
soluble only in acetonitrile, acetone, DMSO, and water,
introducing octyl substituents at the external imidazolium
nitrogens dramatically increased the solubility in DCM and
THF (while, expectedly, reducing the solubility in water). As
we had seen a counterion effect in the activation of benzhydryl
bromide (with the tetrafluoroborate salt being more active than
the triflate one),7f we also exchanged the triflate counterions
contribution becomes more favorable from chloride to iodide,
this effect is counterbalanced by the opposite trend within the
entropy part. The stoichiometry coefficients are close to unity
for all halides, pointing toward a bidentate binding mode of the
XB donor (see Scheme 2).
This binding motive was also confirmed for the solid state, as
an X-ray structural analysis of the complex of m-II-1Me/BPh4
with bromide (generated in situ from CBr4) also showed a
bidentate coordination of the halide (Figure 2).15
In contrast, p-II-1Me/OTf shows a 2:1 stoichiometry in its
complex with chloride (Table 1, no. 4), indicating an
independent monodentate binding mode. The binding constant
is about 1 order of magnitude lower than that of m-II-1Me/OTf
with chloride, and this difference in interaction free energy is
entirely due to a reduced enthalpy contribution. The binding
constants of p-II-1Me/OTf with bromide and iodide are similar
to that of the chloride complex.16 For all following
investigations, we chose m-II-1Me/OTf as our reference system,
focusing mainly on its complex with NBu4Br.
Nature of the Interaction. Next we addressed the obvious
question of whether the binding is actually based on halogen-
bond formation or due to another type of interaction (e.g.,
hydrogen bonds with the protons at C4 and C5 of the
imidazolium moieties). To this end, we tested several
structurally related compounds for their affinity toward
bromide. While the binding of m-BrBr-1Me/OTf with bromide
(Table 1, no. 8) was substantially lower than that of m-II-1Me/
OTf (as is expected from halogen-bond theory), there was no
detectable halide binding for m-HH-1Me/OTf or p-HH-1Me/
OTf under our experimental conditions (Table 1, nos. 10,
11),17,18 ruling out all other types of interactions not involving
the iodine atoms of m-II-1Me/OTf. To make sure the binding
was not simply dependent on a voluminous substituent at
position C2 of the imidazolium moieties, we also tested the bis-
−
−
with PF6 and BPh4 by anion metathesis in methanol. With
these variably modified XB donors in hand, the basis for our
calorimetric measurements was set.
Role of Entropy and First Assessment of Binding.
Table 1 shows binding constants (translating in free energies
ΔG0) and stoichiometry coefficients (n) as well as ΔH0 and
ΔS0 values for various complexes of XB donors 1 with
tetrabutylammonium halides in solution. Quite remarkably, in
almost all complexes investigated the entropic term accounts
for more than 50% of the overall free energy of binding (ΔG0).
Hence, inclusion of the entropy term in these kinds of
investigations is crucial, as data based solely on ΔH0 values will
neglect a substantial proportion of the overall interaction
energy.
The titration of m-II-1Me/OTf and p-II-1Me/OTf with
tetrabutylammonium chloride, bromide, and iodide in acetoni-
trile at room temperature (Table 1, nos. 1−6) allowed a first
assessment of the nature and strength of the charge-assisted
halogen bonds considered in this publication (see Figure 1 for a
typical titration curve).
The association constants for m-II-1Me/OTf with all three
halides are rather similar and in the range of 2.5 × 105−5.2 ×
105 M−1, with the iodide complex being slightly weaker than the
chloride and bromide ones. Interestingly, while the enthalpic
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Table 1. Results of the Isothermal Calorimetric Titrations of Various Halogen-Bond Donors with Halides NBu4X
a
b
c
d
See ref 16. Entry repeated for better comprehensibility. No heat effect was observed in the ITC measurements. In order to get a rough estimate
of the experimental error associated with these measurements, the titration corresponding to entry no. 2 was performed four times (with different
batches of activator and on two different ITC instruments). Binding constants of 1.59 × 105, 2.53 × 105, 4.54 × 105, and 4.76 × 105 mol−1 were
obtained, indicating that only differences in the K value exceeding a factor of 2−3 are to be considered significant.25 The value given here
corresponds to the exact same conditions also employed for entries 1−6 and should therefore offer the best comparability.
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(Table 1, no. 13), while the binding of a simple imidazolium
salt 3 (Figure 3) to bromide is still a bit weaker (Table 1, nos.
14)so either there is an additional hydrogen-bond interaction
in m-HI-2Me/OTf or the second cationic charge of the non-
iodinated imidazolium moiety increases the overall halide
affinity.
Scheme 2. Formation of a Bidentate Halogen-Bond
Complex
a
a
Chemical equation for the titration of m-II-1Me/OTf with NBu4Br in
CH3CN.
Figure 3. Imidazolium derivative 3 and thiourea 4, tested for
comparison with the bidentate XB donors.
We also did not detect any binding with the neutral (non-N-
alkylated) precursor compound of m-1 (i.e., m-II-2), proving
the importance of the cationic charge on the electrophilicity of
the iodine centers (Table 1, no. 15). When testing thiourea 419
for a comparison with a neutral hydrogen-bond donor, we
found a binding constant of approximately 1.5 × 103 M−1
(Table 1, no. 16), about 2 orders of magnitude lower than that
of m-II-1Me/OTf, although related thiourea derivatives have
already been successfully employed in catalyses based on an
“anion binding mechanism”.20
Temperature Dependence and Structural Variations.
In order to investigate the temperature dependence of the
halide binding, we determined the binding parameters of m-II-
1Me/OTf with NBu4Br in acetonitrile in the (experimentally
accessible) temperature range of 10−50 °C (Table 1, nos. 17−
20). From lowest to highest temperature, the binding constant
decreases only by a factor of 2, which means that the free
energy of binding actually becomes slightly more negative (cf.
Figure 4). This is due to both a slightly more favorable
Figure 2. X-ray structural analysis of the complex of bromide with m-
II-1Me/BPh4 (anion omitted for clarity; ellipsoids at 50% probability).
Selected bond lengths [Å] and angles [°]: C1−I1 2.090, C11−I2
2.083, C11−I2−Br1 170.5, C1−I1−Br1 172.1, C11−N3−C7−C6
54.4, C6−C5−N2−C1 −60.0.
methylated compound m-MeMe-1Me/OTf, which again gave
no detectable binding (Table 1, no. 12). These combined
observations provide a very strong indication that with the
present complexes, halogen bonding is indeed the main
interaction responsible for halide binding.
Interestingly, the difference between m-II-1Me/OTf and m-
BrBr-1Me/OTf is mostly based on a significantly less favorable
entropic contribution to ΔG0 for m-BrBr-1Me/OTf. As the
bromo substituents of m-BrBr-1Me/OTf feature a smaller
surface area than the iodo substituents of m-II-1Me/OTf, one
might speculate that the former compound binds less solvent
molecules in solution and consequently releases less solvent
molecules during the binding event − resulting in a less
favorable entropic contribution. To obtain more insight into
the nature of the binding of m-BrBr-1Me/OTf with halides, we
also determined the binding parameters for the chloride and
iodide complexes (Table 1, nos. 7 and 9). The chloride
complex features a similar overall binding strength as the
bromide complex, but the relative enthalpic contribution is
lower, and the relative entropic contribution is higher than in
the bromide case. This trend is similar to the one observed for
the chloride and bromide complexes of m-II-1Me/OTf. The
iodide complex of m-BrBr-1Me/OTf is markedly less favorable
than the chloride and bromide adducts, though, and the
decrease in the free energy of binding, compared to the
bromide complex, is entirely due to a reduced enthalpic
contribution.
Figure 4. Temperature-dependence of the binding of tetrabutylam-
monium bromide to m-II-1Me/OTf (red graph) and m-II-1Me/BPh4
(blue graph) in CH3CN.
enthalpic part and the positive sign of the entropy. The same
trend is seen for m-II-1Me/BPh4 (Table 1, nos. 21−23). The
temperature dependence of halogen bonds had previously been
studied for complexes of C6F5I with metal fluorides in toluene
(or heptane),9d,21a and in all cases the free energy of binding
became less negative for higher temperatures.21b In addition, the
temperature dependence had also been studied in the solid
phase,21c with shorter interaction distances being found at
lower temperatures.
As expected, the binding constant of the mono-iodinated XB
donor m-HI-1Me/OTf lies in the range of that of p-II-1Me/OTf
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The composition of the XB donors should also influence the
strength of the halide binding. First, we replaced the triflate
counteranions in m-II-1Me/OTf with other weakly coordinating
Overall, the dependency of the halogen-bond interaction on
the solvent, though clearly existent, is hard to interpret at this
point.
−
ones, either via the methylating reagent (for BF4 ) or by anion
−
−
CONCLUSIONS
exchange (for PF6 and BPh4 ). We note that for all
counteranions, the respective non-iodinated compounds m-
HH-1Me/(BF4, PF6 or BPh4) again showed no detectable
binding to bromide (Table 1, nos. 24−26). As for the XB
donors themselves, only a very small effect, if at all, is seen for
the different counteranions. The compounds m-II-1Me/PF6
(Table 1, no. 30) and m-II-1Me/BPh4 seem to bind a tad
stronger than m-II-1Me/OTf and m-II-1Me/BF4 (Table 1, no.
29), but this might also still be within the experimental error.
Thus, either the ion pairing in solution is either very weak to
nonexistent or very similar for the different counteranions.22
As a further structural variation, we compared the binding
strengths of the N-methylated and N-octylated compounds m-
II-1Me/OTf and m-II-1Oct/OTf (Table 1, no. 31) as well as m-
II-1Me/BPh4 and m-II-1Oct/BPh4 (Table 1, no. 32). In both
cases, the N-octylated XB donor exhibited a slightly stronger
halide affinity. However, the difference is mainly due to the
entropic part for the triflate salts, and mainly based on the
enthalpic part for the tetraphenylborate salts. Thus, while there
seems to be a notable effect of the N-substituent, it is difficult to
rationalize (and generalize) its influence.
■
In summary, we have conducted isothermal calorimetric
titrations to investigate the halogen-bond strength of cationic
bidentate halogen-bond donors toward halides, using bis-
(iodoimidazolium) compounds 1 as probes. This data is
intended to aid the rational design of halogen-bond donors as
well as the development of halogen-bond-based applications in
solution (e.g., in anion recognition, anion transport, and
catalysis). In all cases examined, the entropic contribution to
the overall free energy of binding was found to be very
important. The binding affinities showed little dependency on
the weakly coordinating counteranions of the halogen-bond
donors, but became slightly stronger with higher temperatures.
We also found a marked influence of different solvents on the
interaction strength. The data presented here should constitute
a valuable benchmark for the development of theoretical
methods concerning the accurate prediction of halogen-bond
interactions in solution.
ASSOCIATED CONTENT
* Supporting Information
■
S
Full experimental details and characterization data, ITC spectra,
and X-ray structural analysis. This material is available free of
charge via the Internet at http://pubs.acs.org.
Solvent Effects. Finally, we considered the influence of
different solvents on the halide binding.23 In addition to
acetonitrile (which features hydrogen-bond acidity αH = 0.09
2
and hydrogen-bond basicity βH = 0.44),24 m-II-1Me/OTf is
2
AUTHOR INFORMATION
Corresponding Author
stefan.m.huber@tum.de
also sufficiently soluble in acetone (αH = 0.04, βH = 0.50),
■
2
2
water/methanol (1/9) (H2O: αH = 0.35, βH = 0.38; MeOH:
2
2
αH2 = 0.37, βH2 = 0.41), ethanol (αH2 = 0.33, βH2 = 0.44), and
DMSO (αH = 0.00, βH = 0.77). Of those solvents, only
Notes
2
2
The authors declare no competing financial interest.
acetone enabled a detectable binding of bromide to the XB
donor (with binding parameters that are very similar to those in
acetonitrile: Table 1, nos. 33−36). In line with these findings, it
had previously been shown9b that hydrogen-bonding solvents
(i.e., water and ethanol) are detrimental to the formation of
halogen bonds. For the N-octylated derivative m-II-1Oct/OTf, a
ACKNOWLEDGMENTS
■
The authors are grateful to the Fonds der Chemischen
Industrie (Liebig Fellowship to S.M.H.), the Deutsche
Forschungsgemeinschaft (DFG), and the Leonhart-Lorenz-
Stiftung for financial support. S.M.W. and F.K. thank the TUM
Graduate School. We thank Prof. Dr. Herbert Mayr (LMU
Munich) for the very generous grant of ITC measurement time
and for interesting discussions. We also thank Prof. Dr.
Thorsten Bach and his group for their great support. L.R.
thanks the Alexander-von-Humboldt Foundation for a
postdoctoral fellowship.
wider range of solvents is available, most notably THF (αH
=
2
0.00, βH2 = 0.51) and CH2Cl2 (αH2 = 0.13, βH2 = 0.05). In the
former solvent, a markedly increased binding constant to
bromide was found (K = 3.3 × 106 M−1, see Table 1, no. 37),
one of the highest binding constants detected so far for
halogen-bond-based complexes.10 An almost identical binding
constant was found in CH2Cl2 (Table 1, no. 39), but in this
case the relative contribution of the entropy is distinctly higher
(and consequently the contribution of the enthalpy is lower)
compared to THF. Additionally, we determined the binding
data for m-II-1Oct/OTf with tetrabutylammonium chloride and
iodide in CH2Cl2 (Table 1, nos. 38 and 40). As was the case for
m-II-1Me/OTf in acetonitrile, we found very similar association
constants for all three halides (with the chloride complex being
perhaps marginally less favorable than the ones of bromide and
iodide). Interestingly, there is again a compensation of two
opposite trends: from the chloride to the iodide complex, the
absolute value of the enthalpy increases while the entropy
decreases. The same trend, albeit not as linear as with the other
two examples, is found for the complexes of m-II-1Oct/OTf
with chloride, bromide, and iodide in acetone (Table 1, nos.
41−43).
REFERENCES
■
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(15) m-II-1Me/(Br,BPh4): colorless needle C38H34BBrI2N4 Mr =
891.20; monoclinic, space group P21/n (No. 14), a = 10.6294(3), b =
14.2671(4), and c = 23.8757(7) Å, β = 102.4990(12)°, V =
3534.95(18) Å3, Z = 4, λ(Mo Kα) = 0.71073 Å, μ = 2.942 mm−1,
ρcalcd = 1.675 g cm−3, T = 123(1) K, F(000) = 1744, θmax = 25.35°, R1
= 0.0150 (6251 observed data), wR2 = 0.0389 (all 6464 data), GOF =
1.093, 417 parameters, Δρmax/min = 0.56/−0.37 e Å−3. Crystallographic
data (excluding structure factors) for the structure reported in this
paper have been deposited with the Cambridge Crystallographic Data
Centre as supplementary publication no. CCDC-859901. Copies of
the data can be obtained free of charge on application to CCDC, 12
Union Rd., Cambridge CB2 1EZ, UK (fax: (+44)1223-336-033; e-
mail: deposit@ccdc.cam.ac.uk). For more details see the Supporting
Information.
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smaller than two, indicating a convolution of the two halide binding
events.
(17) We note that for simple acyclic imidazolium compounds
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́
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(18) In addition to the experiments mentioned in Table 1, we also
titrated m-HH-1Me/OTf (in acetonitrile and acetone) as well as m-
HH-1Oct/OTf (in acetonitrile) with tetrabutylammonium chloride,
bromide, and iodide at 30°. No heat effect was detected in any of these
measurements in acetonitrile. In acetone, heat effects (likely due to
precipitation) were observed, but no binding parameters could be
obtained from the ITC data.
(8) See, for example, ref 5b as well as: (a) Reid, C.; Mulliken, R. S. J.
Am. Chem. Soc. 1954, 76, 3869. (b) Larsen, D. W.; Allred, A. L. J. Am.
Chem. Soc. 1965, 87, 1216. (c) Larsen, D. W.; Allred, A. L. J. Phys.
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(19) See Zhang, Z.; Schreiner, P. R. Chem. Soc. Rev. 2009, 38, 1187
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(9) Selected recent publications: (a) Cabot, R.; Hunter, C. A. Chem.
Commun. 2009, 2005. (b) Sarwar, M. G.; Dragisic, B.; Salsberg, L. J.;
Gouliaras, C.; Taylor, M. S. J. Am. Chem. Soc. 2010, 132, 1646.
(c) Laurence, C.; Graton, J.; Berthelot, M.; El Ghomari, M. J. Chem.
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(10) (a) An even higher binding constant was found for the complex
of 1-methyl-1,3-imidazoline-2-selenone with iodine in dichloro-
methane (K = 2.57 × 106 M−1), see ref 9c. To the best of our
knowledge, as far as neutral halogen bond donors are concerned, this
represents the highest binding constant for a halogen-bond-based non-
covalent complex reported in the literature; (b) During the revision
process of this manuscript, an interesting study was published by Beer
et al., which includes a charge-assisted halogen-bond-based complex
with a binding constant of 3.71 × 106 M−1: Caballero, A.; Zapata, F.;
(20) See ref 14 as well as: (a) Doyle, A. G.; Jacobsen, E. N. Chem.
Rev. 2007, 107, 5713. (b) Birrell, J. A.; Desrosiers, J.-N.; Jacobsen, E.
N. J. Am. Chem. Soc. 2011, 133, 13872. (c) Klauber, E. G.; Mittal, N.;
Shah, T. K.; Seidel, D. Org. Lett. 2011, 13, 2464 and references cited
therein.
(21) (a) Libri, S.; Jasim, N. A.; Perutz, R. N.; Brammer, L. J. Am.
Chem. Soc. 2008, 130, 7842. (b) For the temperature dependence of
the halogen bonding between halotrifluoromethanes and various Lewis
bases in liquid noble gases, see: Hauchecorne, D.; Nagels, N.; van der
Veken, B. J.; Herrebout, W. A. Phys. Chem. Chem. Phys. 2012, 14, 681
and references cited therein. (c) Forni, A.; Metrangolo, P.; Pilati, T.;
Resnati, G. Cryst. Growth Des. 2004, 4, 291 and references cited
therein.
(22) For a recent review on counterion effects in supramolecular
chemistry, see: Gasa, T. B.; Valente, C.; Stoddart, J. F. Chem. Soc. Rev.
2011, 40, 57. For interactions of imidazolium cations with their
counteranions in solution, see e.g.: (a) Avent, A. G.; Chaloner, P. A.;
Day, M. P.; Seddon, K. R.; Welton, T. J. Chem. Soc., Dalton Trans.
1994, 3405. (b) Elaiwi, A.; Hitchcock, P. B.; Seddon, K. R.; Srinivasan,
N.; Tan, Y. M.; Welton, T.; Zora, J. A. J. Chem. Soc., Dalton Trans.
́
White, N. G.; Costa, P. J.; Felix, V.; Beer, P. D. Angew. Chem., Int. Ed.
2012, 51, 1876. (c) Also during the revision process of this
1
manuscript, a H NMR titration study of haloimidazolium salts with
various anions was published: Cametti, M.; Raatikainen, K.;
Metrangolo, P.; Pilati, T.; Terraneo, G.; Resnati, G. Org. Biomol.
Chem. 2012, 10, 1329.
1995, 3467. (c) Kovacevic, A.; Grundemann, S.; Miecznikowski, J. R.;
̈
Clot, E.; Eisenstein, O.; Crabtree, R. H. Chem. Commun. 2002, 2580
and references cited therein.
(11) See ref 9c and references cited therein.
(23) Theoretical study on XBs in solution: Lu, Y.; Li, H.; Zhu, X.;
Zhu, W.; Liu, H. J. Phys. Chem. A 2011, 115, 4467.
(12) In this manuscript, “binding strength” refers to the free energy
ΔG0.
(24) Hydrogen-bond acidity αH and hydrogen-bond basicity βH
2
2
(13) (a) Schmidtchen, F. P. Isothermal Titration Calorimetry in
Supramolecular Chemistry. In Supramolecular Chemistry: from
Molecules to Nanomaterials; Steed, J. W., Gale, P. A., Eds.; Wiley:
Chichester, 2012; pp 275−296; (b) Ball, V.; Maechling, C. Int. J. Mol.
Sci. 2009, 10, 3283.
refer to the scale established by Abraham et al., see: Abraham, M. H.
Chem. Soc. Rev. 1993, 22, 73 and references cited therein.
(25) See also: Sessler, J. L.; Gross, D. E.; Cho, W.-S.; Lynch, V. M.;
Schmidtchen, F. P.; Bates, G. W.; Light, M. E.; Gale, P. A. J. Am. Chem.
Soc. 2006, 128, 12281.
(14) ITC measurements have frequently been employed for the
determination of interaction strengths of supramolecular complexes
and protein−ligand complexes, see e.g.: Koch, C.; Heine, A.; Klebe, G.
J. Mol. Biol. 2011, 406, 700. We note that isothermal calorimetric
titrations simultaneously provide ΔH and ΔS values in one
measurement. For the important role of entropy in supramolecular
interactions, see: (a) Schmidtchen, F. P. Coord. Chem. Rev. 2006, 250,
8512
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