Thermodynamics of Halide Binding to a Neutral ... - ACS Publications

Article Cite This: J. Org. Chem. 2018, 83, 1903−1912

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Thermodynamics of Halide Binding to a Neutral Bambusuril in Water and Organic Solvents Tomas Fiala,∥ Kristina Sleziakova,∥ Kamil Marsalek, Karolina Salvadori, and Vladimir Sindelar* Department of Chemistry and RECETOX, Faculty of Science, Masaryk University, Kamenice 5, 625 00 Brno, Czech Republic S Supporting Information *

ABSTRACT: Driving forces of anion binding in water in contrast to nonpolar environments are of high interest because of their relevance to biology and medicine. Here we report a neutral bambusuril macrocycle (1), soluble in both water and nonpolar solvents due to decoration with 12 polyethylene glycol-based substituents. The new bambusuril has the highest affinity for I− in pure water ever reported for a synthetic macrocycle relying on hydrogen bonding interactions rather than metal coordination or Coulombic forces. Isothermal titration calorimetry (ITC) experiments in nine different solvents, ranging from polar water to nonpolar carbon tetrachloride, provided insight into the forces responsible for halide binding by bambusurils. The different importance of anion solvation and solvent expulsion from the cavity of the macrocycle in various solvents is illustrated by the fact that halide binding in water and chloroform is exclusively driven by favorable enthalpy with an entropic penalty, while in alcohols and nonpolar solvents, both favorable enthalpy and entropy contribute to anion encapsulation. DFT calculations and correlation of thermodynamic data with the solvent Swain acity parameter further underscore the importance of solvent effects on anion binding by bambusurils.

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INTRODUCTION Anions are crucial for existence of living organisms, because many vital functions are based on supramolecular recognition of anions and interaction of negatively charged species with a variety of molecules.1−3 Although these interactions are expected to operate in aqueous solution, some of them take place in rather nonpolar environments. These examples include transport of anions through lipophilic cell membranes or binding of anions in hydrophobic protein pockets.4−7 It is therefore of great interest to investigate the supramolecular interactions which drive anion binding not only in water but also in more nonpolar environments. Many studies were published which used synthetic receptors to evaluate binding affinities toward anions in single solvent systems. However, the studies comparing binding of inorganic anions within synthetic receptors in several organic solvents are rare.8−11 Binding of inorganic anions in pure water is particularly problematic.12 Anions tend to be highly hydrated, and the surrounding hydration shell prevents anions from interacting with a receptor. In addition, neutral receptor molecules are usually poorly soluble in water. For these reasons, anion receptors functioning in water are usually equipped with positive charge(s) in their structure, which aid anion binding through Coulombic interactions and also increase solubility of the receptor in water. Only several synthetic anion receptors lacking positive charge were recently reported to efficiently bind anions in water.13−18 For example, Kubik and co-workers prepared a cyclopeptide receptor19 and its dimer which preferred to bind © 2018 American Chemical Society

sulfate anions over halides even in pure water. The preferential binding of the highly solvated sulfate was explained as a combination of hydrogen bonding interactions between the anion and the receptor and the release of water molecules from the receptor binding pocket. Recently, our group designed bambusuril 2 bearing negative charges (Figure 1) which was soluble in aqueous buffers upon deprotonation of carboxylic functional groups.20 Although negatively charged carboxylate groups occupying both rims of bambusuril were recently shown to repel entering anions,21 this derivative possesses surprisingly high binding affinities toward anions. However, the solubility of bambusurils decorated with carboxylates is poor in a majority of organic solvents, which disabled further comparison of its binding properties in different solvents. Although we were able to prepare bambusuril derivatives with good solubility in nonpolar solvents,22−24 direct correlation of binding characteristics could be affected by dissimilar structural nuances. Here we report the synthesis of a novel bambusuril macrocycle 1 (Figure 1) bearing short oligo(ethylene glycol) substituents with excellent solubility in solvents ranging from nonpolar CCl4 to pure water. This exceptional solubility profile of macrocycle 1 enabled a study of solvent-dependent anion binding by a macrocyclic host. Unlike previously published bambusuril 2, the Received: November 9, 2017 Published: January 26, 2018 1903

DOI: 10.1021/acs.joc.7b02846 J. Org. Chem. 2018, 83, 1903−1912

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The Journal of Organic Chemistry

water over 1 week (Scheme 1b). A less-time-consuming yet more laborious alternative method of anion removal from the 1· Na+BF4− complex, involving a series of simple washes, was also devised (see Experimental Section). A 1H NMR competition experiment with dodecabenzylbambus[6]uril (3)23 provided evidence that the obtained macrocycle 1 was anion-free after the washing process (Figure S13). The macrocycle exhibited >1 mM solubility in every solvent we tested including D2O, DMSO, CD3CN, CD3OD, acetone, THF, CDCl3, toluene, and others (Figure S12). Supramolecular Properties. The formation of supramolecular complexes of bambusuril 1 and inorganic anions was first studied using 1H NMR (Figure S14). In a typical experiment, addition of 1.0 equiv of anion was present in the solution, only one set of signals corresponding to the bound macrocycle was present in the spectra. A similar complexation-induced binding pattern was observed for different anions and solvent systems and was in agreement with the formation of a supramolecular complex of 1:1 stoichiometry with slow binding kinetics on the NMR time scale. This binding pattern is characteristic for other bambusuril complexes in which one anion is included in the center of bambusuril stabilized by 12 C−H···anion interactions (Figure 1b,c).25,26 Despite recent progress in the number of synthetic receptors, binding inorganic anions in water via interactions other than electrostatic is relatively limited. This is because of difficulties to overcome high hydration energies of anions and poor solubility of anion receptors in water. Because bambusuril 1 has outstanding solubility in a variety of solvents, we decided to investigate its binding to anions in water and eight organic solvents listed in Figure 2. Chloride, bromide, and iodide anions were selected for this purpose, as they all interact with 1, share spherical shape, but differ in many properties connected to their

Figure 1. (A) Structures of bambusurils 1−3. Visualization of a bambusuril complex with iodide: (B) side view and (C) top view. R groups are omitted for clarity.

derivative 1 belongs to the small family of neutral anion receptors functioning in pure water.

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RESULTS AND DISCUSSION Synthesis. Bambusuril 1 was synthesized from 2,4-bis(3,6,9trioxadecyl)glycoluril (4). This precursor was prepared in five steps, starting from triethylene glycol monomethyl ether (5) in a 59% overall yield on a multigram scale without chromatographic purification (Scheme 1a). Acid-catalyzed condensation of compound 4 with paraformaldehyde (PFA) in the presence of tetrabutylammonium tetrafluoroborate (TBA+BF4−) provided bambusuril 1·Na+BF4− complex in a 45% yield after separation from the unwanted four-membered macrocycle. The anion-free macrocycle 1 was obtained by continuous washing of a dichloromethane solution of the 1·Na+BF4− complex with Scheme 1a

(A) Synthesis of glycoluril 4; a: TsCl, NaOH, THF, 0 °C to rt, 3 h; b: NaN3, H2O/acetone, reflux, 14 h; c: PPh3, THF, rt, overnight, then H2O, rt, overnight; d: bis(4-nitrophenyl) carbonate, DIPEA, CH2Cl2, rt, 3 h; e: diphenyl carbonate, DIPEA, reflux, 3 h; f: trans-4,5-dihydroxyimidazolidinone (10), HBF4, H2O/MeOH, reflux, 100 min. (B) Synthesis of bambusuril 1; g: (CH2O)n, TBA+BF4−, HBF4, dioxane, reflux, 14 h; h: CH2Cl2, continuous extraction with H2O for 7 days; i: CH2Cl2, 15× wash with brine, then 20× wash with ultrapure water. a

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the surrounding aqueous solvent; therefore, removing its solvation shell is enthalpically unfavorable but has a relatively favorable entropic term. The larger chaotropic halides, on the other hand, exhibit weaker anion−water interactions, resulting in an enthalpy-favored but entropy-disfavored desolvation.30,31 Anion inclusion into bambusuril 1 is aided by C−H···X− hydrogen bonds as we proposed in our previous work.22 As the observed enthalpy of binding is higher for the larger iodide compared to the smaller chloride, we wondered if this trend results from the fact that chloride is too small to be ideally stabilized inside the macrocycle compared to the larger iodide. To answer this question, we performed DFT calculations in vacuo and in two solvents, methanol and chloroform, which were modeled by the implicit C-PCM solvent model32 (see Supporting Information). The results showed that the bambusuril macrocycle is flexible enough to adopt a conformation in which the C−H···X−bonds have an optimal length. The complex stability in vacuo increased in the order iodide < bromide < chloride which corresponds to the increasing charge density of anions. On the other hand, complex stability in methanol and chloroform increased in the order chloride < bromide < iodide, in agreement with experimental results. Considering opposite trends in binding energy obtained from theoretical calculations in vacuo and both experiments and calculations in solvents, we can assume that although supramolecular interactions between anions and the macrocycle positively contribute to the complex stability, they are overcompensated by the anion desolvation. Although the ion pairing equilibrium typically favors separated ions in pure water33 and should not affect anion binding, countercations could potentially alter 1−X− binding by other means, e.g. through weak association with the oligo(ethylene glycol) side chains. To verify that counter-cations do not affect halide binding by bambusuril 1 in water, we measured the association constant for the 1·Br− complex using NaBr, CsBr and TMA+Br−. Indeed, the differences in log Ka values for all three bromide salts were statistically insignificant (Table S2), providing evidence that counter-cations do not play a prominent role in 1−X− binding in water. To summarize binding in pure water, the main driving force of the inclusion of halides inside bambusuril 1 is the enthalpically favorable release of high-energy water molecules from the bambusuril cavity28 accompanied by stabilization of anion through hydrogen bonding upon its inclusion. On the other hand, high enthalpy of anion desolvation, entropic penalties connected to the organization of water molecules released from the anion solvation shell and from the interior of the macrocycle, disfavor the complex formation. Thermodynamic parameters connected with the release of high-energy water are expected to be independent of the anion size, as all included water molecules are expelled from the cavity upon anion inclusion. Thus, differences in overall binding for different halides are caused mainly by differences in their solvation energies. Indeed, the strongly solvated chloride is bound by about 4 orders of magnitude weaker in terms of association constant Ka compared to the loosely solvated iodide. This also explains why binding strength of anions by bambusuril in water follows the Hofmeister series. We decided to evaluate further the formation of the 1·X− complex in pure water in contrast to aqueous phosphate buffer and compare these results with binding properties of the previously reported carboxybenzyl-substituted bambusuril 234 (Figure 1a), which is only soluble in aqueous buffer solutions of

Figure 2. Thermodynamic parameters obtained for binding between bambusuril 1 and three halide anions determined in nine different solvents by ITC at 298.15 K. Tetrabutylammonium halide salts were used in all experiments expect for those in water in which sodium salts were used instead.

size and charge distribution. For example, chloride is known as a relatively strongly solvated anion compared to the weakly solvated iodide. For our investigation, we decided to employ isothermal titration calorimetry (ITC) which provides information about complex stability and furthermore gives insight into the thermodynamic parameters of the host−guest interaction. Most interactions were studied using direct ITC titration methods; however, competition experiments were also employed to evaluate high complex stabilities not suitable to be measured by direct titration (association constants larger than 106 M−1). All performed titrations confirmed formation of complexes between 1 and the halide anion of 1:1 stoichiometry independently on the type of solvent and halide anion. In the following text, thermodynamic parameters of binding between 1 and halide anions in different solvents are discussed. ITC yields only a single set of thermodynamic parameters which is, however, the result of many events which take place upon binding. Our effort is to understand these events and identify those which drive the binding process. Binding in Water. Results summarized in Figure 2 and Table S1 show that binding of halides by bambusuril 1 in pure water was driven enthalpically while the entropic term, −TΔS, was unfavorable for all three anions. These observations are consistent with the nonclassical hydrophobic effect.27−29 Water molecules of high energy encapsulated in the cavity of an anionfree macrocycle 1 are displaced upon anion complexation. Their incorporation into the hydrogen bonded network of water molecules in the bulk solvent is characterized by large negative ΔH which is partially compensated by a positive entropic term, −TΔS, as the released water molecules lose degrees of freedom. Going from Cl− to I−, the enthalpic term becomes more negative, while −TΔS increases. This trend can be explained by anion desolvation. The kosmotropic chloride interacts strongly with 1905


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The Journal of Organic Chemistry neutral or basic pH. While 1·Br− binding is virtually unaffected by substituting pure water for phosphate buffer as the solvent medium (Table S3), the affinity of macrocycle 2 to Br− is significantly higher (log Ka = 5.3) in 100 mM sodium phosphate buffer than in 6 mM buffer (log Ka = 4.5). A similar phenomenon has been previously reported by Gibb et al.35,36 for complexation of adamantane carboxylate and perchlorate by the hydrophobic concavity host octa acid at varying concentrations of sodium salts. The authors attributed this phenomenon to ion-pairing of Na+ with the carboxylate groups of the host, which decreases the net negative charge and thus promotes anion binding. As portals of bambusuril 2 are decorated with carboxylate groups, we expected that observed changes of anion binding caused by buffer concentration are also the result of the ion-pairing. As expected, bambusuril 1 lacking carboxylic groups binds anions with affinity which is independent of the buffer concentration. Binding in Chloroform. Globally evaluating the thermodynamic parameters of 1·X− complex formation (Figure 2, Table S1), two crucial observations can be made: (a) binding enthalpy is favorable in all solvents; (b) the entropic term, −TΔS, is favorable or close to zero in all solvents, except for water and chloroform, in which it is highly unfavorable. In our previous work,26 we attributed the unfavorable entropy of anion binding by dodecabenzylbambus[6]uril in chloroform to enforced conformational change of the macrocycle upon complexation. However, direct comparison of bambusuril binding in a series of solvents, now possible for macrocycle 1, refutes this claim. Conformational change of the bambusuril macrocycle induced by anion inclusion should be solvent-independent. Nevertheless, unfavorable binding entropy is observed only in water and chloroform. Even in carbon tetrachloride, a halogenated solvent similarly polarizable to CHCl3, 1−X− binding was driven both by favorable enthalpy and entropy (Figure 2, Table S1). As discussed above, positive −TΔS for binding in water can be explained by release of high-energy water molecules from the macrocycle’s cavity which join the structured hydrogen-bonding network in the bulk solvent. But for chloroform, an equivalent explanation seems counterintuitive at first: CHCl3 is typically considered as a relatively nonpolar yet polarizable solvent with very little structural organization in the liquid phase.37,38 For these reasons, many experimental results in chloroform as a solvent are considered to be translatable to lipid phases or even the gas phase. Evidence that CHCl3 has certain properties more comparable to water than nonpolar environments is essential to our deeper understanding of this widely used solvent. The perception of chloroform as an unstructured medium was challenged in a recent study by Salzmann et al.39 Their neutron diffraction experiment showed that liquid chloroform exhibits a high level of organization, with over 29% of the solvent molecules participating in polar stacks with collinear dipole orientation. The ordered nature of bulk chloroform can be the reason for the observed similarities in thermodynamic parameters of binding in water and chloroform. To show that expulsion of chloroform molecules from the cavity of bambusuril is a possible explanation for the observed phenomenon, we decided to investigate if chloroform molecules can be encapsulated in the cavity of bambusuril despite their larger size compared to water molecules. While we were unable to crystallize bambusuril 1 under any conditions, most probably due to highly disordered 3,6,9trioxadecyl substituents, we were lucky to obtain single crystals of anion-free dodecabenzylbambus[6]uril (3) from a chloroform solution. The crystal structure revealed that two chloroform molecules can be encapsulated inside the cavity of the

macrocycle (Figure S1). Thus, it seems likely that expulsion of solvent molecules significantly contributes to the complexation event in both water and chloroform. It should be noted that liquid water is still significantly more structured than chloroform, being a tetrahedrally coordinated random network.40,41 Therefore, enthalpic gain and entropic loss from liberating high-energy water is likely to be more pronounced than expelling chloroform from the cavity of the host. However, thermodynamic parameters obtained for halide inclusion in both solvents are similar in terms of enthalpic and entropic contributions and their trends. This indicates that the large enthalpic gain and entropic loss from expelling water molecules must be significantly more compensated by other factors as compared to chloroform. This can be explained by differences between the solvation energies of chloride, bromide, and iodide which are 40, 32, and 20 kJ mol−1, respectively, higher in chloroform than in water.45 Thus, effects contributing both positively (cavity-bound solvent expulsion) and negatively (anion desolvation) to binding of halides to bambusuril 1 are more pronounced in water than in chloroform, but their compensation results in thermodynamic parameters which are qualitatively similar for binding in both solvents. Tighter hydrogen-bonding between CHCl3 and Cl−46 as well as stronger ion-pairing of Cl− salts47 compared to those of Br− and I− give rise to the trend of decreasing ΔH and increasing −TΔS going from 1·Cl− to 1·I− complex formation. Binding in Alcohols. Halide binding by bambusuril 1 was further investigated in a homological series of alcohols, namely methanol, ethanol, and n-propanol. In these alcohols, the binding of Cl− and Br− is driven both by favorable enthalpy and entropy, with a greater contribution from the enthalpic term. 1−I− binding is almost exclusively driven by favorable enthalpy, the entropic term being near neutral. While binding enthalpy grows more negative from chloride to iodide, the entropic term, −TΔS, exhibits the opposite trend. In all alcohols, the most stable complexes are those with iodide, followed by those with bromide and chloride. We may illustrate the observed trends by using data obtained for binding in methanol. Despite relatively high structural organization of methanol molecules in bulk,48 the inclusion of chloride and bromide into bambusuril 1 in this solvent is accompanied by a favorable entropic term. This is in contrast with unfavorable entropy observed for complexation of these anions in water and chloroform. The overall binding entropy is determined by two main components, entropy of anion desolvation and entropy of encapsulated solvent molecule release. These opposing contributions can significantly differ depending on the type of solvent and anion. In methanol, the contribution of anion desolvation entropy prevails over that of solvent molecule release resulting in overall favorable entropy of Cl− and Br− binding. On the contrary, 1−I− binding was characterized by small unfavorable binding entropy as solvation entropy of iodide in methanol is the lowest among all investigated halides49−51 and is not able to overcompensate the entropy of encapsulated solvent molecule release. In the case of water and chloroform, entropy of solvent molecule release is highly dominant, resulting in unfavorable binding entropy for all anions. The driving force of complexation in alcohols is enthalpy due to the concert of high energy solvent molecule release and complex stabilization via multiple hydrogen bonds between methine hydrogen atoms of the macrocycle and the anion. These stabilizing effects are partially compensated by enthalpy of anion desolvation. Chloride has the highest charge density among the 1906


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Table 1. Association Constants (Ka) for Complexes of Bambusuril 1 and Halide Anions Obtained by ITC at 298.15 K and Selected Solvent Parameters Including Swain Acity (A),42 Acceptor Number (AN),43 and Et(30)44 a Ka (M−1)b solvent H2O MeOH EtOH n-PrOH CHCl3 DMSO CH3CN EtOAc CCl4

Swain acity (A) 1.00 0.75 0.66 0.63 0.42 0.34 0.37 0.21 0.09

acceptor number (AN) 54.8 41.5 37.9 37.3 23.1 19.3 18.9 9.3 8.6

−

Br−

Et(30)

Cl

63.1 55.4 51.9 50.7 39.1 45.1 45.6 38.1 32.4

1.2 × 10 4.3 × 103 2.0 × 104 5.3 × 104 2.5 × 104 1.1 × 105 5.6 × 105 1.7 × 106 1.4 × 106 3

I−

1.3 × 10 2.4 × 105 1.0 × 106 1.3 × 106 6.0 × 105 1.8 × 105 1.8 × 106 1.9 × 107 2.9 × 107 5

3.2 × 107 2.0 × 107 9.1 × 107 7.2 × 107 2.0 × 107 1.9 × 105 4.7 × 107 c d

a

Tetrabutylammonium halide salts were used in all experiments expect for those in water in which sodium salts were used instead. bAll measurements were carried out in at least duplicates. The relative standard deviations of log Ka values for all measurements were lower than 4%. cNot measured. Dissolution of TBA+I− in EtOAc resulted in darkening of the solution, presumably oxidation of I− to I3−. The problems persisted even with high-purity grade EtOAc which was degassed prior to the experiment. dNot measured. TBA+I− is poorly soluble in CCl4.

Figure 3. (A) Plot of binding free energy of 1·X− complexes, expressed as log Ka, vs Swain acity parameter.42 A near-linear relationship is observed for all halides when seven out of nine solvents are considered. For the two clear outliers, DMSO and CHCl3, anion desolvation most likely plays a less prominent role in driving the complexation event. (B) Plot of binding selectivity of bambusuril 1 between pairs of halides vs Swain acity parameter. Selectivity is defined as Δlog Ka = log Ka(Br−) − log Ka(Cl−) and analogously for each pair of halides. A near-linear relationship is observed for all halides in solvents other than the polar aprotic DMSO and MeCN.

these ordered molecules become free, which accounts for increase in entropy. This effect is most pronounced in the case of chloride with the highest charge density of all investigated halides, which results in large entropic contribution to binding of this anion in DMSO. The stability of complexes increases in the order chloride < bromide < iodide which is mainly caused by decreasing solvation energy of anions in this series.51 This increase is however significantly less pronounced in DMSO compared to other solvents (Table 1) due to pronounced entropic effects in DMSO as discussed above. Comparison of the Results Obtained in the Different Solvents. As discussed above, the overall binding of anions with bambusuril 1 in all solvent systems is driven by several binding events and processes. The most important ones are proposed to be the events related to release of encapsulated solvent molecules and with desolvation of the anion of interest. To support our hypotheses, we searched for correlations of the obtained thermodynamic parameters with parameters describing solvent−anion interactions. Anion-solvating tendency is measured by the Swain acity parameter (A).42,53 Other parameters, such as the Gutmann−Beckett acceptor number (AN), a measure of solvent’s Lewis acidity,43 and the Dimroth−Reichardt Et(30) parameter, a solvent polarity scale,44 are also relevant to quantifying anion−solvent interactions.11 All these solvent

studied halides. It is therefore more effectively solvated by molecules of methanol compared to, for example, iodide.51 Desolvation of chloride is thus enthalpicaly more unfavorable than in the case of bromide and iodide which translates to the lowest enthalpy of binding measured for chloride. When comparing overall stability of chloride complexes in all alcohols (Table 1), the stability decreases in the order n-propanol > ethanol > methanol. This is in agreement with increasing solvating ability in this alcohol series. Binding in Other Solvents. The last group of solvents included in this study contains DMSO, acetonitrile, EtOAc, and CCl4. Thermodynamics of halide binding in these solvents is characterized by relatively highly favorable entropy although binding enthalpy is usually the main driving force for inclusion. Favorable binding entropy is mainly caused by dominant contribution of desolvation entropy while binding enthalpy accounts for encapsulated solvent molecule release and hydrogen bonding stabilization within the bambusuril cavity. The only exception is binding of chloride in DMSO which is dominated by the entropic factor. We offer the following interpretation for this phenomenon. It was proposed that anions are solvated by DMSO molecules in a multilayer arrangement in which polarized DMSO molecules are ordered due to solvent−solvent dipolar interaction.52 After anion inclusion into the bambusuril cavity, 1907


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selectivity in MeCN compared to what would be expected based on their acity values. Differences in binding behavior in these solvents can be explained by domination of the entropic term over the enthalpic term in halide binding by macrocycle 1 (Figure 2). As discussed above for DMSO, the highly negative entropic term, −TΔS, probably originates from release of solvent molecules from a multilayered solvation shell upon anion inclusion.52 Small enthalpy contribution can be connected to a more favorable encapsulation of these polar solvent molecules in the partially positively charged cavity of the macrocycle, with possible aid from H-bonding stabilization (a similar effect was described previously for a thiourea-based anion receptor9). For a similar reason, the affinity of halides to 1 in DMSO is significantly lower compared to other solvents. This is most pronounced in the case of iodide. Recently, semithiobambusuril was reported to interact with halide anions in DMSO with similar affinities and selectivities,54 which we report here for macrocycle 1 in the same solvent. Thus, semithiobambusurils have very similar binding properties to bambusurils and their selectivity can be expected to be significantly improved when the binding study is performed in a different solvent. We also predict that anion binding by macrocycles from the bambusuril family in other polar aprotic solvents, e.g. dimethylformamide (DMF) and hexamethylphosphoramide (HMPA), will display very low selectivities as well.

parameters are in roughly linear relationship with the free energies, expressed as log Ka, of 1·X− complexes (Figures 3a, S29, and S30, and Table 1), although two out of the nine investigated solvents, DMSO and CHCl3, deviated from this correlation (vide infra). The observed linear relationship indicates that the capacity of the anion to lose its solvation shell is crucial to the overall binding process. Association constants for 1·X − complexes increase with decreasing solvent acity. The largest slope was recorded for chloride, as this anion has the highest charge density. For example, the stability of the 1·Cl− complex in water (A = 1.00) is more than 400 times lower than in acetonitrile (A = 0.37), but there is only 1.5-fold difference comparing binding affinities for the 1·I− complex in the same two solvents. Because anion desolvation is not the only event contributing to different binding thermodynamics of 1·X− complexes in various solvents, it is not surprising that deviations from linearity were observed in the log Ka vs acity plot (Figure 3a). 1−X− binding in chloroform stands out with a significantly lower affinity, especially in the case of the chloride complex, than would be expected considering its acity value (A = 0.42). We believe that this phenomenon is caused by the fact that, in addition to anion desolvation, expulsion of chloroform molecules from the cavity of bambusuril plays an important role in this system, as discussed above in Binding in Chloroform. The fact that the stability of 1· X− complexes in chloroform is lower than expected may imply that expulsion of chloroform molecules does not positively contribute to complex formation but rather decreases stability of 1·X− complexes in this medium. In that case, the entropic penalty for encapsulated chloroform expulsion would have to be higher than the enthalpic gain from this process. Unlike in water where “high-energy water molecules” are expelled, “low-energy chloroform molecules” would be liberated in CHCl3. Binding in chloroform would thus be driven by a concert of anion desolvation and C−H···X− interactions. To gain deeper insight into the possible source of unexpected thermodynamic behavior of 1·X− in chloroform, we first expressed selectivities of bambusuril 1 for one halide over the other as difference in log Ka values between the two corresponding complexes (Δlog Ka). Then we plotted the selectivities vs the Swain acity parameter (Figure 3b, Table S4). Any anion-independent effect on 1−X− binding which significantly affects the log Ka vs acity plot (Figure 3a) should be mitigated in the selectivity plot (Figure 3b). Indeed, a linear relationship between Δlog Ka and acity was found (Figure 3b) in which chloroform no longer stood out as an outlier. This indicates that an anion-independent effect causes the unexpected behavior of 1−X− binding in CHCl3. Because expulsion of encapsulated solvent molecules is expected not to depend on the nature of the guest, endergonic release of chloroform molecules upon binding remains a hot candidate for explaining the thermodynamic behavior of 1·X− complexes in chloroform. Experiments with anions other than halides are underway to deeper probe the nature of this phenomenon. Selectivity of bambusuril 1 between halide anions is significantly reduced in solvents with low acity compared to those with higher A (Figure 3b, Table S4). Thus, an outstanding selectivity of 27 000 for iodide over chloride was found in water (A = 1.00) compared to a significantly lower value of 1400 in less polar but protic n-propanol (A = 0.63), and a modest selectivity of 84 in the polar aprotic acetonitrile (A = 0.37). It should be stated, though, that DMSO and to a certain extent also acetonitrile, both polar aprotic solvents, emerged as outliers in the selectivity plot (Figure 3b). The macrocycle has negligible selectivity for all three halides in DMSO and diminished

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CONCLUSION

We report a multigram scale synthesis of a bambusuril anion receptor 1, which is soluble in a number of solvents with different polarities varying from water to tetrachloromethane. The bambusuril macrocycle is capable of exceptionally strong anion binding which is generally ethalpically driven. Although anions are stabilized inside bambusuril by multiple C−H···X− hydrogen bonding interactions, these are not the main driving force for complex formation, especially in solvents with a higher degree of molecule organization. In the absence of anions, the macrocycle’s cavity is filled with solvent molecules, which are released to the bulk solvent upon anion inclusion. In the case of highly structured solvents (such as water), the release of high energy solvent molecules is the main driving force for anion binding. We further showed that for all halides, the stability of 1·X− complexes decreases with an increasing Swain acity parameter of solvents, indicating an important role of anion solvation. It is important to note that while the contribution from high-energy solvent molecule release is essentially independent of the type of anion, desolvation is significantly more pronounced for the small chloride anion with high charge density compared to the large iodide. Thus, macrocycle 1 is selective for iodide over bromide and chloride in all investigated solvents. The selectivity is the highest in water and decreases with decreasing acity of solvent. Weak iodide solvation is likely the reason why the 1·I− complex stability in water (were the high-energy solvent molecule release is the most pronounced) is comparable or even stronger than in organic solvents. Our study also revealed that this nonclassical solvophobic effect plays a significant role not only in water but also in chloroform and thus supports the recently published results of the structured nature of this solvent.39 Correlation of complex stabilities with the Swain acity parameter provides clues that, unlike in water, expulsion of chloroform molecules from the bambusuril cavity might be the cause of decreased binding strength in this solvent. 1908


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phase was washed with water (100 mL), dried over anhydrous MgSO4, and concentrated in vacuo to give tosylate 6 as a colorless liquid. Yield: 66.0 g (0.207 mol, 97%). 1H NMR (500 MHz, CDCl3): 2.43 (s, 3H, C12H3); 3.35 (s, 3H, C7H3); 3.51 (m, 2H, C6H2); 3.56−3.60 (m, 6H, C3,4,5H2); 3.67 (t, J = 4.9 Hz, 2H, C2H2); 4.15 (t, J = 4.9 Hz, 2H, C1H2); 7.32 (d, J = 8.2 Hz, 2H, C10H); 7.78 (d, J = 8.2 Hz, 2H, C9H). 13C NMR (125 MHz, CDCl3): 21.7 (C12H3); 59.1 (C7H3); 68.8 (C2H2); 69.3 (C1H2); 70.65, 70.66 (C3,4H2); 70.9 (C5H2); 72.0 (C6H2); 128.1 (C9H); 129.9 (C10H); 133.2 (C11); 144.9 (C8). The characterization data match those previously reported.56 Triethylene Glycol Monomethyl Ether Monoazide (7). Procedure A: Compound 7 was prepared according to a literature procedure.57 A

EXPERIMENTAL SECTION

General Methods. All chemicals were purchased from commercial suppliers and used as received except for dry solvents. Commercial solvents (HPLC grade) were dried over activated 3 Å molecular sieves (20% v/v) for at least 3 days. Molecular sieves were activated by heating at 260 °C in vacuum for 24 h. Thin-layer chromatography (TLC) was performed on aluminum plates coated with silica gel (60 Å pore diameter, 200 μm layer thickness, 8.0−12.0 μm particle size) with fluorescent indicator. The analyte was detected by UV light (254 nm) or by staining with acidic ninhydrin, basic KMnO4, or methanolic H2SO4 or in an iodine chamber. Column chromatography separations were performed using Acros Organics silica gel (60 Å pore diameter, 40−60 μm particle size). 1H NMR and 13C NMR spectra were recorded at 30 °C on a Bruker Avance 300 spectrometer (working frequency 300.13 MHz for 1H, 75.48 MHz for 13C) or Bruker Avance 500 spectrometer (working frequency 500.13 MHz for 1H, 75.48 MHz for 13C). Both spectrometers were equipped with the BBFO probe. Chemical shifts were referenced to solvent residual peaks55 and are given in parts per million (ppm), and coupling constants (J) are given in hertz (Hz). The multiplicities of signals are reported as singlet (s), doublet (d), triplet (t), multiplet (m), broad (br), or a combination of these. IR spectra were recorded on a Bruker Alpha FTIR Platinum ATR spectrometer. Characteristic absorptions of major functional groups are reported in reciprocal centimeters (cm−1). Matrix-assisted laser desorption ionization−time-of-flight mass spectra (MALDI-TOF MS) were measured on a MALDI-TOF MS UltrafleXtreme (Bruker Daltonics), and samples were ionized with the aid of a Nd:YAG laser (355 nm) from 2,5-dihydroxybenzoic acid (DHB) or α-cyano-4-hydroxycinnamic acid (HCCA) matrixes. High-resolution mass spectra (HRMS) were obtained using an Agilent 6224 Accurate-Mass TOF, and samples were ionized by electrospray ionization (ESI) or atmospheric-pressure chemical ionization (APCI). Thermodynamic measurements were performed using a VP-ITC microcalorimeter (Microcal, Malvern Instruments). The experiments were carried out at 303.15 ± 0.1 K using HPLC-grade solvents except for measurements in water. Ultrapure water for was obtained using a Barnstead MicroPure, Thermo Scientific system). Salt solutions were automatically injected into the solution of bambusuril 1 present in the calorimeter cell while stirring at 307 rpm. Correction for dilution effects was performed by titration of the salt solution into the pure solvent, and the results were subtracted from the corresponding data. Integrated heat effects were analyzed by nonlinear regression using a single-site model or competitive model (Microcal Origin 7). The experimental data fitted to a theoretical titration curve provided the association constant Ka and the standard binding enthalpy ΔH°. The other thermodynamic parameters, such as standard free energy ΔG° and standard entropy ΔS°, were calculated from the equation: ΔG° = ΔH° −TΔS° = −RT × ln Ka, where T is the absolute temperature and R is the molar gas constant (8.3145 J mol−1 K−1). pH measurements were conducted using an Orion Star A211 pH meter with an Orion Micro glass microelectrode (Thermo Scientific) at room temperature. The pH meter was calibrated using standard Orion Application Solution buffers (Thermo Scientific), pH = 4.01, 7.00, and 10.01. Synthetic Procedures. Triethylene Glycol Monomethyl Ether Monotosylate (6). Compound 6 was prepared according to a literature

solution of tosylate 6 (65 g, 0.20 mol) and sodium azide (26.5 g, 0.408 mol) in acetone (250 mL) and water (50 mL) was heated at reflux overnight. The reaction mixture was concentrated in vacuo to approximately 100 mL, diluted with water (150 mL), and extracted with ether (4 × 150 mL). The combined organic extracts were dried over anhydrous MgSO4 and concentrated in vacuo to provide pure azide 7 as a slightly yellowish liquid. Yield: 37.1 g (0.196 mol, 96%). Procedure B: Alternatively, compound 7 could be prepared directly from triethylene glycol monomethyl ether (5). Diphenyl phosphoryl azide (7.14 g, 25.9 mmol), triethylene glycol monomethyl ether (5, 3.45 mL, 21.6 mmol), 1,8-diazabicyclo[5.4.0]undec-7-ene (DBU, 3.90 mL, 26.1 mmol), and sodium azide (2.97 g, 45.7 mmol) were dissolved in dry dimethylformamide (DMF, 35 mL) and heated under argon at 90 °C for 2 h. The reaction mixture was left to cool to rt, diluted with water (120 mL) and extracted with ether (4 × 75 mL). The combined organic extracts were washed with brine (80 mL), dried over anhydrous MgSO4, and concentrated in vacuo. The crude product was purified by silica gel column chromatography using pentane:ether (9:1 v/v) to CH2Cl2:MeOH (49:1 v/v) as the mobile phase. Compound 7 was obtained as a yellowish liquid. Yield: 3.15 g (16.6 mmol, 77%). 1H NMR (300 MHz, CDCl3): 3.33−3.36 (m, 5H); 3.51 (m, 2H); 3.60−3.65 (m, 8H). 13C NMR (75 MHz, CDCl3): 50.7; 58.9; 70.0; 70.6; 70.6; 70.7; 71.9. The characterization data match those previously reported.57 Triethylene Glycol Monomethyl Ether Monoamine (8). Azide 7 (37.1 g, 0.196 mol) was dissolved in THF (300 mL). Triphenylphos-

phine (56.6 g, 0.216 mol) was added in portions, and the resulting solution was stirred under nitrogen at rt overnight. Water (100 mL) was added, and the mixture was stirred at rt over another night. The solvents were evaporated in vacuo with azeotropic coevaporation with toluene (2 × 100 mL). The crude yellow liquid was treated with cold ether (−20 °C, 400 mL) resulting in crystallization of triphenylphosphine oxide. The side-product was filtered off, and the filtrate was concentrated in vacuo. Water (300 mL) was added, and the aqueous solution was washed with ether (3 × 100 mL) and dichloromethane (100 mL). Water was removed in vacuo with azeotropic coevaporation with acetonitrile (2 × 50 mL), providing amine 8 as a pale-yellow liquid. Yield: 26.7 g (0.164 mol, 83%). 1H NMR (500 MHz, CDCl3): 2.08 (br s, 2H, NH2); 2.86 (t, J = 5.2 Hz, 2H); 3.36 (s, 3H, C7H3); 3.50 (t, J = 5.2 Hz, 2H); 3.53 (m, 2H); 3.60−3.65 (m, 6H). 13C NMR (125 MHz, CDCl3): 41.8; 59.1; 70.4; 70.6; 70.7; 72.0; 73.1. The characterization data match those previously reported.58 N,N′-Bis(3,6,9-trioxadecyl)urea (9). Procedure A: Hünig’s base (DIPEA, 2.7 mL, 16 mmol) and bis(4-nitrophenyl) carbonate (1.47 g,

procedure.56 Triethylene glycol monomethyl ether (5, 35.17 g, 0.214 mol) was dissolved in THF (65 mL), and the solution was cooled to 0 °C. A 6 M aq NaOH (65 mL) solution was added dropwise over 10 min, followed by a dropwise addition of a tosyl chloride (TsCl, 51 g, 0.27 mol) solution in THF (70 mL) over 20 min. The reaction mixture was further stirred at 0 °C for 60 min, left to warm to rt, and stirred for 80 min. The resulting solution was diluted with ether (500 mL) and 1 M aq NaOH (175 mL). After the aqueous layer was separated, the organic 1909


Article


72.1 (CPEGH2); 158.8, 161.0 (CO). IR (ATR): 1694.4 (CO), 3325.9 (N−H). HRMS: calculated Mr for [C18H34N4O8 + H]+, 435.2449; found, 435.2447. Dodecakis(3,6,9-trioxadecyl)bambus[6]uril−NaBF4 Complex (1· NaBF4). A solution of glycoluril 4 (3.27 g, 7.53 mmol), paraformalde-

4.83 mmol) were added to a solution of amine 8 (1.68 g, 10.3 mmol) in dry dichloromethane (20 mL). The reaction mixture was stirred under argon at rt for 3 h. The solvent was evaporated in vacuo, and the crude product was dissolved in 1 M aq NaOH (50 mL). The solution was washed with ether (3 × 50 mL) and extracted with CH2Cl2 (4 × 50 mL). The combined organic extracts were dried over anhydrous MgSO4 and concentrated in vacuo to provide urea 9 as a yellow oil. Yield: 1.70 g (4.82 mmol, quantitative). Procedure B: A solution of amine 8 (25.3 g, 0.155 mol) and diphenyl carbonate (16.27 g, 0.0759 mol) in DIPEA (67.5 mL, 0.388 mol) was heated at reflux for 3 h. The resulting solution was dissolved in 1 M aq NaOH (400 mL), washed with ether (3 × 200 mL), and extracted with CH2Cl2 (4 × 200 mL). The combined organic extracts were dried over anhydrous MgSO4 and concentrated in vacuo to provide urea 9 as a yellowish oil. Yield: 26.2 g (0.0743 mol, 98%). 1H NMR (500 MHz, CDCl3): 3.34−3.38 (m, 10H, C1H2 and C7H3); 3.52− 3.60 (m, 8H, C2H2 and C?H2); 3.62−3.65 (m, 12H, remaining CH2); 5.35 (br s, 2H, NH). 13C NMR (125 MHz, CDCl3): 40.2 (C1H2); 59.0 (C7H3); 70.2; 70.4; 70.5; 70.8; 72.1; 158.7 (CO). HRMS: calculated Mr for [C15H32N2O7 + H]+ − 353.2282, found −353.2287. trans-4,5-Dihydroxyimidazolidin-2-one (10). Compound 10 was prepared by modifying a literature procedure.59 A mixture of urea (104

hyde (0.34 g, 11 mmol), and TBA+BF4− (2.48 g, 7.53 mmol) in dioxane (15 mL) was heated until it came to a boil. A 48% aq HBF4 solution (68.8 μL) was added, and the reaction mixture was heated at reflux for 14 h. The solvent was evaporated in vacuo, and the resulting oil was treated with water (70 mL), resulting in crystallization of excess TBA+BF4−. The salt was filtered off (regenerated yield: 1.36 g, 55%), and the resulting solution was washed with ether (4 × 40 mL) and extracted with CH2Cl2 (4 × 40 mL). The combined organic extracts were dried over anhydrous MgSO4 and concentrated in vacuo to provide 5.17 g of crude brown oil containing mainly macrocycle 1 and the corresponding bambus[4]uril (2:1 m/m) and TBA+BF4−. The product was isolated by silica gel column chromatography using CH2Cl2:MeOH (19:1 to 9:1 v/v) as the mobile phase. Due to ion exchange properties of silica gel, macrocycle 1 was obtained as its Na+ (instead of TBA+) BF4− complex as a colorless oil. Yield: 1.58 g (0.567 mmol, 45%). 1H NMR (300 MHz, CDCl3): 3.36 (s, 36H, C9H3); 3.50−3.78 (m, 132H, C4−8H2 + C3Ha); 3.96 (m, 12H, C3Hb); 5.08 (s, 12H, C2H2); 5.30 (s, 12H, C1H). Dodecakis(3,6,9-trioxadecyl)bambus[6]uril−NaCl Complex (1· NaCl). Macrocycle 1 was fully characterized as the 1·NaCl complex. 1·

g, 1.73 mol) and 40% aq glyoxal solution (100 mL, 0.87 mol) was heated until complete dissolution. Thirty drops of a 1% phenolphthalein indicator solution in EtOH−H2O 7:3 was added followed by 40% aq NaOH until a constant color change. The solution was further heated at 60 °C for 40 min. Every 5 min, the pH of the solution was adjusted with 40% aq NaOH until a constant color change. The reaction mixture was allowed to cool to rt during which a white crystalline solid emerged. The product was filtered, washed with EtOH (100 mL), and dried to obtain 31.3 g of product 10. The filtrate was left to stand at 4 °C for 1 week during which a second portion of product crystallized. The compound was filtered, washed with EtOH (100 mL), and dried to obtain 22.0 g of product 10. The filtrate was concentrated in vacuo to approximately half the original volume and diluted with EtOH (100 mL). The solution was seeded with crystals of pure compound 10 and stirred at rt overnight during which more product crystallized. The compound was filtered, washed with EtOH (250 mL), and dried to obtain 10.9 g of product 10. Total yield: 64.2 g (0.54 mol, 62%). Melting point: 143−144 °C. 1H NMR (500 MHz, DMSO-d6): 4.59 (d, J = 6.6 Hz, 2H, OH); 5.80 (d, J = 6.6 Hz, 2H, CH); 7.03 (br s, 2H, NH). 13C NMR (125 MHz, DMSOd6): 83.9 (CH); 160.3 (CO). The characterization data match those previously reported.59 2,4-Bis(3,6,9-trioxadecyl)glycoluril (4). A suspension of trans-4,5dihydroxyimidazolidin-2-one 10 (17.9 g, 152 mmol) and urea 9 (10.67

NaBF4 (335 mg, 120 μmol) was dissolved in HPLC grade CH2Cl2 (90 mL). The solution was divided into 15 mL aliquots in six centrifuge tubes. Each sample was washed with brine (15 × 15 mL). Separation of layers was extremely slow and had to be achieved by centrifugation. The organic layers were combined and concentrated in vacuo with azeotropic removal of residual water with CH3CN (2 × 25 mL) to provide 1·NaCl as a colorless oil. Yield: 299 mg (109 μmol, 91%). 1H NMR (500 MHz, CDCl3): 3.36 (s, 36H, C9H3); 3.52 (m, 24H, CPEGH2); 3.58−3.69 (m, 96H, CPEGH2 + C3Ha + C4Ha); 3.77 (m, 12H, C4Hb); 3.93 (m, 12H, C3Hb); 5.15 (s, 12H, C2H2); 5.38 (s, 12H, C1H). 13 C NMR (125 MHz, CDCl3): 44.4 (C3H2); 47.5 (C2H2); 59.1 (C9H3); 68.7 (C1H); 69.5 (C4H2); 70.2, 70.3, 70.5, 71.9 (CPEGH2); 159.1, 160.5 (CO). IR (ATR): 1695.5 (CO). HRMS: calculated Mr for [C114H204N24O48 + H]+, 2678.4333; found, 2678.4287. Anion-Free Dodecakis(3,6,9-trioxadecyl)bambus[6]uril (1). Procedure A: 1·NaBF4 was continuously extracted over 1 week with water,

g, 30.3 mmol) in MeOH (100 mL) and water (50 mL) was heated until it came to a boil. A 48% aq HBF4 solution (0.59 mL) was added, and the reaction mixture was heated at reflux for 100 min. The solvents were removed in vacuo, the residue was dissolved in 1 M aq HCl (200 mL) and washed with ether (2 × 100 mL). The product was extracted with CH2Cl2 (4 × 100 mL), dried over anhydrous MgSO4, and concentrated in vacuo to provide glycoluril 4 as a yellowish oil. Yield: 10.4 g (23.9 mmol, 79%). 1H NMR (500 MHz, CDCl3): 3.13−3.18 (m, 2H, C2Ha); 3.37 (s, 6H, C8H3); 3.55−3.73 (m, 20H, C3−7H2); 3.80−3.83 (m, 2H, C2Hb); 5.24 (s, 2H, C1H); 6.38 (br s, 2H, NH). 13C NMR (125 MHz, CDCl3): 43.8 (C2H2); 59.0 (C8H3); 69.0 (C1H); 70.2, 70.4, 70.7, 71.1, 1910


The Journal of Organic Chemistry which was passed through the solution of 1·NaBF4 (500 mg, 179 μmol) in CH2Cl2 (200 mL). The resulting CH2Cl2 solution was concentrated in vacuo with azeotropic removal of residual water with CH3CN (2 × 80 mL) to obtain anion-free 1 as a colorless oil. Yield: 375 mg (134.5 μmol, 75%). Procedure B: 1·NaBF4 (270 mg, 96.8 μmol) was dissolved in HPLC grade CH2Cl2 (90 mL). The solution was divided into 15 mL aliquots in six centrifuge tubes. Each sample was washed with brine (15 × 15 mL) and ultrapure water (20 × 15 mL). Separation of layers was extremely slow and had to be achieved by centrifugation. The organic layers were combined and concentrated in vacuo with azeotropic removal of residual water with CH3CN (2 × 25 mL) to provide anionfree 1 as a colorless oil. Yield: 221 mg (82.5 μmol, 85%). 1H NMR (300 MHz, D2O): 3.42 (s, 36H, C9H3); 3.63 (m, 24H, CPEGH2); 3.68−3.82 (m, 96H, CPEGH2); 3.93 (m, 24H, CPEGH2); 5.23 (s, 12H, C2H2); 5.55 (s, 12H, C1H). To provide evidence that the obtained macrocycle 1 was truly anion-free, a 1:1 mixture of 1 and anion-free Bn12BU[6] (3) was dissolved in CDCl3 and a 1H NMR spectrum was taken (Figure S13). No signal corresponding to 3·BF4− (for procedure A) and 3·Cl− (for procedure B) was observed. Crystallographic Data. Crystal Structure of Anion-Free Bn12BU[6]·2CHCl3 (3·2CHCl3). A saturated solution of anion-free bambusuril 3 in CHCl3 was cooled to −10 °C, resulting in formation of a colorless crystalline solid. X-ray crystallography revealed that the crystals contained bambusuril 3 encapsulating two chloroform molecules in its cavity (Figure S1). Crystallographic data for C70H54N12O6·5CHCl3, Mr = 1599.96, crystal dimensions 0.25 × 0.14 × 0.09 mm, space group P1, a = 14.0083(12) Å, b = 15.8585(10) Å, c = 18.1756(15) Å, α = 98.852(6)°, β = 110.605(8)°, γ = 104.194(6)°, V = 3533.4(5) Å3, Z = 2, ρ = 1.504 g/cm3, μ = 5.836 mm−1. X-ray intensity data were measured at 120 K on a on a Gemini, Oxford Diffraction, with Ultra-Cu collimator (Cu Kα radiation) and CCD detector Atlas. 24449 reflections collected, 8589 unique reflections, 2942 reflections with I > 2σ(I), Rint = 0.1689, θmax = 54.24°, R[F2 > 3σ(F2) = 0.1364, wR(F2) = 0.3320, S = 1.044, parameters = 643, Δρmax = 0.925 eA−3, Δρmin = −0.641 eA−3. Structure solution and refinement was performed by SHELX97.60 CCDC 809241 contains the supplementary crystallographic data for this structure. These data can be obtained free of charge from The Cambridge Crystallographic Data Centre via www.ccdc.cam.ac.uk/data_request/ cif.

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ACKNOWLEDGMENTS

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REFERENCES

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ASSOCIATED CONTENT

The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/acs.joc.7b02846. 1 H, 13C, and 2D NMR spectra, IR spectra, MALDI MS, NMR and ITC titrations, the crystal characterization, and computational details (PDF) CIF data for 3 (CIF) Ball-and-stick model for dodecamethylbambus[6]uril (BU6) (XYZ) Ball-and-stick model for the BU6·Br− complex (XYZ) Ball-and-stick model for the BU6·Cl− complex (XYZ) Ball-and-stick model for the BU6·F− complex (XYZ) Ball-and-stick model for the BU6·I− complex (XYZ)

AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected]. ORCID

Vladimir Sindelar: 0000-0003-0090-5961 Author Contributions ∥

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We thank Dr. Vaclav Havel and Dr. Michal Dusek for providing crystallographic data for this publication. We acknowledge Dr. Petr Kulhanek for DFT calculations. This work was supported by the Czech Science Foundation (13-15576S) and the RECETOX Research Infrastructure (LM2015051 and CZ.02.1.01/0.0/0.0/ 16_013/0001761). We acknowledge the CF Proteomics supported by the CIISB research infrastructure (LM2015043 funded by MEYS CR).

S Supporting Information *

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These authors contributed equally to this work.

Notes

The authors declare no competing financial interest. 1911


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