Solvation Effect on Complexation of Alkali Metal Cations by a Calix[4

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Solvation Effect on Complexation of Alkali Metal Cations by a Calix[4]arene Ketone Derivative Josip Pozar, Ivana Niksic-Franjic, Marija Cvetnic, Katarina Leko, Nikola Cindro, Katarina Piculjan, Ivana Borilovic, Leo Frkanec, and Vladislav Tomisic J. Phys. Chem. B, Just Accepted Manuscript • DOI: 10.1021/acs.jpcb.7b05093 • Publication Date (Web): 14 Aug 2017 Downloaded from http://pubs.acs.org on August 14, 2017

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The Journal of Physical Chemistry

Solvation Effect on Complexation of Alkali Metal Cations by a Calix[4]arene Ketone Derivative

Josip Požar,*,† Ivana Nikšić-Franjić,† Marija Cvetnić,† Katarina Leko,† Nikola Cindro,† Katarina Pičuljan, † Ivana Borilović,† Leo Frkanec,‡ and Vladislav Tomišić*,†

†

Department of Chemistry, Faculty of Science, University of Zagreb, Horvatovac 102a, 10000

Zagreb, Croatia ‡ Department

of Organic Chemistry and Biochemistry, Ruđer Bošković Institute, Bijenička 54,

10000 Zagreb, Croatia

E-mail: [email protected] [email protected] Phone: J. Požar +385 1 46 06 133 V. Tomišić +385 1 46 06 136

ACS Paragon Plus Environment

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ABSTRACT The medium effect on the complexation of alkali metal cations with a calix[4]arene ketone derivative (L)

was

systematically

examined

in

methanol,

ethanol,

N-methylformamide,

N,N-

dimethylformamide, dimethyl sulfoxide, and acetonitrile. In all solvents the binding of Na + cation by L was rather efficient, whereas the complexation of other alkali metal cations was observed only in methanol and acetonitrile. Complexation reactions were enthalpically controlled, while ligand dissolution was endothermic in all cases. A notable influence of the solvent on NaL+ complex stability could be mainly attributed to the differences in complexation entropies. The higher NaL+ stability in comparison to complexes with other alkali metal cations in acetonitrile was predominantly due to a more favorable complexation enthalpy. The 1 H NMR investigations revealed a relative ly low affinity of the calixarene sodium complex for inclusion of the solvent molecule in the calixare ne hydrophobic cavity, with the exception of acetonitrile. Differences in complex stabilities in the explored solvents, apart from N,N-dimethylformamide and acetonitrile, could be mostly explained by taking into account solely the cation and complex solvation. A considerable solvent effect on the complexation equilibria was proven to be due to an interesting interplay between the transfer enthalpies and entropies of the reactants and the complexes formed.


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INTRODUCTION Complexation properties of calixarene derivatives have been extensively investigated during the past several decades.1 This class of compounds has been exploited for different tasks,2,3 and in recent advancements the potentials of calixarene-based compounds as biomimetics,4

drug-deliver y

systems,5 ion chanells,6,7 and nanomaterial components8,9 have been particularly explored. Among the vast variety of prepared calixarenes, those functionalized with carbonyl-containing substitue nts at the lower rim (calixarene ketones, esters or amides) have been recognized as efficient receptors of alkali metal cations.10,11 Such derivatives have found various applications, e.g. as extraction reagents,12–14 catalysts,14 active substances in ion selective electrodes,12,15 fluorescence sensors,12,16,17 and biologically active compounds.4,18 A rather high ability of the above mentioned ligands for cation complexation is a consequence of a well-preorganized binding site which contains electron-rich oxygen donor atoms. The affinities of these macrocycles towards first-group cations have been found dependent on the cation size and type of functionality attached to the carbonyl groups. Calix[4]arene-based ligand s prefer smaller cations (Li+, Na+), whereas larger macrocycles are more suitable for Rb+ and Cs+ complexation.19,20 In general, stability constants of the complexes with tertiary calixarene amides are higher than those with the corresponding secondary amides and simple ketone and ester derivatives.13,19–22 Complexation

is strongly

influenced

by solvation

of reactants

and

complexes.13,19–25 Much lower affinities of calixarene ketones, amides, and esters towards alkali metal cations in methanol compared to acetonitrile have been attributed to intermolecular hydrogen bonding with the solvent molecules.26–28 However, the solubilities of calixarenes in methanol and acetonitrile were comparable in some cases (standard Gibbs energies of macrocycle transfer, Δt G°, among these solvents amounted to only a few kJ mol–1 )21,24,29 As it turned out, differences in


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standard thermodynamic parameters of cation complexation were predominantly a consequence of the favorable complex transfer from methanol to acetonitrile.21,24 This was ascribed, at least partly, to concomitant inclusion of the acetonitrile molecule into the hydrophobic cavity of the complex, 30– 32

as the affinity of macrocycles towards methanol was found to be considerably lower.17,21,22,24

Investigations of calixarene complexation properties in other solvents are less common. The alkali metal cation binding of several low-rim derivatives was studied in benzonitrile.24,29,33 The stability constants obtained in this solvent were quite similar to those determined in acetonitrile. Calorimetr ic investigations indicated this to be consequence of interesting enthalpy-entropy compensation.24,29,33 According to Danil de Namor et al. who studied complexation of Na+ with ketone calix[4]are ne derivative in N,N-dimethylformamide,29 the macrocycle affinity towards the cation in this solvent was more than five orders of magnitude lower than in both acetonitrile and benzonitrile. Detailed thermodynamic analysis of the complexation process revealed that this was due to much stronger free cation solvation in amide on the one hand, and to more favorable solvation of the product in acetonitrile, on the other. One of the most comprehensive thermodynamic investigations of the solvation effect on calix[4]arene binding abilities was that regarding the complexation of Ag+ with a simple lower-rim derivative containing nitrogen and sulfur donor atoms in six solvents (methano l, ethanol, 1-propanol, acetonitrile, benzonitrile and N,N-dimethylformamide).32,34 The macrocycle affinity for Ag+ complexation in simple alcohols was considerably higher when compared to acetonitrile as a result of particularly favorable cation solvation in the latter solvent. On the other hand, solvation of product favored complexation in acetonitrile. A somewhat different trend was noticed in N,N-dimethylformamide.34 Solvation of the complex was more favorable than in acetonitrile, whereas the opposite was observed for the free cation. There have also been several examples of alkali metal cation complexation in mixed solvents. Danil de Namor et al.35 reported


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an interesting dependence of the standard thermodynamic parameters of sodium complexation with a simple ketone derivative on the solvent composition in N,N-dimethylformamide-acetonitr ile mixtures. The authors have correlated the stability of the sodium complex and differences in affinity for acetonitrile molecule inclusion depending on the amount of N,N-dimethylformamide present. More recent thermodynamic and structural investigations, revealed that selective solvation of reactants and products most likely occurs in methanol-dichloromethane

and acetonitrile-

dichloromethane solvent mixtures.17 The brief introduction given above indicates that despite a great progress in the understanding of thermodynamics of calixarene complexation reactions, the data regarding the influence of the solvent on the binding process remained relatively limited. Namely, the binding affinities of lower-rim calix[4]arene derivatives towards a larger number of cations were usually examined in two or three solvents, or complexation with only one ion was explored in a larger number of solvents. In this work we have therefore decided to investigate the alkali metal cation complexatio n with

a

simple

ketone

derivative,

namely,

5,11,17,23-tetra-tert-butyl-26,28,25,27-

tetra(acetoyl)methoxycalix(4)arene (L, Figure 1) in methanol (MeOH), ethanol (EtOH), Nmethylformamide (NMF), N,N-dimethylformamide (DMF), dimethyl sulfoxide (DMSO), and acetonitrile (MeCN). The solvents were chosen taking into account differences in their cation and ligand solvation abilities and the possibility of ligand-solvent and complex-solvent hydrogen bonding established through interactions of calixarene oxygen atoms as proton acceptors and solvent molecules as proton donors. For instance, NMF, DMF, and DMSO solvate the charged species particularly strongly. However, only NMF is capable of forming hydrogen bonds with the ligand L pendant arms. NMF molecules are strongly hydrogen bonded themselves, whereas only strong dipole-dipole interactions are present in DMF and DMSO. Likewise, the EtOH molecules form


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fewer and weaker hydrogen bonds when compared with MeOH. This affects the solvation of the free cations in the mentioned alcohols, which should also hold for the free ligand and the complexes formed. Apart from investigations concerning the thermodynamics of alkali metal cations complexation, we examined inclusion of the solvent molecules into the hydrophobic cavity of the sodium complex in chloroform in order to gain insight whether concomitant inclusion of the solvent molecules accompanied coordination of cations. The previous computational studies of ligand L and its complexes with Cd2+ and Pb2+ in MeCN indicated that this could be expected in the case of this solvent.36 It should be mentioned that binding affinity of L towards alkali metal cations was previously examined in MeCN and MeOH spectrophotometrically. 37 The reported stability constants were in some cases too high to be reliably determined with this method. In addition, no attempt at a more comprehensive thermodynamic study of the corresponding complexation reactions has been made. The ligand extraction properties have also been examined,38–40 and Danil de Namor et al. studied the binding of trivalent cations with the receptor in MeCN.41 In addition, Kriz et al. explored its affinity toward H3 O+ in nitrobenzene.42 The herein reported standard thermodynamic complexation parameters, ligand solubilities, and standard Gibbs energies, enthalpies, and entropies of cation transfer between the explored solvents have enabled detailed interpretation of the solvent effect on the ligand complexatio n properties in methanol, acetonitrile, and in solvents which are not so often used in the investigatio ns of calixarene ionophoric abilities.


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Figure 1. Structure of compound L

EXPERIMENTAL SECTION Materials. Compound L was prepared according to procedure described elsewhere. 37 The solvents methanol (J. T. Baker, HPLC Gradient Grade), ethanol (Sigma-Aldrich,

Spectranal), N-

methylformamide (Sigma-Aldrich, 99 %), N,N-dimethylformamide (Sigma-Aldrich, HPLC Grade), and dimethyl sulfoxide (Sigma-Aldrich, HPLC Grade) were used without further purification. The salts used for the investigation of calixarene complexation were LiClO 4 (Sigma Aldrich 99.99 %), NaClO 4 (Sigma Aldrich 98+ %), KClO 4 (Merck, p.a.), KCl (Sigma-Aldrich), RbCl (Sigma-Aldric h), CsCl (Sigma-Aldrich), CsNO 3 (Sigma, 99.5 %), and RbNO 3 (Sigma, 99.7 %). Due to inertness of the perchlorate anion regarding ion pairing, perchlorates were used whenever these salts were soluble enough in particular solvent to allow reliable determination of the complex stability constants. When this condition was not met, salts with other counterions, e.g. chloride and iodide, were used. In order to check a possible influence of the anion on complexation reaction, some titrations were performed with salts comprising different anions (NaCl and NaClO 4 ).


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Spectrophotometry. Spectrophotometric titrations were carried out at (25.0 ± 0.1) o C by means of a Varian Cary 5 and Agilent Cary 60 spectrophotometers equipped with a thermostatting devices. Spectral changes of L solutions (c0 = 2 × 104 mol dm3 , V0 = 2.3 cm3 ) were recorded upon stepwise addition of alkali metal salt solutions (c = 2 × 103 mol dm3 to 3 × 102 mol dm3 ) into the measuring quartz cell (Hellma, Suprasil QX, l = 1 cm). Absorbances were sampled at 1 nm intervals, with an integration time of 0.2 s. Titrations for each M+/ L (M+ stands for alkali metal cation) system were done in triplicate. The obtained spectrophotometric data were processed using HYPERQUAD 43 program. Calorimetry. Microcalorimetric measurements were performed with an isothermal titration calorimeter Microcal VP-ITC at 25.0 °C. Enthalpy changes were recorded upon stepwise, automatic addition of alkali metal salt solution (c = 2 × 103 mol dm–3 to 3 × 102 mol dm–3 ) to macrocycle solution (c0 = 1 × 104 mol dm–3 to 4 × 104 mol dm–3 ). Blank experiments were carried out in order to make corrections for the enthalpy changes corresponding to the dilution of the alkali metal salt solution in the pure solvent. The dependence of successive enthalpy change on the titrant volume was processed using the Microcal OriginPro 7.0 and OriginPro 7.5 programs. Titrations for each cation/ligand system were repeated three or more times. The calorimeter reliability was checked by carrying out the complexation of barium(II) by 18- crown6 in water at 25.0 °C. The results obtained (log K = 3.74, ΔrH = –30.9 kJ mol–1 ) were in good agreement with the literature values (log K = 3.73, ΔrH = –31.5 kJ mol–1 ).44 Solubility measurements. Saturated solutions of L were prepared by adding excess amounts of the solid to the solvents explored. The obtained mixtures were left in a thermostat at 25.0 °C for several days in order to equilibrate. After the equilibrium had been reached, solutions aliquots were taken for solubility determination. Concentrations of saturated L solutions of at 25.0 °C were


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determined spectrophotometrically by mean of an Agilent Cary 60 spectrophotometer equipped with a thermostatting device. Molar absorption coefficients of the compound were obtained by measuring the absorbances of L solutions of known concentrations. Dissolution enthalpies of the compound were determined from the temperature dependence of solubilities by means of van't Hoff equation (ligand dissolution was found to be very slow). Solvate formation was tested by placing a known amount of L (m = 40 mg) in a desiccator where it was exposed to a saturated atmosphere of investigated solvents for several days. No solvate formation was observed. NMR investigations. NMR spectra were recorded by means of a Bruker Avance III HD 400 MHz/54 mm Ascend spectrometer equipped with a 5 mm PA BBI 1H/D-BB probe head with zgradient and automated tuning and matching accessory. All proton spectra were acquired at 25.0 °C using 64 K data points, spectral width of 20 ppm, recycle delay of 1.0 s and 16 scans. CDCl3 was used as solvent and TMS as internal standard for proton chemical shifts.

1H

NMR titratio ns

were performed by recording spectral changes of L solutions in CDCl3 (c0 ≈ 6 × 103 mol dm3 , V0 = 0.500 cm3 ) upon stepwise addition of MeCN, EtOH, MeOH, NMF, DMF, or DMSO solutio ns in CDCl3 . The dependence of selected proton chemical shifts on reactant concentrations was processed by means of the HypNMR45 program.

RESULTS Complexation investigations. The affinity of compound L towards Li+ in methanol was explored spectrophotometrically and calorimetrically. No absorbance or enthalpy changes were recorded upon addition of LiClO 4 to ligand solutions in high excess (n(Li+) / n(L) ≈ 100). This finding was in disagreement with the previously reported spectrophotometrically determined value for LiL+


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(log K = 2.7).37 Since no complexation of Li+ could be observed by the two experimental methods in the herein presented study, it can be concluded that examined calixarene ketone does not bind this cation in MeOH efficiently. Spectral changes observed upon the addition of NaClO 4 or KCl to the L solution in MeOH were qualitatively similar. The UV spectrum of receptor solutions exhibited a hypochromic shift of its larger part, accompanied by the occurrence of an isosbestic point (Figures 2a and S1a, Supporting Information). Equilibrium constants for reactions of formation of complexes with 1:1 stoichiome tr y (and hence standard reaction Gibbs energies, ΔrG°) were calculated by the least squares non-linear regression analysis of spectrophotometric titration data (Table 1, Figures 2b and S1b, Supporting Information).

A

0.8

(a)

0.6

A282

(b)

0.72 0.68 0.64

0.4

0.60 0.2 0.56 0.0 250

0.52 260

270

280

290

300

310

0

2

4

6

8

10

12

14

+

 / nm

n(K ) / n(L)

Figure 2. a) Spectrophotometric titration of L (c = 1.77 × 10–4 mol dm–3 , V0 = 2.3 mL) with KCl (c = 1.00 × 10–2 mol dm–3 ) in methanol. l = 1 cm;  = 25.0 C. The spectra are corrected for dilution. b) Dependence of absorbance at 282 nm on n(KCl) / n(L) ratio. ■ experimental; ― calculated.

The results of the corresponding microcalorimetric titrations can be seen in Figures 3a, 3b, and S2, Supporting Information. By processing the collected data (Figures 3b and S2b, Supporting Information)

all standard thermodynamic

parameters of complexation


reactions (reaction

10

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enthalpies, entropies, and Gibbs energies) as well as the corresponding equilibrium constants for the binding of Na+ and K + with the investigated calixarene were calculated (Table 1). The stability constants determined by both methods are in good agreement with the values reported by ArnaudNeu et al.37 The standard reaction enthalpy for the complexation of K + is much less favorable, having as a result a lower KL+ stability constant. Standard reaction entropies are negative for both cations, more so in the case of Na+. The complexation of Rb+ and Cs+ was explored spectrophotometrically and calorimetrically. The stability of RbL+ is very low, and the binding of

(a)

52 48 44

(H) / mJ

Cs+ was not observed.

P / W

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


(b) 0.0 -0.2

40

-0.4

36

-0.6

32

-0.8

28 0

25

50

75

100

125

150

175

200

-1.0 0.0

t / min

0.5

1.0

+

1.5

2.0

2.5

n(Na ) / n(L)

Figure 3. a) Microcalorimetric titration of L (c = 1.73  104 mol dm3 , V = 1.42 ml) with NaClO4 (c = 1.91  103 mol dm3 ) in methanol at 25 °C; b) Dependence of successive enthalpy change on n(NaClO4 ) / n(L) ratio. ■ experimental; ― calculated.


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Table 1. Thermodynamic Parameters for Complexation of Alkali Metal Cations with Compound L in Examined Solvents at 25 °C.

solvent

cation log K(ML+) ± SE

MeOH

Na+

H SE kJ mol 1 r

S SE J K 1mol 1 r

5.36 ± 0.06a

–30.6 ± 0.3

5.56 ± 0.01b

–31.73 ± 0.05 –45.8 ± 0.4

3.11 ± 0.03a

–17.7 ± 0.2

3.33 ± 0.01b

–19.01 ± 0.05 –24.3 ± 0.5

Na+

5.77 ± 0.01b

–32.93 ± 0.04 –43.22 ± 0.07 –34.4 ± 0.2

K+

3.32 ± 0.01b

–18.94 ± 0.06 –21.88 ± 0.05 –9.8 ± 0.3

4.27 ± 0.01a

–24.37 ± 0.09

4.36 ± 0.01b

–24.91 ± 0.09 –48 ± 1

3.47 ± 0.02a

–19.8 ± 0.1

3.428 ± 0.001b

–19.56 ± 0.01 –49 ± 1

2.770 ± 0.004a

–15.81 ± 0.03

2.62 ± 0.00b

–14.97 ± 0.07 –50.20 ± 0.03 –118 ± 1

K+

EtOH

G SE kJ mol 1 r

–48 ± 2

–18 ± 2

NMF

Na+

DMF

Na+

DMSO

Na+

MeCN

Li+

7.19 ± 0.02b

–41.0 ± 0.1

–37.6 ± 0.2

11.6 ± 0.8

Na+

9.31 ± 0.02b

–53.1 ± 0.01

–66.7 ± 0.2

–45.6 ± 0.7

K+

5.02 ± 0.01b

–28.7 ± 0.01

–43.0 ± 0.2

–48.2 ± 0.5

Rb+

2.27 ± 0.03b

–12.9 ± 0.9

–27.8 ± 0.7

–53 ± 2

a

–76 ± 4

–98 ± 3

Determined spectrophotometrically. b Determined calorimetrically. SE = standard error of the mean

(N = 3 or 4).


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The complexation of Li+ with compound L in EtOH was explored analogously as in MeOH. The binding of this cation was too weak for reliable determination of complexations thermodyna mic parameters. This can be ascribed to its strongest solvation among the alkali metal cations.46 In contrast, the addition of Na+ to solution of L resulted in a notable hypochromic shift up to equivalence (Figure S3, Supporting Information), indicating strong binding of this cation. Because of the limited sensitivity of spectrophotometry, the stability constant of NaL+(EtOH) could not be determined reliably. Complexation was also studied by means of isothermal titration calorimetr y, which allowed for the determination of all standard complexation thermodynamic parameters for sodium (Table 1 and Figure S4, Supporting Information). The results were, both in terms of magnitude and signs of rH° and rS°, quite similar to those obtained in MeOH (Table 1). The binding of potassium in EtOH was investigated calorimetrically (Figure S5, Supporting Information). The determined ΔrG°, ΔrH°, and ΔrS° values are given in Table 1. The complexatio n of Rb+ and Cs+ could not be explored because of extremely low solubilities of their salts in this alcohol. However, the results obtained in methanol suggest low stabilities of RbL+ and CsL+ in EtOH. Complexation of Li+, K +, and Rb+ was observed in NMF. However, the corresponding stability constants were too low for their reliable determination. On the other hand, the addition of CsCl to ligand solution in considerable excess (up to cation to ligand molar ratios as high as 200) did not result in observable complex formation. Consequently, the standard thermodyna mic parameters of complexation could only be determined for the binding of Na+ with L in this solvent. Spectrophotometric titration of L with Na+ is shown in Figure S6 in Supporting Information, and the stability constant of NaL+ complex computed by processing these data is given in Table 1. The results of the corresponding microcalorimetric titrations can be seen in Figure S7 in Supporting


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Information. Standard reaction enthalpies and equilibrium constants were calculated by the leastsquares non-linear analyses of the calorimetric data (Table 1). The equilibrium constant for the complexation of sodium with the explored macrocycle in this solvent is approximately one order of magnitude lower than in studied alcohols. This is mainly due to the less favorable complexatio n entropy, as complexation enthalpies were quite similar in all studied solvents (Table 1). As an example of the results obtained in DMF, experimental and fitted data corresponding to the titration of L with Na+ are shown in Figure S8 in Supporting Information (the results of the corresponding spectrophotometric measurements are given in Figure S9, Supporting Informatio n). Complexation of Li+, K +, Rb+, and Cs+ in this solvent was not observed. The stability of NaL+ in DMF was lower than in NMF. By inspecting standard complexation parameters listed in Table 1, one can conclude that this is a consequence of notably less favorable standard complexation entropy. The entropy changes accompanying the binding of sodium in DMSO were lowest among the explored solvents. As a result, the stability of NaL+ in this solvent decreased even further in comparison to DMF (Table 1, Figures S10 and S11, Supporting Information). Again, complexatio n of other investigated cations could not be observed. Unlike in previously described solvents, compound L was found to be quite an efficie nt binder of Li+, Na+, and K+ in MeCN (Table 1, Figures S12-S16, Supporting information). Indeed, standard thermodynamic parameters of Na+ complexation could not be determined by direct calorimetric titration because it bound so strongly. The stability constant of NaL+ was obtained by processing the results of competitive experiments in which the Na + cation displaced the potassium cation in the KL+ complex. Likewise, the stability constant of LiL+ was determined by displaceme nt of Li+ in LiL+ with Na+. The ΔrH° values for the binding of lithium and sodium were obtained by direct calorimetric experiments assuming quantitative cation complexation.


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The thermodynamic data listed in Table 1 reveal that the remarkable stability of LiL+ arises from uniquely energetically and entropically advantageous cation coordination. Such behavior was also noticed in the case of other lower-rim calix[4]arene derivatives with carbonyl-containing substituents in MeCN.21,22,24,33 The highest stability (log K = 9.3) of NaL+ in MeCN is due to the most favorable ΔrH° value among all studied complexation reactions. The binding of K+ was also accompanied with negative enthalpy changes, however the measured values were much higher than in the case of sodium. In contrast, ΔrS° values for Na+ and K + complexation were rather simila r. Such combination of standard thermodynamic parameters of complexation resulted in a KL+ stability constant almost four orders of magnitude lower than that of the corresponding sodium complex. The compound L displayed modest affinity for the complexation of Rb+ in MeCN (log K = 2.27, Figure S17, Supporting Information), as a consequence of the least favorable complexation energetics. As expected, the complexation of Cs+ with compound L in MeCN was not observed because it was too large to fit well enough into the ligand cation-binding site.19,20 The difference in the ligand binding affinity for the lithium ion in MeOH and MeCN is strongly affected by quite favorable solvation of this small cation in hydrogen-bonded solvents (Δt G°(Li+, MeCN→MeOH) = –21 kJ mol–1 ).46 On the other hand, the transfer of Na+ cation from MeOH to DMSO is just as favorable (Δt G°(Na+, MeOH→DMSO) = –21 kJ mol–1 )46 as is its transfer from MeCN to MeOH. However, the compound L binds this cation quite efficiently in DMSO (log K = 2.6). Likewise, solvation of Rb+ in acetonitrile is more favorable than in methano l (Δt G°(Rb+, MeOH→MeCN) = –4 kJ mol–1 )46 , yet its complexation was observed in this solvent. In order to explain the influence of the solvent on the thermodynamic parameters of complexation reactions, the differences in solvation of the ligand, the cation and that of the corresponding complex need to be considered.19,21,24,29,32,34 This can be accomplished


by

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determining standard thermodynamic functions of the receptor and complex transfer. The Δt G° for the ligand can be determined from the differences in compound solubilities. The corresponding enthalpy (and hence Δt S°) can be obtained calorimetrically or from their temperature dependence by means of van’t Hoff equation. Standard thermodynamic functions of the complex transfer can be determined using Hess law by combining the standard thermodynamic transfer functions of the reactants with standard complexation parameters obtained in the two solvents. Finally, it should be mentioned that a possible anion effect on the studied reactions was checked by following Na+ complexation processes in MeOH, EtOH, and NMF using both sodium perchlorate and chloride (Figures S18-S20, Supporting Information). These were the only cases where salt solubilities allowed us to utilize more than one counterion. As no significant differe nces in the determined thermodynamic reaction parameters were observed (Table S1), it can be concluded that the influence of anion is negligible or does not exist.

Solubility investigations. The solubilities of ligand L in the investigated solvents are reported in Table S2 in Supporting Information. The standard solution Gibbs energies of the compound, listed in Table 2, were calculated from the solubility data using the equation:

ΔsolG° = –RT lnK° = –RT ln(γLs/c°) ≈ –RT ln(s/c°)

(1)

where s denotes solubility, (c◦ = 1 mol dm−3 ) is a standard concentration, and γL stands for the activity coefficient of the ligand, assumed to be close to unity. Enthalpies of solution (Table 2) were determined from the temperature dependence of solubilities using van’t Hoff equation.


16

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As an example, the s(T) dependence in MeCN is displayed in Figure 4. The corresponding results for other examined solvents are given in Supporting Information (Table S2, Figures S21-S25). The solubilities of L are quite similar in all solvents except in MeCN and DMF, where the compound was the most soluble. Unlike ΔsolG°, ΔsolH° (and therefore ΔsolS°) depends on the solvatio n properties of examined solvents considerably. The dissolution is endothermic in all cases, more strongly so in DMSO and EtOH. On the other hand, standard solution entropies are favorable solely in these two solvents. Similar ΔsolG° values in DMSO and EtOH as compared to MeOH and NMF are hence a consequence of entropy-enthalpy compensation. The highest solubility of the compound in DMF is due to the lowest ΔsolH° (and relatively low, negative ΔsolS°). The ligand solvatio n contribution to standard thermodynamic parameters of alkali metal cation complexation will be discussed in detail in the Discussion section.

-5.7

lnKso -5.8

-5.9

-6.0

-6.1 3.2

3.3

103(T / K)1

3.4

3.5

Figure 4. Temperature dependence of L solubility in acetonitrile. ■ experimental; ― calculated by using van’t Hoff equation.


17


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Page 18 of 42

Table 2. The Standard Ligand Solution Parameters in the Investigated Solvents at 25 °C.

Solvent

104 s / mol dm–3

G kJ mol sol

H kJ mol sol

1

1

S J K mol sol 1

MeOH

4.46

19.1



−8.7

EtOH

3.35

19.3

30.9

39.0

NMF

3.39

19.8

11.0

−29.3

DMF

77.6

12.0

10.5

−5.3

DMSO

3.35

18.7

26.0

24.6

MeCN

27.3

14.6

12.4

−7.5

1

DISCUSSION As already stated, a detailed thermodynamic insight into the observed solvent effects on the described complexation reactions can be obtained by calculating the standard thermodyna mic functions of the reactant and product transfers among the investigated solvents. These can be combined into thermodynamic cycles which enable straightforward interpretation of the solvent influence on the cation complexation thermodynamics. In order to simplify data analysis, one solvent is usually chosen as reference, i.e. as solvent from which standard thermodynamic functio ns of transfer to all other solvents are calculated. We have decided to use methanol for this purpose. The t X° (X = G, H, S) of cations from MeOH to any other solvent (e.g. EtOH) were obtained by combining the corresponding functions of transfer from water: Δt X° (M+, MeOH → EtOH) = Δt X° (M+, H2 O → EtOH) – Δt X° (M+, H2 O → MeOH).

The data, based on Ph4 AsPh4 B convention,

(2)

were taken from refs. 46 and 47. Transfer

thermodynamic parameters for ML+ were calculated using the following expression:


18

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ΔrX° (EtOH) – ΔrX° (MeOH) = Δt X° (ML+, MeOH → EtOH)

(3)

–

Δt X° (M+, MeOH → EtOH)

–

Δt X° (L, MeOH → EtOH).

Such analysis resulted, for example, in the thermodynamic cycle shown in Scheme 1, which explains differences in standard reaction Gibbs energies for the binding of Na+ in MeOH and EtOH.

Scheme 1. Thermodynamic Cycle Explaining the Differences in Standard Complexation Gibbs Energies of Sodium Cation in Several Solvents (S) Relative to Methanol (Most of the Data are Given in Attendant Table).

Na+(MeOH)

tG

+

L(MeOH)

(Na )

Na+(S)

Solvent (S)

tG

+

r

G

31.7 kJ mol

a

kJ mol

1

NaL+(MeOH)

(L)

tG

r

L(S)

rG

1

tG

(Na ) kJ mol 1

b

(NaL )

G NaL+(S)

tG

(L ) kJ mol 1

tG

(NaL ) kJ mol 1

EtOH

−32.9

−6.0

0.2

5.0

NMF

−24.9

−15.0

0.7

−7.5

DMF

−19.6

−18.0

−7.1

−12.9

DMSO

−15.0

−21.0

−0.4

−4.6

MeCN

−53.1

7.0

−4.5

−18.9

a Gibbs b Data

energy of the reaction Na+(S) + L(S) → NaL+(S).

taken from references 46,47.


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Page 20 of 42

One can see that the small differences in the stability of NaL+ in MeOH and EtOH can be explained by comparable cation and complex transfer Gibbs energies and by almost equal ligand solubilities in both alcohols. According to Eq. (3) this results in small differences in NaL+ stabilities in these solvents. The same conclusion can be drawn for the complex of compound L with potassium (Schemes 2, S3 and S4, Supporting Information).

Scheme 2. Thermodynamic Cycle Explaining Differences in Standard Complexation Gibbs Energies of Potassium Cation in MeOH and EtOH.

K+(MeOH)

t

G

+

L(MeOH)

6 kJ mol

K+(EtOH)

+

1 t

G

r

G

0.2 kJ mol

L(EtOH)

r

G

19.01 kJ mol

1

KL+(MeOH)

1 t

18.94 kJ mol

G

6 kJ mol

1

1

KL+(EtOH)

Such behavior is due to an interesting relationship between reactant and product transfer enthalp ies and entropies (Schemes S1 and S2, Supporting Information). The energetically disadvantageo us NaL+ transfer from MeOH to EtOH (Δt H° (NaL+, MeOH → EtOH) = 18.3 kJ mol–1 is compensated by an almost equally unfavorable transfer of the ligand Δt H° (L, MeOH → EtOH) = 14.4 kJ mol–1 ). In contrast, differences in cation solvation contribute to overall differences in complexatio n energetics negligibly. The more exothermic solvation of L in MeOH than in EtOH can be explained, at least partly, by stronger intermolecular solvent-ligand hydrogen bonds in the former solvent. The energetically less favorable solvation of the complex in EtOH is more difficult to account for if small

differences

in

enthalpies

of the free cation

solvation


are taken into

account

20

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(Δt H° (Na+, MeOH → EtOH) = 1.3 kJ mol–1 ).46 The opposite trend in standard transfer entropies of the ligand and complex (Δt S° (NaL+, MeOH → EtOH) = 45.6 J K –1 mol–1 , Δt S° (L, MeOH → EtOH) = 47.7 J K –1 mol–1 ) on the one hand, and the free cation on the other (Δt S° (Na+, MeOH → EtOH = –15.8 J K –1 mol–1 )46 suggests that the Na+ transfer from EtOH to MeOH disrupts the more ordered structure of the latter solvent, whereas the opposite seems to hold for the ligand and the sodium complex. As will be evident from a comparative study of the complex affinity for the inclusion of the investigated solvent molecules into the calixarene hydrophobic cavity (see below), the enthalpically favorable and entropically unfavorable transfer of the complex from EtOH to MeOH is, at least partly, consequence of a more stable adduct with MeOH in comparison to that with EtOH. Contribution of standard transfer enthalpies and entropies of K + and KL+ complex to the differences in complexation Gibbs energies in the two examined alcohols is both qualitatively and quantitatively similar to that for Na+. Actually, the log K(KL+, MeOH) – log K(KL+, EtOH) = 0.11 is almost equal to that for the corresponding sodium complex (log K (NaL+, MeOH) – log K(NaL+, EtOH) = –0.11). As already mentioned, complexation of Li+ in studied alcohols could not be observed. This small cation is particularly well solvated in methanol in comparison for instance with MeCN (Δt G°(Li+, MeOH → MeCN) = 21 kJ mol–1 ).46 The Gibbs energy of lithium cation transfer from MeOH to EtOH amounts to Δt G°(Li+, MeOH → EtOH) = 7 kJ mol–1 (ref. 46) and is comparable to the corresponding sodium and potassium transfers. Consequently, Li+ is strongly solvated in EtOH as well. Considering that the structure of all alkali metal complexes with the investigated receptor should not differ significantly, the Δt G° for LiL+ is expected to be comparable to those obtained in the case of NaL+ and KL+. Consequently, Gibbs energies of the Li+ and LiL+ transfers in equation


21


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(3) again most likely cancel each other out, leading to a similar, low affinity of compound L for Li+ in both EtOH and MeOH. Complexation of Rb+ and Cs+ in MeOH could not be observed. As stated in the Introductio n, despite their less favorable solvation in comparison with other alkali metal cations, they do not fit well into the binding sites of lower-rim calix[4]derivatives containing carbonyl-donor groups. Due to this fact, their binding with such receptors in MeOH can usually be noticed in the case of more basic tertiary-amide-based compounds.48 The thermodynamic

cycle explaining

the differences

in calixarene

derivative

L

complexation properties in MeOH and NMF is shown in Scheme 1. Decrease in NaL+ stability in NMF is obviously due to the differences in the cation and complex transfer Gibbs energies (the Δt G° for ligand transfer is very low). The solvation of both charged species is more favorable in NMF, however notably more so in the case of Na +. This in turn leads to a lower stability of the complex in NMF. The transfer of Na+ from MeOH to NMF is favorable both in terms of enthalpy and entropy (Δt H°(Na+, MeOH → NMF) = –9 kJ mol–1 ; Δt S°( Na+, MeOH → NMF) = 21 J K –1 mol–1 , Schemes S5 and S6, Supporting Information).46,47 This can be explained by a combination of the high dipole moment of NMF (3.86 D in comparison to 1.86 D for MeOH) and its more ordered hydrogen-bonding network in comparison to alcohol. The introduction of charged Na + therefore disrupts the NMF structure more strongly than it does in MeOH, leading to favorable Δt S° from this amide to MeOH. On the other hand, the more pronounced association of NMF molecules should enthalpically favor the solvation in MeOH. Consequently, the exothermic transfer of sodium from MeOH to NMF must be due to stronger interactions between the cation and NMF molecules, which, by large, be consequence of the notable difference in solvent dipole moments. The favorable transfer of NaL+ from MeOH to NMF


22

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(Δt H°(NaL+, MeOH → NMF) = –16.1 kJ mol–1 , Schemes S5 and S6, Supporting information) can be explained in a similar fashion as in the case of free cation. The NMF also solvates the macrocycle more favorably from the enthalpic point of view (Δt H°(L, MeOH → NMF) = –5.5 k J mol–1 ). The described combination of transfer enthalpies eventually leads to similar complexation enthalp ies with sodium in MeOH and NMF. However, standard entropies of complex transfer from methano l to N-methylformamide is considerably negative, hence favoring complexation in methanol. The entropies of reactant transfer are comparable in magnitude but opposite in sign, and therefore cancel each other out in Eq. (3). The low affinity of L towards the rest of alkali metal cations in NMF as compared to MeOH can be rationalized

by

much

more

favorable

free

cation

solvation

in

NMF

(Δt G°(Li+, MeOH → NMF) = –24 kJ mol–1 ; Δt G°(K+, MeOH → NMF) = –16 kJ mol–1 ; Δt G°(Rb+, MeOH → NMF) = –18 kJ mol–1 ; Δt G°(Cs+, MeOH → NMF) = –22 kJ mol–1 ).46 As in the case of NaL+, the expected, more favorable solvation of complexes in NMF (Scheme 1), cannot compensate the thermodynamically advantageous solvation of free alkali metal cations in amide with respect to alcohol. The solubility of receptor L is an order of magnitude higher in DMF than in other examined solvents (Table 2). This makes its transfer from this solvent to all other media unfavorab le. Consequently, the Gibbs energies of cation, ligand, and complex transfer have to be taken into account when explaining the differences in the stability constants determined in DMF and the rest of investigated solvents. The corresponding thermodynamic cycle describing the complexation of Na+ with ligand L in DMF and MeOH is shown in Scheme 1. As can be seen, the considerably lower affinity of the receptor for sodium in DMF is due to a combination of a strongly favorable cation and somewhat less favorable calixarene transfer from methanol to DMF. These two contributio ns


23


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Page 24 of 42

would lower the NaL+ stability constant by approximately five orders of magnitude were it not for the exergonic transfer of the complex to DMF. Once the entropic and enthalpic contributions to Δt G° of reactants and the product are examined (Schemes S7 and S8, Supporting Information) it can be again noticed that the exothermic complex, ligand, and cations transfer from methanol to DMF lead to similar complexation enthalpies in methanol and this amide (Scheme S7, Supporting Information). The situation is somewhat different in the case of Δt S° and consequently ΔrS°. The entropy of complex transfer is strongly negative, (Δt S°(NaL+, MeOH → DMF) = –25.5 J K –1 mol–1 , Scheme S8, Supporting Information) favoring the complexation in methanol. This disadvantageous entropic contribution to ΔrG° in amide is further enhanced by considerably favorable entropy of the free cation transfer (Δt S°(Na+, MeOH → DMF) = 20 J K –1 mol–1 , Scheme S8, Supporting Information)46,47 and far less so by the entropic contribution of the ligand transfer (Δ t S°(L, MeOH → DMF) = 3.4 J K –1 mol–1 , Scheme S8, Supporting Information). Complexation of all other alkali metal cations was not observed. As in NMF, this is largely due to

the

more

favorable

cation

solvation

in

comparison

to

methano l.

(Δt G°(Li+, MeOH → DMF) = –14 kJ mol–1 ; Δt G°(K+, MeOH → DMF) = –20 kJ mol–1 ; Δt G°(Rb+, MeOH → DMF) = –20 kJ mol–1 ; Δt G°(Cs+, MeOH → DMF) = –20 kJ mol–1 ).46 This effect, combined with favorable ligand transfer, can account for quite a low stability of the corresponding complexes in DMF. The reason why the receptor binds Na+ in DMF is because of the most favorable ligand-cation interactions realized in the case of this cation. This can be explained by its almost perfect fit into the binding sites of lower-rim calix[4]arene derivatives.

19,21,24,25,33

Comparison of ligand binding affinities for sodium in NMF and DMF is particular ly interesting because of the presence of intermolecular hydrogen bonding in NMF and its absence in


24

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DMF. The corresponding thermodynamic cycle is given in Scheme 3. One can see that the standard Gibbs energies of complex and cation transfers in Eq. (3) almost cancel each other out. Therefore, differences in NaL+ stabilities are mostly due to favorable transfer of the ligand from NMF to DMF. Consequently, the solvent-solvent interactions and the solvent structure have to be taken into account to explain the observed solvent effects. The receptor L can form intermolecular hydrogen bonds with NMF, whereas no such possibility exists in DMF. One would, therefore, expect more favorable solvation of L in NMF.

Scheme 3. Thermodynamic Cycle Explaining Differences in Standard Complexation Gibbs Energies of Sodium Cation in NMF and DMF. r

Na+(NMF)

t

G

+

3 kJ mol

G

+

1

NaL+(NMF)

L(NMF)

1 t

G

7.8 kJ mol

r

Na+(DMF)

24.9 kJ mol

G

1

19.6 kJ mol

L(DMF)

t

G

5.4 kJ mol

1

1

NaL+(DMF)

In order to form such interactions with the ligand, hydrogen bonds between solvent molecules need to be broken. This seems to be enthalpically quite costly as the solution enthalpies of ligand in NMF and DMF are practically identical (Δt H° (L, NMF → DMF) = –0.5 kJ mol–1 , Scheme S9 in Supporting Information). The standard enthalpies of free cation and the complex transfers from NMF to DMF are negative, though not strongly so (Δt H°(Na+, NMF → DMF) = –2.7 kJ mol–1 (refs. 46 and 47), Δt H°(NaL+, NMF → DMF) = –4.6 kJ mol–1 , Scheme S9 in Supporting Information). As a result of such combination of the Δt H° of reactants and the product formed, the complexatio n


25


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Page 26 of 42

enthalpies in the two solvents are highly comparable. As far as differences in complexation entropies in NMF and DMF are concerned, the entropies of free cation and complex transfers among the solvents are quite low (Δt S°(Na+, NMF→ DMF) = 1.1 J K –1 mol–1 (refs. 46 and 47), Δt S°(NaL+, NMF→ DMF) = 2.8 J K –1 mol–1 , Scheme S10 in Supporting Information). Differe nces in calixarene binding affinity in NMF and DMF are therefore by large a consequence of the entropically

disadvantageous

transfer

of

the

receptor

from

NMF

to

DMF

(Δt S°(L, NMF→ DMF) = 23.9 J K –1 mol–1 , Scheme S10 in Supporting Information). Introductio n of bulky calixarene might have induced the ordering of NMF in a similar fashion as the introductio n of hydrophobic species seems to do in water.49,50 In that case one would expect much more pronounced differences in transfer entropy of the complex. However, the inclusion of NMF into the hydrophobic cavity of the complex was observed in chloroform (see below). In contrast, no formation of the corresponding adduct with DMF molecule could be observed. In the former case the solvent inclusion process leads to more favorable complex solvation in NMF. Among the solvents explored, compound L displays the lowest affinity for Na+ in dimethyl sulfoxide. The standard Gibbs energies of Na+, L, and NaL+ transfer from chosen reference solvent (MeOH) to DMSO (Scheme 1) enabled detailed thermodynamic analysis of the differences in complex stabilities in these solvents. A much lower complexation Gibbs energy in DMSO can, by large, be ascribed to strongly exergonic cation transfer from MeOH to DMSO. In fact, dimethy l sulfoxide

is

the

most

efficient

solvent

for

alkali

metal

cations

(Δt G°(Li+, MeOH → DMSO) = –19 kJ mol–1 ; Δt G°(K+, MeOH → DMSO) = –23 kJ mol–1 ; Δt G°(Rb+, MeOH → DMSO) = –20 kJ mol–1 ; Δt G°(Cs+, MeOH → DMSO) = –22 kJ mol–1 ).46 Taking into account the relatively small Δt G° of NaL+ and L transfers from MeOH to DMSO, and the reaction Gibbs energies for Na+, and K + binding in DMSO, the absence of complexation of all


26

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alkali metal cations except sodium in DMSO is understandable. The most interesting differe nce between DMSO and other “good” free-cation solvents (NMF and DMF) is relatively small Δt G° of NaL+ from MeOH to DMSO. The mentioned transfer is still favorable, but much less so than in the case of NMF and DMF. The thermodynamic cycle explaining differences in standard complexatio n enthalpies (Scheme S11, Supporting Information) indicates that the endothermic ligand transfer from MeOH to DMSO (Δt H°(L, MeOH → DMSO) = 9.5 kJ mol–1 ) largely compensates the exothermic transfer of the free cation between these two solvents. The enthalpy of complex transfer is quite low (Δt H° (NaL+, MeOH → DMSO) = –3.3 kJ mol–1 ). As mentioned before, differences in the standard complexation Gibbs energies in MeOH and all other solvents, with the exception of MeCN, are mainly due to the differences in complexation entropies. An appropriate thermodyna mic cycle explaining differences in ΔrS° in MeOH and DMSO is presented in Scheme S12 in Supporting Information. The data reveal that the notable reduction of the calixarene affinity in DMSO with respect to methanol is a consequence of favorable entropy of cation and ligand transfer from MeOH to DMSO (Δt S°(Na+, MeOH → DMSO) = 41.9 J K –1 mol–1 (ref. 46), Δt S°(L, MeOH → DMSO) = 33.4 J K –1 mol–1 ). The most suitable solvent for alkali metal cation complexation is MeCN. The binding of Li+ , Na+, K+, and Rb+ with ligand L was observed in this solvent. The thermodynamic cycle explaining the differences in complexation Gibbs energies for Na + in this and the reference solvent (MeOH) is given in Scheme 1. Ligand solvation is more exergonic in MeCN, whereas the opposite holds for the free cation solvation. Contribution of the reactant transfers to differences in standard complexation Gibbs energies, by large rule, each other out. The sodium cation in MeCN is hence strongly favored because of the exergonic transfer of the complex from methanol to acetonitrile. This also holds for


27


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Page 28 of 42

the corresponding potassium complex (Scheme 4). The cycles explaining differences in ΔrH° and ΔrS° for both cations are given in Schemes S13-S16 in Supporting Information. As far as sodium is concerned, complexation in MeCN is considerably more exothermic which can be accounted for by exothermic transfer of the complex to acetonitrile (Δt H°(NaL+, MeOH → MeCN) = –17.5 kJ mol– 1 , Scheme

S13, Supporting Information) and somewhat less so due to the corresponding endothermic

transfer of the free cation (Δt H°(Na+, MeOH → MeCN) = 7.4 kJ mol–1 .46 Interactions of the solvent molecules

with

the

receptor

are

more

favorable

in

MeCN

than

in

MeOH

(Δt H°(L, MeOH → MeCN) = –4 kJ mol–1 Scheme S13, Supporting Information). Differences in standard complexation entropies in these solvents are quite low because of the low entropies of transfers of the complex, sodium cation, and the ligand. The situation is somewhat different in the case of potassium cation. The complex with this cation is again more stable in MeCN because of its exergonic transfer from methanol to this solvent. However, the K+ solvation is highly similar in both solvents, leading to less pronounced differenc es in standard complexation Gibbs energies as compared to those observed for the sodium cation. The relative contribution of standard transfer enthalpies and entropies to differences in Δ rG° for potassium complexation (Schemes S15 and S16, Supporting Information) is greatly different in comparison to sodium (Schemes S13 and S14, Supporting Information). Transfer of the potassium complex

is

significantly

more

exothermic

than

that

of

the

sodium

complex

(Δt H°(KL+, MeOH → MeCN) = –26.6 kJ mol–1 ). In contrast, the Δt H° of free cation transfer from MeOH to MeCN is negative, which favors its complexation in MeOH. As far as complexatio n entropies are concerned, the entropically disadvantageous KL+ transfer from MeOH to MeCN lowers its stability in the latter solvent. This is partly counterbalanced by a negative entropy of K + transfer from the reference solvent (Δt S°(K +, MeOH → MeCN) = –6.4 J K –1 mol–1 ).46 One can


28

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conclude that differences in the calixarene affinities towards Rb + in acetonitrile and methanol are again consequence of the exergonic complex transfer from alcohol to acetonitrile. N amely, the transfers of Rb+ (Δt G°(Rb+, MeOH → MeCN) = –4 kJ mol–1 )46 and that of the ligand (Δt G°(L, MeOH → MeCN) = –4.5 kJ mol–1 ) from the reference solvent to MeCN are favourable. A notably more favorable solvation of alkali metal cation complexes with other carbonylcontaining lower-rim calix[4]arene derivatives in MeCN as compared to MeOH was previous ly reported.21,24 This was ascribed, at least partly, to the inclusion of the solvent molecule in the macrocycle hydrophobic cavity. Namely, the affinity for MeCN inclusion into the aromatic basket of the free calixarene was found to be much lower as a result of its flattened

cone

conformation.21,22,24,36,51 In contrast, the corresponding cation complexes adopted the square cone conformation, suitable for the solvent molecule inclusion. 21,22,24,32,36,51,52

Scheme 4. Thermodynamic Cycle Explaining Differences in Standard Complexation Gibbs Energies of Potassium Cation in MeOH and MeCN.

r

K+(MeOH)

t

G

+

2 kJ mol

G

+

1

KL+(MeOH)

L(MeOH)

1 t

G

4.5 kJ mol

r

K+(MeCN)

19.0 kJ mol

G

1

28.7 kJ mol

L(MeCN)

t

G

16.2 kJ mol

1

1

KL+(MeCN)

To examine interactions of the investigated calixarene sodium complex with the acetonitrile molecule, calorimetric (Figure S26a, Supporting Information), spectrophotometric (Figure S26b, Supporting Information), and 1 H NMR (Figures 5a and 5b) titrations of this species with MeCN in


29


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CDCl3 were carried out. The binding of MeCN with the NaL+ complex could be clearly observed. By processing the corresponding calorimetric, spectrophotometric, and 1 H NMR titration data, the stability constants of NaLMeCN + adduct in CDCl3 were determined (Table S3, Supporting Information). Apart from that, the standard complexation enthalpy and entropy for the process of solvent inclusion were assessed calorimetrically (Table S3, Supporting Information). The formatio n of NaLMeCN + inclusion complex was entirely enthalpically driven. The interactions realized between the electron-rich aromatic rings and the methyl- group protons of MeCN were quite favorable. On the other hand, the immobilization of MeCN inside the hydrophobic receptor cavity resulted in negative entropy changes. The adduct stability in chloroform was therefore relative ly low. However, in pure acetonitrile the extent of the solvent-molecule inclusion process should be much higher and that could result in almost stoichiometric adduct formation. It should be noted that the use of mixed solvents can influence chemical shifts of the compound protons even without specific solvent-solute interactions. However, the shape of the 1 H NMR titration curves (Figures 5, S27, and S28, Supporting Information), as well as the fact that the data were quite satisfactorily processed by means of the model corresponding to the formation of 1:1 calixarene-solvent molecule adduct, suggests that such influence is not significant under the conditions used. Moreover, the values of equilibrium constant for the reaction of inclusion of the MeCN molecule in the NaL+ complexes in chloroform determined by three different methods (spectrophotometry, NMR, and calorimetry) are in excellent agreement (Table S3, Supporting Information), which would not be the case if the “nonspecific” solvent effect on the NMR spectra would be significant.


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(a)

n(MeCN) / n(NaL+) 160 80 30

d

16 a

a

10 6

c2

 (tert-butyl-H)/ppm

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


(b) 1.20

1.18

1.16

4 2

1.14 a

c2

d

0

1.12 7.20 7.16 7.12 7.08 7.04 3.44 3.40 3.36 3.32

1.20 1.16 1.12 1.08

0

30

60

90

120

150

180

n(MeCN) / n(NaL+)

 /ppm

Figure 5. a) 1 H NMR titration of NaL+ (c = 6.02 × 10–3 mol dm–3 , V0 = 0.5 mL) with MeCN (c = 0.594 mol dm–3 ) in CDCl3 ,  = 25.0 C. b) Dependence of chemical shift of tert-butyl protons on n(MeCN) / n(NaL+) ratio. ■ experimental; ― calculated.

The explored calixarene derivative exhibited lower affinity for inclusion of MeOH and NMF. Unfortunately, the high dilution enthalpies of MeOH(CHCl3 ) and NMF(CHCl3 ) prevented calorimetric investigations of the corresponding inclusion processes. Likewise, changes in UV spectra of NaL+(CHCl3 ) solution upon NMF and MeOH addition were quite low. The corresponding binding constants were therefore determined by processing 1 H NMR titration data (Figures S27 and S28, Supporting Information). The affinity of the sodium complex cavity for DMSO and DMF was even lower (Table S3, Supporting Information). Differences in the stability of NaLMeOH+ and NaLEtOH+ were related to the enthalpically favorable and entropically unfavorable transfer of NaL+ from EtOH to MeOH (Schemes S4 and S3, Supporting Information). Likewise, the Δt G° for NaL+ transfer from MeOH to EtOH and from NMF to DMF would probably have been lower were it not for the formation of solvent adducts of different stability.


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CONCLUSION The presented study of the solvation effects on the affinity of the calix[4]arene derivative towards alkali metal cations revealed it to be an efficient binder of these ions in MeCN, somewhat less so in MeOH and EtOH, whereas its complexation affinity in NMF, DMF, and DMSO was found to be much lower. In all investigated solvents the ligand formed the most stable complex with sodium. The highest stability of NaL+ in MeCN was a consequence of the most favorable complexatio n enthalpy. The strongly exothermic coordination of this cation could be ascribed to the compatibility of the receptor binding site and cation sizes.17,19,21,24,25 The situation was somewhat different in the case of smaller lithium cation. The considerable stability of its complex with L in MeCN was a consequence of favorable enthalpy and entropy changes. Such combination of thermodyna mic complexation parameters is typical for Li+ in this solvent and can be rationalized by its very strong solvation.21,22,24,33 The stabilities of KL+(MeCN) and RbL+(MeCN) were expectedly much lower in comparison to Li+ and Na+, and the corresponding complexation reactions were complete ly enthalpically driven. Exothermic transfer of the K+ and Na+ complexes from methanol to acetonitrile contributed most strongly to differences in their stabilities in these solvents. This was in agreement with the previously reported studies,21 and could be explained, at least partly, by the enthalpica lly favorable inclusion of MeCN molecules inside the hydrophobic cavity of the calixarene complexes. The remarkably low ligand affinity for Li+ in MeOH compared to MeCN was consequence of its stronger solvation in this alcohol (Δt G°(Li+, MeOH → MeCN) = 21 kJ mol–1 ).46 The stabilities of the sodium and potassium complexes in ethanol and methanol were comparable owing to enthalpy-entropy compensation. Standard enthalpies and entropies of the ligand and the complex transfers from MeOH to EtOH were considerably large. However, their relative contributions to the corresponding transfer Gibbs energies were similar in magnitude but


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opposite in sign, so they largely ruled each other out. This could be accounted for by considering differences in the extent and strength of hydrogen-bonding in the investigated alcohols and the formation of more stable complex-solvent adducts with methanol. Much less favorable interactio ns of Rb+ and Cs+ with the receptor binding site than those realized with smaller cations and weaker solvation of alkali metal cation complexes in the investigated alcohols in comparison to MeCN led to their quite low stability in MeOH and EtOH. The moderate affinity of L for sodium cation in NMF, DMF, and DMSO was a result of particularly favorable solvation of free cations in these solvents. Namely, the NaL+ transfer from methanol to NMF, DMF, and DMSO was less exergonic than of the corresponding Na+ transfer, particularly in the case of DMSO. The opposite was true for the transfer of complexes from acetonitrile to any of these solvents, despite the fact that free cations are considerably less strongly solvated in MeCN. This was due to the already mentioned inclusion of the solvent molecule in the calixarene complex basket, which was definitely most pronounced in the case of acetonitrile. The very low affinity of the receptor for all other investigated cations in amides and dimethyl sulfoxide was by large the result of less favorable cation-ligand interactions in comparison with sodium. As already mentioned, the best fit of this cation into the binding site of calix[4]are ne derivative could sufficiently compensate the far more favorable free cation solvation in these solvents in comparison to MeCN and the alcohols. From the above considerations, it is obvious that the solvation effect on the binding affinity of the explored calix[4]arene derivative towards alkali metal cations is indeed remarkable. Comprehensive comparative analysis of the herein reported thermodynamic data enabled us to gain a very deep insight into the factors governing this effect in solvents with quite different properties. As a final conclusion, we believe that the results presented in this work can help in understandin g,


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and even anticipating, the solvent influence on the complexation equilibria in general, particular ly those involving macrocyclic compounds.

Supporting Information Additional results of spectrophotometric, microcalorimetric, and NMR titrations. Temperature dependence of ligand L solubilities in the investigated solvents and the corresponding van’t Hoff plots. Thermodynamic cycles for complexation reactions in different solvents expressed in terms of enthalpies and entropies.

ACKNOWLEDGEMENTS This work has been fully supported by Croatian Science Foundation under the project IP-2014-097309 (SupraCAR).


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Table of Contents Graphic


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Solvation Effect on Complexation of Alkali Metal Cations by a Calix[4

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