J Solution Chem (2010) 39: 835–848DOI 10.1007/s10953-010-9550-9

Thermodynamics of Complexation of Alkali MetalCations by a Lower-Rim Calix[4]arene Amino AcidDerivative

Josip Požar · Tajana Preocanin · Leo Frkanec ·Vladislav Tomišic

Received: 17 November 2009 / Accepted: 26 January 2010 / Published online: 15 June 2010© Springer Science+Business Media, LLC 2010

Abstract Complexation of alkali metal cations with 5,11,17,23-tetra-tert-butyl-26,28,25,27-tetrakis(O-methyl-D-α-phenylglycylcarbonylmethoxy)calix[4]arene (L) in methanoland acetonitrile was studied by means of direct and competitive microcalorimetric titra-tions at 25 °C. The thermodynamic parameters of complexation reactions showed that allthe reactions investigated were enthalpically controlled. In both solvents the reaction en-thalpy was most favorable for Na+ binding with L leading to the highest affinity of the ex-amined calix[4]arene derivative towards this cation. The solubilities (and consequently thesolution Gibbs energies) of the ligand were determined, as were the corresponding solutionenthalpies and entropies. No significant difference was observed between the solution ther-modynamic quantities of L in the two solvents, whereas the transfer of complex species frommethanol to acetonitrile was found to be quite favorable. The interactions of solvent mole-cules with the free and the complexed ligand were investigated by 1H NMR spectroscopy. Itwas concluded that in both cases inclusion of an acetonitrile molecule into the hydrophobiccavity of L occurred, which significantly affected the cation complexation in this solvent.The thermodynamic data were discussed regarding the structural properties of the ligand, thefree and the complexed cations as well as the solvation abilities of the solvents examined.In this respect, the specific solvent-solute interactions and the intramolecular NH · · ·O=Chydrogen bonds at the lower rim of L were particularly addressed.

Keywords Calixarenes · Complexation · Microcalorimetry · Thermodynamics ·Solvation · Hydrogen bonds

Electronic supplementary material The online version of this article(doi:10.1007/s10953-010-9550-9) contains supplementary material, which is available to authorizedusers.

J. Požar · T. Preocanin · V. Tomišic (�)Laboratory of Physical Chemistry, Department of Chemistry, Faculty of Science, University of Zagreb,Horvatovac 102a, 10000 Zagreb, Croatiae-mail: vtomisic@chem.pmf.hr

L. FrkanecLaboratory of Supramolecular and Nucleoside Chemistry, Department of Organic Chemistryand Biochemistry, Ruder Boškovic Institute, Zagreb, Croatia

836 J Solution Chem (2010) 39: 835–848

Fig. 1 Structure of L

1 Introduction

Over the past several decades calixarenes have been recognized as suitable parent com-pounds for the synthesis of a wide variety of ionophores and molecular receptors [1–3].Their main advantages as host templates are the relatively simple synthesis and the possi-bility of both lower- and upper-rim functionalization. However, the targetted preparation ofcalixarene derivatives, which could selectively bind ions or neutral molecules, presupposes adetailed thermodynamic understanding of all the factors contributing to the hosting process.Knowledge of thermodynamic state functions of complexation reactions, i.e. �rG

◦, �rH◦

and �rS◦, and the accurate thermodynamic information concerning solutes solvation (host,

guest and complex), are required for explaining the ligand selectivity for particular species aswell as for understanding the solvent effects involved [4–7]. Consequently, this knowledgesignificantly contributes to a better comprehension of the structure-reactivity relationships,which can be of great importance in targeted molecular design.

The calix[4]arene investigated in the present paper, namely 5,11,17,23-tetra-tert-butyl-26,28,25,27-tetrakis(O-methyl-D-α-phenylglycylcarbonylmethoxy)calix[4]arene (L,Fig. 1) [8], belongs to a class of calixarene derivatives known as peptidocalixarenes [9]. Thederivatives of this kind, carrying amino acid or peptide substituents at the lower or uppercalixarene rim, are used as receptors for cations, anions and neutral guests, and have beenrecognized as particularly interesting biomimetic compounds [10].

The N -linked amino acid calix[4]arene derivative L with secondary-amide-containingsubstituents at the lower rim was shown to be an effective receptor for some alkali metalcations [8, 11]. As for many other calixarene derivatives, its hosting properties are enhancedby the presence of bulky tert-butyl groups at the upper rim which stabilize the cone con-formation of the macrocycle thus making it suitable for cation binding. However, com-pound L was found to have a lower affinity towards alkali metal cations as compared totertiary amide derivatives. This was ascribed to the presence of intramolecular NH · · ·O=Chydrogen-bonding in L. A remarkable solvent effect on the cation complexation equilibriawas observed, the stability constants of ML+ (M stands for alkali metal) complexes beingconsiderably higher in acetonitrile than in methanol.

This work is focused on a microcalorimetric study of complexation of alkali metal cationsby L in methanol and acetonitrile at 25 °C. The obtained standard reaction Gibbs energies,enthalpies and entropies, as well as solution thermodynamic and transfer parameters, have

J Solution Chem (2010) 39: 835–848 837

provided a deeper insight into the complexation processes. The medium effect on the re-actions studied is particularly addressed with emphasis on the specific solvent-ligand andsolvent-complex interactions.

2 Experimental

2.1 Materials

Compound L was prepared according to the procedure described elsewhere [8]. The sol-vents, methanol, MeOH (Merck, Uvasol) and acetonitrile, MeCN (Merck, Uvasol; Riedel-de Haen, spectranal) were used without further purification. The salts for the calorimetrictitrations were perchlorates (LiClO4, NaClO4, KClO4, Merck, p.a.) and nitrates (RbNO3,Merck, p.a.; CsNO3, Merck, puriss.). Because of the low solubility of KClO4 in methanol,KCl (Merck, p.a.) was used in the experiments conducted in this solvent.

2.2 Solubility Measurements

Saturated solutions of L were prepared by adding an excess amount of the solid substanceto the solvent. The obtained mixtures were left in a thermostat at 25 °C for several daysin order to equilibrate. Aliquots of solutions were then taken for the determination of thesolubility. The concentrations of saturated solutions of L in MeCN and MeOH at 25.0 °Cwere determined spectrophotometrically by means of a Varian Cary 5 spectrophotometerequipped with a thermostatting device. Calibration curves were obtained by measuring theabsorbances of L solutions of known concentrations.

Solvate formation was tested by placing a known amount of L (m = 40 mg) in a sealeddesiccator where it was exposed, for several weeks, to a saturated atmosphere of acetoni-trile or methanol. Solvent uptake was then checked by weighing the samples. No solvateformation was observed.

2.3 Calorimetric Measurements

Microcalorimetric experiments were performed by means of an isothermal titrationcalorimeter CSC 4200 ITC; Calorimetry Sciences Corporation, at 25.0 °C. The calorimeterwas calibrated electrically and chemically by means of the standard reaction of protonationof THAM(aq) with HCl(aq) [12]. The reliability of the microcalorimeter was additionallychecked by carrying out the complexation of barium(II) by 18-crown-6 in aqueous mediumat 25 °C. The results obtained (log10 K = 3.77, �rH = −31.0 kJ·mol−1) were in goodagreement with the literature values (log10 K = 3.77, �rH = −31.4 kJ·mol−1 [13]). Calori-metric data were processed using the Titration Bindworks and OriginPro 7.5 programs.

2.3.1 Calorimetric Titrations

The calorimeter reaction cell was filled with L solution (V = 1.3 cm3;c = 1 × 10−3 mol·dm−3) in the appropriate solvent. The enthalpy changes were recordedupon stepwise automatic additions (5 min intervals) of alkali-metal salt solution (c =1 × 10−2 mol·dm−3 or 4 × 10−2 mol·dm−3) from a 250 µL Hamilton syringe. Blank experi-ments were performed in order to make corrections for the enthalpy changes correspondingto the dilution of the alkali metal salt solution in the pure solvent.

838 J Solution Chem (2010) 39: 835–848

In the case of the Li+ complexation with L in acetonitrile the corresponding equilib-rium constant was higher than 106, so competitive calorimetric titrations were performedwhere the lithium ion in the complex was replaced by the sodium ion. A solution ofNaClO4 (c = 1 × 10−2 mol·dm−3) was placed in the syringe and added into a solu-tion containing the lithium complex with the macrocycle (c(L) = 1 × 10−3 mol·dm−3,c(Li+) = 1 × 10−2 mol·dm−3). The obtained enthalpy changes were corrected for the heatof dilution of the NaClO4 solution.

Mathematical deconvolution was used in both direct and competitive complexation ex-periments to shorten the period needed to attain thermal equilibrium between additions. Allmeasurements were done in triplicate.

2.3.2 Solution Enthalpy Determinations

Enthalpies of solution of L in both solvents were determined calorimetrically. The samplecell was filled with a suspension of L(s) in a saturated solution of the ligand in MeCN orMeOH. The pure solvent was then added from a syringe and the corresponding enthalpychanges were measured. The amount of L dissolved was calculated from the solubility ofthe ligand in the appropriate solvent. The solutions of (uncharged) ligand were assumed tobehave almost ideally, and consequently, the determined solution enthalpies were supposedto be close to their standard values.

2.4 1H NMR Measurements

NMR spectra were taken on a Bruker AV600 spectrometer (δ in ppm relative to Me4Si as aninternal standard).

3 Results and Discussion

As an example, a thermogram obtained by titration of L with NaClO4 in methanol at 25 °C isshown in Fig. 2a. The stepwise addition of NaClO4 resulted in exothermic enthalpy changes.The standard reaction enthalpy and the equilibrium constant (hence the standard reactionGibbs energy) for the complexation of Na+ with the macrocycle were calculated by a least-squares non-linear regression analysis of calorimetric titration data (Fig. 2b). The standardreaction entropy was calculated from the reaction enthalpy and Gibbs energy (Eq. 1):

�rG◦ = �rH

◦ − T �rS◦

(1)

The obtained thermodynamic parameters, as well as those corresponding to reaction of K+with L in methanol (determined following the same procedure), are given in Table 1.

Addition of LiClO4, RbNO3 or CsNO3 into a methanolic calixarene solution did not re-sult in measurable enthalpy changes, indicating that no observable complexation took place,or that the value of the complexation enthalpy was close to zero. However, the latter couldprobably be ruled out as previous conductometric and spectrophotometric experiments [11]suggested that the above mentioned cations did not form complexes with L in methanol toany significant extent under similar conditions.

Calorimetric titrations of L with NaClO4 and LiClO4 in acetonitrile provided informa-tion about the enthalpies of complexation reactions. However, the corresponding equilib-rium constants were too high to be determined by direct calorimetric titrations. As already

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Fig. 2 (a) Microcalorimetrictitration of L(c = 1.01 × 10−3 mol·dm−3),V = 1.3 mL with NaClO4(c = 9.97 × 10−3 mol·dm−3) inmethanol; t = 25 °C;(b) Dependence of successiveenthalpy change on volume ofNaClO4 solution.� experimental; – calculated

Table 1 Thermodynamic parameters for complexation of alkali metal cations with L in methanol at 25 °C

Cation log10K± SE (�rG◦/kJ·mol−1) ± SE (�rH

◦/kJ·mol−1) ± SE (�rS◦/J·K−1·mol−1) ± SE

Li+ –a –a –a –a

Na+ 4.29 ±0.04 −24.5 ± 0.2 −43.4 ± 0.1 −64 ± 1

K+ 2.39 ±0.02 −13.64 ± 0.09 −35 ± 1 −73 ± 4

Rb+ –a –a –a –a

Cs+ –a –a –a –a

aNo complexation was observed

SE = standard error of the mean (N = 3)

mentioned, in order to determine the stability constant of the LiL+ complex, Li+/Na+ dis-placement experiments were conducted. The NaL+ stability constant was determined previ-ously by direct potentiometry [11], and was used in processing the competitive calorimetric

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Table 2 Thermodynamic parameters for complexation of alkali metal cations with L in acetonitrile at 25 °C

Cation log10K ± SE (�rG◦/kJ·mol−1) ± SE (�rH

◦/kJ·mol−1) ± SE (�rS◦/J·K−1·mol−1) ± SE

Li+ 6.47 ± 0.03 −36.9 ± 0.2 −27.4 ± 0.8 32 ±3

Na+ 7.66a −43.7 −59.7 ± 0.5 −53 ± 2

K+ 4.32 ±0.03 −24.6 ± 0.2 −40.3 ± 0.5 −52 ± 2

Rb+ −b –b –b –b

Cs+ −b –b –b –b

aDetermined by potentiometry [11]

bNo complexation was observed

SE = standard error of the mean (N = 3)

Table 3 Thermodynamic parameters of solution of L in acetonitrile and methanol at 25 °C

Solvent 103·s/mol·dm−3 �solG◦/kJ·mol−1 �solH

◦/kJ·mol−1 �solS◦/J·K−1·mol−1

Acetonitrile 2.72 14.64 8.99 −19.0

Methanol 3.45 14.05 7.51 −21.9

titration data. The results of the described measurements are presented in Table 2, whichalso includes the thermodynamic parameters for the reaction of K+ with L in acetonitrileobtained by direct microcalorimetry.

As was the case in methanol, no measurable heat effects were observed upon the additionof RbNO3 or CsNO3 into acetonitrile L solutions. These findings are in accord with theresults of previous spectrophotometric and conductometric investigations, from which it wasconcluded that the larger Rb+ and Cs+ cations did not fit in the investigated calix[4]areneion-binding site formed by four ether and four amide carbonyl oxygen atoms.

It is worth mentioning that the values of stability constants determined in this workare in good agreement with those obtained previously by means of spectrophotometrictitrations [11]. In both cases the possible effect of ion-pairing reactions (M+ + A− andML+ + A−; A− denotes Cl− or ClO−

4 ) was assumed to be negligible.The values of standard thermodynamic functions of solution of L in acetonitrile and

methanol are reported in Table 3. Standard solution Gibbs energy was calculated from thesolubility data using the equation:

�solG◦ = −RT lnK◦ = −RT ln(γLs/c◦) ≈ −RT ln(s/c◦) (2)

where s denotes solubility, c◦ (= 1 mol·dm−3) is a standard concentration, and γL stands forthe activity coefficient of the ligand which is assumed to be close to unity.

Enthalpies of solution (�solH◦) were determined by means of calorimetric experiments

in which the solvent was added into the saturated solution of L containing the suspendedsolid ligand, as described in the Experimental section. As an example, the thermogram ofdissolution of L in methanol is displayed in Fig. 3. In both solvents the ligand dissolutionreaction was found to be endothermic. Standard solution entropies (�solS

◦) were calculatedfrom �solG

◦ and �solH◦ values.

The data presented in Tables 1 and 2 indicate that all the investigated complexation reac-tions are enthalpy driven. The enthalphic contribution to the reaction Gibbs energy is most

J Solution Chem (2010) 39: 835–848 841

Fig. 3 Calorimetric curve for thedissolution of L in methanol;0.2 mL of solvent added into thesaturated solution of L containingsuspended solid ligand at 25 °C

favorable for Na+ binding with L in both solvents leading to the highest affinity of the ex-amined calix[4]arene derivative towards this cation. The reaction entropy is positive only inthe case of the Li+ complexation with L in acetonitrile, for which the −�rH

◦ value is alsothe lowest. This could be explained, at least partly, by a stronger solvation of the Li+ cationcompared to the other alkali metal ions. Like sodium, the lithium cation fits well into thecalix[4]arene ion-binding site, so its binding to the oxygen atoms of the ligand should beenergetically quite favorable. However, in order to undergo complexation the cation mustrelease directly and indirectly bound solvent molecules, and this process is enthalphicallymost demanding for the smallest and strongest most strongly solvated Li+ cation. On theother hand, desolvation results in a gain in entropy which is greater if the cation is smallerbecause of the more pronounced structuring of the surrounding solvent. This entropy gainseems to be a decisive factor in determining the sign of the reaction entropy in the case ofLiL+ formation. Thus despite the less favorable reaction enthalpy, it leads to a considerablyhigher stability constant for the LiL+ than for the KL+ complex in acetonitrile. The situa-tion is quite opposite when the reaction is carried out in methanol, i.e. although L stronglybinds Li+ in acetonitrile, in methanol no complexation can be observed. This differencecan serve as an example which clearly emphasizes the influence of the free cation solvationon the equilibrium of the complexation reaction. As a hydrogen bonding solvent, methanolstrongly solvates small cations [14], thus making the substitution of its molecules by lig-and binding sites thermodynamically unfavorable. The entropy gained by the desolvation ofLi+(MeOH) should be greater than that of Li+(MeCN) but is still insufficient to compensatethe energy required for the overall reaction.

It is interesting to note that from the thermodynamic point of view there is no significantdifference between the solvation of compound L in methanol and that in acetonitrile. Thesolubilities and derived solution Gibbs energies, solution enthalpies and entropies of the lig-and are similar for the two solvents (Table 3). As a consequence, the values of correspondingthermodynamic functions of transfer are rather low (Table 4). This means that the consider-able difference in ML+ stability constants in MeOH and MeCN (Tables 1 and 2) cannot beexplained by the difference in solvation of L, although that was previously suggested to bea very important factor to consider [11].

Even though the free cation solvation plays an important role in the complexationprocess, this phenomenon cannot completely account for the observed difference in complex

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Table 4 Thermodynamic functions of transfer of L and its complexes with K+ and Na+ from methanol toacetonitrile at 25 °C

Species �tG◦/kJ·mol−1 �tH

◦/kJ·mol−1 �tS◦/J·K−1·mol−1

L 0.6 1.5 3.0

NaL+ −13.2 −7.2 20.1

KL+ −12.5 −7.5 16.8

stabilities in the solvents under consideration (e.g. K+ ion is even slightly better solvated inacetonitrile than in methanol, the Gibbs energy of transfer of this cation from MeOH toMeCN being −2.1 kJ·mol−1 [15]). At least in some cases the decisive factor should there-fore be the solvation of the complex species formed. Transfer thermodynamic parameters(�tX(M+L)) that support this conclusion can be calculated by the following equation:

�rX(MeCN) − �rX(MeOH) = �tX(M+L,MeOH → MeCN)

− �tX(M+,MeOH → MeCN)

− �tX(L,MeOH → MeCN) (3)

where index r denotes a complexation reaction and X stands for enthalpy, entropy or Gibbsenergy.

The �tG◦(M+, MeOH → MeCN) and �tH

◦(M+, MeOH → MeCN) values for Na+and K+ were calculated by combining the corresponding functions of transfer of cationsfrom water to methanol or acetonitrile:

�tX(M+,MeOH → MeCN) = �tX(M+,H2O → MeCN) − �tX(M+,H2O → MeOH)

(4)

These data, based on Ph4AsPh4B convention, were taken from [15].The transfer Gibbs energies and enthalpies of NaL+ and KL+ complexes obtained by

means of Eq. 3, as well as entropies of transfer computed from �tG◦ and �tH

◦ values, arepresented in Table 4 and demonstrate that in both cases the transfer of ML+ from methanolto acetonitrile is quite favorable. Such a strong stabilization of the complexes in acetonitrileseems to be the predominant factor in determining the solvent effect on the complexation ofNa+ and K+ with L. It should also be noted that the values of both transfer enthalpies andentropies (and consequently those of Gibbs energies) for NaL+ and KL+ are rather similar.This could suggest that there is no significant difference in the solvation of the two species ina given solvent (the same was proposed on the basis of conductivity data in the case of NaL+and LiL+ complexes in acetonitrile [11]). This finding supports our previous assumption(made without having the necessary thermodynamic data) that the main reason why theNa+/K+ selectivity, i.e. ratio KNaL

+/KKL+, is more than ten times higher in acetonitrile

than in methanol (Tables 1 and 2) lies in the difference in the solvation of free cations inthese solvents.

As already mentioned, there are no significant differences between the solution thermo-dynamic parameters of L in MeOH and MeCN (Table 3). Since MeOH with its proton-donorand proton-acceptor abilities is expected to interact rather strongly with the pendant armsof L via intermolecular hydrogen bonds formation, one would expect these differences tobe to some extent larger. To examine the interactions of the investigated macrocycle and

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Fig. 4 1H NMR spectra of compound L in CDCl3 at two different temperatures

its sodium complex with the solvent molecules in more detail, we have undertaken severalNMR experiments and the obtained results are described below.

The 1H NMR spectrum of L in CDCl3 was reported in [8]. As it consisted of a single setof signals in the temperature range (−20–40) °C, it was concluded that the ligand existed in arigid cone conformation. Seeking the difference between its conformations in uncomplexedand complexed forms, in this work we recorded the CDCl3 spectrum of L at a significantlylower temperature, i.e. −55 °C. The obtained spectrum showed two quite broad signals formost of the protons (Fig. 4), indicating the C2v flattened cone conformation of L. Fromthis finding it can be concluded that at room temperature a fast (at the NMR time scale)C2v � C2v interconversion occurs [16]. This is in accordance with the molecular structureof L in the solid state [8], and is of importance for the binding of the guest molecule in theupper-rim hydrophobic calixarene cavity (see below).

Comparison of 1H NMR spectra of L solutions in CDCl3, CD3CN and MeOD (TableS1 in Electronic supplementary material) reveals that, among the other differences, the sig-nals of the upper-rim aromatic protons are significantly shifted downfield on going fromCDCl3 (δ = 6.66, 6.64) and CD3OD (δ = 6.77, 6.76) to CD3CN (δ = 7.01–6.99). Thisshift can be taken as an indication of interaction of the acetonitrile molecule with thecalix[4]arene upper-rim aromatic rings [17, 18]. Inclusion of MeCN into the hydrophobiccavity of calix[4]arene derivatives both in the solid state [19, 20] and in the solution [17, 18,21–25] has been well documented in the literature. To corroborate such L–MeCN interac-tion, 1H NMR titration of the macrocycle solution in CDCl3 with acetonitrile was carriedout. However, the chemical shift of CH3CN protons in the MeCN/L molar ratio range 0.1–2showed no difference with respect to the corresponding signal of acetonitrile in deuteratedchloroform (δ = 2.0). In addition, no significant change of any of the chemical shifts ofL protons was observed during the titration (Fig. S1 in Electronic supplementary mater-ial). Likewise, in the case of titration with CH3OH, the methyl protons signal was found atδ = 3.49 independently of the MeOH/L molar ratio (Fig. S2 in Electronic supplementarymaterial). The fact that the complexation of acetonitrile with L was not observed in chlo-roform could be accounted for by assuming that in this solvent, under the stoichiometric

844 J Solution Chem (2010) 39: 835–848

conditions applied, the extent of MeCN inclusion was rather low. Consequently, the for-mation of the L–MeCN complex could not be detected. Nevertheless, the above-mentionedchemical shift of the ligand aromatic protons in acetonitrile solution strongly suggests theenthalpically favorable inclusion of the MeCN molecule into the hydrophobic cavity ofL [22]. Another favorable process taking part in acetonitrile solutions is the formation ofcircular intramolecular NH· · ·O=C hydrogen bonds. The presence of these bonds in the cal-ixarene derivative studied has been undoubtedly confirmed by Frkanec et al. [8], and can bededuced from the relatively high chemical shift of the amide NH protons which is approx-imately the same in CDCl3 and CD3CN solutions (Table S1 in Electronic supplementarymaterial). Contrary to the latter solvents, in methanol, which is a strong hydrogen bondingsolvent (and thus competes for NH and C=O amide sites), the formation of intramolecularH-bonds is expected to occur to a much lesser extent.

The 1H NMR titration of the NaL+ solution in CDCl3 with acetonitrile clearly indi-cated the binding of MeCN with the complex (Fig. S3 in Electronic supplementary mate-rial). At the lowest MeCN/L ratio examined, the signal of acetonitrile protons appeared atδ = 1.81 ppm (0.19 ppm upfield shift), and then gradually shifted towards the position ofthe MeCN signal in CDCl3. Besides, the signals corresponding to the upper-rim aromaticand tert-butyl protons exhibited downfield shifts with increasing MeCN concentration. Theobvious higher affinity of the complex species toward MeCN compared to the free ligandcan be explained by taking into account the conformational difference, i.e. the complex withC4 cone conformation, which is stabilized by binding of the cation, is better preorganized toaccept the acetonitrile molecule into its hydrophobic cavity [21, 26].

As suggested by Danil de Namor at al. [17, 27, 28], the interaction of acetonitrile withboth the free and the complexed ligand can (at least partly) account for the complexation andsolution thermodynamic parameters determined in this work. By considering the inclusionof the MeCN molecule into the hydrophobic cavity of L, one can partly explain the smalldifference between the L solubilities in the examined solvents, as well as between the relatedthermodynamic solution quantities (Table 3) corresponding to these solvents. On the otherhand, the stronger binding of MeCN by the complex species could be responsible for thefavorable transfer of NaL+ and KL+ complexes from MeOH to MeCN, and this is in turnreflected in the higher stabilities of the complexes in the latter solvent relative to the formerone.

As stated in the Introduction, the stability constants of ML+ species in both MeOH andMeCN are considerably lower than those corresponding to the complexes of alkali metalcations with tertiary amide calix[4]arene derivatives [11, 29–31]. Since we have alreadyestablished that the herein examined complexation reactions are enthalpy controlled, weshall make an attempt to explain the lower affinity of the secondary amide derivative Ltowards M+ ions by considering the available complexation reaction enthalpies. As canbee seen in Table 5, the −�rH

◦ values for the reactions of Li+, Na+ and K+ with p-tert-butylcalix[4]arene tetrakis(diethylamide) [29] are appreciably higher than those given herein Tables 1 and 2.

There are several factors which could be responsible for this observation: the hydrophiliccharacter of the binding cavity, i.e. basicity of the carbonyl oxygen atoms [30], solvation ofthe free and complexed forms of the ligand [4–6], and formation of inter- and intramolecularhydrogen bonds [30, 32]. However, it is interesting to note that a drop in the ethalphic sta-bility of ML+ complexes with respect to the corresponding complexes comprizing tertiaryamide calix[4]arene derivative is substantially larger in MeCN than in MeOH (Table 5). Thatcould be at least partly attributed to the presence of above mentioned circular intramolecularNH· · ·O=C hydrogen bonds in L. These bonds obviously need to be disrupted through an

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Scheme 1 Thermodynamic cycles expressed in terms of Gibbs energies for complexation of Na+ and K+with L in methanol and acetonitrile

enthalpically demanding process to allow a change in orientation of the amide groups intoa position favorable for binding the cation. This effect, which does not play a role in thecase of a tertiary amide derivative (no intermolecular hydrogen bonds present), is expectedto be much more pronounced in acetonitrile than in methanol leading to a larger differencein �rH

◦ values in the former solvent relative to the latter.

4 Conclusion

The results described in this paper clearly show how composite the medium effect on thecomplexation reactions is. Notwithstanding the importance of solvation of all reactants andproduct(s), solvation of some of the particular species involved can play a major role indetermining the complexation equilibrium. As an example that quite strongly supports thisstatement could be the huge difference between the stabilities of the LiL+ complex in MeOHand MeCN (Tables 1 and 2). That can be mostly accounted for by the stronger solvation ofthe Li+ cation in the former solvent. Thus, in this case the decisive factor is presumably thesolvation of the free metal ion.

Contrary to our earlier conclusion [11], here we have established that from the thermody-namic point of view there is no significant difference between the solvation of L in MeOHand that in MeCN. The transfer Gibbs energies given in thermodynamic cycles presented inScheme 1 show that in the case of reaction of Na+ and L, cation solvation favors its com-plexation in acetonitrile as compared to methanol, whereas for K+ there is a slight oppositeeffect. However, the predominant factor leading to the higher stabilities of the complexesNaL+ and KL+ in acetonitrile is the more favorable solvation of the complex species in thissolvent (Scheme 1). That is proposed to be a consequence of the specific solvent-complexinteraction, i.e. the inclusion of acetonitrile molecule into the complexed ligand upper-rimhydrophobic cavity.

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Table 5 Comparison of complexation enthalpies for the reactions of alkali metal cations with L and tertiaryamide derivative L1a in methanol and acetonitrile; t = 25 °C

�rH◦/kJ·mol−1

MeOH MeCN

Na+ K+ Li+ Na+ K+

L1b −50.6 −42.4 −55 −79 −64

L −43.4 −35.5 −27.4 −59.7 −40.3

�(�rH◦)c −7.2 −6.9 −27.6 −19.3 −23.7

aL1, p-tert-butylcalix[4]arene tetrakis(diethylamide)

bData from [29]c�(�rH

◦) = �rH◦(L1) − �rH

◦(L)

Such inclusion is believed to occur also in the case of the free ligand, but to a lowerextent. The reason for higher affinity of the ML+ complex towards MeCN molecules liesin the better preorganization of its upper-rim cavity, i.e. in the fact that in solution it existsin a rigid cone conformation, whereas the free macrocycle L experiences interconversionbetween two flattened cone conformations.

Comparison of the enthalpies of complexation of alkali metal cations by L with thosecorresponding to the reactions of these cations with tertiary amide calix[4]arene derivativeL1 shows that the enthalpic stabilities of ML+ complexes are lower in both acetonitrile andmethanol (Table 5). Moreover, the differences in reaction enthalpies are much larger in thecase of acetonitrile as a solvent. That is suggested to be due to the presence of intramolecularNH· · ·O=C hydrogen-bonding in the herein investigated ligand, which is more likely tooccur in acetonitrile than in methanol.

Acknowledgements This work was supported by the Ministry of Science, Education and Sports of theRepublic of Croatia (projects 119-1191342-2960, 119-1191342-2961 and 098-0982904-2912). The authorswarmly thank Professor Angela F. Danil de Namor for helpful discussions.

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