J. Phys. Chem. 1995,99, 16781-16785
16781
Thermodynamics of Alkyl p-tert-Butylcalix(4)arenetetraethanoates. 2. Solution Studies of Calix(4)arene Esters and Lithium and Sodium Ethylcalix(4)arene Ester Complexes: Coordination Data in the Solid State Angela F. Danil de Namor,*'t Elisabeth Gil,? Margot A. Llosa Tanco,? David A. Pacheco Tanaka,' Lupe E. Pulcha Salazar,? Ronald A. Schulz,$ and Jianji Wangt Laboratory of Thermochemistry, Department of Chemistry, and Department of Chemical and Processing Engineering, University of Surrey, Guildford, Surrey GU2 5XH, U.K. Received: July 17, 1995'
Solubility data for ethyl p-tert-butylcalix(4)arenetetraethanoate and its sodium perchlorate complex in acetonitrile at 298.15 K are used to derive the standard solution Gibbs energies of these species in this solvent at 298.15 K. The thermochemical behavior of alkyl (methyl, ethyl, and n-butyl) p-tert-butylcalix(4)arene esters is discussed on the basis of 'H NMR data in these solvents. Standard enthalpies of solution of the sodium and lithium perchlorate complexes of the tetraethyl ester derivative in acetonitrile and in benzonitrile at 298.15 K are first reported. A considerable change is observed in the enthalpic stability of the free ligand relative to its sodium perchlorate complex, in these solvents. From enthalpy data for the host, the guest, and the resulting complex the differences previously observed in the standard enthalpies of complexation of lithium and sodium with calix(4)arene esters in these solvents are explained. Combination of solution and previously reported complexation data leads for the first time in the area of calixarene chemistry to the calculation of the thermodynamic parameters associated with the coordination process referred to reactants and product in the solid state. The importance of these data is emphasized.
The thermodynamics of complexation of alkyl (methyl, ethyl, n-butyl) p-tert-butylcalix(4)arenetetraethanoates,RCalix(4), and alkali-metal cations, M+, in acetonitrile and in benzonitrile was investigated at 298.15 K, and from the results reported in the preceding paper,' a quantitative assessment (based on stability constant new data) was made on the selective recognition of these ligands for these cations in these solvents. It was also established that (i) complex stability as reflected by the Gibbs energy is generally enthalpy controlled. (ii) Entropy data referred to the process M+(g)
+ RCalix(4)(s) -M+RCalix(4)(s)
(1)
suggest appreciable differences between the solvation of the free ligand and its metal-ion complex: M+RCalix(4). So far, the thermodynamic interest in the area of macrocyclic chemistry in general and calixarenes in particular tends to end with the complexation process. This statement is corroborated by the lack of solution thermodynamic studies on calixarene derivatives (nonelectrolytes) and their metal-ion complexes (electrolytes). Undoubtedly, ion-solvent and ligand-solvent interactions play a crucial role in complexation processes involving macrocyclic ligands and metal cations in solution. As far as calix(4)arene derivatives and their metal-ion complexes are concerned, investigations on their solution properties are particularly relevant since, unlike cryptands or crowns, these macrocycles possess an additional hydrophobic region able to interact with neutral species. Therefore, for the purpose of useful interpretation of the complexation process in a given medium, solution studies of these ligands and their compounds should be carried out in this medium. In this paper thermodynamic aspects of the solution process involving alkyl p-tert-butylcalix(4)arenetetraethanoates in acetonitrile and in
' Department of Chemistry. @
Department of Chemical and Processing Engineering. Abstract published in Advance ACS Abstracts, October 15, 1995.
0022-365419512099-16781$09.0010
benzonitrile at 298.15 K are discussed, and 'H Nh4R data in these media are reported. On the basis of the stability of the metal-ion complexes,' the lithium and sodium perchlorates of ethyl p-tert-butylcalix(4)arenetetraethanoatehave been isolated and thermodynamically characterized in these solvents. The main aims of this paper are (a) to gain further understanding of the complexation of calix(4)arene esters and metal cations in acetonitrile and in benzonitrile from interpretation of the solution thermodynamics of the host, the guest, and the resulting complex; (b) to derive for the first time the thermodynamics associated with the coordination process (referred to reactants and product in their solid state) from complexation and solution data.
Experimental Section Chemicals. Acetonitrile and benzonitrile (both from Aldrich) were purified as described elsewhere.' Lithium and sodium complexes of ethyl p-tert-butylcalix(4)arene tetraethanoates were prepared by the addition of the calixarene ester to a nonaqueous solution containing the appropriate lithium or sodium (in excess) perchlorate. The mixture was heated at 40 "C during 30 min, after which distilled water was added until complete precipitation of the white solid occurred. It was then filtered. The solid was washed several times with water, and then it was recrystallized from absolute ethanol. Microanalysis was carried out at the University of Surrey. For the sodium perchlorate complex of EtCalix(4), found percentages (C, 64.23; H, 7.27) are in good agreement with calculated values (C, 64.56; H, 7.23). Similar agreement was found for the lithium perchlorate complex of EtCalix(4) (calculated percent: C, 65.23; H, 7.33; experimental percent: C, 65.29; H, 7.23). Solubility Measurements. Saturated solutions of the ligand or its complexes were prepared by adding an excess amount of the solid to the solvent. The mixtures were left for several days in a thermostat at 298.15 f 0.01 K until equilibrium was 0 1995 American Chemical Society
16782 J. Phys. Chem., Vol. 99, No. 45, 1995 reached. Aliquots of the saturated solution were taken and analyzed gravimetrically. Analyses were performed in triplicate. Blank experiments were carried out in all cases. Solvate formation was tested by placing small quantities of the solid in open vessels over the appropriate solvent placed at the bottom of a closed desiccator.2 In this way a saturated atmosphere of the solvent was ensured. Uptake of the solvent was checked by weighing the samples from time to time. Thermochemical Measurements. Enthalpies of solution at 298.15 K were measured with the Tronac 550 calorimeter designed by Izatt and Chri~tensen.~The reliability of the equipment was checked by measuring the enthalpy of solution of THAM (tris(hydroximethy1)aminomethane) in an aqueous solution of hydrochloric acid (0.1 M).4 A value of 29.75 f 0.03 kJ mol-' was obtained, which is in close agreement with that reported in the literature5 for this reaction using the same calorimetric equipment. Enthalpies of solution were measured by breaking glass ampules containing different amounts of the appropriate solute (ligand or metal-ion complex) in 50 cm3 of the solvent. The total heats observed were corrected for the heat associated with the breaking of empty ampules in the appropriate solvent. Standard enthalpies of solution are the value at c = 0 from a plot of AH versus c"*, where c is the final concentration of the compound in the appropriate solvent. 'H NMR Measurements. 'H NMR measurements were carried out with an AC-300 NMR spectrometer at a spectral frequency (SF) of 300.135 MHz, a spectral width (SW)-of 4504 ppm, a pulse width of 2.0, and delay time of 1.60 s; an acquisition time (AQ) of 1.819 s and a line broadening (LB) of 0.55 were applied. The appropriate solutions (calix(4)arene esters) were placed in 5 mm NMR tubes using TMS as the internal reference to measure the spectrum of the ligand
Results and Discussion
(a) Solution Thermodynamics of Calix(4)arene Esters and Their Complexes at 298.15 K. Solubilities and Derived Gibbs Energies. Calix(n)arene esters are essentially nonelectrolytes, and therefore, the solution Gibbs energy can be calculated from solubility measurements of these compounds in the appropriate solvent at a given temperature. Thus, the solubility of ethyl p-tert-butylcalix(4)arenetetraethanoate in acetonitrile at 298.15 K was found to be 1.12 x lo-* M of saturated solution. From this value, the standard Gibbs energy A,G" (11.14 kJ mol-') was calculated. However, due to the extensive solvation of this ligand when exposed to an atmosphere of benzonitrile, the standard Gibbs energy of ethyl p-tert-butylcalix(4)arenetetraethanoate in this solvent could not be calculated. Nevertheless, these observations are useful since it can be concluded (although in qualitative terms) that benzonitrile is a better solvating medium for this ligand than acetonitrile. Additional useful information obtained is that related with the solubility of the sodium perchlorate complex of the tetraethyl ester derivative at 298.15 K in acetonitrile. Thus, a value of 6.25 x M of saturated solution was obtained for the solubility of [Na+EtCalix(4)]C104- in acetonitrile. Conductance measurements had shown that this electrolyte is fully dissociated in this solvent, and therefore, no correction for ion-pair formation is requiredS6Using the extended Debye-Hiickel equation for the calculation of the mean molar ionic activity coefficient, the standard Gibbs energy of solution (ASGO = 16.11 kJ mol-') was calculated. As pointed out previously,' interpretation based solely on Gibbs energy data can be misleading, and therefore, we proceeded with calorimetricmeasurements to derive the standard
Danil de Namor et al. enthalpies of solution. In addition, some of the limitations found in the determination of Gibbs energies are not applicable in the calculation of enthalpies. As far as the ligands are concerned and to assist with the interpretation of enthalpy data, 'H NMR studies were carried out, and this is discussed below. Standard Enthalpies and 'H NMR Studies. As pointed out previously,6the strength of ligand-solvent interactions are best reflected in the enthalpies of transfer of the macrocyclic ligand between two solvents. Thus, on the basis of transfer enthalpy data for ethyl p-terf-butylcalix(4)arenetetraethanoate from methanol (a solvent which does not show specific interactions with the ligand) to acetonitrile and 'HNMR studies of this ligand in C D F N , it was c o n c l ~ d e dthat ~ ~acetonitrile ~ interacts with the hydrophobic cavity of the ethyl ester, producing an "allosteric effect" which preorganizes better the hydrophilic cavity to interact with metal cations in this solvent than in methanol. These statements were supported by the availability of X-ray diffraction studies which demonstrate that acetonitrile sits in the hydrophobic cavity of a similar derivative, p-tert-butylcalix(4)arenetetracarbonate, forming a 1:1 complex with it.'o The availability of enthalpy of solution data for the tetraalkyl esters in acetonitrile and in benzonitrile at 298.15 K led us to proceed with 'H NMR studies in these solvents involving other calixarene ester derivatives. Chloroform does not show specific interactions with the esters, and therefore this solvent was used as reference. Details are given in Table 1. The 'H NMR spectra of the alkyl p-tert-butylcalix(4)arenes in acetonitrile and in benzonitrile are typical of the cone conformation observed for these ligands in CDC13, as reflected by the splitting pattem for the ArCHzAR protons (a pair of doublets with A6 values of 1.55 ppm for the methyl and ethyl ester and 1.54 ppm for the n-butyl ester in acetonitrile; A6 values in benzonitrile are 1.67, 1.69, and 1.70 ppm for the methyl, ethyl, and n-butyl ester, respectively. The most significant change observed for the esters in acetonitrile relative to chloroform is found in the downfield shifts of the aromatic protons, suggesting that similar interactions to those found for the ethyl derivative are observed for the methyl and n-butyl calix(4)arene esters in acetonitrile. No significant changes are found in the 'H NMR data in moving from the methyl to the n-butyl ester. Therefore, the decrease in the endothermic character of the reaction reflected in the ASH" value (20.80 kJ mol-') for the n-butyl relative to the ethyl (ASH" = 22.67 kJ mol-') in acetonitrile must be attributed to a slightly lower energy requirement to break the lattice of the former relative to the latter. These results lend support to the statement made in the preceding paper that the stability enhancement observed on cation complexation when moving from the methyl to the ethyl or n-butyl may be mainly attributed to an increase in the electron-donating effect of the alkyl group rather to an increase in the specific solvent-cavity interaction. As far as benzonitrile is concerned, judging from the transfer enthalpy from acetonitrile to benzonitrile (AtHo= -8.64 kJ and from solubility data in these solvents (discussedabove), a greater interaction occurs with the latter relative to the former solvent. This is also reflected in the I NMR spectra of calix(4)arene esters in benzonitrile. Unlike acetonitrile, significant downfield shifts are observed in this solvent, particularly for the exo and endo protons of the methylene groups attached to the aromatic rings. The downfield shifts observed for H-1 and H-4, although less pronounced, are significant. It was initially thought that benzonitrile may interact with the hydrophobic cavity of the ligand. However, computer-modeling calculations suggest that the possibility of stacking interactions between the solvent and the benzene rings outside the cavity are not excluded. The most
J. Phys. Chem., Vol. 99, No. 45, I995 16783
Alkyl p-tert-Butylcalix(4)arenetetraethanoates
TABLE 1: Proton CKemical Shws (6) and A6 Values of Alkyl (Methyl, Ethyl, nlutyl) p-tert-Butylcalix(4)arenetetraethanoatesin Acetonitrile and Benzonitrile at 298.15 K (Reference Solvent: CDC13)
benzonitrile 1.11 6.99 3.22 4.77 4.77 3.72
0.04
0.21 0.03. -0.06 -0.04 -0.04
acetonitrile
benzonitrile
acetonitrile
6 (ppm) Ad (ppm)" 6 (ppm) Ad (ppmY 6 (ppm) Ad (ppm)b 6 (ppm) Ad (PpmY 1.19 0.12 1.12 0.05 1.19 0.12 1.12 0.05 7.01 0.24 7.00 0.23 3.43 0.24 3.24 0.06 3.42 0.23 3.23 0.04 5.12 0.27 4.78 -0.08 5.09 0.26 4.78 -0.07 4.89 0.09 4.77 -0.04 4.89 0.08 4.76 -0.04 0.00 4.18 -0.03 4.13 3.67 -0.09 4.18 -0.03 1.24 -0.04 1.63 0.00 1.26 -0.02 1.37 0.01 0.00 0.93
benzonitrile
3.44 5.14 4.94 4.12 1.54 1.31 0.87
0.26 0.28 0.13 -0.01 -0.09 -0.05 -0.06
a Chemical shifts (ppm) for I in CDCl3: dH(1) = 1.07; dH(2) = 6.78; dH(3) = 3.19(exo) and 4.83(endo); dH(4) = 4.81; and dH(5) = 3.76. Chemical shifts (ppm) for 11in CDC13: 6H(1) = 1.07; dH(2) = 6.77; dH(3) = 3.19(exo) and 4.85(endo); dH(4) = 4.80; dH(5) = 4.21; and dH(6) = 1.28. Chemical shifts (ppm) for 111 in CDC13: dH(1) = 1.07; dH(2) = 6.77; dH(3) = 3.18(exo) and 4.86 (endo); dH(4) = 4.81; dH(5) = 4.13; dH(6) = 1.63; dH(7) = 1.36; and dH(8) = 0.93. The chemical shift of H(2) in benzonitrile could not be detected due to the overlap with the solvent peaks. Ad values are the differences found for the chemical shifts in acetonitrile and/or benzonitrile relative to chloroform.
TABLE 2: Enthalpies of Solution of Ethyl p-tert-Butylcalix(4)arenetetraethanoate Lithium and Sodium Perchlorate Complexes in Acetonitrile and in Benzonitrile at 298.15 K [LifEtCalix(4)]C104acetonitrile benzonitrile dmol dm-3 A J f M mol-' dmol dm-3 AJfM mol-' 4.06 x 10-4 -9.51 -9.83 8.95 x 10-4 -9.64 1.81 x 10-3 -9.31 1.36 x 10-3 2.40 x 10-3 -11.41 -6.73 3.18 x 10-3 2.77 x 10-3 -12.05 3.61 x 10-3 -9.79 -9.84 3.89 x 10-3 -9.80 5.12 x 10-3 A J P = - 10.42 f 1.23 AJP = -9.17 f 1.36 kT mol-' kT mol-' ~~~~
~
[NafEtCalix(4)]C104acetonitrile benzonitrile AJfM mol-' clmol dm-) AJfM mol-' dmol dm-3 7.28 x 10-4 -11.49 2.77 x 10-4 -34.52 1.34 x 10-3 -12.81 -37.26 5.59 x 10-4 1.47 x 10-3 -12.71 -33.37 1.25 x 10-3 -35.09 1.60 x 10-3 -12.65 1.57 x 10-3 2.16 x 10-3 - 12.83 -35.29 1.97 x 10-3 -32.18 3.07 x 10-3 -13.73 2.61 x 10-3 AQIIO = -12.70 f 0.72 A J P = -34.62 f 1.74 kT mol-' kT mol-' relevant conclusion that can be drawn from these results is that as far as calixarene derivatives are concerned, it is important to determine not only the extent of ligand solvation but also the possible sites of interaction between the ligand and the solvent. Regarding the metal-ion complexes in these solvents, standard enthalpies of solution of two new electrolytes, lithium ([Li+EtCalix(4)]ClO4-) and sodium ([Na+EtCalix(4)]C104-) ethyl p-tert-butylcalix(4)areneperchlorates, in acetonitrile and in benzonitrile at 298.15 K, are listed in Table 2. Since these are the first data ever reported on solution enthalpies involving calixarene-derivedelectrolytes, AJI values at each concentration are also given. As the data are independent of the electrolyte concentration, an average is taken as the standard enthalpy of solution of the electrolyte in the appropriate solvent.
It seems appropriate to compare these data with those for the uncomplexed electrolytes in acetonitrile (AJP NaC104 = -17.28 kJ mol-'; A J P LiC104 = -43.26 kJ mol-')I2 and in benzonitrile (AJP NaC104 = -10.42 kJ mol-'; ASH"LiC104 = -27.32 kJ mol-')'2 at 298.15 K. Standard enthalpies of solution are the result of two processes, namely, the breakage of the crystal lattice (endothermic) and the solvation (exothermic). For electrolytes containing the same anion, it is expected that as the size of the cation increases from M+ to M+EtCalix(4), the energetic requirement to overcome the crystal lattice will be lower, and therefore, it is not surprising to find that, for sodium, the dissolution process is enthalpically more favored for the electrolyte containing the complexed cation than for the metal-ion perchlorate. However, the fact that, for a given salt, the crystal lattice contribution to the solution enthalpy is the same strongly suggests that the solvation of [Na+EtCalix(4)] is greater in acetonitrile than in benzonitrile. Furthermore, the solvation of the complexed sodium cation is greater than the sodium ion. These statements are best reflected in the standard enthalpies of transfer. A@, from acetonitrile (reference solvent) to benzonitrile for Na+ (3.77 kJ mol-'), [Na+EtCalix(4)] = 17.91 kJ mol-' (data based on the P b A s P b B c~nvention),'~ as well as that for the ligand (-8.64 W mol-'). The most striking feature of these results is the dramatic changes observed in enthalpic stabilities for the free ligand to that of the inclusion complex in these solvents. These data unambiguously demonstrate that the selective solvation of the ethyl ester derivative in benzonitrile relative to acetonitrile is reversed for the complexed cation, and this may be attributed to conformational changes that the ligand undergoes upon complexation with sodium in acetonitrile. In fact, the increase in exothermic behavior previously shown in the complexation of sodium and EtCalix(4) in acetonitrile relative to the same process in benzonitrile (A,HO(MeCN) A,HO(PhCN) -19 kJ mol-') is almost entirely due to the higher enthalpic stability of [Na+EtCalix(4)] in the former relative to the latter solvent, as reflected in the AtW value for this cation shown above. This is further corroborated by the
16784 J. Phys. Chem., Vol. 99,No. 45, 1995
Danil de Namor et al.
TABLE 3: Enthalpies of Coordination (W mol-') of'Ethy1p-tert-Butylcalix(4)areneLithium and Sodium Perchlorates at 298.15 K _ _ _ _ ~ ~
-43.26" -27.32'
A P EtCalix(4) 22.64' 14.03'
NaC104
ASH" EtCalix(4)
LiC104 MeCN PhCN
- 17.28"
MeCN PhCN a
- 10.42'
22.67' 14.03'
[Li+EtCalix(4)]C104-9.17d - 10.42d
[Na+EtCalix(4)]C104-34.62d -12.7od
AcW Li++EtCalix(4) (solution)
AcmniH' LiC104+EtCalix(4) (solid state)
-48.78' -57.20' A X Na++EtCalix(4) (solution)
-60.23 -60.07 NaC104+EtCalix(4) (solid state)
-69.20' -50.70'
-29.19 -34.39
Acmrdw
'
Reference 5. Reference 12. Reference 6. From Table 2. e Reference 1
greater loss in entropy observed for the complexation of Naf and EtCalix(4) in acetonitrile (&So = -85.1 J K-'mol-') with respect to the same process in benzonitrile ( 4 s " = -25.1 J K-'mol-'). Inspection of enthalpy data for lithium contrasts with those for sodium. Thus, the higher enthalpic stability (A@ more negative) observed in the dissolution of LiC104 (see above) relative to [LifEtCalix(4)]C104- strongly suggests that the former is more solvated than the latter in these solvents, since the energy requirements to break the crystal lattice (endothermic process) for LiC104 is expected to be greater than for the complexed salt. Transfer enthalpies from acetonitrile to benzonitrile for Lif (A@ = -13.18 kJ mol-'), [LifEtCalix(4)] (A,W = -4.34 kJ mol-'), and [EtCalix(4)] (AtW = -8.64 kJ mol-') show that the value for the free ligand is very similar to that of the complexed cation, and if so by inserting the appropriate enthalpies in the relationship between complexation AcH and transfer data AtH expressed by eq 2,
e1~ewhere.I~ As far as calixarenes are concerned, information regarding the thermodynamics of coordination in the solid state is nonexistent, and this is now discussed. Gibbs Energy of Coordination. The availability of solution and complexation data allows the calculation of the Gibbs energy of coordination, AcWrdGO, referred to reactants and product in their solid (sol) state. By inserting the appropriate standard Gibbs energies (ASGO for NaC104 in acetonitrile at 298.15 K is that from the literat~re)~ in the following thermodynamic cycle,
As@ = 2.51 kJ mol-'
I
4,& ,z1 I
A,GO = 11.14 kJ mol-'
Na+(s) + CIO4-(s) + Ekalix(4)(s)
ekJ=
A,@ = 16.11 kJ mol-'
[Na+EtCalix(4)](s)+ CI04-(s)
(3)
AcHo(PhCN) - AcHo(MeCN) = AJP[Li+EtCalix(4)] - AJP[EtCalix(4)] - AJP(Lif) (2) a much clearer interpretation of the factors contributing to the variation in enthalpic stability previously observed in the complexation of Li+ and EtCalix(4) in benzonitrile (AcW = -57.2 kJ mol-') relative to acetonitrile (AcW = -48.78 kJ mol-') can be made. Given that the AtHo of the complex is largely canceled by that for the free ligand, it can be concluded that the difference observed in the AcW values in these solvents is largely influenced by the changes in the solvation energy of lithium in these media. This may partially explain the gain in entropy observed for this cation upon complexation with the calixarene ester in acetonitrile (AcSo = -44.9 J K-'mol-') relative to benzonitrile (AJ" = -86.7 J K-'mol-') since its desolvation is expected to be greater in the former than in the latter solvent. These studies reveal that, unlike cryptands, where the solvation of the cation is normally dominant in dipolar aprotic medial4 in determining the variation in complex stability, in calixarene chemistry, depending on the system under study, the solvation energies of the macrocycle, the cation, and the complex need to be considered in the interpretation of cation complexation processes involving these ligands. (b) Thermodynamics of the Coordination Process in the Solid State. Undoubtedly, structural information in the solid state is best obtained from X-ray crystallographic studies. However, the isolation of suitable crystals for these purposes can be a lengthy process, and often this cannot be successfully achieved. We have recently shownI5that for systems involving 1-aza-12-crown-4 and lithium salts, the enthalpy of coordination is a suitable reporter of specific interactions in the solid state. Further applications of coordination data have been discussed
AcoordGofor this system is calculated. It should be noted that this value does not differ significantly from the Gibbs energy associated with the complexation process in solution, although the two sets of data are not strictly comparable given that in the solid state the electrostatic interaction between cation and anion may be significant, particularly if the cation is not fully shielded by the ligand. Although this may not be the case for lithium and sodium, as shown in the preceding paper. However, the availability of literature data for the AcoordGoof sodium perchlorate with cryptand 222 (-67.4 kJ mol-I)I6 allows us to conclude that the stability of the sodium ethyl ester calix(4)areneperchlorate complex is lower by about 21 kJ mol-' from that of sodium cryptate. Enthalpies and Entropies of Coordination. Equation 3 expressed in terms of enthalpy allows the calculation of the enthalpies of coordination, AcoordH0.for lithium and sodium perchlorates and the calixarene ester in their solid state. A,W values for sodium and lithium perchlorates and for ethyl p-tertbutylcalix(4)arenetetraethanoate in acetonitrile and benzonitrile are those from the literature.'* Details are given in Table 3. Since for a given metal-ion salt and for a given ligand the thermodynamic data referred to the process in the solid state should be the same, independent of the solvent from which this is derived, the calculation of ACoordWoffers a suitable means of checking the accuracy of complexation data previously reported' and solution enthalpies shown in this paper. Considering the number of steps involved in the derivation of coordination data, a good agreement between the values derived from the two solvents is found. Thus, average values of -60.15 and -31.93 kJ mol-' are taken for the enthalpies of coordination of lithium and sodium perchlorates with EtCalix(4),respectively. The trend observed for the enthalpies of coordination of lithium and sodium perchlorates with the calix(4)arene ester
J. Phys. Chem., Vol. 99,No. 45, 1995 16785
Alkyl p-tert-Butylcalix(4)arenetetraethanoates derivative is somehow similar to that observed for the coordination of cryptand with sodium (AcoordHO= -59.2 kJ mol-') and potassium (AcoordH"= -36.0 kJ mol-') in that the stability of enthalpic terms in the solid state is lower for the complex, which shows the higher enthalpic stability in solution (AJP of K+ 222 is higher than Na' 222 in most solvents). It seems that the higher the interaction between the metal cation and the ligand, the lower its electrostatic interaction with the anion. A possible way of visualizing this statement is by considering that the more stable complexes are best shielded by the ligand, and as a result, the electrostatic cation-anion interaction becomes weaker. This is also reflected in the entropies of coordination of sodium perchlorate and the tetraethyl ester (Acoordso= 48.1 J K-' mol-') and the corresponding sodium cryptate salt (AcoordSO = 28.0 J K-' mol-'), which are positive, and as previously stated,I6 this may be mainly attributed to a weaker cation-anion interaction in the sodium macrocyclic salt in the solid state relative to the stronger electrostatic cation-anion interaction between sodium and perchlorate in the ionic solid.
+
+
Final Remarks From solution studies, we therefore conclude that these results call upon further investigationsregarding the solvation properties of calixarene derivatives and their metal-ion complexes in different solvents. Although the use of these macrocycles as sequestering agents for metal cations from aqueous solutions into nonaqueous media has been explored," these studies demonstrate the important role of the solvent in processes involving calixarene derivatives. Undoubtedly, the choice of the solvent in metal-extraction processes will be greatly facilitated by the availability of the thermodynamic and kinetic data for these systems. As far as the coordination process is concemed, this paper reports data that can be combined with solution parameters involving host, guest, and complex in low-dielectric media to gain information regarding the binding process in these media.I8 It is well established that in solvents of low permitti~ity'~ extensive ion-pair formation occurs, and therefore, any attempt to determine directly the thermodynamics associated with these processes is likely to lead to misrepresentative data. We are now proceeding with solution studies of lithium and sodium calix(4)arene ester electrolytes containing different anions, aiming to assess the anion effect in the coordination process.*O
Acknowledgment. The financial support given by the European Commission, DG-MI, ISC, to D.A.P.T., L.E.P.S., and J.W. to carry out research at the Thermochemistry Laboratory is gratefully acknowledged. E.G. thanks the National University of the South, Bahia Blanca, Argentina, for sabbatical leave. References and Notes (1) Danil de Namor, A. F.; Gil, E.; Llosa Tanco, M. A,; Pacheco Tanaka, D. A.; Pulcha Salazar, L. E.; Schulz, R. A.; Wang, J. J . Phys. Chem. 1995, 99, 16776. (2) Bax, D.; de Ligny, C. L.; Renynse, A. C. Recl. Trav. Chim. 1972, 91, 125. (3) Christensen, J. J.; Izatt, R. M.; Hansen, L. D. Rev. Sci. Znstrum. 1965, 36, 779. (4) Irving, R. J.; Wadso, I. Acta Chem. Scand. 1964, 28, 195. ( 5 ) Ghousseini, L. Ph.D. Thesis, University of Surrey, 1985. (6) Danil de Namor, A. F.; Cabaieiro, M. C.; Vuano, B. M.; Salomon, M.; Pieroni, 0. I.; Pacheco Tanaka, D. A.; Ng, C. Y . ;Llosa Tanco, M. A.; Rodriguez, N. M.; CArdenas Garcia, J. D.; Casal, A. R. Pure Appl. Chem. 1994, 66, 435. (7) Danil de Namor, A. F.; Traboulssi, R.; Lewis, D. F. V. J . Am. Chem. Soc. 1990, 112, 8442. (8) Danil de Namor, A. F.; Apaza de Sueros, N.; McKervey, M. A,; Bmett, G.; Amaud Neu, F.; Schwing-Weill, M. J. J . Chem. Soc., Chem. Commun. 1991, 1546. (9) Danil de Namor, A. F. Pure Appl. Chem. 1993, 65, 193. (10) McKervey, M. A.; Seward, E. M.; Ferguson, G.; Ruhl, B. L. J . Org. Chem. 1986, 51, 3551. (1 1) Shinkai, S.; Fujimoto, K.; Otsuka, T.; Ammon, H. L. J . Org. Chem. 1986, 51, 3551. (12) Danil de Namor, A. F.; Ghousseini, L. J . Chem. SOC., Faraday Trans. 1985,81,781. Danil de Namor, A. F.; Berroa de Ponce, H. J . Chem. Soc., Faraday Trans. 1988, 84, 1671. (13) Cox, B. G.; Hedwig, G. R.; Parker, A. J.; Watts, D. W. Aust. J . Chem. 1974, 27,477. (14) Danil de Namor, A. F. J . Chem. SOC., Faraday Trans. 1988, 84, 244 1. (15) Danil de Namor, A. F.; Llosa Tanco, M. A.; Salomon, M.; Ng, C. Y . J . Phys. Chem. 1994, 98, 11796. (16) Cox, B. G.; Schneider, H. Pure Appl. Chem. 1989, 61, 171. (17) Penin, R.; Harris, S. In Topics in Inclusion Phenomena. Calixarenes. A Versatile Class of Macrocyclic Compounds; Vicens, J., Bohmer, V., Eds.; Kluwer Academy: Dordrecht, 1990. (18) Danil de Namor, A. F.; Llosa Tanco, M. A.; Ng, C. Y.; Salomon, M. Pure Appl. Chem. 1995, 67, 1095. (19) Femandez Prini, R. Physical Chemistry of Organic Solvent Systems; Covington, A. K., Dickinson, T., Eds.; Plenum Press, New York, 1973. (20) Danil de Namor, A. F.; Llosa Tanco, M. A.; Pulcha Salazar, L. E.; Salomon, M. Work in progress, 1995. Jp951974U
