J. Phys. Chem. 1995,99, 16776-16780
16776
Thermodynamics of Calix(4)arene Esters. 1. Complexation of Alkyl p-tert-Butylcalix(4)arenetetraethanoatesand Alkali-Metal Cations in Acetonitrile and in Benzonitrile Angela F. Danil de Namor,**tElisabeth GilJ Margot A. Llosa Tanco,? David A. Pacheco TanakaJ 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@
Using silver electrodes a double competition reaction involving cryptands and calixarene esters has been for the first time used for the potentiometric determination of stability constants of highly stable complexes of calixarene esters and metal cations (lithium and sodium) in acetonitrile and in benzonitrile at 298.15 K. Corresponding data for less stable complexes were derived from thermochemical measurements using titration calorimetry (macro and micro). Thus, Gibbs energies, enthalpies, and entropies of complexation of alkyl (methyl, ethyl, and n-butyl) p-tert-butylcalix(4)tetraethanoates and alkali-metal cations in acetonitrile and in benzonitrile at 298.15 K are reported. The results show that as the electron-donating effect of the ligand increases in moving from the methyl to the ethyl and to a lesser extent to the tert-butylcalix(4)arene ester, making the carbonyl oxygen more electronegative, their interaction with metal cations increases. The implication of these finding on the selective recognition of alkali-metal cations by calix(4)arene esters in these solvents is discussed. It is concluded that in both solvents (i) metal ion-ligand stability (hence Gibbs energy) is enthalpically controlled and (ii) entropies of calix(4)arenate formation reflect marked differences between the solvation of the ligands relative to their metal-ion complexes.
The design of ligands with selective properties for a particular neutral or ionic species is one of the main targets in synthetic macrocyclic chemistry. This field of research, greatly motivated by the discovery of the crown ethers by Pedersen in 1967,' has presented the physical chemist with a number of interesting and challenging questions regarding the factors which contribute to selective host-guest complexation in solution processes. If a thermodynamic approach is considered, information concerning the enthalpy and entropy contributions to the stability of the complexes as reflected in the Gibbs energy should be obtained since these data provide valuable information. Thus, hostguest binding energies as well as solvation energies of the ligand, metal cation, and complex will predominantly contribute to the magnitude of AH', while the entropy of complex formation is largely dependent upon the number of reacting species, the reaction medium, and conformational changes resulting from complexation. A direct implication of the above statements is the need of accurate thermodynamic data, the derivation of which is largely dependent on the methodology employed. Calixarenes, an interesting class of macrocyclic ligands, are obtained from the base-catalyzed reaction between parasubstituted phenol and formaldehyde.* Functionalization of the low rim of parent calix(n)arenes has led to a variety of derivative^.^-^ Particularly interesting are the calix(4)arene esters 1. X-ray crystallographic and NMR studies (mainly in CDC13) have shown that these compounds are in a fixed cone conformation in the solid state and in s~lution.~.'Therefore, these macrocyclics are unusual in that they contain both hydrophilic and hydrophobic regions. Thus, calix(4)arene esters have (i) a hydrophobic envelope formed between the aromatic rings able to enter interaction with neutral species and (ii) a hydrophilic ring able to interact with metal cations. In addition,
' Department of Chemistry.
* Department of Chemical and Processing Engineering. @
Abstract published in Advance ACS Abstracts, October 15, 1995
0022-3654/95/2099- 16776$09.00/0
1
R.Methy1, Ethyl, n Butyl
x
o
OR
these ligands have found interesting analytical and industrial applications.8 There have been a number of reports on stability constants of calixarene esters and metal cations mainly in methanol and a c e t ~ n i t r i l eat~ .298.15 ~ K. Most of the data were derived from spectrophotometric measurements, with only few stability constants determined by either calorimetric or conductimetric titrations.I0 It may be noted that on the basis of spectrophotometric data, a selectivity index for calixarene esters and univalent metal cations has been attempted.8 However, it should be emphasized that absorbance (W spectrophotometry), heat (Calorimetry),and conductance (conductimetric titrations) effects are proportional to concentration, and therefore, the determi0 1995 American Chemical Society
J. Phys. Chem., Vol. 99, No. 45, 1995 16777
Thermodynamics of Calix(4)arene Esters nation of stability constants based on these measurements are less versatile than those dependent upon the logarithm of activities." Among the latter and for highly stable complexes, the use of the silver-silver ion electrodes in potentiometric titrations has proved to be one of the most sensitive methods for accurate potential measurements even for cases in which competitive processes are involved.' I Thus, SchneiderI2 and extensively used the potentiometric method to measure the activity of free silver ions in the presence of a competitive equilibrium involving the relevant metal cation in order to calculate the stability constants of metal-ion cryptates in a wide variety of nonaqueous solvents. The derivation of stability constants by this method requires equilibria data for the following processes: (i) the complexation of silver and cryptand (Cry) (eq 1) and (ii) the competition between the silver cryptate and the relevant cation M+ (eq 2) in a given solvent (s).
(Aldrich) was dried with CaC12 and distilled from P2O5 under reduced pressure.I6 The middle fraction was collected and redistilled after refluxing the solvent for a few hours. The amount of water in the solvent did not exceed 0.01%. Potentiometric Titrations. Potentiometric titrations were carried out in a Camlab potentiometer, using the electrochemical cell suggested by Schneider and c o - w ~ r k e r swith ~ ~ ~a' silver~ silver ion reference electrode consisting of a silver wire introduced in a solution of silver perchlorate or nitrate (0.01 M). To keep the ionic strength constant, a solution of TEAP (0.05 M) was used at the indicator electrode. Both electrodes were separated by a salt bridge containing 0.05 M of TEAP in the appropriate solvent. A schematic representation of the electrochemical cell is given: Et,NC10, AgIAg', X M 11 0.05 M 11 0.01 M, Ag+/Ag
It is evident from eq 1 that to apply this method to other systems, the ligand must be able to form stable complexes with silver. However, calixarene esters are poor complexing agents for silver in acetonitrile and benzonitrile. Therefore, for the determination of stability constants of metal cations and these ligands, the potentiometric method described above has been expanded by introducing a second competition reaction (eq 3) involving calixarene esters (R Calix). M+CT(S)
+ R Calix(s) -M+R Calix(s) + cry(s) K2
(3)
Thus, combination of eqs 1, 2, and 3 leads to the calculation of the stability constant of the metal-ion calixarenate (eq 4). (4) In this paper, stability constants of alkali-metal cations and various calix(4)arene esters (R = methyl, ethyl, and n-butyl) in acetonitrile (MeCN) and benzonitrile (PhCN) at 298.15 K derived from the double competitive potentiometric method (Li+ and Na+) and from titration calorimetry (K+ and Rbf) are reported. From these data, the standard Gibbs energies associated with the complexation process are calculated. The implication of these results on the selective recognition of calix(4)arene esters for metal cations is quantitatively evaluated. Combination of Gibbs energies with enthalpies measured calorimetrically yields the entropies of complexation of alkali-metal cations and the various calix(4)arene esters in these solvents. The results, interpreted in terms of the ligand, the solvent, and the cation, are compared with previous work involving cryptands in these media. Experimental Section Chemicals. Kryptofix 222 and 22 (Fluka) and silver, sodium, potassium, and rubidium perchlorates (Aldrich) were dried under vacuum for several days before use. Lithium perchlorate (Aldrich) was twice recrystallized from water. Tetra-n-ethylammonium perchlorate (TEAP) was prepared from tetraethylammonium hydroxide and perchloric acid. The product was recrystallized from methanol, washed with ether, and dried under vacuum. Methyl, ethyl, and n-butyl p-tert-butylcalix(n)arenetetraethanoates were synthesized as detailed elsewhere.I4 Acetonitrile (Aldrich) was refluxed and distilled from calcium hydride.I5 The water content of the solvent, checked by Karl Fischer titration, was not more than 0.02%. Benzonitrile
(5)
The electrodes were kept at constant temperature (298.15 & 0.05 K) using a thermostat bath. A computational simulation program was used to select the appropriate concentrations of ligand and metal-ion salt to be used for these measurements. Titrations were carried out until an excess of titrant was added. To carry out these titrations, a solution of silver nitrate or perchlorate was added into the cell containing TEAP (25 m3; 0.05 M) solution in the appropriate solvent. At least 10 additions were performed, and the data were used to calculate the standard electrode potential of the reference cell. This was followed by the addition of the appropriate cryptand to a solution containing silver perchlorate in the solvent. Potential readings were taken after each addition, and the data were used to calculate the stability constant of the silver cryptate complex (eq 1). Then, in the same experiment, additions of the metalion salt in the required solvent were made. The potential data were used to determine the equilibrium constant for the process represented by eq 2. Finally, a solution of the appropriate alkyl p-tert-butylcalix(4)arenetetraethanoate in the nonaqueous solvent was added (at least 10 additions) to the mixture contained in the electrochemical cell. The potentials reflecting the decrease of the silver activity (as the process given by eq 3 took place) were taken. These data were used to calculate the equilibrium constant for the process represented by eq 3. Thermochemical Measurements. Enthalpies of complexation of metal cations and alkyl p-tert-butylcalix(4)arenetetraethanoate esters were determined by classical calorimetry using a Tronac 450 titration calorimeter andor by titration microcalorimetry. For the latter, microcalorimetric experiments at 298.15 K were carried out using the titration vessel of the 2277 Thermal Activity Monitor. The vessel was filled with 2.8 cm3 of a solution of the metal-ion perchlorate in the appropriate solvent (concentrations used are dependent on salt solubility), and calixarene esters (concentrations used were about 10 times higher than that of the metal-ion salt) were injected (-15 injections; 0.015-0.025 cm3 for each run) from a 0.5 cm3 gas-tight Hamilton syringe, attached to a computer-operated syringe drive, at 5-min intervals. In cases where ligand solubility was relatively low, this was placed in the calorimetric vessel and the metal-ion salt was placed in the syringe. Blank experiments were carried out in all cases to account for heat of dilution effects resulting from the addition of either ligand or metal-ion salt to the solvent contained in the calorimetric vessel. A dynamic correction based on Tian's eq~ation".'~was used to calculate the integrals from the microcalorimetric curve. The reliability of the equipment was checked by carrying out the
16778 J. Phys. Chem., Vol. 99, No. 45, 1995 -280
QP -300
% a *X
-320
340
-360
-380
-400 V o l u m e of MeCalix(4) s o h a d d e d (mil
Figure 1. Experimental ( x ) and calculated (0) values for the potentiometric titrations of sodium cryptate 22 with methyl p - f u r butylcalix(4)arenetetraethanoate in benzonitrile at 298.15 K (eq 3 in the text).
standard reactions suggested in the 1iterat~re.I~ The procedure used for the Tronac 450 is described elsewhere.*O Results and Discussion Potentiometric Determination of Equilibria Data for Highly Stable Complexes. Given that the double competitive potentiometric reaction using silver electrodes is first reported for the determination of stability constants of metal cations and macrocyclic ligands (which are essentially nonselective for silver), a representative example for the process represented by eq 3 is given. Thus, Figure 1 shows the potentiometric titration curve obtained by the addition of methyl p-tert-butylcalix(4)arenetetraethanoate to a solution containing sodium 22-cryptate in benzonitrile at 298.15 K. The experimental data shown in this figure reflect that as the amount of ester added increases, the equilibrium position is shifted as to favor the release of the metal cation (from its cryptate complex) to form the metal-ion calixarenate complex, allowing the free cryptand to interact with the silver ion. This is unambiguously demonstrated by the decrease observed in the silver ion activity as the reaction proceeds. Figure 1 also shows calculated data obtained from the results from the simulation program. Excellent agreement between calculated and experimental data is observed. Thus, Table 1 lists stability constants (expressed as log K,) of alkyl p-tert-butylcalix(4)arenetetraethanoates (methyl; MeCalix(4); ethyl; EtCalix(4); n-butyl; n-BuCalix(4)) and lithium and sodium cations (eq 4) in acetonitrile and in benzonitrile at 298.15 K derived from the double competitive potentiometric method described above. Data for alkali-metal cryptate (cryptand 222; 222 or cryptand 22; 22) in these solvents required for the evaluation of log K, for calixarene esters and metal cations in media are also reported in Table 1. The standard deviations of the data are also included in Table 1. Standard deviations in log K , for the process represented by eq 1 are calculated from s = [Z(xi - x)2/(n- l)Il2. For processes 2 and 3, the standard deviations shown in the log K , values are the square root of the sum of the individual variances. For comparison purposes, published data for cryptand 222 and these metal cations in a~etonitrile'~ and in benzonitrile20at 298.15 K are included in Table 1. Except for lithium and cryptand 222 in benzonitrile, good agreement is found between these two sets of data. As far as cryptand 22 is concemed, log K, values for silver and sodium in benzonitrile at 298.15 K are first reported. The data shown in Table 1 for lithium and calixarene esters reflect that stability constants for these systems
Danil de Namor et al. determined by different methods are in close agreement. However, as the magnitude of log Ksincreases to values which are outside of the scope of techniques such as spectrophotometry,' calorimetry, and to some extent conductimetry,I0considerable differences are often found between the results obtained by these methods and those derived from potentiometry. This statement is illustrated by the data for the Na+-EtCalix(4) complex in acetonitrile, where the value obtained by spectrophotometry differs by more than 2 log units from that obtained potentiometrically. Similar comments apply to the same system in benzonitrile as well as for the calorimetric value obtained for Na+-MeCalix(4) in this solvent. However, the agreement found between the log K , values determined by the double competitive potentiometric method using two different cryptands reinforces previous statements regarding the versatility of this method for the determination of equilibria data involving highly stable complexes.' I We therefore conclude that this new approach can be applied to other systems in which the first ligand has a high affinity for silver while the second ligand has low or no affinity for this cation. For the particular case of highly stable metal-ion complexes of calix(n)arene esters, the outcome of this research leads us to state that, so far, this is the most suitable method available at present. For the less stable metal-ion complexes, titration calorimetry has been used for the derivation of stability constant data. An additional advantage of this methodology relative to others currently used for the determination of log Ks values is that enthalpy data are also obtained from thermochemical measurements. In the following section we report log K , values for the complexation of less stable metal-ion complexes of calix(4)arene esters and the thermodynamics associated with the complexation of alkali-metal cations with these ligands at 298.15 K. Titration Calorimetry. Derived Thermodynamic Parameters of Complexation. Table 2 lists stability constants (log K,) for the complexation of potassium and rubidium with the various calix(4)arene esters in acetonitrile and for potassium and these ligands in benzonitrile at 298.15 K obtained from titration calorimetry. In Table 2 are also included derived Gibbs energies, enthalpies, and entropies for the complexation of alkali-metal cations and calixarene esters in acetonitrile and in benzonitrile (eq 4) at 298.15 K. In cases where stability constant data derived from different methods are available (see Table 1) and reasonable agreement is found between these values, an average of log K, is taken for the calculation of the Gibbs energy of complexation. All standard enthalpies reported in Table 2 have been measured calorimetrically. Due to the much higher sensitivity of titration microcalorimetry relative to classical calorimetry, the A,Ho values for ethyl p-tert-butylcalix(4)arenetetraethanoateand metal cations (Na+, K+, Rb+) slightly differ from those previously reported.21 The most relevant aspect to highlight is that no heat was detected in the microcalorimeter for the reaction between cesium ions and calix(4)arene esters in acetonitrile. It may be correctly argued that this is not an indication that these ligands are unable to interact with this cation given that if the heat associated with these processes is close to 0 kJ, calorimetry is not a suitable reporter of molecular events for these systems. On the other hand, attempts to reproduce the spectrophotometric value reported in the literature for the tetraethyl ester and cesium in acetonitrile at 298.15 K (log K , = 2.8)7 were unsuccessful since no significant changes were observed in the spectra by the addition of cesium salts to solutions of calix(4)arene esters in this solvent. The availability of thermodynamic parameters of complex-
J. Phys. Chem., Vol. 99, No. 45, 1995 16779
Thermodynamics of Calix(4)arene Esters
TABLE 1: Stability Constants of Silver, Sodium, and Lithium Cryptates and Lithium and Sodium Alkyl (Methyl, Ethyl, n-Butyl) ptert-Butylcalix(4)arenetetraethanoates in Acetonitrile and Benzonitrile at 298.15 K log Ks acetonitrile Ag+222 Ag+22 Li+222 Na+222 Na+22 LPMeCalix(4) Na'MeCalix(4) LPEtCalix(4) Na+EtCalix(4) Li+n-ButCalix(4) Na+n-ButCalix(4) a
9.01 f 0.01
8.92"
6.98 f 0.05 9.72 f 0.14
6.97" 9.63"
5.61 f 0.06 6.97 f 0.14 6.10 f 0.19 7.53 f 0.15 6.21 f 0.05 7.67 f 0.22
benzonitrile
6.40b 5.80b
10.00 f 0.03 9.45 f 0.01 8.18 f 0.04 11.20 f 0.04 5.66 f 0.03 5.63 f 0.06 6.80 f 0.04 5.49 f 0.22 7.57 f 0.30 6.09 f 0.05 7.56 f 0.05
6.1OC 7.82'
9.97" 9.14" 11.60" 5.27' 6.76 Ifr. 0.04d
5.51e
6.17' 7.32 f 0.16d
References 13 and 20. Reference 7. Conductimetric method. Using cryptand 22. Reference 10.
TABLE 2: Thermodynamic Parameters for the Complexation of Alkali-Metal Cations and Alkyl (Methyl, Ethyl, n-Butyl) p-tert-Butylcalix(4)arenetetraethanoates in Acetonitrile and in Benzonitrile at 298.15 K Methyl p-tert-Butylcalix(4)arenetetraethanoate cation
log Ks
A,G", kT mol-'
A,Ho, kT mol-'
A,So, J K-' mol-'
AsolvSod, J K-'mol-'
A,$", J K-I mol-'
-252.3 -214.2 -187.0 - 174.0
-271.7 -292.0 -246.4 -164.1
-237.7 -218.0 -205.2
-291.1 -226.0 -225.1
-252.3 -214.2 -187.0 -174.0
-297.2 -299.3 -263.1 -213.5
-237.7 -218.0 -205.2
-324.4 -242.9 -215.8
-252.3 -214.2 -174.0
-288.7 -294.8 -225.2
-237.7 -218.0 -205.2
-311.3 -245.6 -220.1
Acetonitrile
f 0.80 f 0.50 f 0.70 f 0.02
Li+ Na+ K' Rb+
5.61 f 0.03 6.97 & 0.02 4.01 f 0.03 2.25 f 0.04
-32.02 f 0.07 -39.79 f 0.05 -22.89 f 0.07 - 12.84 f 0.09
-37.80 -63.00 -40.63 -9.89
Li+ Na+
5.45 f 0.03 6.78 f 0.04 2.70 f 0.07"
-31.11 f 0.07 -38.70 f 0.10 -15.41 f 0.17'
Benzonitrile -47.02 f 0.29" -53.4 -41.08 f 1.23" -8.0 -21.34 f 0.68" - 19.9a
K+
-19.4 -77.8 -59.4 9.9
Ethyl p-tert-Butylcalix(4)arenetetraethanoate Acetonitrile -48.78' -44.9 -69.20 f 0.96 -85.1 -45.75 f 0.45 -76.1 -23.34 f 1.36 -39.0
6.20 f 0.05 7.68 f 0.08 4.04 f 0.03 2.05 f 0.03 (1 .90)b
-35.39 f 0.12 -43.81 f 0.20 -23.06 f 0.07 - 11.70 f 0.07
5.49 f 0.22 7.57 f 0.02 3.51 f 0.03
-31.34 f 0.55 -43.27 f 0.05 -20.04 f 0.07
K+
6.21 f 0.01 7.67 f 0.03 2.05 f 0.03
n-Butyl p-tert-Butylcalix(4)arenetetraethanoate Acetonitrile -35.45 f 0.02 -46.30 f 1.00 -36.4 -43.78 f 0.07 -67.80 f 1.00 -80.6 -11.67 f 0.20 -26.91 f 1.54 -51.0
Li+ Na+ K+
6.09 f 0.02 7.44 & 0.03 3.48 f 0.06
-34.76 f 0.05 -42.47 f 0.07 -19.86 f 0.15
Li+ Na+
K+ Rb+
Li+ Na+
K+
Li+ Na+
a
Benzonitrile -57.20 f 1.80 -86.7 -50.70 f 1.10 -24.9 -23.21 f 0.86 -10.6
Benzonitrile -56.70 f 0.78 -73.6 -27.6 -50.70 & 1.33 -14.9 -24.30 f 0.33
Reference 10. Reference 7. Reference 21. Reference 20.
ation of alkali-metal cations and various calix(4)arene esters in two solvents prompts us to discuss these data in terms of the ligand, the solvent, and the cation. As far as the ligands are concerned, the results show that as the electron-donating effect of the alkyl group increases in moving from the methyl to the ethyl and the n-butyl derivatives, making the carbonyl oxygen of the ester more electronegative, their interaction with metal cations increases. Thus, the highest stability enhancement is brought about by replacing the methyl by the ethyl ester and to a lesser extent in moving from the latter to the n-butyl derivative. It seems appropriate at this stage to discuss the implications of these results on the selectivity of alkyl p-tert-butylcalix(4)arene tetraethanoates for alkali-metal cations in these solvents. Thus, by taking the ratio of stability constants2*in the appropriate solvent, a quantitative assessment on the selective properties of a given ester for a metal cation with respect to another can be made. Thus, in acetonitrile, the
methyl ester is more selective for sodium than for lithium, potassium, and rubidium by factors of about 23, 912, and 5.2 x lo4,respectively. Substitution of methyl by ethyl or n-butyl groups increases the ability of these ligands to selectively recognize sodium with respect to lithium (-30), potassium (2.4 x lo3), and rubidium (4.6 x lo5). In benzonitrile, the most dramatic change observed with respect to acetonitrile is that found for sodium relative to potassium (-1.2 x 104 for methyl, -1.1 x lo4 for ethyl, and 9 x lo3 for n-butyl). It it therefore concluded that in acetonitrile, calix(4)arene esters are able to recognize alkali-metal cations far more selectively than it was originally ~ t a t e d . ~ As far as the solvent is concerned, stability constants (hence, Gibbs energies) listed in Table 2 show that in both acetonitrile and benzonitrile the same selectivity pattern is observed, with the net result of maximum stability for sodium with a monotonic decrease from sodium to rubidium in acetonitrile and to
Danil de Namor et al.
16780 J. Phys. Chem., Vol. 99, No. 45, 1995 potassium in benzonitrile. In terms of enthalpy, the contribution of this term to the Gibbs energy of the process is dominant relative to the entropic contribution to the extent that, in acetonitrile, the pattem found in enthalpies is almost mirrored in the Gibbs energies. Thus, the greatest entropy losses are found for the complexes of highest stabilities, as reflected in the data shown in Table 2 for sodium and the various calix(4)arene esters in acetonitrile. In benzonitrile, the enthalpy term is again dominant; however, there is a definite size effect as far as enthalpies are concerned, with the highest stability for lithium and the lowest for potassium. To interpret the higher stability observed for these ligands and most of these cations in acetonitrile relative to benzonitrile, it seems relevant to consider previous thermodynamic studies carried out with other macrocycles and these cations in these solvents. Thus, for ligands like crypt and^^^.^^.^^ it was found that the most highly solvated is the cation in a given medium; the weakest is its interaction with cryptands in that medium. Thus, single-ion-transfer Gibbs energies, AtGo, for alkali-metal cations from acetonitrile to benzonitrile20based on the P h A s P b B convention25show that the latter solvent is a poorer solvator for these cations than the former. As a result, benzonitrile offers a more suitable complexation medium for cryptands and alkali-metal cations than acetonitrile. Quite clearly these concepts do not always apply to processes involving calixarene esters and metal cations in these solvents, as illustrated in the data shown in Table 2. Solvation changes upon complexation are best reflected in the entropy term. Therefore, in an attempt to analyze the cation effect upon complexation with these ligands, their entropies of solvation, Aso~vSo (see Table 2) in these solvents are considered.20 The total replacement of solvent molecules by the binding sites of the ligand in the coordination sphere of the cation will reverse the trend in the entropies of complexation relative to that of solvation. In fact, this trend is only observed for lithium and sodium in acetonitrile. For potassium and rubidium, both A,S" and A,,I,S" follow the same trend. This is also true for all cations considered in benzonitrile. These observations seem to indicate that for the larger cations in acetonitrile (K+ and Rbf) and for most cations in benzonitrile either the cation is not fully desolvated upon complexation with calixarene esters or even more likely that in the metal-ion complexes the cations are able to enter direct interaction with the solvent. A rather interesting picture emerges by considering the entropies of calixarenate formation, A&?", referred to the process represented in eq 6. M+(g)
+ RCalix(4)(s) -M+RCalix(4)(s)
(6)
which involves the transfer of the cation from the gas phase to the solvated ligand to give the solvated metal-ion calix(4)arenate. The large variations found in the A&" values shown in Table 2 suggest that, unlike metal-ion crypt ate^,^^ considerable differences may be found in the solvation of the ligand relative to the metal-ion complex. Final Remarks From these investigations, the following is concluded: (i) The use of silver electrodes in the double potentiometric method applied for the first time in the determination of highly stable complexes of calixarene esters and metal ions offers a new approach for the determination of stability constants for systems in which the first ligand is highly selective for silver while the second shows little or no selectivity for this cation. (ii) The determination of stability constant data is by no means a trivial process, and therefore, complex stability needs to be checked by several methods since this provides the basis to assess the selective recognition of macrocyclic ligands for metal
cations and provides a guide to the complexes selected for preparative isolation of solids. (iii) Although the thermodynamics of the complexation process reveals the important role played by the solvation of the host, guest, and resulting complex in these processes, a detailed study requires the thermodynamic characterization of the individual species which participate in the binding process in the appropriate solvent. Acknowledgment. The financial support given by the European Commission, DG-XII, 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. The authors thank Prof. S. GlHb, Department of Chemistry, University of Warsaw, Poland, for useful discussions regarding the double competition potentiometric method. References and Notes (1) Pedersen, C. J. J. Am. Chem. Soc. 1967, 89, 2495. (2) Gutsche, C. D.; Muthukrishnan, R. J . Org. Chem. 1978, 43,4905. (3) Arduini, A,; Pochini, A,; Reverber, S.; Ungaro, R. J . Chem. Soc., Chem. Commun. 1984, 981. (4) Chang, S.; Cho, I. Chem. Lett. 1984, 477. ( 5 ) McKervey, M. A.; Seward, E. M.; Ferguson, G.; Ruhl, B.; Harris, S. J. J. Chem. Soc., Chem. Commun. 1985, 388. (6) Ungaro, R.; Pochini, A. In Tropics in Inclusion Phenomena. Calixarenes. A Versatile Class of Macrocyclic Compounds: Vicens, J., Bohmer, V., Eds.; Kluwer Academy: Dordrecht, 1990. (7) Schwing, M. J.; McKervey, M. A. In Topics in Inclusion Phenomena. Calixarenes. A Versatile Class of Macrocyclic Compounds; Vicens, J., Bohmer, V., Eds.; Kluwer Academy: Dordrecht, 1990. (8) Pemn, R.; Harris, S. In Topics in Inclusion Phenomena. Calixarenes. A Versatile Class of Macrocyclic Compounds; Vicens, J., Bohmer, V., Eds.; Kluwer Academy: Dordrecht, 1990. (9) Amaud-Neu, F.; Barrett, G.; Cremin, S.; Deazy, M.; Ferguson, G.; Harris, S . J.; Lough, A. J.; Guerra, L.; McKervey, M. A,; Schwing-Weill, M. J.; Schwinte, P. J . Chem. Soc., Perkin Trans. 2 1992, 1119, and therein references. (10) Danil de Namor, A. F.; Cabaleiro, M. C.; Vuano, B. M.; Salomon, M.; Pieroni, 0. I.; Pacheco Tanaka, D. A,; Ng, C. Y.; Llosa Tanco, M. A,; Rodriguez, N. M.: CBrdenas Garcia, J. D.; Casal, A. R. Pure Appl. Chem. 1994, 66, 435. (1 1) Cox, B. G.; Schneider, H. Coordination and Transport Properties of Macrocyclic Compounds in Solution; Elsevier: New York, 1992. (12) Gutknecht, J.; Schneider, H.; Stroka, J. Inorg. Chem. 1978, 17, 3326. (13) Cox, B. G.; Garcia Rosas, J.; Schneider, H. J . Am. Chem. Soc. 1981, 103, 1384. (14) Danil de Namor, A. F.; Sueros Velarde, F. J. To be submitted. (15) Perrin, D. D.; Armarego, W. L. F. Purification of Laborarory Chemicals, 2nd ed.; Pergamon Press: New York, 1980. (16) Danil de Namor, A. F.; Berroa de Ponce, H. J . Chem. Soc., Faraday Trans. 1987, 83, 1569. (17) Bastos, M.; Hagg, S.; Lonnbro, P.; Wadso, I. J. Biochem. Biophys. Methods 1991, 23, 255. (18) Randzio, S. L.; Suurkuusk, J. In Biological Microcalorimetry; Beezer, A. E., Ed.; Academic Press, London, 1980. (19) Briggner, L. E.; Wadso, I. J . Biochem. Biophys. Methods 1991, 22, 101. (20) 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. (21) Danil de Namor, A. F.; Apaza de Sueros, N.; McKervey, M. A,; Barrett, G.; Amaud Neu, F.; Schwing-Weill, M. J. J. Chem. SOC.,Chem. Commun. 1991, 1546. (22) Abraham, M. H.; Danil de Namor, A. F.; Schulz, R. A. J . Chem. Soc., Faraday Trans. I 1980, 76, 869. (23) Danil de Namor, A. F.; Ghousseini, L. J . Chem. Soc., Faraday Trans. 1984, 80, 2349. (24) Danil de Namor, A. F. J. Chem. SOC.,Faraday Trans. 1988, 84, 2441. (25) Cox, B. G.; Hedwig, G. R.; Parker, A. J.; Watts, D. W. Aust. J. Chem. 1974, 27,477. JP95 19732
