Complexation of Calixarene Derivatives and Lanthanide Cations in

8018

J. Phys. Chem. B 2001, 105, 8018-8027

Complexation of Calixarene Derivatives and Lanthanide Cations in Nonaqueous Media Angela F. Danil de Namor* and Olga Jafou Laboratory of Thermochemistry, Department of Chemistry, School of Physics & Chemistry, UniVersity of Surrey, Guildford, Surrey GU2 5XH, United Kingdom ReceiVed: March 26, 2001; In Final Form: May 30, 2001

The complexing properties of two lower rim calix(4)arene derivatives, namely, 5,11,17,23-tetrakis(1,1dimethylethyl)-25,27-bis[2-(methylthio)ethoxy]-26,28-bis[2-(diethylamine)ethoxy]calix(4)arene (1a) and p-tertbutylcalix(4)arene tetradiisopropylethanoamide (1b) toward lanthanide(III), scandium, and yttrium cations in acetonitrile and in N,N-dimethylformamide at 298.15 K were investigated. 1H NMR complexation experiments established the presence of interactions between the hydrophilic cavity of these macrocycles and these metal cations and revealed the active sites of complexation of these ligands. Conductance measurements were used to (i) establish the concentrations at which the lanthanide trifluoromethanesulfonate salts are fully dissociated 3:1 electrolytes in these solvents and (ii) determine the composition of the metal-ion complex in these solvents. Titration microcalorimetry was used to derive the thermodynamics of complexation of these macrocycles and lanthanide(III) cations in acetonitrile and N,N-dimethylformamide at 298.15 K. No reliable thermodynamic data could be obtained from classical calorimetry due to the slow kinetics observed in the complexation of these calixarene derivatives and these cations in these solvents. Stability constants of 1a were also determined by the competitive potentiometric method using the silver electrodes. Excellent agreement was found between the data derived from calorimetry and those derived by potentiometry. For all the systems investigated, the complexation process between these cations and these ligands was enthalpically controlled. Enthalpy-entropy compensation effects were observed in the complexation of 1a and the different lanthanide(III) cations in acetonitrile and in N,N-dimethylformamide, as suggested by the absence of significant variations in the Gibbs energies of complexation in each case. As far 1b is concerned, a selective behavior was observed for this ligand and the various cations in acetonitrile with the highest stabilities found for gadolinium and europium. The enthalpic and entropic contributions to the Gibbs energy associated with these processes are analyzed. Final conclusions are given.

Introduction Parent calixarenes and their chemically modified derivatives bearing substituents with hard donor atoms at the lower rim have shown to interact with lanthanide cations which are hard cations. Harrowfield and co-workers1-6 have reported the crystal structures of a number of lanthanide complexes with p-tertbutylcalix(n)arenes (n ) 4, 6, 8). The binuclear complex of europium with the cyclic octamer isolated by Harrowfield and co-workers7 from N,N-dimethylformamide was the first lanthanide complex with a parent phenolic calixarene whose structure was established by X-ray crystallography. A number of research articles have been reported in the literature aiming to investigate the extraction properties of calix(4)arene derivatives for lanthanide cations from water to nonaqueous media. A recent article in this area reporting the lanthanide solvent extraction capabilities of a partially substituted calix(4)arene amide derivative as well as the structure determinations of some of its lanthanide(III) complexes has been published by Beer and co-workers.8 Thus, the structures of lanthanum, samarium, ytterbium, and lutetium complexes of this ligand determined by X-ray diffraction studies show the lanthanide cation fully encapsulated in an eight-coordinated polar environment. Their conclusions regarding solvent extraction of lanthanide cations from water to dichloromethane indicated that under optimal experimental conditions (e.g. pH) their ligand exhibited generally very high percentage extract-

abilities for most cations. A recent review article from our group9 clearly indicates that most data reported on the interactions of calixarene derivatives and lanthanide cations are limited to the determination of the stability constants which, in fact have been scarcely reported, while enthalpy and entropy data for these systems are practically absent from the literature. However, to proceed with the thermodynamics of calixarene derivatives and lanthanide cations in nonaqueous solvents, knowledge of the speciation in solution is required. As far as lanthanide(III) cation salts are concerned, few studies have been reported in nonaqueous solvents (some of which are conflicting) which suggest the presence of anion coordination to lanthanide(III) cations in these solvents with certain anions. Thus, Bu¨nzli and co-workers10-14 have investigated the interaction of lanthanide(III) cations and various anions in organic solvents. Using 139La NMR and FT-IR measurements10 these authors showed the presence of significant inner-sphere interactions between La(III) and various anions in methanol at 298 K. These interactions result in the formation of complexes of the type LaXn(3-n)+ and increase in the order

CF3SO3- = ClO4- < Br- < Cl- < NO3-

(1)

Other studies carried out by Bu¨nzli and co-workers include investigations of anhydrous solutions of lanthanide(III) perchlorates in acetonitrile,11,12 europium perchlorate and nitrate in N,N-dimethylformamide13 as well as lanthanide(III) tri-

10.1021/jp011124g CCC: $20.00 © 2001 American Chemical Society Published on Web 07/31/2001

Calixarene Derivatives and Lanthanide Cations

J. Phys. Chem. B, Vol. 105, No. 33, 2001 8019

fluoromethanesulfonates in propylene carbonate.14 In the case of anhydrous Ln(III) perchlorate solutions (Ln ) lanthanide) in acetonitrile (all solutions prepared were 0.05 mol dm-3), these authors showed that both monodentate and bidentate perchlorate anions are in the inner coordination sphere resulting in complicated equilibria between species which differ by the number and type of coordinated perchlorate anions and probably the number of coordinated acetonitrile molecules. Two changes in the average coordination number across the lanthanide series were shown to occur, one between La and Pr and the other at Gd. As far as europium perchlorate and nitrate in N,Ndimethylformamide are concerned, Bu¨nzli et al.,13 demonstrated the absence of inner-sphere interactions between ClO4- and Eu3+ ions. In the case of Eu(NO3)3 in the same solvent, three different species in equilibrium were identified in the concentration range studied. Choppin and co-workers15 have investigated the ability of the trifluoromethanesulfonate anion to coordinate with tervalent lanthanides in anhydrous acetonitrile. FT-IR measurements have shown the presence of two main species in solution at the range of concentrations studied (10-33 mmol dm-3), the equilibrium of which was suggested to be represented by eq 2

Ln(CF3SO3)2+ + CF3SO3 - f Ln(CF3SO3)3

(2)

The same group carried out conductance measurements on two solutions (1.0 and 10.0 mmol dm-3) of Yb(III) triflate at 298 K. The value obtained for the molar conductance at the lower concentration (136 S cm2 mol-1) is in agreement with the accepted value for 1:1 electrolytes, whereas for the more concentrated solution, the molar conductance (58 S cm2 mol-1) suggested the presence of even greater association. These findings are in disagreement, however, with the data reported by Seminara and Rizzarelli.16 According to their results, coordination of trifluoromethanesulfonate anions with lanthanide cations does not occur. It is quite clear from the above discussion that the nature of the ionic species formed in solution is dependent on the concentrations used. Therefore, a systematic study in the appropriate solvent must be carried out prior to the determination of the thermodynamics associated with the complexation process involving these cations. The aims of this paper are as follows: (i) To establish the presence or absence of interactions between two calixarene derivatives (1a 1b) and the lanthanide cations as well as the sites of interaction in acetonitrile (MeCN) and in N,N-dimethylformamide (DMF) at 298 K using 1H NMR studies whenever possible.

(ii) To determine the concentration range at which the lanthanide(III) cations are predominantly in solution and to establish the composition of the complexes formed in the appropriate solvent.

(iii) To characterize thermodynamically the complexation process involving calixarene derivatives (1a 1b) and lanthanide cations in dipolar aprotic media using potentiometry and titration calorimetry (macro and micro). Some of the data reported in this paper also include cations such as scandium and yttrium. Although the radius of the former is considerably smaller than those of any of the lanthanides, its behavior is intermediate between that of aluminum and that of lanthanides.17 As far as yttrium is concerned, its size is similar to gadolinium and terbium. Its complexation chemistry therefore closely resembles that of the late-middle lanthanides and is generally found in Nature with the lanthanides.17,18 Experimental Procedure Chemicals. Scandium (99%), yttrium (98%), lanthanum (99.9%), praseodymium (98%), neodymium (98%), samarium (98%), europium (98%), gadolinium (98%), terbium (98%), dysprosium (98%), holmium (98%), erbium (98%), ytterbium (99.9%), lutetium (98%) trifluoromethanesulfonate salts, and silver perchlorate (99.9%) were all purchased from Aldrich Chemical Co. These were dried over P4O10 under vacuum for several days before use. Potassium chloride, Fisher Chemical Company, was twice recrystallized from distilled water and dried in a drying pistol at 393 K for 3 days before use. Tris(hydroxymethyl)aminomethane, ultrapure grade 99.9% from Aldrich and tetra-n-butylammonium perchlorate (TBAP) (electrochemical grade g 99%, Fluka), were used without any further purification. The macrocycles 1a and 1b were synthesized at the Thermochemistry Laboratory according to the procedures reported in the literature.19-21 They were dried in a drying pistol at 8090 °C for several days before use. Acetonitrile,22 HPLC grade from Fisher was refluxed in a nitrogen atmosphere and distilled from calcium hydride. Only the middle fraction of the solvent was collected. Its water content, as determined by the Karl Fisher titration method, was not more than 0.02%. N,N-dimethylformamide,22 HPLC grade from Fisher was dried over 4 Å molecular sieves and subsequently distilled under reduced pressure. The middle fraction of the solvent was used. Its water content determined by the same method as described for acetonitrile was less than 0.02%. Deuterated chloroform (CDCl3), acetonitrile (CD3CN), N,Ndimethylformamide (DMF-d7) and tetramethylsilane (TMS) were purchased from Aldrich Chemical Co. 1H NMR Measurements. 1H NMR measurements were recorded at 298 K using a Bruker AC-300E pulsed Fourier transform NMR spectrometer. Typical operating conditions for routine proton measurements involved ‘pulse’ or flip angle of 30°, spectral frequency (SF) of 300.135 MHz, delay time of 1.60 s, acquisition time (AQ) of 1.819 s and line broadening of 0.55 Hz. Solutions of the samples in question were prepared in the appropriate deuterated solvent and then placed in 5 mm NMR tubes using TMS (tetramethylsilane) as the internal reference. The complexing behavior of calix(4)arene derivatives toward tervalent cations was studied by adding a metal-ion salt of known concentration (lower than 10-3 mol dm-3) into the NMR tube already containing a known amount of the ligand in question (concentration range from 9.5 × 10-4 to 1.1 × 10-3 mol dm-3) dissolved in the appropriate solvent. Stepwise additions of varying amount of the metal cation salt were undertaken until chemical shift changes ceased. Proton chemical shifts of the free and complex ligand were recorded. All measurements were carried out at 298 K.

8020 J. Phys. Chem. B, Vol. 105, No. 33, 2001

Danil de Namor and Jafou

TABLE 1: 1H NMR Chemical Shift Changes (ppm) of 1a with the Addition of Y(CF3SO3)3 in CD3CN at 298 K mole ratio (Y3+/L)

ArH

Hax

Heq

free ligand 0.09 0.18 0.27 0.36 0.44 0.53 0.62 0.71 0.80 0.89 1.01 1.15 1.33

7.00, 6.67 -0.03, +0.02 -0.07, +0.04 -0.10, +0.06 -0.13, +0.08 -0.14, +0.09 -0.15, +0.10 -0.16, +0.10 -0.16, +0.10 -0.16, +0.10 -0.15, +0.10 -0.15, +0.10 -0.15, +0.10 -0.15, +0.10

4.20 -0.01 -0.02 -0.04 -0.05 -0.06 -0.07 -0.08 -0.10 -0.11 -0.12 -0.11 -0.12 -0.12

3.04 +0.00 +0.01 +0.02 +0.03 +0.04 +0.05 +0.05 +0.06 +0.07 +0.07 +0.07 +0.07 +0.07

O-CH2CH2S O-CH2CH2N O-CH2CH2S O-CH2CH2N NCH2CH3 SCH3 NCH2CH3 3.95 -0.01 -0.02 -0.04 -0.04 -0.05 -0.05 -0.06 -0.06 -0.06 -0.06 -0.06 -0.06 -0.06

3.63 +0.04 +0.09 +0.14 +0.19 +0.27 +0.27 +0.31 +0.34 +0.37 +0.39 +0.39 +0.39 +0.39

Conductance Measurements. For these measurements, a Wayne-Kerr Autobalance Universal Bridge, type B642 was used. For the determination of the cell constant of the conductivity cell, an aqueous solution of KCl (0.1000 mol dm-3) was added by steps to the deionized water containing cell. The cell was kept in a thermostated bath at 298.15 K. The subsequent conductances following each injection were then recorded after allowing a sufficient time to ensure the system was properly mixed and the cell had attained the temperature of the bath. The conductance of the water was measured in advance and subtracted from each conductance change recorded. The corresponding molar conductances (S cm2 mol-1) were calculated from the equation of Lind, Zwolenik, and Fuoss.23 The molar conductance of KCl was used to calculate corresponding values of the specific conductance and then the cell constant. Conductance measurements at different concentrations (9.0 × 10-6 - 3 × 10-4 mol dm-3) of tervalent cation salts as trifluoromethanesulfonate were measured in acentonitrile and in N,N-dimethylformamide at 298.15 K in order to establish the concentration limit at which the tervalent cation salts were predominantly behaving as 3:1 electrolytes in solution. For conductometric titrations, the cell was filled with the metal-ion salt solution (concentration range from 9.0 × 10-6 to 1.2 × 10-5 mol dm-3) prepared in the appropriate solvent (∼20 cm-3), reweighed, sealed and left in a thermostated bath at 298.15 K to reach thermal equilibrium, while dry nitrogen was passed through the solution. After thermal equilibrium was reached, the ligand solution (concentration range from 8.5 × 10-4 to 1.0 × 10-3 mol dm-3) was added in steps, using a hypodermic syringe. All measurements were taken at 298.15 K. Determination of the Stability Constant from Potentiomety. The method based on a competitive reaction between the tervalent cation and an auxiliary cation (Ag+) as described elsewhere24-27 was used. Calorimetric Titrations. For macrocalorimetric titrations, the Tronac 450 calorimeter was used as an isoperibol titration calorimeter. The reproducibility of the apparatus was checked by (i) calibrating the buret and (ii) chemical calibration using as standard reaction the enthalpy of protonation of an aqueous solution of tris(hydroxymethyl)aminomethane (THAM) in hydrochloric acid at 298.15 K. For this purpose, an aqueous solution of THAM placed in the buret was incrementally titrated into the reaction vessel containing an aqueous solution of hydrochloric acid (50 mL). The value of -47.51 ( 0.15 kJ ¨ jelund mol-1 is in excellent agreement with that reported by O and Wadso¨28 (-47.49 ( 0.04 kJ mol-1) using an LKB reaction calorimeter. For microcalorimetric titrations, the four channel heat conduction microcalorimeter (Thermometric 2277) designed by

3.06 -0.04 -0.08 -0.12 -0.15 -0.17 -0.18 -0.20 -0.21 -0.21 -0.22 -0.21 -0.22 -0.22

2.75 +0.06 +0.14 +0.19 +0.31 +0.37 +0.44 +0.50 +0.56 +0.61 +0.65 +0.65 +0.65 +0.65

2.39 +0.06 +0.12 +0.21 +0.29 +0.36 +0.37 +0.48 +0.56 +0.61 +0.65 +0.66 +0.66 +0.66

2.01 -0.00 -0.03 -0.05 -0.04 -0.05 -0.05 -0.05 -0.05 -0.05 -0.05 -0.05 -0.05 -0.05

0.82 +0.03 +0.06 +0.10 +0.13 +0.15 +0.18 +0.20 +0.23 +0.25 +0.26 +0.26 +0.26 +0.26

t-butyl 1.06, 0.81 -0.02, +0.01 -0.04, +0.03 -0.07, +0.05 -0.08, +0.06 -0.10, +0.07 -0.10, +0.08 -0.10, +0.08 -0.11, +0.07 -0.11, +0.07 -0.10, +0.07 -0.10, +0.07 -0.10, +0.07 -0.10, +0.07

Suurkuusk and Wadso¨29 was used. Electrical and chemical calibrations were carried out to check the reliability of the equipment.30,31 For calorimetric titrations using the Tronac 450 calorimeter, a solution of the lanthanide salt (concentration range from 9.5 × 10-5 to 3.0 × 10-4 mol dm-3) in the appropriate solvent was placed in the buret. The macrocyclic ligand in question (50 cm3, concentration range from 1.0 × 10-4 to 1.0 × 10-3 mol dm-3) in the same solvent was pipetted into the reaction vessel. The whole system was then immersed in a thermostated water bath at 298.15 K and allowed to reach thermal equilibrium. The titrating solution was subsequently added at a fixed buret delivery rate and the time of each addition was recorded. A chart recorder was used to monitor the reaction. An electrical calibration experiment was carried out after every titration experiment. Blank experiments were carried out in all cases to account for the heat of dilution effects resulting from the addition of the metal-ion salt solution into the appropriate solvent. In a typical microcalorimetric titration experiment, the reaction vessel was charged with 2.8 cm3 of the macrocycle solution (concentration range from 7.5 × 10-5 to 1.0 × 10-3 mol dm-3) in the appropriate solvent. The lanthanide salt solution (concentration range from 9.5 × 10-5 to 1.3 × 10-4 mol dm-3) was subsequently injected incrementally using a 0.5 cm3 gastight motor driven Hamilton syringe. In each titration experiment about 15 injections were made at time intervals of 1 h. Corrections for enthalpy of dilution of the titrant in the solvent were carried out in all cases. A computer program for the calculation of log Ks and ∆cH was developed at the University of Surrey.32 Results and Discussion 1H

NMR Titrations of Calixarene Derivatives and Tervalent Cations. 1H NMR titrations were carried with both ligands (1a) and (1b) and tervalent cations in CD3CN at 298 K. As far as the former ligand is concerned, 1H NMR titrations were also carried out in DMF-d7 at 298 K. 1H NMR data for 1a and tervalent cations in CD3CN and DMF-d7 at 298 K are now discussed. 1H NMR Titrations of 1a and Tervalent Cations in CD CN 3 at 298 K. Table 1 shows the chemical shift changes with respect to the free ligand observed in the titration of 1a with yttrium in CD3CN. Only the protons of the nitrogen substituents are deshielded upon complexation, while the sulfur substituents move upfield. The most marked downfield shifts seem to be those experienced by the methylene protons on either side of nitrogen, followed by the methylene protons adjacent to the oxygen in the OCH2CH2N system. The ligand therefore uses



TABLE 2: Summary of 1H NMR Chemical Shift Changes (ppm) for 1a Metal-Ion Complexes with Respect to the Free Ligand in CD3CN at 298 K proton

∆δ(ppm) Sc3+ complex

∆δ(ppm) Y3+ complex

∆δ(ppm) La3+ complex

∆δ(ppm) Eu3+ complex

ArH Hax Heq O-CH2CH2S O-CH2CH2N O-CH2CH2S O-CH2CH2N NCH2CH3 SCH3 NCH2CH3 t-butyl

-0.19, +0.13 -0.11 +0.09 -0.06 +0.42 -0.23 +0.66 +0.66 -0.07 +0.27 -0.13, +0.10

-0.15, +0.10 -0.12 +0.07 -0.06 +0.39 -0.22 +0.65 +0.66 -0.05 +0.26 -0.10, +0.07

-0.13, +0.08 -0.12 +0.07 -0.05 +0.38 -0.20 +0.65 +0.65 +0.00 +0.27 -0.09, +0.06

-0.19, +0.14 -0.11 +0.10 -0.07 +0.43 -0.24 +0.67 +0.67 -0.01 +0.28 -0.13, +0.10

TABLE 3: 1H NMR Chemical Shift Changes (ppm) of 1a with the Addition of Y(CF3SO3)3 in DMF-d7 at 298 K mole ratio (Y3+/L)

ArH

Hax

Heq

free ligand 0.13 0.26 0.43 0.60 0.77 0.90 1.07 1.24 1.63 2.29

7.24, 6.63 -0.01, +0.01 -0.01, +0.01 -0.01, +0.02 -0.01, +0.02 -0.02, +0.02 -0.02, +0.02 -0.02, +0.02 -0.02, +0.02 -0.02, +0.03 -0.02, +0.03

4.43 -0.00 -0.00 -0.00 -0.00 -0.00 -0.00 -0.00 -0.00 -0.00 -0.00

3.26 +0.00 +0.00 +0.00 +0.00 +0.00 +0.00 +0.00 +0.00 +0.00 +0.00

O-CH2CH2S O-CH2CH2N O-CH2CH2S O-CH2CH2N NCH2CH3 SCH3 NCH2CH3 4.20 -0.00 -0.00 -0.00 -0.00 -0.00 -0.01 -0.01 -0.01 -0.01 -0.01

3.82 +0.02 +0.02 +0.03 +0.03 +0.03 +0.04 +0.04 +0.04 +0.04 +0.04

its N, O, and S donor sites in order to complex the cation, although the nitrogen atoms seem to interact more strongly. As the complexation process proceeds, the two aromatic and tertbutyl signals move progressively closer and indeed remain very close together at the end of the titration. The fact that these signals move to a similar resonance position indicates the formation of a more symmetrical ‘cone’ conformation as complexation with the metal cation forces the pendant arms to move closer. This is also supported by the conformational changes which occur in the bridging methylene protons on addition of the yttrium salt. Thus, the axial protons are shielded while the equatorial protons are deshielded, resulting in a shift difference between the pair of doublets of 0.98 ppm which resembles a more symmetrical parent-like ‘cone’ conformation33 when compared with 1.16 ppm in the free ligand. After the 1:1 ratio is reached, all chemical shifts virtually are nonexistent. As far as the 1H NMR titrations for 1a and scandium, lanthanum and europium in CD3CN are concerned, similar patterns are observed to those discussed for yttrium as shown in Table 2. In DMF-d7, the interaction of 1a with yttrium gives rise to small shift changes in the proton signals relative to those of the free ligand in this solvent as shown in Table 3. During the complexation titration experiment all the protons on the amine lower rim substituents are deshielded, while the methylethylthio lower rim protons move upfield, similar to that observed in the same titration experiment carried out in CD3CN. However, it appears that only the nitrogen and oxygen donor atoms of the amine pendant arms provide the sites of complexation in DMFd7. As far as this cation and this ligand is concerned, the 1H NMR results do not provide any evidence that the methylethylthio pendant arms participate significantly in the complexation process, since no considerable chemical shift changes are recorded for the protons adjacent to the sulfur donor atoms in this solvent. The conformational changes described above for the same titration experiment in CD3CN also take place here although to a lesser extent. Thus, the two aromatic and two

3.24 -0.00 -0.00 -0.00 -0.00 -0.01 -0.01 -0.01 -0.01 -0.01 -0.01

2.93 +0.10 +0.11 +0.12 +0.12 +0.13 +0.13 +0.14 +0.14 +0.14 +0.14

2.64 +0.02 +0.03 +0.04 +0.05 +0.05 +0.05 +0.06 +0.06 +0.06 +0.06

2.27 0.00 -0.00 -0.00 -0.00 -0.00 -0.00 -0.00 -0.00 -0.00 -0.00

1.06 +0.01 +0.01 +0.02 +0.02 +0.02 +0.02 +0.03 +0.03 +0.03 +0.03

t-butyl 1.32, 0.90 -0.01, +0.01 -0.01, +0.01 -0.01, +0.01 -0.01, +0.01 -0.01, +0.01 -0.01, +0.01 -0.01, +0.02 -0.01, +0.02 -0.01, +0.02 -0.02, +0.02

TABLE 4: 1H NMR Chemical Shifts (ppm) of 1b and Its Metal-Ion Complexes in CD3CN at 298 K proton ArH Hax Heq CH2CON (O)NCH(CH3)2 NCH(CH3)2 (O)NCH(CH3)2 NCH(CH3)2 t-butyl

δ(ppm) ∆δ(ppm) ∆δ(ppm) ∆δ(ppm) free ligand Y3+complex La3+complex Lu3+complex 7.11 5.35 3.22 5.00 3.95 3.51 1.36 1.13 1.17

+0.49 +0.09 +1.64 -1.23 +0.11 +0.09 +0.15 +0.14 +0.05

+0.50 +0.05 +1.88 -1.14 +0.05 +0.08 +0.18 +0.11 +0.05

+0.49 +0.07 +1.67 -1.24 +0.18 +0.09 +0.13 +0.15 +0.05

tert-butyl signals become closer together as shown by the decrease in the shift difference between each pair of signals shown in Table 3, although this movement is less pronounced in the case of this solvent. Further conformational changes are also observed in the bridging methylene protons during the complexation experiment. The axial protons are again shielded while the equatorial protons become less shielded, as in CD3CN. A direct consequence of that is that there is hardly any shift difference between each pair of doublets of the ArCH2Ar system, from the free to the complex ligand. This is consistent with the overall small chemical shift changes that accompany the complexation process in this solvent, indicating that less pronounced interactions take place in this medium. 1H NMR Titrations of 1b and Tervalent Cations in CD3CN at 298 K. Table 4 lists the 1H NMR chemical shifts of 1b and its yttrium, lanthanum and lutetium complexes in CD3CN at 298 K. Comparison of the 1H NMR spectrum of the free ligand with those of the Y3+ and Lu3+ complexes shows considerable changes in most signals. The most significant observation is the presence of two sets of aromatic signals (a pair of doublets) replacing the singlet of the equivalent aromatic protons in the 1H NMR spectrum of the free ligand. This may suggest that the aromatic rings adopt a 2-fold ‘cone’ symmetry in the presence of these metal cations. In the case of the La3+

8022 J. Phys. Chem. B, Vol. 105, No. 33, 2001 complex of 1b however, the macrocycle maintained its 4-fold symmetry because the 1H NMR spectrum shows only a singlet for the aromatic protons. Beer and co-workers34 have reported the crystal structures of a series of transition metal complexes of an analogous derivative known as p-tert-butylcalix(4)arene tetraethanoamide whose structure is very close to 1b (ethyl groups replace the isopropyl ones in 1b). They found that the size of the hydrophilic cavity defined by the eight oxygen atoms (four ethereal and four carbonyl) is such that it fits the stereochemical requirements of the larger cations such as K+ and Pb2+ (cation sizes, 1.38 and 1.19 Å, respectively, for coordination number 6).35 It would be rather difficult for the cavity to adjust its size so as to allow small transition metal-ions to fit in with eight equivalent bond lengths, mainly due to the relative inflexibility of the ether oxygen atoms which are unable to approach each other. Unlike these, the carbonyl oxygens are more flexible and therefore more able to move closer in order to interact with the smaller transition-metal cations. Thus, cations such as Zn2+, Fe2+ and Cu2+ (cation sizes 0.74, 0.61, and 0.73 Å respectively for coordination number 6)35 interact strongly with the four carbonyl oxygen atoms while the interactions between these cations and the four ether oxygen atoms are much weaker. A distortedoctahedral structure was found for Ni2+ (cation size 0.69 Å for coordination number 6)35 where this cation is placed close to only six oxygen atoms (three ethereal and three carbonyl). Similarly, in these studies, the difference in the symmetry of the metal-ion complexes revealed by 1H NMR spectroscopy could reflect the size cation effect. Thus, La3+ is first in the lanthanide series, and therefore, the biggest cation with an ionic radius of 1.06 Å,36 able to form a 4-fold symmetric complex with 1b. Lu3+ is last in the series and therefore the smallest cation with an ionic radius of 0.85 Å36 forming a 2-fold symmetric complex with 1b. The same situation was observed for Eu3+ and Y3+ whose ionic radii are 0.95 and 0.88 Å, respectively.36 Further conformational changes which occur as complexation with metal-ions take place are also observed in the bridging methylene protons. In all cases, both axial and equatorial protons are deshielded which results in a considerable decrease in the shift difference between the pair of doublets from the free to the complex ligand from 2.12 to 0.57 and 0.52 ppm in the cases of yttrium and lutetium, respectively. As previously stated33 the ∆δ(ppm) values between the axial and equatorial protons in calix(4)arenes depend on their conformation and these are generally ca. 0.5 ( 0.1 ppm for a system in the ‘flattened cone’ conformation similar to the values observed above. A direct consequence of the conformational changes of the ligand upon complexation is that the donor atoms on the pendant arms of the nearly flattened aromatic rings are able to interact more strongly with the metal cation. The largest chemical shift change is observed for the equatorial protons which move upfield. The next most significant shift is observed for the methylene protons of the acetamide groups CH2CON which exhibit high upfield shifts of 1.23, 1.14, and 1.24 ppm for Y3+, La3+ and Lu3+ respectively. These may indicate that the amide and possibly the ethereal oxygens are strongly involved in the complexation process. This upfield shift can be explained by assuming that the CO groups move inward the hydrophilic cavity toward the metal cation upon complexation which brings the CH2 protons toward the exterior of the hydrophilic cavity where they experience a shielding effect in the presence of the aromatic rings. The nitrogen donor atoms seem to be less involved in the complexation since the adjacent terminal protons exhibit


Figure 1. Plot of Λm against c1/2 for Yb(CF3SO3)3 solutions in acetonitrile at 298.15 K.

small shift changes. In all the above 1H NMR titration experiments, the 1:1 stoichiometry was clearly demonstrated because all the signals remained virtually unchanged after the salt/ligand ratio reached the unity value. To (i) determine the concentration range to be safely used and ensure that lanthanide(III) triflates are present as 3:1 electrolytes in the appropriate solvent (ii) establish the composition of the metal-ion complexes, conductance measurements were performed and these are now discussed. Solution Properties of Lanthanide(III) Trifluoromethanesulfonate Salts in Acetonitrile and in N,N-Dimethylformamide at 298.15 K. Preliminary measurements of the molar conductivities of acetonitrile solutions containing lanthanide(III) triflates suggested that at higher concentrations, the Λm decreases considerably reflecting that possibly some interaction takes place between lanthanide(III) cations and the triflate anion. However, at the initial points of the curves (corresponding to the salt concentrations used in the titration experiments, see Experimental Section) reasonable linear relationships were obtained. A representative example of this statement is shown in Figure 1 where the molar conductances, Λm, are plotted against the square root of the concentration for ytterbium triflate solutions in acetonitrile at 298.15 K. It is therefore evident from the straight line observed that ion-pair formation does not play a significant role in the range of electrolyte concentrations studied (below 1.0 × 10-4 mol dm-3). This is also corroborated by the conductometric titrations discussed below which show that Λm values are within the accepted range for 3:1 electrolytes in acetonitrile (above 340 S cm2 mol-1) and in N,N-dimethylformamide (above 200 S cm2 mol-1).37 These were performed with the aim of establishing the composition of the metal-ion complexes in these solvents and are discussed in the following section. Conductometric Titrations of Lanthanide(III) Cations and Calix(4)arene Derivatives in Dipolar Aprotic Media at 298.15 K. Measurements of the variation of electrical conductance with the concentration of salt and ligand can be used to assess the interactions taking place, determine the strength and stoichiometry of the complex as well as its stability constant.26 In this work, these measurements were used (i) to establish the composition of the complex and (ii) to assess qualitatively the strength of the complex from the shape of the conductometric titration curve. Thus, for very little or no complexation, plots were characterized by a slight slope and without any indication of a change in slope at any given mole ratio. For complexes of moderate stability, plots with a well-defined change in curvature at the stoichiometry of the reaction were obtained. Plots formed by two straight lines intersecting at the stoichiometry of the reaction were indicative of strong complexation.


Figure 2. Conductimetric titration curve for the addition of 1a to Nd(CF3SO3)3 in acetonitrile at 298.15 K.

Figure 3. Conductimetric titration curve for the addition of Tb(CF3SO3)3 to 1a in acetonitrile at 298.15 K.

Conductometric titrations of lanthanide(III) trifluoromethanesulfonate salts (neodymium, samarium, gadolinium, dysprosium and terbium) with 1a, in acetonitrile at 298.15 K demonstrated that the curves are a result of a combination of two linear segments with a well-defined change in curvature at the 1:1 stoichiometry which suggest the formation of relatively strong complexes in this solvent. A representative example is shown in Figure 2 where data for the titration of Nd(CF3SO3)3 with 1a are given. Because the electrolyte was placed in the conductance vessel a decrease in the conductivity on addition of the ligand would be expected due to the large size of the complexed metal-ion compared with the free solvated ion. The opposite observation suggests that the lanthanide cation is much more solvated than the new complexed cation, therefore less mobile. An increase in electrolyte conductance has also been observed by Danil de Namor and co-workers38,39 on the addition of crown ethers (1-aza-12-crown-4 and 15-crown-5) to lithium salts (LiBF4 and LiCF3SO3) in acetonitrile and propylene carbonate. This conductivity enhancement was also attributed to the fact that lithium coronand electrolytes were less solvated in these solvents than the lithium salts, a fact which was further verified by solution thermochemical studies of these electrolytes in these solvents. The behavior of these systems in the reverse experimental order, i.e., placing the ligand in the conductometric vessel and following by the stepwise addition of the electrolyte in question was explored. Thus, the conductometric titration curve for the addition of Tb(CF3SO3)3 to 1a in acetonitrile is shown in Figure 3. As expected the conductivity of the initial solution (neutral ligand) is close to zero. As the cation is added an expected increase in the conductivity is observed. A reasonably clear change in the slope of the curve at a metal/ligand ratio 1


Figure 4. Conductimetric titration curve for the addition of 1a to La(CF3SO3)3 in N,N-dimethylformamide at 298.15 K.

demonstrates again that a ligand unit interacts with a cation unit. However, this is not the recommended way of performing a conductometric titration because an increase in conductance is bound to occur due to the addition of the electrolyte solution and therefore it is usually difficult to draw any conclusions regarding complex formation in this case. To demonstrate the medium effect on the complexation of 1a and these cations, conductometric titrations were performed in N,N-dimethylformamide (DMF) at 298.15 K. A representative example is given in Figure 4 where the changes in conductance upon complexation of 1a and La3+ against the ligand: metal cation molar ratios in DMF are relatively small compared to the changes observed in acetonitrile. No significant breaks in the curvature of the plots are observed indicating the formation of low stability complexes. In this solvent, the molar conductance decreases as the metal-ion complex is formed which may indicate that the complexes are better solvated than the free salts and therefore less mobile resulting in a decrease in conductance. Thermodynamics of Complexation. (i) Thermodynamics of Complexation of 1a and TerValent Cations in Acetonitrile at 298.15 K. The complexation between 1a and tervalent cations in acetonitrile was investigated at 298.15 K. For this purpose, titration microcalorimetry, classical macrocalorimetry and competitive potentiometry using silver electrodes were used. While the latter technique was used to determine the stability constants, Ks, calorimetry (macro and micro) was used for the determination of Ks (expressed as log Ks) and the enthalpy of complexation, ∆cH°. To prove the validity of the nonlinear minimization algorithm used, an exhaustive search of the configuration space of the problem was carried out in each case. According to this, the error function, U, was evaluated for all possible discrete values of Ks within a given range, as illustrated in Figure 5 for the microcalorimetric titration of 1a and erbium in acetonitrile at 298.15 K. Figure 6 shows a microcalorimetric plot for the titration of the same system at the same temperature in which the potential is plotted against time. This graph is typical of a 1:1 complexation reaction with an intermediate value of stability constant. Because on addition of an excess of the metal-ion solution to the calorimetric vessel containing the ligand solution no significant heat was generated, it was assumed that 1:1 complexes were formed. This is further reinforced by the availability of 1H NMR and conductance measurements. We have also evidence from the conductometric titration curves that in the range of concentrations used, the metal-ion complexes are fully dissociated in this solvent. Thermodynamic data for these systems therefore refer to the process described by eq 3

M3 + (s) + 1a(s) f M3 + 1a(s)

(3)


Figure 5. Assumed values of log U against log KS for the microcalorimetric titration of 1a with erbium in MeCN at 298.15 K.

Figure 6. Plot of potential Vs. time for the microcalorimetric titration of 1a with erbium in acetonitrile at 298.15 K.

Danil de Namor and Jafou Stability constants (log Ks), derived standard Gibbs energies, ∆cG°, enthalpies, ∆cH° and entropies, ∆cS°, for the complexation of 1a and tervalent cations derived from titration microcalorimetry are listed in Table 5. Also included in this table are stability constant data obtained from potentiometry. We found no agreement between these data and those obtained from classical macrocalorimetry. This is due to the fact that slow kinetics were observed in the complexation of these cations and this ligand in acetonitrile as shown by conductance, potentiometric as well as calorimetric studies possibly due to the slow removal of the solvation shell of the highly charged lanthanide(III) cations. Therefore, classical calorimetry was not suitable for measurements involving these systems and this was clearly reflected in the relatively large errors in the thermodynamic parameters measured by classical macrocalorimetry. Titration microcalorimetry has the capacity of measuring heats associated with slow processes.31,40 Comparison of data derived from microcalorimetry with those measured potentiometrically shows that good agreement is found between the two sets of data derived from the two independent methods. Thus, these data were combined and the average log Ks values were used in order to calculate ∆cG° for these systems as shown in Table 5. ∆cH° and ∆cS° values are also shown in this table. A general analysis of the thermodynamic parameters shows that the complexation process is favored in terms of enthalpy (∆cH° < 0) but not in terms of entropy (∆cS° < 0) in all the above systems. Therefore, the complexation process is enthalpically controlled. The data in Table 5 show that similar ∆cG° values are obtained for 1a and the different cations studied in acetonitrile. A linear relationship was in fact found when ∆cH° values were plotted against ∆cS° values as shown in Figure 7. Therefore, these data suggest the presence of enthalpy-entropy compensation effects possibly due to solvent reorganization upon complexation. These effects have been discussed in detail by Grunwald and Steel.41 The slope of this plot representing the experimental temperature (298 K) is 287 K, the correlation coefficient is 0.993 while the intercept of the fitted line, ∆cG°, is -30.03 ( 3.32 kJ mol-1. According to Choppin,42 a good linear relationship is often found

TABLE 5: Thermodynamic Parameters of Complexation of 1a and Lanthanide(III) Cations in Acetonitrile at 298.15 K Determined by Titration Microcalorimetry and Potentiometry cation Sc3+ Y3+ La3+ Pr3+ Nd3+ Sm3+ Eu3+ Gd3+ Tb3+ Dy3+ Ho3+ Er3+ Yb3+ a

4.49 ( 0.03a 4.42 ( 0.06b 4.39 ( 0.05a 4.45 ( 0.07b 4.59 ( 0.01a 4.55 ( 0.09b 4.66 ( 0.02a 4.59 ( 0.04b 4.59 ( 0.02a 4.09 ( 0.04a 4.19 ( 0.09b 4.80 ( 0.07a 4.84 ( 0.07b 5.07 ( 0.05a 5.05 ( 0.05b 3.68 ( 0.02a 3.65 ( 0.08b 3.58 ( 0.02a 3.89 ( 0.01a 3.93 ( 0.07a 3.98 ( 0.06b 4.39 ( 0.07a 4.30 ( 0.09b

} } } } } } } } } }

∆cG° (kJ mol-1)

∆cH° (kJ mol-1)

∆cS (J K-1 mol-1)

4.46 ( 0.03

-25.43 ( 0.15c

-164.2 ( 0.5a

-465

4.42 ( 0.04

-25.23 ( 0.23c

-94.4 ( 0.3a

-232

4.57 ( 0.01

-26.09 ( 0.06c

-188.3 ( 0.9a

-544

4.62 ( 0.02

-26.37 ( 0.10c

-172.5 ( 0.6a

-490

4.14 ( 0.04

-26.20 ( 0.11a -23.63 ( 0.21c

-179.1 ( 0.6a -164.5 ( 0.8a

-513 -473

4.82 ( 0.05

-27.51 ( 0.28c

-142.1 ( 0.9a

-384

5.06 ( 0.04

-28.88 ( 0.20c

-146.4 ( 0.5a

-394

3.67 ( 0.02

-20.95 ( 0.11c

-187.2 ( 0.9a

-558

3.96 ( 0.05

-20.46 ( 0.09a -22.20 ( 0.06a -22.60 ( 0.26c

-206.5 ( 0.7a -143.1 ( 0.4a -170.9 ( 0.9a

-624 -406 -497

4.35 ( 0.06

-24.83 ( 0.31c

-156.5 ( 0.2a

-442

log Ks

Microcalorimetric data. b Potentiometric data. c Calculated from the average of log Ks values (calorimetric and potentiometric).



Figure 7. Plot of ∆cH° Vs ∆cS° for the complexation of lanthanide(III) cations by 1a in acetonitrile at 298.15 K. Figure 9. Variation of log KS with the ionic radius for some lanthanide(III) complexes of 1b in acetonitrile at 298.15 K.

Figure 8. Enthalpy and entropy complexation changes of 1a and tervalent cations in acetonitrile as a function of the ionic size.

TABLE 6: Thermodynamic Parameters for the Complexation of Lanthanide(III) Cations and 1b in Acetonitrile at 298.15 K by Titration Microcalorimetry cation

log Ks

Sc3+

3.75 ( 0.03 4.96 ( 0.05 3.67 ( 0.06 3.86 ( 0.06 5.09 ( 0.04 5.21 ( 0.10 5.24 ( 0.09 4.66 ( 0.06 4.69 ( 0.06 4.63 ( 0.09 4.29 ( 0.10 3.78 ( 0.04

Y3+ Pr3+ Nd3+ Sm3+ Eu3+ Gd3+ Tb3+ Dy3+ Ho3+ Er3+ Yb3+

∆cG° (kJ mol-1)

∆cH° (kJ mol-1)

-21.40 ( 0.17 -100.7 ( 0.5 -28.31 ( 0.28 -70.8 ( 0.2 -20.95 ( 0.34 -108.9 ( 0.7 -22.03 ( 0.34 -90.8 ( 0.7 -29.05 ( 0.23 -86.4 ( 0.3 -29.74 ( 0.57 -72.8 ( 0.4 -29.91 ( 0.51 -79.1 ( 0.6 -26.60 ( 0.34 -101.8 ( 0.4 -26.77 ( 0.34 -69.1 ( 0.3 -26.43 ( 0.51 -95.8 ( 0.5 -24.49 ( 0.57 -81.2 ( 0.2 -21.58 ( 0.23 -83.7 ( 0.3

∆cS° (J K-1 mol-1) -266 -143 -295 -231 -192 -144 -165 -252 -142 -233 -190 -208

between the values of the enthalpy and the entropy of complexation for the same ligand and the different lanthanides. The variation in these two thermodynamic parameters (∆cH° and ∆cS°) with the cation size is quite striking as shown in Figure 8 which is a plot of the ∆cH° against the cation radii. This plot shows several breaks with points of maximum and minimum exothermicity which are compensated by more and less favorable entropies respectively along the series. (ii) Thermodynamics of Complexation of 1b and TerValent Cations in Acetonitrile at 298.15 K. The complexation of 1b and tervalent cations was investigated in MeCN at 298.15 K using titration microcalorimetry and classical (macro) calorimetry. With the latter technique the slow kinetics of the process impeded us to derive reliable thermodynamic data. Therefore, the data reported in Table 6 are the result of titration microcalorimetry. Inspection of Table 6 shows that the complexation process involving this ligand and the various lanthanide cations can be differentiated in terms of stability constant. Despite the

Figure 10. Variation of ∆cH° with the ionic radius for some lanthanide(III) complexes of 1b in acetonitrile at 298.15 K.

small variation in the size of lanthanide cations as assessed by their ionic radii (r) (from Pr3+; r ) 1.01 Å to Yb3+; r ) 0.86 Å, the size variation is 0.15 Å),36 1b is able to interact selectively with these cations. Indeed, there is a change in stability of about 1.6 log units between the weakest (log Ks Pr3+1b ) 3.67) to the strongest (log Ks Gd3+1b ) 5.24 = log Ks Eu3+1b ) 5.21) lanthanide complex. In fact, a plot of log Ks values against the ionic radii (Figure 9) shows a selectivity peak similar to that observed for this ligand and alkali-metal cations in this solvent. It should be noted that among the latter cations, this ligand shows the highest selectivity for sodium27 in acetonitrile at 298.15 K. In fact the ionic radius of this cation (0.99 Å)43 is very close to Gd3+ (0.94 Å) and Eu3+ (0.95 Å)36 which are, as reflected in Figure 9, the cations which form the most stable complexes with 1b. Judging solely on stability constants (hence Gibbs energies of complexation) one may be prompted to conclude that the behavior of 1b for tervalent cations is similar to that observed for this ligand and related lower rim calix(4)arene derivatives as well as other macrocyclic ligands and alkali-metal cations. However, when the thermodynamic origin of complex stability as assessed from the enthalpy and entropy associated with these processes is considered, dramatic differences are observed. Thus, for alkali-metal cations and some calix(4)arene derivatives,9 the complex stability results from maximum exothermicity accompanied by the greatest entropy loss, whereas for lanthanide(III) cations and 1b the behavior is strikingly different. Thus, Figure 10 shows the variation of enthalpy of some lanthanide(III) complexes of 1b in acetonitrile with the cation radii. From Pr3+ to Tb3+ the pattern followed is very similar to that shown in Figure 8 for 1a and these cations in this solvent. However this is not the case from Tb3+ to Yb3+. Among the cations



TABLE 7: Thermodynamic Parameters for the Complexation of Some Lanthanide(III) Cations and 1a in DMF at 298.15 K cation

log Ks

∆cG° (kJ mol-1)

∆cH° (kJ mol-1)

∆cS° (J K-1 mol-1)

Sc3+ Y3+ Eu3+ Gd3+ Tb3+ Yb3+

3.19 ( 0.10 2.73 ( 0.08 3.32 ( 0.08 3.28 ( 0.09 3.28 ( 0.08 3.51 ( 0.08

-18.21 ( 0.56 -15.58 ( 0.46 -18.95 ( 0.45 -18.72 ( 0.50 -18.72 ( 0.45 -20.03 ( 0.46

-84.3 ( 0.4 -6.3 ( 0.5 -19.7 ( 0.3 -19.4 ( 0.6 -6.8 ( 0.2 -16.3 ( 0.5

-222 31 -3 -2 40 13

investigated the maximum exothermicity (and the greater loss of entropy) is observed for the largest cation (Pr3+). There is an enthalpy decrease in moving from Pr3+ to Dy3+ accompanied by an entropy gain. From Dy3+ to Yb3+ the pattern observed is quite irregular. (iii) Thermodynamics of 1a and TerValent Cations in N,NDimethylformamide at 298.15 K. The medium effect on the complexation process was investigated by the determination of thermodynamic data for 1a and various tervalent cations in N,Ndimethylformamide (see Table 7). Complex formation between a macrocyclic ligand and a metal cation in solution involves a competition between the ligand and the solvent for the free cation.44 If the metal cation is well solvated in a given solvent (strength of solvation depending upon the nature of the solvent and the cation concerned) it will be more reluctant to enter complexation with the ligand. The binding properties of the ligand therefore depend very much on the medium where the complexation process takes place. Dipolar aprotic solvents containing oxygen atoms with localized negative charges (i.e. N,N-dimethylformamide or dimethylsulfoxide) are protophilic solvents and therefore are expected to solvate cations better than acetonitrile (protophobic dipolar aprotic solvent).26,45 However, a careful interpretation of the medium effect requires knowledge on the differences in solvation of these metal cations as well as those for the ligand and the metal-ion complex and these need to be further investigated. The similarity of ∆cG° values observed for the systems studied in N,N-dimethylformamide leads to reasonable linear relationships when plotting the ∆cH° values against ∆cS° values for the complexation of 1a with these cations. The intercept value for this plot is -18.36 ( 0.70 kJ mol-1 while the slope is 298 K. Accordingly, these data provide good examples of enthalpyentropy compensation effects similarly to what was observed for the complexation of lanthanide(III) cations and 1a in acetonitrile. Inspection of Table 7 reveals that there is an entropy gain in N,N-dimethylformamide with respect to acetonitrile (Table 5) associated with the complexation reaction between lanthanide cations and 1a. Final Conclusions From the above discussion, the following is concluded: 1. 1H NMR complexation experiments have demonstrated the following conclusions: (i) Scandium, yttrium, lanthanum, and europium interact with 1a forming 1:1 complexes in CD3CN. The active sites of complexation are provided by the nitrogen and oxygen donor atoms, while chemical shift changes experienced by the protons adjacent to the sulfur atoms are less pronounced. (ii) The interactions of yttrium, lanthanum, europium, and lutetium with 1b in CD3CN gave rise to 1:1 complexes. The large chemical shift changes observed for the protons of the acetamide groups suggested that both amide and possibly ethereal oxygens participate in the complexation process. The oxygen donor atoms are believed to provide the

main sites of complexation, whereas the nitrogen atoms play a secondary role. The formation of 2-fold symmetric cones was observed in the cases of yttrium, europium and lutetium complexes, whereas the lanthanum complex adopted a 4-fold symmetry, as indicated by the signals of the aromatic protons giving two sets of peaks in the former case, while being equivalent in the latter. This was attributed to the difference in cation size. (iii) Weak interactions take place between 1a and yttrium in DMF-d7. Only the nitrogen and oxygen donor atoms of the amine pendant arms participate in the complexation since no significant shift changes were observed for the protons adjacent to the sulfur atoms. 2. Conductance data have shown the following: (i) Welldefined breaks in the titration curves demonstrated the formation of 1:1 complexes between 1a and lanthanide cations (Nd3+, Sm3+, Gd3+, Dy3+, and Tb3+) in acetonitrile at 298.15 K. (ii) No information could be obtained regarding complex stoichiometries for the interactions between 1a and two lanthanide cations (La3+ and Yb3+) in N,N-dimethylformamide at 298.15 K, since no significant breaks were observed in the conductance curves. (iii) In the concentration range used throughout this work in these solvents, the predominant ionic species in solution are the tervalent cations. 3. The thermodynamic data reported in this paper show the following: (i) Tervalent cations form 1:1 complexes in acetonitrile with 1a and 1b. (ii) Ligands 1a and 1b form complexes of similar stability constants in acetonitrile, whereas the enthalpic stability is higher for 1a relative to 1b in this solvent. Although 1b has a higher number of donor atoms all of which are efficient active sites for lanthanide cations (oxygen and nitrogen) relative to 1a, the presence of the isopropyl groups in its lower rim (which are bulkier than the methyl and ethyl groups in 1a) contribute enthalpically less favorably to the complex stability in this case. The difference in enthalpic stability could be explained on the basis of a higher degree of ‘preorganization’ of 1a relative to 1b prior to complexation. As far as the latter ligand is concerned, X-ray diffraction studies of an analogous macrocycle, namely p-tert-butylcalix(4)arene tetraacetamide and on the assumption that the ligand keeps the same structure in solution as in the solid state, have shown that a rotation of the CO groups inward the cavity toward the metal cation is necessary. This is due to the fact that the X-ray structure of the uncomplexed ligand46 shows two opposed carbonyls having diverging orientations. This CO movement is reflected in the large upfield shift observed by the CH2CON protons in this work (see Table 4) as well as similar studies reported by Ungaro and co-workers46 involving p-tert-butylcalix(4)arene tetraacetamide and alkali-metal cations. Among other factors, the consequences of the energy cost required for achieving converging orientations of all binding sites toward the metal cation are reflected in the less negative ∆cH° values observed for 1b (see Table 6) in acetonitrile. The more positive entropies of complexation, ∆cS°, found for 1b relative to 1a may indicate (a) a greater desolvation of 1b upon complexation which leads to an increase in disorder because more free solvent molecules are released, (b) greater solvation of the 1a complexes relative to 1b complexes since this would result in a decrease of the disorder of the system, hence more negative ∆cS° values in the first case. This assumption seems rather reasonable because 1a complexes are expected to be more exposed to solvent molecules relative to the corresponding ones of 1b due to the lower degree of interactions between the soft sulfur atoms and hard lanthanide cations which is in fact what was observed in the 1H NMR complexation experiments.

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Complexation of Calixarene Derivatives and Lanthanide Cations in

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