J. Phys. Chem. B 1998, 102, 7839-7844
7839
The Medium Effect on the Thermodynamics of Complexation of 5,11,17,23-Tetrakis-(1,1-dimethylethyl)-25,27-bis(methylthioethoxy)26,28-bis[(diethylamine)ethoxy]calix(4)arene and the Silver Cation Angela F. Danil de Namor,* Mariel L. Zapata-Ormachea, and Robert G. Hutcherson Laboratory of Thermochemistry, Department of Chemistry, School of Physical Sciences, UniVersity of Surrey, Guildford, Surrey GU2 5XH, UK ReceiVed: April 7, 1998; In Final Form: July 6, 1998
The interaction of 5,11,17,23-tetrakis-(1,1-dimethylethyl)-25,27-bis(methylthioethoxy)-26,28-bis[(diethylamine)ethoxy]calix(4)arene and the silver cation was investigated by 1H NMR in CD3OD and CD3CN at 298 K. The results suggest that the hydrophilic cavity of the ligand hosts the metal cation through interaction with the different donor atoms (O, N, S). Conductometric titrations of Ag+ and this ligand in methanol and in ethanol indicate that the stoichiometry of the metal ion-ligand complex is 1:1. The thermodynamics of complexation of 1a and silver in six different solvents (methanol, ethanol, propan-1-ol, N,N-dimethylformamide, acetonitrile, and benzonitrile) at 298.15 K was derived from titration calorimetry. Stability constants were also checked by direct potentiometry using silver electrodes. Standard Gibbs energies of solution of 5,11,17,23tetrakis(1,1-dimethylethyl)-25,27-bis(methylthioethoxy)-26,28-bis[(diethylamine)ethoxy]calix(4)arene derived from solubility measurements in different solvents are used to calculate transfer Gibbs energies from acetonitrile to various solvents. The medium effect on the complexation process is assessed in terms of Gibbs energies taking into account the differences in solvation of the free and the complex metal cation and the ligand in the appropriate solvents. Representative examples are given to demonstrate that the medium alters the stability of the complex in a selective manner which is dependent upon the solvation changes that the reactants and the product undergo upon complexation in the various solvents.
Introduction The field of calixarene chemistry has been the subject of several books,1-3 reviews,4,5 and research publications. The possibility of substituting the phenolic hydrogens of p-tertbutylcalix(4)arenes by different functional groups has generated a variety of macrocycles which are characterized by the presence of two cavities, a hydrophobic, able to interact with neutral species and a hydrophilic which is known to interact with metal cations. One of the factors that affect the complexing properties of these ligands for metal cations is the nature of the donor atoms present in the introduced functional groups.1-4 An area of increasing interest is the search for ligands with selective properties for soft metal cations. We have previously reported6,7 the synthesis and characterization of a series of lower rim functionalized calix(4)arenes containing aliphatic and alicyclic tertiary amines. The attachment of these functional groups provides a suitable arrangement for interaction with toxic metal cations (Hg2+, Cd2+, Pb2+) while cations such a Na+, K+, and Ba2+ are discriminated against. Recently, new lower rim p-tertbutylcalix(4)arene derivatives containing two tertiary (aliphatic and alicyclic) and two methyl thioether pendant groups alternately arranged have been synthesized and characterized by our group.8 X-ray crystallographic studies on one of these derivatives9 have shown that these macrocycles have a distorted “cone” conformation. The aromatic moieties containing the aminoethoxy substituents are parallel to each other while those containing the methylthio substituents adopt a flattened conformation. Information regarding the conformation of these compounds in solution was obtained from 1H and 13C NMR studies in Cl3CD and CD3OD at 298 K. Preliminary studies
demonstrated that these ligands are able to complex selectively Hg2+, Ag+, Cd2+, Pb2+, and Cu2+ in methanol. The thermodynamics of calixarene chemistry has been recently reviewed by our group.10 As far as the complexation of lower rim calix(4)arene derivatives and metal cations is concerned most investigations have been carried out in methanol, acetonitrile, and benzonitrile as the reaction media.11-13 However, there are not systematic studies in a variety of solvents of moderate dielectric constant where ions rather than ion pairs are the predominant species in solution. This paper concerns 1H NMR, conductance, and thermodynamic studies of 5,11,17,23-tetrakis(1,1-dimethylethyl)-25,27-bis(methylthioethoxy)-26,28-bis[(diethylamine)ethoxy]calix(4)arene (1a) and the silver cation in solution at 298.15 K.
Since the differences in the thermodynamic parameters of complexation, ∆cPo(Po) Go, Ho, So), of 1a and Ag+ in one solvent (s1, reference solvent) relative to another solvent (s2) are dependent on the transfer parameters, ∆tPoof the metal cation, M+, the ligand, 1a, and the metal ion complex M+1a as shown in eq 1
S1089-5647(98)01778-7 CCC: $15.00 © 1998 American Chemical Society Published on Web 09/09/1998
7840 J. Phys. Chem. B, Vol. 102, No. 40, 1998
∆cPo(s2) - ∆cPo(s1) ) ∆tPo(M+ 1a)(s1fs2) ∆tPo(M+)(s1fs2) - ∆tPo1a(s1fs2) (1) These parameters are considered in terms of Gibbs energies in order to assess quantitatively the effect of the solvation changes of the reactants and the product on the stability of complex formation in this system. Experimental Section Chemicals. A two-step procedure was used to synthesize 1a which consisted of treatment of p-tert-butylcalix(4)arene with 2-chloromethylethyl sulfide in acetonitrile at 363 K to give the 1,3-bismethyl ethyl thioether derivative. Treatment of this derivative with 2-diethylaminoethyl chloride hydrochloride using 18-crown-6 as a transfer catalyst in THF/DMF medium at 363 K yielded 1a as detailed elsewhere.9 This ligand was recrystallized from a methanol-dichloromethane solvent mixture. Anal. Calcd % for C62H94O4S2N2: C, 74.80; H, 9.52; N, 2.81. Found: % C, 74.79; H, 9.92; N, 2,77. N,N-Dimethylformamide (Fischer), tetrahydrofuran (Aldrich), hexane (Fischer), ethyl ethanoate (Fischer), acetonitrile (Fischer), dichloromethane (Fischer), ethanol (Hayman), methanol (HPLC grade, Fischer), chloroform (Fischer), nitrobenzene (Fluka), benzonitrile (Aldrich), butan-1-ol (Fischer), and propylene carbonate (Aldrich) were purified as described elsewhere.11,14 Silver(I) perchlorate (Aldrich) and silver(I) nitrate (99.8%, Rose chemicals) were used. These were dried over phosphorus pentoxide under vacuum for several days. Tetra-n-butylammonium perchlorate (99% electrochemical grade) (Fluka) was used without further purification. Complexation Studies by 1H NMR. 1H NMR measurements were recorded at 298 K on a Brucker 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. The NMR standard used was TMS. 1H NMR titration experiments to assess the interaction of 1a with the metal cation in the appropriate solvent were carried out by injecting a solution of the silver salt (6 × 10-3 mol dm-3) into a NMR tube containing a known concentration of the ligand (1 × 10-3 mol dm-3), predissolved in the deuterated solvent. Stepwise additions of silver cation solutions were made until no further shift changes were observed in the spectrum. Proton shifts of the free and complexed ligand were noted in parts per million. Conductance Measurements. Conductance measurements at 298.15 K were carried out with a Wayne-Kerr autobalance universal bridge, type B642. A solution containing the silver cation salt (concentration range from 1 × 10-4 to 3 × 10-4 mol dm-3) in the appropriate solvent was titrated with a solution of the ligand 1a (concentration range from 1 × 10-3 to 2 × 10-3 mol dm-3) in the same solvent. A plot of molar conductances against the ligand-metal ion concentration ratio, C1a/CAg+, was used to determine the stoichiometry of the complex. Titration Calorimetry. Stability constants (log Ks) and enthalpies, ∆cHo, of complexation of 1a and the silver cation at 298.15 K were determined by classical calorimetry using a Tronac 450 titration calorimetry (originally designed by Christensen and Izatt15) or by microcalorimetric titrations. The latter titrations were performed with the 2277 Thermal Activity Monitor. Given the low solubility of the ligand in acetonitrile,
Danil de Namor et al. a solution of the calixarene derivative (concentration range from 8 × 10-4 to 1 × 10-3 mol dm-3) was placed in the vessel and the silver salt (concentration range from 2.8 × 10-2 to 3.2 × 10-2 mol dm-3) in the syringe. The procedure used has been described elsewhere.11 For the macrocalorimetric determinations a solution of the calixarene in the appropriate solvent (concentration in MeOH ) EtOH ) PrOH ) DMF, 1.0 × 10-2 mol dm-3; PhCN 2.0 × 10-2 mol dm-3) was titrated in a 50-mL solution of the metal ion salt (concentration range from 2 × 10-4 to 8 × 10-4 mol dm-3) in the same solvent. The technique used has been described elsewhere.16 The uncertainties of the measurements are 2%. Potentiometric Titrations. Potentiometric titrations were used to determine stability constants of the silver cation with 1a at 298.15 K. A silver-silver ion reference electrode consisting of a silver wire, introduced in a solution of silver perchlorate (3 × 10-3 mol dm-3), was used in an electrochemical cell as that suggested by Schneider and co-workers.18,19 A schematic representation of the electrochemical cell is shown,
Ag/Ag+, X M//0.05 M (But4NClO4)//0.01 M, Ag+/Ag Tetra-n-butylammonium perchlorate (TBAP) was used to maintain the ionic strength (0.05 mol dm-3) of the solution constant. All solutions [silver nitrate (3 × 10-3 mol dm-3) and ligand (MeOH ) PrOH 2 × 10-3 mol dm-3; MeCN, 1 × 10-3 mol dm-3)] were prepared in the appropriate solvent and the salt bridge was filled with the solution of tetra-n-butylammonium perchlorate in the same solvent (0.05 mol dm-3). For the calibration of the electrode, the silver perchlorate solution was titrated into the sample cell containing TBAP (25 cm3, 0.05 mol dm-3) solution in the appropriate solvent. These data were used to calculate the standard electrode potential of the reference cell. In all cases, Nernstian behavior was observed. The second titration involved the formation of the silvercalixarene complex. The ligand solution was added to the silver solution and potential changes recorded. The data were used to calculate the stability constant of the metal ion complex. Measurements were carried out in duplicate. Solubility Measurements. Saturated solutions of 1a were prepared by adding an excess amount of the solid to the solvent. The mixtures were left in a thermostat at 298.15 K until equilibrium was reached. Aliquots of the samples were taken and analyzed gravimetrically by triplicate. Separate blank experiments were carried out to ensure the absence of any involatile material in the pure solvent. Solvate formation was checked by exposing the solid to an atmosphere of the solvent for several days.17 Results and Discussion 1H
NMR titration of 1a and Silver in CD3OD and in CD3CN at 298.15 K. The results of the 1H NMR titration of 1a with Ag+ (using perchlorate as the counterion) in a protic solvent (CD3OD) and in a dipolar aprotic solvent (CD3CN) at 298.15 K are reported in Tables 1 and 2, respectively. Chemical shift changes with respect to the free ligand in the appropriate solvent are reported as a function of the metal ion/ligand ratio. As far as the free ligand is concerned it is observed that a solvent change from CD3CN to CD3OD causes the resonance frequency of the methylene protons adjacent to nitrogen in the OCH2CH2N groups to downfield shift (∆δ ) 0.16 ppm) whereas that for the NCH2CH3 groups appears to be affected to a less extent
Medium Effect on Complexation
J. Phys. Chem. B, Vol. 102, No. 40, 1998 7841
TABLE 1: 1H NMR Titration of 1a with Silver in Deuterated Methanol at 298.15 Ka-c 6
9
5
8 7
O
10
12
S
2
O
11
4 3
N 2 1
mole ratio [Ag+]/[1a]
1
2
3
4
free ligand
1.12
2.67
3.11
3.89
0.25
+0.02
+0.03
+0.01
+0.04
0.50
+0.03
+0.08
+0.02
+0.08
0.75
+0.04
+0.10
+0.02
+0.10
1.00
+0.05
+0.12
+0.02
+0.13
1.26
+0.05
+0.14
+0.02
+0.15
1.51
+0.05
+0.14
+0.02
+0.16
2.01
+0.06
+0.15
+0.03
+0.16
2.51
+0.08
+0.18
+0.04
+0.23
3.01
+0.08
+0.19
+0.05
+0.25
a
5 8
6 9
7
7.11 6.59 0 +0.01 -0.01 +0.03 -0.01 +0.04 -0.1 +0.04 -0.01 +0.03 -0.01 +0.03 -0.01 +0.03 -0.01 +0.06 -0.02 +0.07
1.29 0.89 -0.01 +0.01 -0.02 +0.02 -0.02 +0.03 -0.02 +0.03 -0.01 +0.03 -0.01 +0.02 -0.01 +0.03 -0.01 +0.04 -0.04 +0.05
4.37 3.20 0 0 0 +0.02 0 +0.04 -0.01 +0.06 -0.01 +0.07 -0.01 +0.07 -0.01 +0.08 -0.01 0.11 0 +0.11
10
11
12
4.15
3.20
2.22
+0.10
+0.02
+0.05
+0.15
+0.07
+0.12
+0.21
+0.10
+0.17
+0.23
+0.15
+0.24
+0.27
+0.19
+0.30
+0.28
+0.19
+0.30
+0.29
+0.20
+0.31
+0.29
+0.21
+0.33
+0.30
+0.21
+0.33
10
11
12
4.14
3.25
2.21
0
0
+0.02
+0.01
+0.01
+0.05
+0.01
+0.03
+0.07
+0.01
+0.03
+0.08
+0.01
+0.04
+0.10
+0.02
+0.05
+0.11
+0.03
+0.05
+0.12
+0.02
+0.06
+0.13
+0.02
+0.06
+0.13
Chemical shifts (δ) and changes in chemical shifts (∆δ) in ppm. b Ppm (0.01. c +, downfield shift; -, upfield shift.
TABLE 2: 1H NMR Titration of 1a with Silver in Deuterated Acetonitrile at 298.15 Ka-c mole ratio [Ag+]/[1a]
1
2
3
4
free ligand
1.02
2.59
2.95
3.82
0.25
+0.01
+0.02
+0.01
+0.04
0.50
+0.01
+0.03
+0.02
+0.07
0.75
+0.01
+0.04
+0.03
+0.10
1.00
+0.01
+0.05
+0.03
+0.12
1.25
+0.01
+0.06
+0.04
+0.14
1.50
+0.01
+0.06
+0.04
+0.15
2.00
+0.02
+0.07
+0.05
+0.17
2.50
+0.02
+0.07
+0.05
+0.18
3.00
+0.02
+0.07
+0.05
+0.18
a
5 8
6 9
7
7.20 6.87 -0.02 0 -0.04 0 -0.06 0 -0.07 0 -0.08 0 -0.09
1.25 1.01 -0.02 0 -0.03 +0.01 -0.03 +0.01 -0.04 +0.01 -0.05 +0.01 -0.05 +0.01 -0.06 +0.02 -0.06 +0.02 -0.06 +0.02
4.40 3.21 -0.01 +0.01 -0.02 +0.02 -0.03 +0.02 -0.03 +0.02 -0.03 +0.02 -0.04 +0.03 -0.04 +0.03 -0.04 +0.03 -0.04 +0.03
-0.10 0 -0.11 0 -0.11
Chemical shifts (δ) and changes in chemical shifts (∆δ) in ppm. b Ppm (0.01. c +, downfield shift; -, upfield shift.
(∆δ ) 0.08 ppm). This is likely to result from hydrogen bond formation between the amino nitrogens and methanol. As far as the 1H NMR spectrum in CD3OD is concerned, large shift changes in the proton signals relative to the free ligand (chemical shift for 1a in CD3OD are reported Table 1) are recorded. Both donor atoms (N and S) appear to be involved in the complexation of this ligand with this cation. During the process of complexation, all the protons of the amine and methyl thioether pendant arms at the lower rim are deshielded. The conformational changes observed previously upon protonation
of the ligand9 do not take place as the two aromatic and the two tert-butyl signals remain almost stationary. Only a minimum amount of conformational movement of the ligand appears to be required to complex this cation. The sulfur donor atoms seem to bind more strongly to the cation since the adjacent protons are deshielded to a greater extent than those adjacent to the nitrogen atoms. The latter, although deshielded may also experience field effects which account for the lower than expected shift changes. This is particularly true for the methylene protons adjacent to the nitrogen which show a lower
7842 J. Phys. Chem. B, Vol. 102, No. 40, 1998
Figure 1. Conductometric titration of silver perchlorate and 1a in (i) methanol and (ii) ethanol at 298.15 K.
downfield change in shift than the methylene protons next to the oxygen, suggesting that the oxygen atoms are also involved in the complexation of this cation with this ligand. These results are concomitant with those found in the X-ray structure of an analogous derivative containing pyridine pendant arms where the silver cation is found coordinated to both the pyridine nitrogen and the phenolic oxygens.20 As far as the 1H NMR titration of 1a and Ag+ in CD3CN is concerned, the results in Table 2 indicate that a deshielding effect is mainly shown by the oxygen methylene protons of the amino groups as well as by the methyl groups adjacent to the sulfur atom. On the other hand smaller changes are exhibited by the resonances arising from the remaining protons of the pendant groups. These observations seem to suggest that in CD3CN the cation exerts a lower tendency to coordinate with the amine nitrogens. Conductance Measurements. Variation of the electrical conductance with the concentration of the metal ion salt and the ligand were used to determine the composition of the metal ion complex in a variety of solvents. Representative examples are those given in Figure 1 where the conductometric titration curves (plots of molar conductance, Λm, against the ligand: metal cation mole ratio) for silver perchlorate and 1a in methanol (i) and in ethanol (ii) are shown. As far as the data in methanol are concerned, a sharp break is observed in the titration curve when the ligand:metal cation reaches unity, indicating the formation of a highly stable 1:1 complex. The decrease in conductance observed with the addition of the ligand to the methanolic solution containing the silver perchlorate reflects the lower mobility of the metal ion complex relative to that of the free cation. This may be attributed to the larger size of the former relative to the latter. The pattern found in ethanol (Figure 1, ii) is similar to that found in MeOH although the break at 1:1 ligand:metal cation ratio is less sharp in the former solvent. The results from conductance measurements reveal that (i) the macrocycle interacts with the silver cation to give a 1:1 stoichoimetry complex in these solvents and (ii) the strength of complexation of 1a and Ag+ is greater in MeOH than in EtOH.
Danil de Namor et al. To obtain quantitative information on the interaction of 1a with Ag+ in different solvents, we proceeded with calorimetric and potentiometric titrations. The aim is to investigate the thermodynamics associated with the complexation process and to assess the medium effect on the interaction of this cation and this ligand and these are now discussed. Thermodynamics of Complexation. Stabilility constants (log Ks) and derived Gibbs energy, ∆cG°; enthalpy ∆cHo and entropy changes, ∆cSo of complexation of 1a and the silver cation in methanol, MeOH; ethanol, EtOH; propan-1-ol, 1-PrOH; N,N-dimethylformamide, DMF; acetonitrile, MeCN; and benzonitrile, PhCN, at 298.15 K are reported in Table 3. Also included in this table are the standard deviation of the data. Conductance measurements of silver perchlorate in MeOH,21 MeCN,22 DMF,21 and PhCN23 at 298.15 K have shown that these electrolytes are fully dissociated in these solvents. As far as EtOH and 1-PrOH are concerned we are not aware of any conductance data reported for AgClO4 in these solvents. Therefore, calorimetric measurements were carried out at different electrolyte concentrations in order to investigate that no process other than complexation takes place. Since no variation in the log Ks and ∆cHo values were observed, it is reasonable to assume that ions rather than ion pairs are the predominant species in solution in the concentration range used for these measurements. On the basis of the experimental facts, the data reported in Table 3 are referred to the process
Ag+(s) + 1a(s) f Ag+ 1a(s)
(2)
Although the magnitude of the stability constants reported in Table 3 are well within the scope of titration calorimetry, these data were also checked by direct potentiometry using silver electrodes (see Table 3). Again no variations in log Ks values were observed by altering the electrolyte concentration. Good agreement is found between the data derived from the two independent methods. Therefore, an average of the calorimetric and potentiometric log Ks values is used to calculate the standard Gibbs energies of complexation of these systems in the various solvents. Data in Table 3 indicate that in all cases, the complexation process is enthalpy controlled and takes place with a loss of entropy. The medium effect on the complexation of 1a and the silver cation is clearly reflected since significant variations are found in the log Ks values. The strength of complexation follows the sequence,
MeOH > EtOH > PrOH > DMF > PhCN > MeCN As far as MeOH and EtOH are concerned the trend found is the same as that observed from conductance measurements. However, the most dramatic changes are observed in the enthalpy and entropy contributions to the Gibbs energy of the process. Thus, the stability in enthalpic terms is as follows,
PhCN = MeOH > DMF > EtOH = PrOH > MeCN Therefore the lower stability of Ag+ and 1a in PhCN relative to MeOH is due to the greater entropy loss in the complexation process when the medium is PhCN relative to MeOH, since the enthalpy associated with this process in both solvents is about the same. This is indeed not the case for the alcohols. Thus, the higher stability in MeOH relative to the higher alcohols is enthalpy controlled. Indeed, in moving from MeOH to PrOH, no dramatic entropy changes are observed. In an attempt to elucidate the factors which contribute to the selective complexation behavior in one solvent relative to another, eq 1 in terms of Gibbs energies is considered. Transfer
Medium Effect on Complexation
J. Phys. Chem. B, Vol. 102, No. 40, 1998 7843
TABLE 3: Stability Constants (log Ks) and Derived Standard Gibbs Energy, Enthalpy, and Entropy Changes of Complexation of 1a and Silver in Various Solvents at 298.15 K solventa MeCN DMF PhCN
log Ks
}
3.46 ( 0.16b 3.35 ( 0.01c 4.16 ( 0.07 4.00 ( 0.04 5.92 ( 0.09b 5.81 ( 0.04c 4.68 ( 0.05b 4.62 ( 0.11c 4.58 ( 0.04
} }
MeOH EtOH 1-PrOH
∆cG°/kJ mol-1
∆cH°/kJ mol-1
∆cS°/J K-1 mol-1
-19.44 ( 0.91
-23.16 ( 1.65
-12.5
-23.75 ( 0.38 -22.83 ( 0.23
-39.39 ( 0.95 -44.85 ( 0.73
-52.4 -73.8
5.87 ( 0.09
-33.51 ( 0.51
-44.39 ( 1.12
-36.5
4.65 ( 0.08
-26.54 ( 0.28
-36.34 ( 0.28
-32.9
-26.14 ( 0.24
-36.92 ( 1.11
-36.2
3.41 ( 0.12
a
Abbreviations used: acetonitrile, MeCN; N,N-dimethylformamide, DMF; benzonitrile, PhCN; methanol, MeOH; ethanol, EtOH; propan-1-ol, 1-PrOH. b Calorimetric value. c Direct potentiometry using silver electrodes.
TABLE 4: Solubilities and Derived Gibbs Energies of Solution of 1a in Various Solvents at 298.15 K. Transfer Gibbs Energies from Acetonitrile to Other Solvents solventsa
solubility/mol dm-3
∆sG°/kJ mol-1
∆tG°(MeCNf(s)/ kJ mol-1
MeCN MeOH EtOH 1-BuOH PhCN CH2Cl2 PhNO2 DMF PC THF Hex EtAc 1,2 DCE
(3.30 ( 0.04) × 10-3 (1.16 ( 0.03) × 10-2 (2.07 ( 0.03) × 10-2 (1.36 ( 0.04) × 10-1 too solubleb too solubleb too solubleb (1.68 ( 0.01) × 10-2 too insoluble too solubleb too solubleb too solubleb too solubleb
+14.16 +11.05 +9.61 +4.95
0 -3.11 -4.55 -9.21
+10.13
-4.03
a Solvent abbreviations are as follows: acetonitrile, MeCN; methanol, MeOH; ethanol, EtOH; butanol, 1-BuOH; benzonitrile, PhCN; dichloromethane, CH2Cl2; nitrobenzene, PhNO2; N,N-dimethylformamide, DMF; propylene carbonate, PC; tetrahydrofuran, THF; hexane, Hex; ethyl acetate, EtAc; 1,2-dichloroethane, 1,2-DCE. b Solvate formation when exposed to an atmosphere of the solvent.
Gibbs energies of the silver cation from water to various solvents have been reported in the literature.24-26 These data, based on the Ph4AsPh4B convention, can be used for the calculation of ∆tG° values from acetonitrile (reference solvent) to various solvents. However, there are not ∆sG° (hence ∆tG°) data available for this ligand. Therefore, we proceeded with solubility measurements of 1a in a variety of solvents at 298.15 K and these are now discussed. Solubility and Derived Solution Gibbs Energies. Standard Transfer Gibbs Energies from Acetonitrile to Other Solvents. The solubility of 1a in various solvents at 298.15 K is reported in Table 4. No quantitative data could be obtained for this ligand in benzonitrile, chloroform, dichloromethane, ethyl acetate, hexane, 1,2-dichloroethane, nitrobenzene, and tetrahydrofuran due to the high solubility of this ligand in these solvents. Solvate formation was detected when the solute was exposed to a saturated atmosphere of each of these solvents. On the other hand, the solubility of 1a in propylene carbonate was very low, and therefore, it was not possible to obtain quantitative information in this solvent. From solubility data, standard Gibbs energies of solution, ∆sG° (standard state, 1 mol dm-3), were calculated. Since for a given solute the contribution of the crystal lattice to the ∆sG° value is the same, the variation observed in the solution Gibbs energies results from the difference in the solvation of this ligand in these solvents. This is best reflected in the transfer Gibbs energies, ∆tG°. These are calculated by using acetonitrile as the reference solvent. Gibbs energies for this ligand, which is essentially a nonelec-
trolyte, show that its transfer from one solvent to another alters the equilibrium position in a selective manner where differences of up to =9 kJ mol-1 are observed. Thus, the solvation sequence in terms of ∆tG° is PhCN, EtAc, Hex, THF, PhNO2, CH2Cl2, 1,2 DCE > 1-BuOH > EtOH > DMF > MeOH > MeCN. Two points should be stressed regarding the data in Table 4. (i) The fact that this ligand undergoes solvate formation in a variety of solvents underlines the possibility that the use of these solvents for synthetic purposes may lead to the isolation of their complexes rather than the pure ligand. (ii) The data in Table 4 provide further experimental evidence that this derivative is highly solvated in low donor number (DN) solvents (1,2 dichloroethane, DN ) 0; nitrobenzene, DN ) 4.4; PhCN, DN ) 11.9)27 and, therefore, attempts to correlate the solvation of calixarene derivatives with the Gutmann donor number of the solvents should be ruled out as discussed by Danil de Namor and co-workers.10 Interpretation of Gibbs Energy Complexation Data. The availability of standard Gibbs energies of complexation in various solvents (Table 3) and corresponding data for the transfer of the silver cation26,28,29 and the ligand (Table 4) from a reference solvent to another allows (by rearrangement of eq 1) the calculation of the standard transfer Gibbs energies of the metal ion complex (data based on the Ph4AsPh4B convention).26 The individual processes involved in eq 1 are best illustrated via a thermodynamic cycle expressed in terms of Gibbs energies. Ag+(s1) + 1a (s1) ∆tG°
∆cG°
∆tG°
Ag+(s2) + 1a (s2)
Ag+1a (s1) ∆tG°
∆cG°
(3)
Ag+1a (s2)
With the aim of analyzing the various factors (metal cation, ligand, metal-ion complex) which contribute to the medium effect [∆cGo(s2) - ∆cGo(s1)] (eq 1) on the complexation process, three systems are considered where acetonitrile is the reference solvent and s2 ) MeOH, EtOH, DMF. It should be emphasized that the difference between ∆cGo values in the two solvents is independent of any extrathermodynamic convention since it is clear from eq 1 that the anion contribution (Ag+ and Ag+1a) is canceled out. This equation implies that the optimum conditions for higher complex stability in s2 (∆cGo more negative) relative to s1 is fulfilled by a solvent (s2) which offers a better solvating medium for the metal-ion complex [negative ∆tG° (s1 f s2) and a poorer solvating medium for the ligand and the free metal cation [positive ∆tG° (s1 f s2)] than the reference solvent (s1).
7844 J. Phys. Chem. B, Vol. 102, No. 40, 1998 By inserting the appropriate quantities in the thermodynamic cycle, this analysis is carried out in the acetonitrile-methanol system. Ag+(MeCN) + 1a (MeCN)
30.12 kJ mol–1
-19.44 kJ mol–1
-3.11 kJ mol–1
Ag+(MeOH) + 1a (MeOH)
-33.51 kJ mol–1
Ag+1a (MeCN)
12.94 kJ mol–1
Ag+1a
(MeOH)
Hence, it follows from eq 1 that the higher stability of the complex in MeOH relative to MeCN is controlled by the lower solvation of the cation in the former solvent relative to the latter. Indeed, the higher solvation of the ligand and the lower solvation of the metal ion complex in MeOH with respect to MeCN contribute unfavorably to the stability of the complex in this solvent relative to acetonitrile. In the acetonitrile-ethanol solvent system, the contribution Ag+(MeCN) + 1a (MeCN)
28.03 kJ mol–1
Ag+(EtOH)
-19.44 kJ mol–1
-4.55 kJ mol–1
+ 1a (EtOH)
-26.54 kJ mol–1
Ag+1a (MeCN)
16.18 kJ mol–1
Ag+1a (EtOH)
of the cation to the higher stability in EtOH relative to MeCN still predominates. However, the ∆(∆cGo) becomes less negative in moving from MeCN-MeOH to MeCN-EtOH. This is due to the more unfavorable contributions involved in the transfer of the free metal cation and the ligand from MeCN to EtOH relative to those involving the transfer from the same reference solvent to methanol. Quite a different picture emerges when the acetonitrile-N,Ndimethylformamide solvent system is considered. Unlike for the systems involving the alcohols, the ∆tG° value for the free Ag+(MeCN) + 1a (MeCN)
4.65 kJ mol–1
Ag+(DMF)
-19.44 kJ mol–1
-4.03 kJ mol–1
+ 1a (DMF)
-23.75 kJ mol–1
Ag+1a (MeCN)
-3.69 kJ mol–1
Ag+1a (DMF)
cation from MeCN to DMF is relatively small indicating that these solvents do not offer a differentiating solvation media for this cation. Since the ligand contribution does not favor complexation in DMF, the slightly higher complex stability found in this solvent is attributed to the both higher solvation of the metal ion complex in DMF relative to MeCN and the lower solvation of the silver cation in DMF relative to MeCN. The above examples unambiguously demonstrate that the medium alters the stability of the complex in a selective manner which is not only dependent on the solvation of the free cation (as it is often assumed) but it is the result of the solvation changes that the reactants and the product undergo upon complexation in the various solvents.
Danil de Namor et al. We therefore conclude that this is the first thermodynamic study (i) involving this ligand and (ii) in the field of calixarene chemistry in which the medium effect on cation complexation processes is shown in a wide variety of solvents. Furthermore, the results demonstrate that in assessing the medium effect on the stability of complex formation, quantitative information regarding the solvation changes that the reactants and the product undergo upon complexation is required. Similar analysis in terms of enthalpy and entropy (where more pronounced changes are expected) is now in progress.30 Acknowledgment. The authors acknowledge the financial support provided by the EU, DGI, and DGX11 (contract 932038 AR). References and Notes (1) Gutsche, C. D. In Monographs in Supramolecular Chemistry: Stoddart, J. F., Ed.; Royal Society of Chemistry: London, 1989. (2) Calixarenes. A Versatile Class of Macrocyclic Compounds; Vicens, J., Bo¨hmer, V., Eds.; Kluwer Academic Publishers: Dordrecht, The Netherlands, 1991. (3) Calixarenes 50th AnniVersary CommemoratiVe Volume, Eds. Vicens, J., Asfari, Z., Harrowfield, J. M., Eds.; Kluwer Academic Publishers: Dordrecht, The Netherlands, 1994. (4) Gutsche, C. D. Aldrichimica Acta 1995, 28, 3. (5) Bo¨hmer, V. Angew. Chem. Int. Ed. Engl. 1995, 34, 713. (6) Danil de Namor, A. F.; Sueros Velarde, F. J.; Cabaleiro, M. C. J. Chem. Soc., Faraday Trans. 1996, 92, 1731. (7) Danil de Namor, A. F.; Hutcherson, R. G.; Sueros Velarde; F. J.; Piro, O. E.; Castellano, E. E. J. Chem. Soc., Faraday Trans., in press. (8) Danil de Namor, A. F.; Hutcherson, R. G.; Sueros Velarde; F. J.; Zapata-Ormachea, M. L.; Pulcha Salazar, L. E.; Al Jammaz, I.; Al-Rawi, N. Pure Appl. Chem., in press. (9) Danil de Namor, A. F.; Hutcherson, R. G.; Sueros Velarde, F. J.; Alvarez Larena, A.; Brianso´, J. L. J. Chem. Soc., Perkin Trans 1 1998, 2933. (10) Danil de Namor, A. F.; Cleverley, R. M.; Zapata-Ormachea, M. L. Chem. ReV., in press. (11) 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, 16770, 16781 and references therein. (12) Danil de Namor, A. F.; Zapata-Ormachea, M. L.; Jafou, O.; Al Rawi, N. J. Phys. Chem. 1997, 101, 6772. (13) Arnaud Neu, F.; Fanni, S.; Guerra, L.; McGregor, W.; Ziat, K.; Schwing-Weill, M. J.; Barrett, G.; McKervey, M. A.; Marrs, D.; Seward, E. M. J. Chem. Soc., Perkin Trans. 1995, 113. (14) Perrin, D. D.; Armarego, W. L. F. Purification of Laboratory Chemicals, 3rd ed.; Pergamon Press: New York, 1988. (15) Christensen, J. J.; Izatt, R. M.; Hansen, L. D. ReV. Sci. Instrum. 1965, 36, 779. (16) 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. (17) Ligny, C. L.; Bax, D.; Alfenar, M.; Elfenrek, M. G. L. Recl. TraV. Chim. Pays-Bas. 1969, 88, 1183. (18) Cox, B. G.; Garcia Rosas, J.; Schneider, H. J. Am. Chem. Soc. 1981, 103, 1384. (19) Gutknecht, J.; Schneider, H.; Stro¨ka, H. J. Inorg. Chem. 1978, 17, 3326. (20) Danil de Namor, A. F.; Piro, O. E.; Pulcha Salazar, L. E.; Aquilar Cornejo, A. F.; Al Rawi, N.; Castellano, E. E.; Sueros Velarde, F. J. J. Chem. Soc., Faraday Trans., in press. (21) Prue, J. E.; Sherrington, P. J. Trans. Faraday Soc. 1961, 57, 1795. (22) Yeager, H. L.; Kratochvil, B. J. Phys. Chem. 1969, 73, 1963. (23) Danil de Namor, A. F., unpublished results. (24) Ling, H. C. Ph.D. Thesis, University of Surrey, 1981. (25) Ghousseini, L. Ph.D. Thesis, University of Surrey, 1985. (26) Cox, B. G.; Hedwig, G. R.; Parker, A. J.; Watts, D. W. Aust. J. Chem. 1974, 27, 477. (27) Gutmman, V. The donor-acceptor approach to molecular interactions; Plenum Press: New York, 1978; p 20. (28) Johnson, M.; Persson, I. Inorg. Chim. Acta 1987, 127, 25. (29) Ling Ph.D. Thesis, University of Surrey. (30) Danil de Namor, A. F.; Zapata-Ormachea, M. L.; Hutcherson, R. G., work in progress.
