J. Phys. Chem. B 1997, 101, 1643-1648
1643
Solution Thermodynamics of Ba(II)-Cryptand 222 in Various Solvents and Derived Coordination Data in the Solid State Angela F. Danil de Namor* and Dorota Kowalska Laboratory of Thermochemistry, Department of Chemistry, UniVersity of Surrey, Guildford, Surrey GU2 5XH, U.K. ReceiVed: August 20, 1996; In Final Form: NoVember 7, 1996X
Enthalpies of complexation of barium and cryptand 222 in N,N-dimethylformamide, dimethyl sulfoxide, acetonitrile, and propylene carbonate at 298.15 K determined calorimetrically are reported. The stability constant for this system in acetonitrile at the same temperature was measured by competitive potentiometric titration using silver electrodes. Gibbs energies of complexation in a variety of solvents are analyzed in terms of corresponding data for the transfer of ligand and free and complexed cation. Standard enthalpies of solution of barium perchlorate in N,N-dimethylformamide, dimethyl sulfoxide, acetonitrile, and propylene carbonate and for barium-cryptand 222 perchlorate in water, methanol, N,N-dimethylformamide, and dimethyl sulfoxide measured calorimetrically at 298.15 K are reported. The data show that in these solvents the former electrolyte is more solvated than the latter. Single-ion values of transfer enthalpies for the free and complexed cations are calculated on the basis of the Ph4AsPh4B convention. These data and those previously reported for the transfer enthalpy of cryptand 222 among these solvents are used to explain the variation observed in the complexation enthalpies of Ba2+ and this ligand in the various solvents. Coordination enthalpies for the process in the solid state derived from six different solvents are for the first time reported. The agreement found between the values derived from independent measurements reflect the reliability of complexation and solution data reported in this paper. The relationship between complexation and solvation entropy of barium is discussed taking into account previous thermodynamic data on univalent cations and cryptand 222. It is concluded that the effective shielding effect of the ligand for univalent cations weakens considerably for barium and breaks completely for some of the lanthanide cations as assessed from thermochemical data on these systems.
Previous investigations have shown that cation solvation plays a predominant role in the complexation of univalent cations with cryptand 222 in dipolar aprotic media.1-5 Thus, the linear correlation established in terms of entropies between cation solvation and complexation in these media6 provided a strong indication that the cation undergoes complete desolvation when entering the cavity of the ligand. On the other hand, solution thermodynamic studies on lanthanide and lanthanide cryptate electrolytes in dipolar aprotic solvents (acetonitrile and propylene carbonate) showed no net change in the transfer enthalpies of the free compared to the complexed cation.7 These data led to the suggestion that as far as these tervalent cations are concerned these solvents are able to recognize the cation in its free or complexed form. Regarding alkaline-earth metal cations, most efforts have been devoted to determine the stability constants of their complexes with cryptand 222.8 Enthalpy data are scarce,9 and there is hardly any information concerning the solution properties of metal-ion(II) cryptates except that related to the thermochemical characterization of alkaline-earth cryptates in acetonitrile and in propylene carbonate. One of the driving forces for the synthetic developments in the area of macrocyclic chemistry since the late sixties when Pedersen discovered the crown ethers10 is the design of ligands able to bind selectively neutral or ionic species. A direct implication of this statement is the need of accurate thermodynamic data, the derivation of which is strongly dependent on (i) the methodology employed and to a great extent (ii) the correct characterization of the process to which the data are referred to, particularly, in ion complexation processes in lowpermittivity media where ion-pair formation may take place. X
Abstract published in AdVance ACS Abstracts, February 1, 1997.
S1089-5647(96)02568-0 CCC: $14.00
In this paper, enthalpies of complexation of barium(II) and cryptand 222 in various solvents, N,N-dimethylformamide (DMF), dimethyl sulfoxide (Me2SO), acetonitrile (MeCN), and propylene carbonate (PC) determined calorimetrically at 298.15 K are reported. The data combined with corresponding Gibbs energies derived from stability constants reported in the literature8 (except in acetonitrile) yield the entropy of complexation of this cation and cryptand 222 in these solvents. This is followed by calorimetric measurements of the solution enthalpies of barium cryptand perchlorates in protic (water and methanol) and dipolar aprotic solvents (DMF, Me2SO) at 298.15 K and corresponding data for barium perchlorate in DMF, Me2SO, MeCN, and PC at the same temperature. The thermochemical behavior of the barium cryptand electrolyte is discussed with respect to that of the free salt in these solvents. Single-ion values of transfer enthalpies are calculated on the basis of the Ph4AsPh4B convention. Complexation and solution data for the ligand (222) and the free and complexed salt [Ba2+222(ClO4-)2] are combined to derive for the first time the enthalpy of coordination, ∆coordH°, referred to the process (eq 1) where reactants and product are in the solid (sol.) state.
Ba(ClO4)2(sol.) + 222(sol.) f [Ba2+222](ClO4-)2(sol.) (1) The coordination enthalpy for this system is discussed and compared with corresponding data for the process involving KClO4 and this ligand. Experimental Section (i) Chemicals. Kryptofix 222 (Fluka), tetra-n-butylammonium perchlorate (Fluka), silver perchlorate (Aldrich), and © 1997 American Chemical Society
1644 J. Phys. Chem. B, Vol. 101, No. 9, 1997 barium perchlorate (Aldrich) were dried under vacuum for several days before use. Barium-cryptand 222 perchlorate was prepared from barium perchlorate and cryptand 222. The product was recrystallized from methanol and dried under vacuum. Microanalysis was carried out at the University of Surrey. Anal. Calcd (%): C, 30.33; H, 5.09; N, 3.93. Found: C, 30.39; H, 5.38; N 3.94. (ii) Purification of Solvents. Methanol (Fisher, HPLC grade) was used without further purification. Acetonitrile (Fisher, HPLC grade) was refluxed and distilled from calcium hydroxide. Dimethylformamide (Fisher, AR) and dimethyl sulfoxide (Aldrich 99%) were left over dried 4 Å molecular sieves for 48 h before use. Propylene carbonate (Aldrich 99%) was left overnight over dried 4 Å molecular sieves. It was then distilled under reduced pressure. The water content of the solvents determined by Karl Fisher titration was not higher than 0.02% for methanol, acetonitrile, and propylene carbonate. For dimethyl sulfoxide and N,N-dimethylformamide, the water content did not exceed 0.05%. (iii) Determination of Enthalpies of Solution. Enthalpies of solution of barium and barium cryptand 222 perchlorates in water and in the freshly purified solvents (methanol, MeCN, DMF, Me2SO, and PC) were determined calorimetrically at 298.15 K using the Tronac 450 calorimeter. Measurements in water were carried out in alkaline medium using tetra-nbutylammonium hydroxide in order to avoid the hydrolysis of the ligand. The accuracy of the instrument was periodically checked by using the standard reaction of tris(hydroxymethyl)aminomethane (THAM) with HCl (0.01 mol dm-3) suggested by Irving and Wadso¨.12 All heats were corrected for the heat of breaking of empty ampules in the appropriate solvent which was determined experimentally. (iv) Determination of Enthalpies of Complexation. The enthalpies of complexation of barium and cryptand 222 were determined by titration calorimetry using a Tronac 450 calorimeter. The experiments were carried out at 298.15 K. The reliability of the equipment was tested by using the reaction of protonation of THAM with hydrochloric acid suggested by Wilson and Smith.13 For these experiments, the calorimetric vessel was filled with 50 cm3 of a solution of barium perchlorate in the appropriate solvent, and cryptand 222 in solution (concentrations used were about 10 times higher than those of barium perchlorate) was injected from a syringe. Blank experiments were carried out to account for heat of dilution effects resulting from the addition of the ligand to the solvent contained in the calorimetric vessel. (v) Stability Constant Measurements. The stability constant of Ba2+ and cryptand 222 in acetonitrile was determined by potentiometric titrations. The measurements were carried out in a Camlab potentiometer, using the electrochemical cell suggested by Schneider and co-workers14 with a silver-silverion reference electrode consisting of a silver wire introduced in a solution of silver perchlorate. To keep a constant ionic strength, a solution of tetra-n-butylammonium perchlorate (TBAP) was used at the indicator electrode. Both electrodes were separated by a salt bridge containing 0.05 mol dm-3 of TBAP in the appropriate solvent. Experiments were carried out at 298.15 K. The solution of the silver perchlorate (0.01 mol dm-3) was added (15 additions) into the cell containing TBAP solution (25 cm3; 0.05 mol dm-3) in the appropriate solvent. Potential readings were taken after each addition, and the data were used to calculate the standard electrode potential of the reference cell. Then, in the same experiment, a solution of cryptand 222 (2.5 × 10-3 mol dm-3) in acetonitrile was added into the cell containing a solution of silver perchlorate in the
Danil de Namor and Kowalska TABLE 1: Thermodynamic Parameters of Complexation of Barium and Cryptand 222 in Various Solvents at 298.15 K solvent
log Ks
water 9.45a MeOH 12.89b
-53.97a -73.58
DMF 8.04b Me2SO 6.36b PC 17.10b MeCN 17.90 ( 0.12c a
∆cG°/kJ mol-1 ∆cH°/kJ mol-1 ∆cS°/K-1 mol-1
-45.92 -36.30 -97.61 -102.18
-59a -71b -68.9b -50.60 ( 0.30c -47.78 ( 0.28c -103.43 ( 3.10c -108.84 ( 1.30c -108.80b
-16.9a +12.0b -15.7c -38.5c -19.5c -22.3c
Reference 15. b Reference 8. c This work, see text.
appropriate solvent. The data were used to calculate the stability constant of the silver cryptate complex (Ks) in acetonitrile. Titrations were carried out until an excess of titrant was added. Finally, additions of barium perchlorate (0.01 mol dm-3) in this solvent were made to the same electrochemical cell. Potential readings were used to calculate the equilibrium constant (K1) for the competition process between silver cryptate and Ba2+ in acetonitrile. Combination of these data (Ks and K1) led to the calculation of the stability constant of barium cryptate in acetonitrile. Results and Discussion Thermodynamics of Complexation. Standard enthalpies of complexation, ∆cH°, of Ba2+ and cryptand 222 determined calorimetrically at 298.15 K in DMF, Me2SO, MeCN, and PC are reported in Table 1. These data are referred to the process
Ba2+(s) + 222(s) f Ba2+222(s)
(2)
where s denotes the solvent. The standard deviation of the data (σ) was calculated from
σ)
(
)
∑(xi - xj)2 1/2 n-1
(3)
Errors are given as twice the σ values. Previous literature data8,15 for this system in water, methanol, and acetonitrile at 298.15 K determined calorimetrically are also included in Table 1. This table shows that our value for the enthalpy of complexation of Ba2+ and cryptand 222 in acetonitrile is in excellent agreement with the data previously reported for this system in this solvent. Since there is no quantitative information regarding the stability constant of Ba2+ and cryptand 222 in acetonitrile at 298.15 K, the log Ks value obtained by competitive potentiometric titration using silver electrodes is reported in Table 1 together with literature data for this system in the various solvents.8,15 Combination of Gibbs energies (∆cG°) and enthalpies yields entropies of complexation (∆cS°) of Ba2+ and cryptand 222 in DMF, Me2SO, PC, and MeCN. These data analyzed in terms of Gibbs energies reflect that the strength of complexation is greater in MeCN and PC and lower in DMF and Me2SO than in H2O and in MeOH. In fact, the degree of complexation is strongly dependent on the nature of the solvent to the extent that the stability pattern follows the sequence
MeCN > PC > MeOH > H2O > DMF > Me2SO (4) The relationship between cation solvation and complex stability in the various solvents can be investigated by considering single-ion values for the transfer Gibbs energy, ∆tG°, of Ba2+ from water to these solvents since these data reflect the changes in cation solvation from one solvent to another. Transfer data for multivalent ions are scarce in relation to corresponding data for univalent cations. This is not surprising given the complexities involved in the derivation of solution
Solution Thermodynamics of Ba(II)-Cryptand 222
J. Phys. Chem. B, Vol. 101, No. 9, 1997 1645
TABLE 2: Gibbs Energies of Transfer of Cryptand 222 and Its Barium Cryptate Cation from Water to Nonaqueous Solvents at 298.15 K transfer
∆tG°(Ba2+222)a/kJ mol-1
∆tG°[222]b/kJ mol-1
H2O f MeOH H2O f Me2SO H2O f MeCN
+3.52 -2.83 +15.41
+4.73 +4.60 +6.32
a
See text. b From refs 1, 6, and 17.
Gibbs energies, ∆sG°, of fully dissociated salts involving multivalent ions, particularly in nonaqueous media. However, Ahrland16 reported ∆tG° values for barium from water to MeOH (18.4 kJ mol-1), Me2SO (-25.1 kJ mol-1), and MeCN (57.3 kJ mol-1) at 298.15 K. These data based on the Ph4AsPh4B convention reflect that Ba2+ is selectively solvated by these solvents in the following sequence
Me2SO > H2O > MeOH > MeCN
(5)
In fact, the trend observed in cation solvation (eq 5) is just the opposite to that observed in complex stability (eq 4). In order to assess whether or not a quantitative correlation can be found between the different ∆cG° values in these solvents and the ∆tG° values of Ba2+ from a reference solvent (s ) H2O or MeCN) to another, identity 6 (observed for this ligand and univalent cations in dipolar aprotic solvents6) is tested. Thus
∆cG°(s1) - ∆cG°(s2) ) ∆tG°[Ba2+](s1fs2)
(6)
By inserting the appropriate ∆cG° (Table 1) and ∆tG°[Ba2+] (see above) values into eq 6, (s1 f s2, H2O f MeOH; ∆tG° ) 18.4 kJ mol-1, ∆∆cG° ) 19.61 kJ mol-1; H2O f Me2SO; ∆tG° ) -25.1 kJ mol-1, ∆∆cG° ) -17.67 kJ mol-1; H2O f MeCN; ∆tG° ) 57.3 kJ mol-1, ∆∆cG° ) 48.21 kJ mol-1; MeCN f Me2SO; ∆tG° ) -82.4 kJ mol-1, ∆∆cG° ) -65.88 kJ mol-1 ), it is shown that while identity 6 is followed in the watermethanol solvent system, this does not strictly apply to other systems. Expansion of eq 6 as to involve the contribution due to the transfer of the ligand and the metal-ion complex leads to eq 7.1-6
∆cG°(s1) - ∆cG°(s2) ) ∆tG°[Ba2+](s1fs2) + ∆tG°[222](s1fs2) - ∆tG°[Ba2+222](s1fs2) (7) If identity 6 is not held for the Ba2+-cryptand 222 system, it follows that
∆tG°[222](s1fs2) * ∆tG°[Ba2+222](s1fs2)
(8)
The high stability of the Ba2+222 complex led us to isolate the perchlorate salt of the complexed cation. However, the high solubility of this salt in these solvents impeded the derivation of transfer Gibbs energies of [Ba2+222](ClO4)2, and consequently, we were unable to calculate single-ion values of [Ba2+222] from water to other medium by this method. These data were calculated from eq 7 using ∆tG° values for cryptand 2221,6,17 from water to the various solvents. Comparison of these data with those for the complexed cation (Table 2) shows that in the water-methanol system, both ligand and cation are slightly better solvated in water than in methanol. For others, including the transfer process between two dipolar aprotic solvents, data for the ligand differ from that for the complexed cation as expressed in eq 7. We therefore conclude that, in terms of Gibbs energies, the degree to which barium interacts
with cryptand 222 is mainly dependent on the solvation of both the free and complexed cations in most solvents. Allthough the Gibbs energy determines the position of equilibrium of the process, this thermodynamic function is essentially the combination of changes in enthalpy and entropy. Thus, Table 1 shows that the process represented by eq 2 is enthalpically controlled. The limitations discussed above in terms of Gibbs energies are not applicable to the determination of the solution enthalpies of free and complexed barium salts in these solvents. In order to analyze the variations in enthalpy observed in the complexation of barium and cryptand 222 in these solvents, enthalpies of solution of barium and barium cryptate salts in the appropriate solvent at 298.15 K determined calorimetrically are now discussed. Enthalpies of Solution. Enthalpies of solution of barium perchlorate in DMF, Me2SO, MeCN, and PC at 298.15 K at various ionic strenghts (I1/2) are listed in Table 3. In cases where variations in enthalpy were observed with changes in the ionic strength of the solution, the standard value, ∆sH°, was taken as the value at I ) 0 from a plot of ∆sH against I1/2. When no variations in ∆sH were observed with the ionic strength of the solution, ∆sH° is the average value of these measurements. Errors in enthalpy values are expressed as twice the standard deviation of the data. Enthalpies of solution of barium-cryptand 222 perchlorate were determined in water, MeOH, DMF, and Me2SO at 298.15 K. Standard enthalpies of solution of [Ba2+222](ClO4)2 in H2O and Me2SO are the average of the ∆sH values given in Table 3 since no significant changes are observed with changes in the ionic strength of the solution. In order to facilitate the discussion, ∆sH° values for Ba(ClO4)2 and [Ba2+222](ClO4)2 in the various solvents are summarized in Table 4. Also included in this table are literature data for Ba(ClO4)2 in water18 and MeOH19 at 298.15 K and for the complexed salt in MeCN and PC at the same temperature.9 As far as the solution enthalpies of barium perchlorate are concerned, the data shown in Table 4 reflect that except for water the dissolution of this salt in these solvents leads to exothermic heats implying that the solvation process (exothermic) predominates over the crystal lattice (endothermic) process. For a given compound, the latter contribution to the solution enthalpies in the various solvents is the same. Therefore, the differences observed in the ∆sH° values are due to the different degree of solvation of Ba(ClO4)2 in these solvents. The slightly endothermic character of the dissolution process in water must be associated with the large energy required to break the highly ordered structure of water relative to other solvents. As far as ∆sH° values for [Ba2+222](ClO4)2 are concerned, the solvation process predominates in DMF and Me2SO while the dissolution of this salt in water, MeOH, MeCN, and PC is endothermic (see Table 4). In moving from Ba(ClO4)2 to [Ba2+222](ClO4)2, there is a considerable increase in cation size, and therefore, the energy required to break the solid is expected to be lower for the latter relative to the former. Therefore, on the assumption that in a given solvent Ba2+undergoes the same degree of solvation as [Ba2+222], the solution enthalpy of Ba(ClO4)2 would be expected to be more endothermic than that for [Ba2+222](ClO4)2. The results shown in Table 4 indicate that this is not the case. The endothermic (or lower exothermic) character of the dissolution process involving the cryptate salt relative to Ba(ClO4)2 suggests that the [Ba2+222] cation is less solvated than the free cation. An implication of these findings is that either partially or totally cryptand 222 exerts a shielding effect on the cation. This is best reflected in the ∆tH° of the dissociated electrolytes Ba(ClO4)2 and [Ba2+222](ClO4)2 from water to the various solvents shown in Table 5. For Ba(ClO4)2,
1646 J. Phys. Chem. B, Vol. 101, No. 9, 1997
Danil de Namor and Kowalska
TABLE 3: Enthalpies of Solution of Barium and Barium Cryptand 222 Perchlorates in Various Solvents at 298.15 K in kJ mol-1 Ba(ClO4)2 N,N-dimethylformamide
dimethyl sulfoxide
acetonitrile
I1/2
∆sH
I1/2
∆sH
0.043 0.046 0.061 0.086 0.113 0.139
-110.23 -113.62 -112.66 -110.69 -113.58 -110.69
0.042 0.051 0.061 0.071 0.091 0.131
-104.46 -102.83 -104.29 -103.10 -103.09 -104.18
∆sH° ) -111.91 ( 3.10
∆sH° ) -103.66 ( 1.46
I1/2
propylene carbonate ∆sH
0.034 -21.72 0.049 -19.28 0.062 -18.53 0.076 -17.93 0.089 -15.25 0.102 -14.91 0.128 -11.89 ∆sH° ) -24.77 ( 0.56
I1/2
∆sH
0.028 0.030 0.049 0.061 0.065 0.073
-23.02 -22.55 -23.31 -23.86 -23.50 -23.55
∆sH° ) -23.30 ( 0.92
[Ba222](ClO4)2 water 1/2
I
methanol ∆sH
0.034 29.70 0.041 29.66 0.065 29.91 0.083 30.21 0.089 29.37 0.095 30.54 ∆sH° ) 29.90 ( 0.84a a
I1/2
N,N-dimethylformamide I1/2
∆sH
0.024 16.62 0.029 16.11 0.043 14.47 0.045 15.04 0.056 14.54 0.068 14.76 ∆sH° ) 17.23 ( 0.58b
∆ sH
0.029 -19.01 0.037 -17.20 0.048 -16.94 0.055 -16.38 0.092 -15.64 0.106 -15.17 ∆sH° ) -19.14 ( 0.66b
dimethyl sulfoxide I1/2
∆ sH
0.028 -11.86 0.041 -11.70 0.049 -11.87 0.055 -11.34 0.065 -11.31 0.083 -11.31 ∆sH° ) -11.57 ( 0.56a
Average value. b Extrapolated value at I ) 0.
TABLE 4: Standard Enthalpies of Solution of Barium Perchlorate and Barium-Cryptand 222 Perchlorate in Various Solvents at 298.15 K ∆sH°/kJ mol-1 water MeOH DMF Me2SO MeCN PC a
Ba(ClO4)2
[Ba222](ClO4)2
4.18a -59.83b -111.91c -103.66c -24.77c -23.30c
29.90c 17.23c -19.14c -11.57c 11.93d 11.77d
Reference 18. b Reference 19. c This work. d Reference 9.
the data reflect that the stability in enthalpic terms follows the sequence
DMF > Me2SO > MeOH > MeCN = PC > H2O (9) This trend is altered for [Ba2+222](ClO4)2 since the data indicate that the enthalpic stability is
DMF > Me2SO > PC = MeCN > MeOH > H2O
(10)
A clearer picture emerges by partitioning the transfer parameters for the dissociated electrolytes into single-ion quantities. For this purpose, transfer enthalpies of the perchlorate anion from H2O to the various solvents are used.11,20 The data are based on the Ph4AsPh4B convention.11 Thus
∆tH°[Ba2+](H2Ofs) ) ∆tH°Ba(ClO4)2(H2Ofs) 2∆tH°[ClO4-](H2Ofs) (11) and
∆tH°[Ba2+222](H2Ofs) ) ∆tH°[Ba222](ClO4)2(H2Ofs) 2∆tH°[ClO4-](H2Ofs) (12) Details are included in Table 5. The ∆tH°(Ba2+) from water to methanol at 298.15 K is in good agreement with the value of -59.2 kJ mol-1 based on the same convention reported by Ahrland.16 Such agreement is not found in the transfer of this
cation from water to Me2SO (∆tH° ) -78.5 kJ mol-1) and to MeCN (∆tH° ) -8.5 kJ mol-1). Comparison of transfer data for Ba2+ relative to [Ba2+222] shows that in moving from the free to the complexed cation, the medium effect is greatly reduced, particularly, in transfers to MeOH, Me2SO, and DMF. This effect is not too significant in transfers to acetonitrile and propylene carbonate since these solvents are poor cation solvators. These results are interesting and strongly suggest that in [Ba2+222] the cation is shielded by the ligand. To assess the factors which contribute to the differences observed in the enthalpies of complexation of Ba2+ and cryptand 222 in the various solvents (Table 1), ∆tH° values for Ba2+ (Table 5) are considered. These data show that except for methanol, the higher the enthalpic stability of this cation is in a given solvent (more negative ∆tH° value in Table 5), the lower (less exothermic) its enthalpy of complexation with cryptand 222 in that solvent (Table1). However, identity 6 tested in terms of enthalpy shows that, including or excluding protic solvents (H2O, MeOH), this does not hold for the Ba2+-cryptand 222 system. Therefore, taking into account eq 7, it follows that
∆tH°[Ba2+222](s1fs2) * ∆tH°[222](s1fs2)
(13)
The different thermochemical behavior of cryptand 222 relative to the complexed cation is corroborated by comparing the transfer enthalpies of the free ligand1,2,17 and those of [Ba2+222] (Table 5). Provided that water is excluded, the ∆tH° for cryptand 222 between nonaqueous solvents is approximately equal to 0. For pairs of solvents of similar properties, it is also found that ∆tH°[Ba2+222](s1fs2) = 0. Such is the case for transfers involving DMF and Me2SO (dipolar protophilic solvents) or PC and MeCN (dipolar protophobic solvents). However, this is not universally observed. Therefore, it is concluded that expression 13 holds in terms of Gibbs energies and enthalpies. Consequently, the differences observed in the enthalpic stability of the Ba2+-cryptand 222 complex with solvent variation is most clearly shown in eq 7 (expressed in enthalpies). The results for nonaqueous solvents show that the different ∆cH° values (Table 1) resulting from medium effects depend on the balance between the transfer enthalpies of the free and complexed cation. When data in water are included, the transfer enthalpy of cryptand 222 contributes significantly
Solution Thermodynamics of Ba(II)-Cryptand 222
J. Phys. Chem. B, Vol. 101, No. 9, 1997 1647
TABLE 5: Standard Enthalpies of Transfer of Barium and Barium Cryptand 222 Perchlorates from Water (Reference Solvent) to Various Solvents. Derived Single-Ion Values of Ba2+ and [Ba2+222] from Water Based on the Ph4As Ph4B Convention at 298.15 K ∆tH°/kJ mol-1 transfer
Ba(ClO4)2
[Ba222](ClO4)2
Ba2+ a
[Ba2+222] a
[222]b
H2O f H2O H2O f MeOH H2O f DMF H2O f Me2SO H2O f MeCN H2O f PC
0 -64.01 -116.09 -107.84 -28.95 -27.48
0 -12.67 -49.04 -41.47 -17.97 -18.13
0 -58.99 -70.91 -69.34 6.11 5.40
0 -7.65 -3.86 -2.97 17.09 14.75
0 58.97 59.28 60.20 59.22 57.73
a For these calculations, ∆tH°ClO4-(H2Ofs) in kJ mol-1 based on the Ph4As Ph4B convention were as follows: s, MeOH ) -2.51; DMF ) -22.59; Me2SO ) -19.25; MeCN ) -17.53; PC ) -16.44. (Ref 20). b References 1, 6, and 17.
TABLE 6: Enthalpies of Coordination (kJ mol-1) of Ba(ClO4)2 and Cryptand 222 in the Solid State at 298.15 K Derived from Solution and Complexation Data in Six Different Solvents (See Text) from data in
∆coordH° a
H2O MeOH DMF Me2SO PC MeCN
-109.5 -113.5 -108.9 -104.4 -104.0 -112.6 average ) -108.8 ( 4.0
a For these calculations, ∆ H° values (kJ mol-1) for cryptand 222 in s water, MeOH, DMF, Me2SO, PC, and MeCN are -24.76, 33.51, 34.52, 35.44, 34.46, and 32.97, respectively. References 1 and 17.
to the ∆(∆cH°) in MeCN or PC relative to water. For MeOH, the contribution of the free cation from water to this solvent (exothermic) almost completely cancels that of the ligand (endothermic). It is the higher enthalpic stability of [Ba2+222] in MeOH relative to water that increases the exothermic character of the complexation process in the former relative to the latter solvent. However, the contributions of the cryptand and cryptate are not enough to overcome the higher enthalpic stability of Ba2+ in DMF and Me2SO relative to water. As a result, the complexation enthalpy is lower (less exothermic) in these dipolar aprotic solvents than in water. The discussion above may be related to previous work on the thermodynamics of cation complexation process involving 222. For univalent cations, the charge-dipole interactions between cation and solvent (dipolar aprotic solvents) are weak to the extent that the cryptand and its metal-ion cryptates are characterized by almost the same transfer parameters.1-6 This was attributed to the effective shielding exerted by the ligand on the cation which leaves the solvent unable to distinguish between free and complexed ligand. However, this effect seems to weaken considerably for a bivalent cation and breaks completely by the inclusion of a tervalent cation in the ligand cavity since for the latter the transfer enthalpy of the free multicharged ion was found to be almost identical to that of the complexed cation.7 Enthalpies of Coordination. For the calculation of enthalpies of coordination, ∆coordH°, referred to the process in the solid (sol.) state, eq 1, ∆sH° values for cryptand 222,1,17 barium, and barium cryptate salts (Table 6) in a given solvent are combined with the complexation enthalpy of this cation and this ligand (Table 1) in the same solvent. For this purpose, the following equation is used
∆coordH° ) ∆sH°Ba(ClO4)2(s) + ∆sH°222(s) + ∆cH°(s) - ∆sH°[Ba222](ClO4)2(s) (14) The magnitude of the ∆coordH° for this system should be the same, independent of the solvents from which this is derived.21
TABLE 7: Hydration and Solvation Entropy of Barium and Complexation Entropies in Water, Methanol, Dimethyl Sulfoxide, and Acetonitrile at 298.15 K in J K-1 mol-1 solvent H2O MeOH Me2SO MeCN a
∆hydS° a
∆solvS° a
∆cS° b
-313.3 -202.1 -225.4
-16.9 +12.0 -38.5 -22.3
-53.7
Reference 16. b From Table 1.
Therefore, these calculations offer a suitable means to check the reliability of the solution data. Enthalpies of coordination for the process represented by eq 1 derived from six different solvents, each one containing four sets of experimental data are listed in Table 6. Excellent agreement is found between the data derived from independent measurements.Therefore, an average is taken for the ∆coordH° of the barium salt and cryptand 222. While in the solvents considered, the complexation process is free from anion effects as it should be, the magnitude of ∆coordH° is strongly dependent on the nature of the anion.22 Therefore, for useful comparison, the ∆coordH° of KClO4 (a salt containing a cation of similar size to Ba2+and the same anion) and cryptand 222 is considered (∆coordH° = -39 kJ mol-1).23 Specific charge effects are reflected in the coordination process. Despite that Ba2+ and K+ have similar size, the high charge density of the former relative to the latter leads to a greater enthalpic stability for the salt containing the bivalent relative to that involving the univalent cation. We therefore reinforce previous statements22 that coordination enthalpies are suitable reporters of interactions taking place in the solid state. Entropy Data. Entropies of complexation are strongly dependent on the number of species involved, solvation, and conformational changes taking place in the process. The combination of metal cation and ligand to give the complex is expected to lead to entropy losses. This is observed for most solvents, except methanol (Table 1). However, cation and ligand desolvation upon complexation leads to entropy gain. The relationship between complexation and cation solvation can be explored in terms of entropy. Solvation entropy data for Ba2+ in MeOH, Me2SO, and MeCN recorded in Table 7 were calculated from combination of hydration, ∆hS° 16 (Table 7), and transfer, ∆tS°, entropies (eqs 15 and 16). ∆hS°
Ba2+(g) 98 Ba2+(H2O) ∆tS°
Ba2+(H2O) 98 Ba2+(s)
(15) (16)
Inspection of Table 7 shows that, excluding water, a linear relationship between solvation and complexation entropies in MeOH, Me2SO, and MeCN (straight line; cc ) 0.986) is found reflecting that the greater the loss of entropy of Ba2+ upon solvation the greater its desolvation upon complexation
1648 J. Phys. Chem. B, Vol. 101, No. 9, 1997
Danil de Namor and Kowalska
(∆cS° more positive). However, the data do not fit into the correlation (eq 17) previously demonstrated for univalent cations and cryptand 222 in dipolar aprotic solvents6
∆cS° ) constant - ∆solvS°
(17)
since the entropy of cryptate formation, ∆cfS°, referred to the process ∆cfS°
M2+(g) + 222(s) 98 M2+222(s)
(18)
is not constant for this system. This is in accord with the conclusions derived from Gibbs energy and enthalpy data that, in the complexation of this cation with cryptand 222, the medium effect leads to differences in the thermodynamics of complexation of this system and, in nonaqueous solvents, that are not only due to the desolvation of the free cation but also to the solvation changes of the metal-ion cryptate in these solvents. In water, the correlation between hydration and complexation breaks down completely in this solvent (see Table 7). This could be attributed to water-ligand interactions. Final Remarks This paper demonstrates the following: (i) In assessing the complexation of macrocycles and guest species, the solution thermodynamics of the chemical entities participating in the process should be investigated. (ii) Although solvent-cation interactions are weak in cryptate complexes of univalent cations in dipolar aprotic media since the solvents are unable to distinguish between the free and complexed ligand, these interactions become stronger with an increase in cation charge to the extent that, as previously shown, for some of the lanthanides,7 the solvent is able to recognize selectively the tervalent ion in its free or complexed form. However, for these tervalent cations, investigations were limited to MeCN and PC, and these should be extended as to cover a wide range of dipolar aprotic solvents. (iii) Coordination enthalpies are suitable reporters of interactions in the solid state where charge effects are clearly reflected. In addition, these provide an useful mean to check the reliability of thermodynamic data.
Acknowledgment. The authors thank the TEMPUS programme (JEP. 7891) for a scholarship to D. Kowalska and Professors A. Hulanicki and S. Glab (University of Warsaw, Poland), Dr. A. Emsley (University of Surrey), and Dr. J. L. Liebenguth (EHICS, Strasbourg, France) for arranging it. This work was supported by BST 532/2/96. References and Notes (1) Danil de Namor, A. F.; Ghousseini, L. J. Chem. Soc., Faraday Trans. 1 1984, 80, 2843; 1985, 81, 781; 1986, 82, 3275. (2) Danil de Namor, A. F.; Ghousseini, L.; Lee, W. H. J. Chem. Soc., Faraday Trans. 1 1985, 81, 2495. (3) Danil de Namor, A. F.; Ghousseini, L.; Hill, T. J. Chem. Soc., Faraday Trans. 1 1986, 82, 349. (4) Danil de Namor, A. F.; Fernandez Salazar, F.; Greenwood, P. J. Chem. Soc., Faraday Trans. 1 1987, 83, 1569. (5) Danil de Namor, A. F.; Fernandez Salazar, F. J. Chem. Soc., Faraday Trans. 1 1988, 84, 3539. (6) Danil de Namor, A. F. J. Chem. Soc., Faraday Trans. 1988, 84, 2441. (7) Danil de Namor, A. F.; Ritt, M. C.; Schwing-Weill, M. J.; Arnaud Neu, F. J. Chem. Soc., Faraday Trans. 1990, 86, 89. (8) Izatt, R. M.; Pawlak, K.; Bradshaw, J. S. Chem. ReV. 1991, 91, 1721. (9) Danil de Namor, A. F.; Hill, T.; Walker, R. A. C.; Contreras Viguria, E.; Berroa de Ponce, H. J. Chem. Soc., Faraday Trans. 1 1988, 84, 255. (10) Pedersen, C. J. J. Am. Chem. Soc. 1967, 89, 2495, 7017. (11) Cox, B. G.; Hedwig, G. R.; Parker, A. J.; Watts, D. W. Aust. J. Chem. 1974, 27, 477. (12) Irving, R. J.; Wadso¨, I. Acta Chem. Scand. 1964, 18, 195. (13) Wilson, E. W., Jr.; Smith, D. F. Anal. Chem. 1969, 41, 1903. (14) Gutknecht, J.; Schneider, H.; Stroka, J.; Inorg. Chem. 1978, 17, 3326. (15) Kauffmann, E.; Lehn, J. M.; Sauvage, J. P. HelV. Chim. Acta. 1976, 59, 1099. (16) Ahrland, S. Pure Appl. Chem. 1982, 54, 1451. (17) Abraham, M. H.; Danil de Namor, A. F.; Schulz, R. A. J. Chem. Soc., Faraday Trans. 1980, 76, 869. (18) Abraham, M. H.; Ah Sing, E.; Danil de Namor, A. F.; Hill, T.; Nasehzadeh, A.; Schulz, R. A. J. Chem. Soc., Faraday Trans. 1 1978, 74, 359. (19) Drakin, S. I.; Chang, Y. M. Zh. Fiz. Khim. 1964, 38, 2800. (20) Ghousseini, L. Ph.D. thesis, University of Surrey, 1985. (21) Danil de Namor, A. F.; Llosa Tanco, M. A.; Salomon, M.; Ng, J. C. Y. J. Phys. Chem. 1994, 98, 11796. (22) Danil de Namor, A. F.; Ng, J. C. Y.; Llosa Tanco, M. A.; Salomon, M. J. Phys. Chem. 1996, 100, 14485. (23) Danil de Namor, A. F. Unpublished results.
