J. Phys. Chem. 1996, 100, 14485-14491
14485
Thermodynamics of Lithium-Crown Ether (12-crown-4 and 1-Benzyl-1-aza-12-crown-4) Interactions in Acetonitrile and Propylene Carbonate. The Anion Effect on the Coordination Process Angela F. Danil de Namor,*,† Joe C. Y. Ng,† Margot A. Llosa Tanco,† and Mark Salomon‡ Laboratory of Thermochemistry, Department of Chemistry, UniVersity of Surrey, Guildford, Surrey GU2 5XH, and Power Sources DiVision, U.S. Army ARL, Fort Monmouth, New Jersey 07703 ReceiVed: February 20, 1996; In Final Form: May 20, 1996X
Titration microcalorimetry in nonaqueous media (acetonitrile and propylene carbonate) has been used for the determination of stability constants (log Ks) and enthalpies of complexation of lithium and crown ethers (12-crown-4 and 1-benzyl-1-aza-12-crown-4 at 298.15 K. To ensure that the data are referred exclusively to the complexation process, salts containing highly polarizable anions (hexafluoroasenate, tetrafluoroborate, trifluoromethanesulfonate, and perchlorate) are used as sources for lithium. From stability constants and standard enthalpies, standard Gibbs energies and standard entropies are calculated. In propylene carbonate, a correlation is found between the stability of the lithium crown complex and the increase in conductance of the complexed relative to the free cation. Eight new lithium coronand salts of 12-crown-4 and 1-benzyl-1aza-12-crown-4 were isolated. Standard enthalpies of these salts and crown ethers in acetonitrile and propylene carbonate at 298.15 K measured calorimetrically are used to explain (i) the higher molar ionic conductivities, observed for lithium coronand relative to lithium electrolytes, and (ii) the effect of the solution properties of ligand, free, and complexed cation in the binding of these ligands with lithium in these solvents. Enthalpies of coordination first reported show the anion effect in the process involving reactants and product in their pure physical state. The strength of cation-anion interaction follows the sequence ClO4- > CF3SO3- > AsF6- > BF4-.
Introduction The chemistry of crown ethers has been extensively discussed in several books and research publications.1-3 Most thermodynamic studies on cation complexation involving macrocyclic ligands are limited to the binding process. Indeed, efforts aiming to characterize the properties of electrolytes resulting from the complexation of metal ion salts and crown ethers in particular and macrocycles in general are very limited. The importance of fundamental thermodynamics to explain the conductivity enhancement observed by the addition of crown ethers (1-aza-12-crown-4 (1A12C4) and 15-crown-5 (15C5)) to lithium electrolyte solutions in acetonitrile (AN) and in propylene carbonate (PC) has been previously discussed by our group.4,5 However, it is important to rationalize the effect of the thermodynamic stability of the lithium crown ether complex on the observed increase of electrolyte conductance in these solvents. To proceed with this evaluation, titration microcalorimetry was used to derive the stability constants (log Ks) and the enthalpy of complexation (∆cH°) of related crown ethers such as 12-crown-4 (12C4) and 1-benzyl-aza-12-crown-4 (BA12C4) and lithium in AN and in PC. The sensitivity of the microcalorimetric technique is much higher than that of classical titration calorimetry. It is with the latter that most thermodynamic data available on cation-macrocycle complexation reactions in difference solvents have been derived. As far as 12C4 is concerned, the interaction of this ligand with lithium has been studied by Buschmann,6 Anet et al.7 and others mainly in methanol where the formation of 2:1 complexes in this solvent has been reported, although 1:1 complexes are known to be formed in other solvents.8 We are not aware of †
University of Surrey. U.S. Army ARL. X Abstract published in AdVance ACS Abstracts, July 15, 1996. ‡
S0022-3654(96)00519-9 CCC: $12.00
any detailed thermodynamic studies on these crown ethers and lithium salts in these solvents involving the thermochemical characterization of 12C4 and BA12C4 or indeed their lithium coronand salts in these solvents. Studies of this kind are important since these provide information regarding the solvation of the reactants and the product participating in the complexation process. Another important aspect of these studies is that the availability of complexation and solution data enables the calculation of the enthalpy of coordination, ∆coordH°, referred to reactants and product in their pure physical state:
M+X-(sol) + CE (sol or l) f M+CEX- (sol)
(1)
In eq 1, M+X-, CE, and M+CEX- denote the metal-ion salt, the crown ether, and the metal-ion coronand salt, respectively, in the solid (sol) or liquid (l) state. To our knowledge, the role of the anion in the coordination process involving lithium and these crown ethers has not been investigated. In fact, most studies on the anion effect of complexation reactions involving crown ethers or cryptands are referred to as solution processes.9-12 Due to the distinctive features of crown ethers (holes) relative to cryptands (cavities), the anion effect on the coordination process may be quite significant. This paper attempts to establish whether or not enthalpies of coordination are suitable reporters of these effects. Therefore in the following sections we discuss the following: (i) The thermodynamics of complexation of lithium and crown ethers (Figure 1) in AN (MeCN) and PC (PC), solvents currently used in battery technology at 298.15 K. To ensure that the data are exclusively referred to the complexation process, lithium salts containing highly polarizable anions (hexafluoroarsenate, tetrafluoroborate, trifluoromethanesulfonate, and perchlorate) are used. Again these are the electrolytes commonly used in lithium batteries. The effect of complex stability on the conductance © 1996 American Chemical Society
14486 J. Phys. Chem., Vol. 100, No. 34, 1996
Danil de Namor et al. TABLE 1: Microanalysis Data for Lithium Coronand (12C4 and BA12C4) Salts calcd
Figure 1. (a) 12-crown-4 (12C4). (b) 1-Benzyl-1-aza-12-crown-4 (BA12C4).
enhancement of lithium coronand relative to lithium salts is assessed from the ionic molar conductivities of the complex relative to the free cation. (ii) The thermochemical characterization of the crowns, lithium, and eight new lithium coronand salts in AN and PC. (iii) The anion effect on the enthalpies of coordination (eq 1). Experimental Section (i) Chemicals. Lithium hexafluoroarsenate(V) (LiAsF6) (99%), lithium tetrafluoroborate (LiBF4) (99%), lithium trifluoromethanesulfonate (LiCF3SO3) (96%), and lithium perchlorate (LiClO4) (99%) were purchased from Aldrich. LiAsF6, LiBF4, and LiClO4 were left in a desiccator containing CaCl2 for several days before use. LiCF3SO3 was recrystallized from an acetone/toluene (1:4) mixture. Then, it was dried at 60 °C under low pressure for several days before use. 12C4, BA12C4, and 18-crown-6 (18C6) from Aldrich were used without further purification. Acetonitrile (HPLC) from Fision Chem Co. and PC from Aldrich were used as described elsewhere.4 Tris(hydroxymethyl)aminomethane (THAM) was twice recrystallized from a water:methanol (50:50) mixture. The crystals were washed with methanol and dried at room temperature for 24 h. Then these were stored in a vacuum desiccator for 3 days. (ii) Titration Calorimetric Measurements. The fourchannel heat conduction microcalorimeter (Thermometric 2277) designed by Suurkuusk and Wadso¨13 was used to determine the complexation of crown ethers and lithium salts in AN and in PC at 298.15 K. The equipment was calibrated chemically by determining the standard enthalpy of complexation of 18C6 and Ba2+ in water at 298.15 K using slow and fast titration modes. Thus, values for log Ks ) 3.72 ( 0.05, ∆cH° ) -31.71 ( 0.25 kJ mol-1 (slow titration), and log Ks ) 3.72 ( 0.05, ∆cH° ) -31.71 ( 0.31 kJ mol-1 (fast titration14) were obtained. These are in good agreement with the values (log Ks ) 3.77 ( 0.01, ∆cH° ) -31.42 ( 0.20 kJ mol-1) reported in the literature for the Ba2+-18C6 system in water at this temperature.15 Another chemical calibration used was the determination of the standard enthalpy of protonation of THAM in an aqueous solution of hydrochloric acid (0.1 mol dm-3) at 298.15 K. A value of -47.56 ( 0.14 kJ mol-1 is again in good agreement with the reported value of -47.49 ( 0.04 kJ mol-1 reported by Ojelund and Wadso¨.16 Microcalorimetric titrations of lithium salts and crown ethers in nonaqueous solvents were carried out at 298.15 ( 0.01 K. Thus, the calorimetric vessel was charged with the lithium salt solution in the appropriate solvent (2.8 cm3). The crown ether solution (0.07-0.08 mol dm-3) was incrementally added from a 500 µL gas-tight motor-driven Hamilton syringe. All calorimetric measurements were performed by the slow titration procedure at 298.15 K. In each titration experiment, about 15 injections (20-25 µL each) were made at time intervals of 1 h. Separate dilution experiments were carried out. The standard enthalpy change and the stability constant (log Ks) were calculated using a minimization program.17 (iii) Preparation of Lithium Coronand Salts. Solid complexes of cyclic polyethers were prepared by dissolving sto-
lithium coronand salt
%C
%H
[Li12C4]AsF6 [Li12C4]BF4 [Li12C4]CF3SO3 [Li12C4]ClO4 [LiBA12C4]AsF6 [LiBA12C4]BF4 [LiBA12C4]CF3SO3 [LiBA12C4]ClO4
25.83 35.59 32.54 34.00 39.06 50.17 45.61 48.46
4.33 5.97 4.85 5.71 5.03 6.46 5.50 6.24
found %N
%C
%H
%N
3.04 3.90 3.32 3.77
25.42 35.62 32.18 34.08 38.60 50.08 45.66 48.77
4.36 6.12 4.87 5.80 5.07 6.58 5.48 6.38
2.87 3.87 3.17 3.77
ichiometric amounts of the crown ether and the appropriate lithium salt in methanol. The solvent was carefully removed by evaporation, and the solid residue was dried under low pressure for several days. Microanalyses were carried out at the University of Surrey, and these are reported in Table 1. (iv) Calorimetric Measurements for the Determination of Solution Enthalpies. For the determination of enthalpies of solution of the crown ethers, lithium, and lithium coronand salts in AN and in PC, the TRONAC 450 calorimeter originally designed by Christensen and Izatt18 was used. The reliability of the equipment was checked by using the standard reaction of THAM in a 0.1 mol dm-3 aqueous solution of hydrochloric acid suggested by Irving and Wadso¨ in 1964.19 The value of -29.58 ( 0.02 kJ mol-1 for the standard enthalpy of solution of THAM in HCl at 298.15 K is in close agreement with that reported by Ghousseini20 (∆sH ) -29.76 ( 0.02 kJ mol-1) using the same equipment. For the determination of the enthalpies of solution of crown ethers, lithium, and lithium coronand salts in AN and in PC at 298.15 ( 0.01 K, sealed ampules containing accurately weighed amounts (to 0.0001) of the appropriate compound were broken in a volume of solvent (50 cm3) placed in the calorimetric vessel. The observed heat change was in all cases corrected for the heat of breaking of the empty ampule in the solvent. In all cases, an electrical calibration was performed by introducing a known quantity of electrical heat approximately equal to the energy change of the system due to the chemical reaction. Results and Discussion (i) Thermodynamics of Complexation. There are several techniques currently used for the determination of stability constant data for complexation reactions involving metal cations and macrocyclic ligands.3 However, the suitability of the technique is largely dependent on the magnitude of the stability constant. Since trial experiments showed that the log Ks values of these crown ethers and lithium are within the scope of titration microcalorimetry, this was the technique adopted. One of the main advantages of titration calorimetry is that not only the enthalpy associated with the complexation process can be derived but also the stability constant can be evaluated. Using the well-known thermodynamic relationship
∆cG° ) ∆cH° - T∆cS°
(2)
the entropy of complex formation, ∆cS° can be calculated. Thus, Table 2 lists stability constants and derived Gibbs energies, enthalpies, and entropies for the complexation of crown ethers (12C4 and BA12C4) and lithium salts containing different anions in AN and in PC at 298.15 K. For comparison purposes, data for 1A12C4 and lithium salts previously reported in these solvents4 are also included in Table 2. The individual errors in K and ∆cH° expressed as twice the standard deviation of the mean were calculated as detailed elsewhere.4
Thermodynamics of Lithium-Crown Ether Interactions
J. Phys. Chem., Vol. 100, No. 34, 1996 14487
TABLE 2: Stability Constants and Derived Gibbs Energies, Enthalpies, and Entropies of Complexation of Crown Ethers (12-crown-4 and 1-Benzyl-1-aza-12-crown-4) and Lithium (using salts containing different anions) in Acetonitrile and in Propylene Carbonate at 298.15 Ka electrolyte
a
macrocycle
∆cG° (kJ mol-1)
log Ks
Acetonitrile -18.4 ( 0.7 -19.8 ( 1.1 -20.1 ( 1.6 -18.9 ( 0.6
∆cH° (kJ mol-1)
∆cS° (J K-1 mol-1)
-22.78 ( 1.81 -21.66 ( 1.26 -21.35 ( 1.58 -21.87 ( 1.38
-14.7 -6.2 -4.6 -10.0
Li+AsF6Li+BF4Li+CF3SO3Li+ClO4-
12C4 12C4 12C4 12C4
3.23 ( 0.14 3.46 ( 0.18 3.52 ( 0.28 3.31 ( 1.14
Li+AsF6Li+BF4Li+CF3SO3-
1A12C4 1A12C4 1A12C4
4.23 ( 0.34b 4.24 ( 0.12b 4.23 ( 0.86b
-24.2 ( 0.8b -24.2 ( 0.4b -24.2 ( 0.6b
-18.84 ( 1.70b -19.91 ( 0.48b -18.69 ( 1.06b
Li+AsF6Li+BF4Li+CF3SO3Li+ClO4-
BA12C4 BA12C4 BA12C4 BA12C4
4.25 ( 0.08 4.30 ( 0.02 4.31 ( 0.04 4.31 ( 0.08
-24.3 ( 0.5 -24.6 ( 0.1 -24.6 ( 0.5 -24.6 ( 0.5
-27.14 ( 1.16 -27.54 ( 1.06 -27.44 ( 0.66 -28.50 ( 0.34
-9.5 -9.9 -9.3 -13.1
Li+AsF6Li+BF4Li+CF3SO3Li+ClO4-
12C4 12C4 12C4 12C4
2.81 ( 0.08 2.79 ( 0.06 2.84 ( 0.08 2.81 ( 0.08
-17.70 ( 0.08 -17.71 ( 0.54 -17.05 ( 0.90 -15.29 ( 0.08
-5.7 -6.1 -2.8 2.4
Li+AsF6Li+BF4Li+CF3SO3-
1A12C4 1A12C4 1A12C4
3.67 ( 1.06b 3.69 ( 0.44b 3.87 ( 0.84b
-21.0 ( 6.0b -21.1 ( 2.5b -22.1 ( 0.4b
-14.78 ( 1.64b -14.63 ( 0.64b -15.08 ( 0.66b
20.9b 21.7b 23.5b
Li+AsF6Li+BF4Li+CF3SO3Li+ClO4-
BA12C4 BA12C4 BA12C4 BA12C4
4.08 ( 0.10 4.39 ( 0.04 4.59 ( 0.08 4.32 ( 0.08
-23.3 ( 0.1 -25.1 ( 0.1 -26.2 ( 0.5 -24.7 ( 0.5
-24.59 ( 0.62 -24.98 ( 0.42 -24.70 ( 0.06 -23.30 ( 1.24
-4.3 0.4 5.0 4.7
Propylene Carbonate -16.0 ( 0.5 -15.9 ( 0.4 -16.2 ( 0.5 -16.04 ( 0.2
18.0b 14.4b 18.5b
Literature data for related systems are included. b Reference 4.
For these systems, the data fit into a model which corresponds to the formation of a 1:1 metal ion:ligand stoichiometry, and therefore, the data reported in Table 2 refer to the following process in solution (s)
Li+ (s) + CE (s) f Li+CE (s)
(3)
The composition of lithium complexes of 12C4 and BA12C4 investigated by conductance measurements gave a break at ligand metal ion concentration ratios equal to unity, which verify that, in these solvents, 1:1 metal ion-crown complexes are only formed. There is not good agreement between the log Ks values for lithium and 12C4 in AN at 298.15 K found in the literature. Thus, log Ks values of 4.25, 3.40, and 3.14 have been reported.5 We ourselves previously reported a value of 3.91 for the stability constant of lithium and 12C4 in this solvent.5 This value was derived from titration macrocalorimetry. However, the greater sensitivity of the microcalorimetric system used in this work leads us to suggest that the values reported in this paper are characterized by higher accuracy than the previously reported value. Excellent agreement is found between the data for lithium and 12C4 using the perchlorate salt (log Ks ) 2.90 ( 0.04) and for lithium and BA12C4 using tetrafluoroborate as the anion (log Ks ) 4.33 ( 0.02) in PC obtained from conductance data5 and those given in Table 2 for these systems. We consider that the differences in the thermodynamic data shown in Table 2 when different lithium salts are used are indeed very small and within the experimental error; thermodynamic data are independent of the anion. Therefore, the average value of 3.38 for log Ks of lithium and 12C4 in AN at 298.15 K is in good agreement with the value of 3.40 reported in the literature.8 The driving force for the synthetic developments in the field of macrocyclic chemistry is the design of ligands with selective properties for metal cations. Therefore, it is important to discuss these results in terms of the factors which contribute to the selective complexation of macrocycles and metal cations.
Among these are the nature of the donor atoms, the effect of the substituent in the crown ether, and the nature of the solvent. The thermodynamic data reflect the effect of the donor atoms of the ligand in the complexation process involving these solvents. Thus, the more polarizable nitrogen atoms afford a much stronger dipolar interaction as reflected in the higher stabilities of 1A12C4 and BA12C4 with lithium in both solvents relative to 12C4. Major features are observed in the enthalpy and entropy contributions. Thus, the complexation of a tertiary amine such as BA12C4 with lithium is characterized by a higher enthalpic stability than that for 1A12C4 which may be attributed to a lower enthalpy of ligand desolvation (endothermic process) in the former relative to the latter ligand. This is also corroborated by the more favorable entropy changes (positive) observed for the binding of lithium to 1A12C4 relative to its benzyl derivative in these solvents which are typical of processes in which strong desolvation occurs upon complexation. The thermodynamic data show that, in PC, the stability of these crowns for lithium follows the sequence BA12C4 > 1A12C4 > 12C4. However, in AN, within the uncertainties given, BA12C4 = 1A12C4 > 12C4 Ionic molar conductances λ° of lithium coronand cations in PC at 298.15 K (λ°(Li+12C4) ) 10.02 S cm2 mol-1; λ°(Li+1A12C4) ) 10.24 S cm2 mol-1; λ°(Li+BA12C4) ) 11.99 S cm2 mol-1)5 are higher than that for the free cation (λ°(Li+) ) 7.86 S cm2 mol-1)5 in the same solvent and at the same temperature. The data reflect that the highest enhancement in conductances is found for Li+BA12C4 in this solvent. In fact, there is a correlation between the stability of the lithium crown ether complex in PC (log Ks) and the increase in conductance (∆λ+° ) λ+°(Li+CE) - λ+°(Li+)) of the complexed cation relative to the free cation in this solvent. These results may have interesting implications in the selection of new lithium coronand electrolytes to be tested in battery technology. To assess the solvent effect on the complexation process, we proceeded with the thermochemical characterization of crown
14488 J. Phys. Chem., Vol. 100, No. 34, 1996
Danil de Namor et al.
TABLE 3: Enthalpies of Solution of 12-crown-4 and 1-Benzyl-1-aza-12-crown-4 in Acetonitrile and in Propylene Carbonate at 298.15 K acetonitrile -3
propylene carbonate -1
∆sH (kJ mol )
c (mol dm )
c (mol dm-3)
∆sH (kJ mol-1)
12C4 -3.68 1.57 × 10-2 -3.06 4.09 × 10-3 4.60 × 10-3 -3.59 4.68 × 10-2 -2.99 5.45 × 10-3 -3.66 7.76 × 10-2 -3.10 5.90 × 10-3 -3.56 1.08 × 10-1 -3.12 7.12 × 10-3 -3.72 1.83 × 10-1 -3.01 -1 ∆sH° ) -3.64 ( 0.07 kJ mol ∆sH° ) -3.06 ( 0.06 kJ mol-1 BA12C4 2.49 × 10-3 1.71 2.18 × 10-2 -3.59 2.69 × 10-3 1.80 3.26 × 10-2 -3.38 3.30 × 10-3 1.71 3.63 × 10-2 -3.33 3.44 × 10-3 1.78 3.98 × 10-2 -3.01 4.43 × 10-3 1.65 6.24 × 10-2 -3.32 ∆sH° ) 1.73 ( 0.06 kJ mol-1 ∆sH° ) -3.33 ( 0.21 kJ mol-1
TABLE 4: Standard Enthalpies of Transfer of 12-crown-4, 1-Aza-12-crown-4 and 1-Benzyl-1-aza-12-crown-4 from Propylene Carbonate to Acetonitrile at 298.15 K crown ether
∆tH° (PCfMeCN) (kJ mol-1)
12C4 1A12C4 BA12C4
-0.58 -0.07a 5.05
aReference
4.
ethers and their lithium complexes in AN and PC, and this is discussed in the following section. (ii) Standard Enthalpies of Solution and Transfer. Limitations encountered in the derivation of standard Gibbs energies of solution, ∆sG°, of 12C4 and BA12C4 and their lithium complexes in AN and PC are mainly concerned with the high solubilities of these compounds in these solvents. Since these compounds appear to form solvates in the solid phase, the results cannot be used in Gibbs energy calculations. In the determination of standard enthalpies of solution, ∆sH°, these limitations are not applicable. Thus, enthalpies of solution of 12C4 and BA12C4 in AN and in PC at 298.15 K at various concentrations (molar scale) are reported in Table 3. Since there is hardly any variation in the ∆sH values of the appropriate crown with changes in concentration, the standard enthalpies of solution, ∆sH°, of 12C4 and BA12C4 in AN and in PC at 298.15 K are the average of the data given for each crown in each solvent. Errors in the ∆sH° values are given as twice the standard deviation of the mean. Table 3 shows that, in most cases, the dissolution process takes place with a release of energy (exothermic process) except for BA12C4 in AN where ∆sH° is positive. The crystal lattice contribution to the solution enthalpy can be removed by considering the standard enthalpy of transfer, ∆tH°, of these macrocycles from a reference solvent (PC) to AN.
∆tH°(PCfMeCN) ) ∆sH°(MeCN) - ∆sH°(PC)
(4)
Equation 4 reflects the difference in solvation of the solute in these two solvents. Data for 12C4 and BA12C4 are reported in Table 4. For comparison purposes, the ∆tH° value for 1A12C4 previously reported4 is also included in this table. The ∆tH° values reveal that as far as 12C4 and 1A12C4 are concerned not much difference is found in the solvation of these ligands in these solvents. However, BA12C4 is better solvated in PC than in AN. This information is now considered to assess the effect of the ligand solvation in its complexation with metal cations in AN and PC. If the desolvation of the ligand was
likely to make a predominant contribution to the ∆cH° values (Table 2) for each particular ligand and Li+ in these two solvents, the enthalpic stability of 12C4 (and 1A12C4) with this cation would be approximately the same in both solvents. On the other hand, for BA12C4, the process of complexation would be expected to be enthalpically more stable in AN than in PC. Although this pattern is observed for the latter ligand in these solvents, the consistent decreases in the ∆cH° values listed in Table 2 (from 3 to 5 kJ mol-1) in moving from AN to PC led us to assess whether this effect is due to cation desolvation upon complexation or indeed to a greater solvation of the lithium coronand complex in AN than in PC. These two aspects are now discussed. As far as the free cation is concerned, the higher enthalpic stability of lithium in AN relative to PC reflected in the single-ion transfer enthalpy of this cation between these solvents (∆tH° Li+, (PCfMeCN) ) -2.42 kJ mol-1,20 data based on the Ph4As Ph4B convention at 298.15 K20) shows that Li+ is better solvated in AN than in propylene carbonate and, therefore, the contribution of the desolvation enthalpy (endothermic process) would be greater in the former solvent than in the latter. As a result, ∆cH° values would be expected to be enthalpically more stable in PC than in AN. This trend is not observed in the data listed in Table 2. Therefore, it is important to gain information regarding the solvation of lithium coronand complexes in these solvents. To evaluate this effect we proceeded with the solution enthalpies of these electrolytes in AN and in PC. Table 5 lists the enthalpies of solution of lithium hexafluoroarsenate and eight new lithium coronand salts [lithium 12crown-4 hexafluoroarsenate (Li12C4AsF6), lithium 12-crown-4 tetrafluoroborate (Li12C4BF4), lithium 12-crown-4 trifluoromethanesulfonate (Lil2C4CF3SO3), lithium 12-crown-4 perchlorate (Lil2C4ClO4), lithium 1-benzyl-1-aza-12-crown-4 hexafluoroarsenate (LiBA12C4AsF6), lithium 1-benzyl-1-aza12-crown-4 tetrafluoroborate (LiBA12C4BF4), lithium 1-benzyl1-aza-12-crown-4 trifluoromethanesulfonate (LiBA12C4CF3SO3), and lithium 1-benzyl-1-aza-12-crown-4 perchlorate (LiBA12C4ClO4)] in AN and in PC 298.15 K as a function of the electrolyte concentration (molar scale). The importance of reporting the dependence of ∆sH values with concentration cannot be overemphasized particularly in cases where concentration dependence is observed. For electrolytes which do not show systematic variations in ∆sH with changes in their concentration, the standard enthalpies of solution, ∆sH°, are the average value5 of the data given in Table 5. For electrolytes showing systematic variations in ∆sH with the concentration of the electrolyte, the standard enthalpy of solution, ∆sH°, is the intercept at c ) 0 of a plot of ∆sH against the square root of the concentration (c1/2) (for 1:1 electrolytes provided that only one electrolyte is present in solution, the molar concentration equals the ionic strength of the solution). To discuss these data, standard enthalpies of solution of lithium common salts and lithium coronand salts in MeCN and in PC are summarized in Table 6. Also included in this table are data previously reported for lithium 1A12C4 salts containing different anions in these solvents at the same temperature.4 Enthalpies of solution are the result of two contri-butions: (i) the crystal lattice enthalpy (endothermic process) and (ii) the solvation enthalpy (exothermic process). For electrolytes containing the same anion, it is expected that, as the size of the cation increases in moving from lithium to lithium coronand ions, the energy associated with process i decreases (less endothermic). However, the results in Table 6 show that the dissolution of common lithium salts in these solvents is an exothermic process (solvation predominates) while
Thermodynamics of Lithium-Crown Ether Interactions
J. Phys. Chem., Vol. 100, No. 34, 1996 14489
TABLE 5: Enthalpies of Solution of Lithium and Lithium Coronand Electrolytes in Acetonitrile and in Propylene Carbonate at 298.15 K as a Function of the Electrolyte Concentration (molar scale) Acetonitrile c (mol dm-3)
∆sH (kJ mol-1)
c (mol dm-3)
∆sH (kJ mol-1)
LiAsF6 1.23 × 10-3 -18.35 3.73 × 10-3 -18.40 5.82 × 10-3 -18.55 9.43 × 10-3 -18.50
Li12C4AsF6 2.62 × 10-3 6.62 3.23 × 10-3 7.72 5.49 × 10-3 7.05 6.18 × 10-3 7.35 7.93 × 10-3 7.09
∆sH° ) -18.45 ( 0.06 kJ mol-1 a
∆sH° ) 7.17 ( 0.39 kJ mol-1 a
LiBA12C4AsF6 3.14 × 10-3 8.64 6.92 × 10-3 8.88 6.93 × 10-3 8.86 7.13 × 10-3 8.61 8.90 × 10-3 8.75
LiBA12C4BF4 1.59 × 10-3 10.66 1.96 × 10-3 9.29 2.20 × 10-3 9.33 3.14 × 10-3 9.03 5.80 × 10-3 8.80
∆sH° ) 8.75 ( 0.12 kJ mol-1 a
∆sH° ) 11.33 ( 0.58 kJ mol-1 b
c (mol dm-3)
∆sH (kJ mol-1)
c (mol dm-3)
∆sH (kJ mol-1)
∆sH (kJ mol-1)
c (mol dm-3)
Li12C4BF4 1.88 × 10-3 6.08 2.35 × 10-3 6.26 2.50 × 10-3 6.13 2.84 × 10-3 6.03 3.47 × 10-3 6.31 -3 6.18 × 10 7.35 7.93 × 10-3 7.09 ∆sH° ) 4.74 ( 0.23 kJ mol-1 b
Li12C4CF3SO3 1.54 × 10-3 14.51 2.37 × 10-3 15.93 3.51 × 10-3 17.30 4.96 × 10-3 15.33 5.85 × 10-3 15.84 -3 7.51 × 10 15.22
Li12C4ClO4 1.56 × 10-3 1.30 2.57 × 10-3 1.24 3.96 × 10-3 1.14 5.58 × 10-3 1.44 6.64 × 10-3 1.23
∆sH° ) 15.69 ( 0.94 kJ mol-1 a
∆sH° ) 1.27 ( 0.18 kJ mol-1 a
LiBA12C4CF3SO3 1.98 × 10-3 19.74 2.12 × 10-3 18.40 2.45 × 10-3 19.18 2.49 × 10-3 18.10 2.69 × 10-3 19.12 3.74 × 10-3 19.01 -3 5.83 × 10 19.14 ∆sH° ) 18.96 ( 0.54 kJ mol-1 a
LiBA12C4ClO4 1.42 × 10-3 9.34 1.59 × 10-3 9.58 1.72 × 10-3 9.94 2.09 × 10-3 8.94 3.61 × 10-3 8.88 3.63 × 10-3 8.99 -3 4.88 × 10 9.05 ∆sH° ) 10.23 ( 0.35 kJ mol-1 b
Propylene Carbonate c (mol dm-3)
∆sH (kJ mol-1)
Li12C4AsF6 2.15 × 10-3 3.24 × 10-3 4.74 × 10-3 8.02 × 10-3 1.10 × 10-2
11.97 10.99 10.23 8.97 7.92
∆sH° ) 14.92 ( 0.13 kJ mol-1 b LiBA12C4AsF6 1.25 × 10-3 12.47 1.33 × 10-3 12.53 1.53 × 10-3 12.95 2.18 × 10-3 12.10 3.14 × 10-3 11.15 4.14 × 10-3 12.00 ∆sH° ) 13.91 ( 0.49 kJ mol-1 b a
c (mol dm-3)
∆sH (kJ mol-1)
c (mol dm-3)
∆sH (kJ mol-1)
Li12C4BF4 2.01 × 10-3 9.79 3.02 × 10-3 9.04 3.19 × 10-3 10.29 3.26 × 10-3 9.87 3.31 × 10-3 9.33 3.32 × 10-3 9.94 -3 4.28 × 10 9.21 5.75 × 10-3 9.65 ∆sH° ) 9.64 ( 0.42 kJ mol-1 a
Li12C4CF3SO3 1.40 × 10-3 22.34 2.85 × 10-3 19.94 2.87 × 10-3 22.72 3.09 × 10-3 19.88 3.23 × 10-3 19.14 5.58 × 10-3 21.54 -3 7.21 × 10 19.65
LiBA12C4BF4 1.36 × 10-3 13.52 1.45 × 10-3 13.44 1.58 × 10-3 13.30 1.67 × 10-3 13.29 2.65 × 10-3 13.65 3.60 × 10-3 14.19 -3 5.07 × 10 14.18 ∆sH° ) 12.35 ( 0.17 kJ mol-1 b
LiBA12C4CF3SO3 1.37 × 10-3 23.41 1.64 × 10-3 22.15 1.87 × 10-3 22.18 2.35 × 10-3 23.00 2.41 × 10-3 21.98 3.27 × 10-3 22.39 -3 5.44 × 10 22.42 ∆sH° ) 22.50 ( 0.51 kJ mol-1 a
∆sH° ) 20.74 ( 1.43 kJ mol-1 a
c (mol dm-3)
∆sH (kJ mol-1)
Li12C4ClO4 3.81 × 10-3 4.56 × 10-3 6.13 × 10-3 7.28 × 10-3 1.12 × 10-2
5.43 5.39 5.53 5.04 4.77
∆sH° ) 6.51 ( 0.18 kJ mol-1 a LiBA12C4ClO4 1.69 × 10-3 8.66 1.95 × 10-3 10.90 2.62 × 10-3 9.82 2.92 × 10-3 9.62 3.79 × 10-3 9.13 4.64 × 10-3 9.55 ∆sH° ) 9.62 ( 0.75 kJ mol-1 a
Average value. b Value at c ) 0 from a plot of ∆sH against c1/2.
the reverse is true for lithium coronand salts. Therefore, the thermochemical characterization of salts containing the same anion leads to the conclusion that in AN and in PC the lithium coronand is less solvated than the lithium salts, which may explain the higher conductances observed for lithium coronand relative to common lithium electrolytes, in these solvents. Availability of standard enthalpies of solution of electrolytes in two solvents allows the calculation of the transfer enthalpies of these electrolytes, ∆tH°, from PC (reference solvent) to AN. These data are also reported in Table 6. The results show that, in most cases, electrolytes are enthalpically more stable in AN than in PC. To assess the anion and cation contributions to the transfer enthalpy of these electrolytes, single-ion values are calculated using previously reported5 ∆tH° data from PC f MeCN for AsF6- (0.64 kJ mol-1), BF4- (0.97 kJ mol-1), CF3SO3- (-0.62 kJ mol-1), and ClO4- (3.35 kJ mol-1)20 based on the Ph4AsPh4B convention21 at 298.15 K. Values of -6.05 and -3.36 kJ mol-1 (averages of data given in Table 6) were obtained for the enthalpy of transfer of Li+12C4 and Li+-
BA12C4, respectively. These data do not differ significantly from the ∆tH° value of -3.02 kJ mol-1 from PC f MeCN previously reported5 for 1A12C4. This information clearly reflects that these cations are enthalpically more stable in MeCN than in PC. In fact the ∆tH° values given above are well within the stability increases observed in the enthalpies of complexation of these crown ethers and lithium in MeCN relative to PC listed in Table 2. These findings reveal that in assessing the medium effect on the complexation of metal cations and macrocycles not only cation and ligand desolvations should be considered but equally important is the solvation of the metal-ion complex since this also contributes to the overall process of binding in solution. Another interesting aspect of the higher enthalpic stability of lithium coronand cations in MeCN relative to PC is that related to the electrolyte conductances of electrolytes containing these cations. Thus, the enhancement in conductance of lithium coronand electrolytes relative to common lithium salts is likely to be greater in PC (less solvated) than in MeCN. Some
14490 J. Phys. Chem., Vol. 100, No. 34, 1996
Danil de Namor et al.
TABLE 6: Standard Enthalpies of Solution of Lithium and Lithium Coronand Salts in Acetonitrile and in Propylene Carbonate at 298.15 K
electrolyte
∆sH° (kJ mol-1), acetonitrile
∆sH° (kJ mol-1), propylene carbonate
∆tH° (kJ mol-1), PC f MeCN
LiAsF6 LiBF4 LiCF3SO3 LiClO4 [Li12C4]AsF6 [Li12C4]BF4 [Li12C4]CF3SO3 [Li12C4]ClO4 [Li1A12C4]AsF6 [Li1A12C4]BF4 [Li1A12C4]CF3SO3 [LiBA12C4]AsF6 [LiBA12C4]BF4 [LiBA12C4]CF3SO3 [LiBA12C4]ClO4
-18.45 -14.57a -15.59a -43.26b 7.17 4.74 15.69 1.27 3.72a 0.85a 14.66a 8.75 11.33 18.96 10.23
-15.14a -15.55a -12.50a -41.11c 14.92 9.64 20.74 6.51 4.84a 3.73a 18.74a 13.91 12.35 22.50 9.62
-3.31 0.98 -3.09 -2.15 -7.75 -4.90 -5.05 -5.24 -1.12 -2.88a -4.08a -5.16a -1.02a -3.54a 0.61a
a
TABLE 7: Enthalpies of Coordination of Lithium Salts and Crown Ethers (12-crown-4 and 1-Benzyl-1-aza-12-crown-4) at 298.15 K [Li+CE]X[Li+12C4]AsF6[Li+12C4]BF4[Li+12C4]CF3SO3[Li+12C4]ClO4[Li+BA12C4]AsF6[Li+BA12C4]BF4[Li+BA12C4]CF3SO3[Li+BA12C4]ClO4-
∆coordH° (kJ mol-1)a
average ∆coordH° (kJ mol-1)a
-50.97 -50.05 -44.86 -45.18 -56.83 -53.23 -70.08 -67.61 -52.92 -56.76 -51.82 -55.61 -60.47 -62.71 -79.41 -78.44
-50.51 -45.02 -55.03 -68.84 -54.84 -53.71 -61.59 -78.92
a For these calculations, enthalpies of complexation, ∆cH° are the average of the values given in Table 2 for the various anions.
Reference 1. b Reference 24. c Reference 23.
preliminary results obtained by us appear to corroborate this statement.22 (iii) Enthalpies of Coordination. There is now enough information to calculate the enthalpies of coordination, ∆coordH°, associated with the process referred to product and reactants in their pure physical state (eq 1). These data can be calculated from
LiAsF6(sol) + 12C4(l) f [Li12C4]AsF6(sol)
(6)
LiBF4(sol) + 12C4(l) f [Li12C4]BF4(sol)
(7)
Enthalpies of coordination for processes involving BA12C4 are higher (more exothermic) than those involving 12C4. This pattern is also found in the enthalpies of complexation of lithium and these crown ethers in solution (see Table 2). However, the results shown in Table 7 strongly reflect the anion effect on the coordination process. Thus, coordination values follow the sequence ClO4- > CF3SO3- > AsF6- > BF4-. Among the anions, ClO4- and CF3SO3- are known to have some coordinating ability. Thus, X-ray crystallographic studies reveal that oxyanions such as perchlorate can interact with the metal cation through the oxygen atoms in several ways.25,26 Also there is structural information that the trifluoromethanesulfonate anion interacts with metal cations almost exclusively through the oxygen atoms.27 However, fluoride adducts of strong Lewis fluoroacids such as BF4- and AsF6- are characterized by their low basicity and non-coordinating abilities,28 and therefore, their contribution to increase the enthalpic stability of the coordination process (more exothermic) is likely to be relatively small. These statements are corroborated by the ∆coord H° results shown in Table 7, where data involving AsF6- and BF4- anions are enthalpically less stable (more positive) than corresponding data involving ClO4- and CF3SO3- anions.
LiCF3SO3(sol) + 12C4(l) f [Li12C4]CF3SO3(sol)
(8)
Final Remarks
LiClO4(sol) + 12C4(l) f [Li12C4]ClO4(sol)
(9)
∆coordH° ) ∆sH°(CE) + ∆sH°(LiX) + ∆cH° - ∆sH°[(Li+CE)X-] (5) Some of the advantages in deriving data for the coordination process (eq 1) have been previously discussed.4 Standard enthalpies of solution of the appropriate crown ether (Table 3), lithium,4,23 and lithium coronand salts (Table 6) and corresponding data of complexation (Table 2) in each solvent are inserted into eq 5 to calculate the enthalpies of coordination for the following processes
LiAsF6(sol) + BA12C4(l) f [LiBA12C4]AsF6(sol) (10) LiBF4(sol) + BA12C4(l) f [LiBA12C4]BF4(sol)
(11)
LiCF3SO3(sol) + BA12C4(l) f [LiBA12C4]CF3SO3(sol) (12) LiClO4(sol) + BA12C4(l) f [LiBA12C4]ClO4(sol) (13) Table 7 reports the enthalpies of coordination of lithium coronand salts (eqs 6-13) derived from data in AN and in PC. As previously stated,4,24 for a given salt and crown, ∆coordH° should be the same independent of the solvent from which it is derived. This statement is corroborated by the data shown in Table 7, where a good agreement is found between the ∆coordH° derived from two different solvents.
From the above discussion, it is concluded that, (i) Having established that lithium coronands are less solvated than common lithium salts, and as a result the conductances of the former are greater than that for the latter, the conductance enhancement of lithium coronand salts relative to common lithium salts in PC is related to the affinity of the crown ether for lithium. Thus, the most stable lithium-crown ether complex in this solvent shows the highest increase in conductance as corroborated by the pattern found in the limiting ionic conductances which mirrors the sequence observed in complex stability
Li+BA12C4 > Li+1A12C4 > Li+12C4 (ii) Transfer enthalpies for the species participating in the binding process provide important information regarding the medium effect on the complexation of crown ethers and metal ions. (iii) Enthalpies of coordination are good reporters of events taking place when reactants and product are in their pure physical state. This statement is corroborated by the differences
Thermodynamics of Lithium-Crown Ether Interactions observed in the enthalpies associated with the coordination process by altering the anion constituent of the lithium salt. The strength of interaction follows the sequence
ClO4- > CF3SO3- > AsF6- > BF4Acknowledgment. The authors thank the European Research Office of the U.S. Army and ARPA for financial support. References and Notes (1) Pederson, C. J. J. Am. Chem. Soc. 1967, 89, 2495. (2) Gokel, G. W. Crown Ethers and Cryptands; Stoddart, J. F., Ed.; The Royal Society of Chemistry: London, 1991. (3) Cox, B. G.; Schneider, H. In Coordination and Transport Properties of Macrocyclic Compounds in Solution; Elsevier Science Publishers: New York, 1992. (4) Danil de Namor, A. F.; Llosa Tanco, M. A.; Salomon, M.; Ng, J. C. Y. J. Phys. Chem. 1994, 98, 11796. (5) Danil de Namor, A. F.; Llosa Tanco M. A.; Ng J. C. Y.; Salomon, M. Pure Appl. Chem. 1995, 67, 1095. (6) Buschmann, H. -J. J. Soln. Chem. 1987, 16, 181. (7) Anet, F. A. L.; Krane, J.; Dale, J.; Daasvatn, K.; Kristiansen, P. O. Acta Chem. Scand. 1973, 27, 3395. (8) Hopkins, H. P.; Norman, A. B. J. Phys. Chem. 1980, 84, 309. (9) Takaki, U.; Hogen Esch, T. E.; Smid, J. J. Am. Chem. Soc. 1971, 93, 6760. (10) Nakamura, K. J. Am. Chem. Soc. 1980, 102, 7846. (11) Za´vada, J.; Pechanec, V.; Kocia´n, O. Collect. Czech. Chem. Commun. 1983, 48, 2509.
J. Phys. Chem., Vol. 100, No. 34, 1996 14491 (12) Za´vada, J.; Pechanec, V.; Zajı´cek J.; Stibor, I.; Vitel, A. Collect. Czech. Chem. Commun. 1985, 50, 1184. (13) Suurkuusk, J.; Wadso¨, I. Chem. Scripta 1982, 20, 155. (14) Bastos, M.; Ha¨gg, S.; Lo¨nnbro, P.; Wadso¨, I. J. Biochem. Biophys. Methods 1991, 23, 255. (15) Briggner, L.-E.; Wadso¨, I. J. Biochem. Biophys. Methods 1991, 22, 101. (16) Ojelund, G.; Wadso¨, I. Acta Chem. Scand. 1968, 22, 2691. (17) Ebert, K.; Ederer, H. F.; Isenhour, I. L. O. Computer Applications in Chemistry; VCH: Weinheim, Germany, 1989. (18) Christensen, J. J.; Izatt, R. M.; Hansen, L. D. ReV. Sci. Instrum. 1965, 36, 779. (19) Irving, R. J.; Wadso¨, I. Acta Chim. Scand. 1964, 11, 189. (20) Ghousseini, L. F. Ph.D. Thesis, Surrey University, 1984. (21) Cox, B. G.; Hedwig, G. R.; Parker, A. J.; Watts, D. W. Aust. J. Chem. 1974, 27, 477. (22) Danil de Namor, A. F.; Llosa Tanco, M. A. Unpublished results. (23) Criss, C. M. In Physical Chemistry of Organic SolVent Systems; Covington, A. K., Dickinson, T., Ed.; Plenum Press: London, 1973. (24) 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, 16781. (25) Gouda, N. M. N.; Naikar, S. B.; Reddy, G. K. N. AdV. Inorg. Chem. Radiochem. 1984, 28, 255. (26) Lawrance, G. A. Chem. ReV. 1986, 86, 17. (27) Strauss, S. H. Chem. ReV.; 1993, 93, 927. (28) Huheey, J. E.; Keiter, E. A.; Keiter, R. L. Inorganic Chemistry: Principles of Structure and ReactiVity, 4th ed.; Harper Collins College Publishers, 1993.
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