J. Phys. Chem. 1994, 98, 11796- 11802
11796
Thermodynamic, Structural, and Conductance Studies of Lithium Coronand Electrolytes Relevant to Lithium Battery Technology Angela F. D a d de Namor,*'t Margot A. Llosa TancoJ Mark Salomon? and Joe C. Y. Ngt Laboratory of Thermochemistry, Department of Chemistry, University of Surrey, Guildford, Surrey GU2 5XH, U.K.,and US Army ARL, Power Sources Division, Fort Monmouth, New Jersey 07703 Received: June 7, 1994; In Final Form: August 16, I994@
Gibbs energies, enthalpies, and entropies of coronand macrocycles (15-crown-5 and 1-aza- 12-crown-4) and lithium salts containing highly polarizable anions (hexafluoroarsenate, tetrafluoroborate, and trifluoromethanesulfonate) in acetonitrile and in propylene carbonate at 298.15 K are reported. These titration calorimetry studies are accompanied by I3C and 'HNh4R measurements in acetonitrile (CD3CN). On the basis of the stability of the complexes, six new coronand electrolytes were isolated. The thermochemical behavior of these electrolytes as assessed from the standard enthalpies of solution is compared with corresponding data for common lithium salts. Interpretation of these data reveals that the new coronand electrolytes are much less solvated by these solvents than the latter. The implications of these results on the conductivity enhancement observed by the addition of 15-crown-5 and 1-aza-12-crown-4 to nonaqueous lithium solutions are demonstrated leading to the conclusion that the use of these electrolytes in lithium batteries shows promise. Enthalpies of coordination referred to reactants and products in their pure physical state for these systems are first reported. A notable feature of the data is the higher enthalpic stability observed for the coordination of l-aza-12crown-4 and lithium trifluoromethanesulfonate with respect to that of the same ligand and other lithium salts. The observed enhancement of stability is within the range expected for hydrogen bond formation likely to be attributed to a specific interaction between the hydrogen atom of the aza crown ligand and the anion. Further investigations in this area are suggested with particular emphasis on fundamental research where more efforts should be geared to overcome some of the problems encountered in lithium battery technology.
Introduction The solution chemistry of lithium electrolytes has received considerable attention in recent years as reflected by the large number of research publications and review articles in this area,' most of which emphasize its wide range of biologicalZ and i n d u ~ t r i a l ~applications. ,~ For the latter, there are major intemational efforts competing in the development of rechargeable lithium batterie~,~.~ particularly those based on the "rocking chair" concept7-10 which offer an environmentally friendly and safsr alternative to metallic lithium and the more conventional lead acid and nickel cadmium systems. Whether metallic lithium or the appealing rocking chair system is considered, a major problem in the development of these systems is that related to the electrolyte solution." For metallic lithium battery systems, the required use of an inert (nonprotic) solvent generally of medium or very low permittivity results in a number of problems relating to solubilities and low conductivities due to the formation of nonconducting species such as ion pairs and, in some cases, dimers. Rocking chair systems are essentially concentration cells, in which lithium is the only ion transported and where the metallic lithium is replaced by lithiated carbon or the intercalate LiTiS2.1° While these systems suffer from the same problems encountered with metallic lithium battery systems, they offer an additional advantage in that more aggressive aprotic solvents such as acetonitrile can be used. Recent approaches to solution of the electrolyte problems discussed above involved the use of macrocyclic ligand^;'^-'^ as a result of the developments of the last 27 years a large variety University of Surrey.
* us Army Alu.
+
@
Abstract published in Advance ACS Abstracts, October 1, 1994.
of nonelectrolytes (neutral ligand^)'^-'^ are available. The complexation of metal cations with neutral macrocyclic ligands results in an even wider range of new electrolytes, but the thermodynamic properties of these new electrolytes have not been carefully in~estigated,'~ nor has their use in lithium battery systems been fully explored. Although we are currently working with various macro cycle^,^^,^^ the ligands selected for this paper are 15-crown-5 (15C5) and 1-aza-12-crown4 (l-A-l2C4), and the electrolytes considered are those containing cations of relatively large size (lithium coronands) and highly polarizable anions such as hexafluoroarsenate, As&-; tetrafluoroborate, BF4-; and trifluoromethanesulfonate, CF3S03-. However, the isolation of these electrolytes requires thermodynamic data for the complexation process (A,G", A S , and AJ") involving crown ethers and lithium ions in the appropriate solvents. Solution data for the ligand and the complexed and uncomplexed salts in these solvents are of crucial importance to gain knowledge about the solvation properties of these species in solution.20 Hence, this paper focuses on the energetics of (a) the complexation process involving 15-crown-5 and 1-aza- 12crown-4 and different lithium salts [Li+X-] in two solvents, namely, acetonitrile,MeCN, and propylene carbonate, PC.Thus, titration calorimetry was used to determine the stability constant (log Ks)of the binding process as well as the enthalpy associated with this reaction. Calorimetric studies are complemented by 13C and 'H NMR measurements in CD3CN. (b) the solution process, A@, for the macrocyclic ligands (15C5 and 1-A12C4), the lithium, and the new lithium coronand electrolytes [Li+CE]X- in these solvents. (c) the coordination process, A,,,JP', of these salts and the appropriate crown ether referred
0022-365419412098-11796$04.50/0 0 1994 American Chemical Society
Electrolytes Relevant to Lithium Battery Technology to the process where reactants and products are in their pure physical state. Finally, representative examples are given to demonstrate the effect of crown ethers on the electrolyte conductivity and their possible applications in lithium battery technology.
Experimental Section (a) Chemicals. LiCF3S03 (Aldrich, 98%) recrystallized from an acetone/toluene (1:4) mixture was dried at 60 "C under low pressure for several days. LiAsF6 and LiBF4 (both from Fluka) were dried at 50-60 "C. 15-crown-5 and 1-aza-12-crown-4 (Fluka) were purified as suggested in the literature.*' Acetonitrile (Fisons, HPLC grade; 99.9%) was first refluxed with CaH2 for several hours and then distilled. Only the middle fraction was collected, and its water content was determined by Karl Fischer titration and found to be less than 0.02%. The solvent conductivity measured by a Wayne-Kerr Autobalance Universal Bridge, type B 642 was 1 x S-cm-' at 298.15 K. Propylene carbonate (Aldrich, 99%) was stored overnight over dried 4A molecular sieves, and then it was distilled under reduced pressure. The water content of the solvent (0.015%) and its conductivity (1 -4 x 3cm-I) were determined as described above. (b) Titration Calorimetry. Calorimetric titrations were carried out in the Tronac 450 calorimeter, which is the commercial version of the solution calorimeter originally designed by Christensen et al.22 The reliability of the equipment was tested by using the reaction of protonation of THAM [tris(hydroxymethy1)aminomethanel with hydrochloric acid suggested by Wilson and Smith.23 The value obtained (A,W = -47.49 f 0.62 kJ*mol-') was in excellent agreement with that reported in the literature (-47.49 kJm~l-').*~ For enthalpy of complexation measurements, a solution of the crown ether (33.5 x m ~ l - d m - ~in) the appropriate solvent was placed in the burette. Then, 50 mL of the appropriate electrolyte (8 x 10-4-1 x m ~ l - d m - ~was ) pipetted into the reaction vessel and placed in the calorimetric tray. The whole system was immersed in a thermostatic water bath at 298.15 f 0.01 K and allowed to reach thermal equilibrium. Then, the solution of the crown ether was added for a given time at a fixed burette delivery rate (BDR). A chart recorder was used to monitor the reaction. After every titrant addition, an electrical calibration was carried out. Corrections for enthalpy of dilution of the titrant in the solvent were carried out in all cases. (c) 'H and 13C NMR Studies. ' H NMR Measurements. Measurements were carried out in a Brucker A-C-300 NMR spectrophotometer at a spectral frequency (SF)of 300.135 MHz, delay time of 1.60 s, and acquisition time (AQ) of 1.819 s, and a line broadening of 0.55 was applied. Solutions of crown ethers in CD3CN (1 -5 x m~l-dm-~ were ) placed in 5 mm NMR tubes using TMS as the internal reference to measure the spectrum of the ligand. Then, additions of lithium salt solutions in CD3CN (1-2 m ~ l - d m - ~were ) made to obtain the 'H NMR spectrum of the lithium complex. AU NMR measurements were carried out at 298 K. 13CNMR Measurement. For these measurements, SF,75.469; SW, 301.15 MHz; pulse width, 350, delay time, 0.279 s; acquisition time, 0.721 s; and line broadening, 0.55 were used. Crown ethers solution in CD3CN (1-2 x mol~dm-~) placed in 10 mL NMR tubes were used to measure the spectra of the ligand with T M S as internal reference. For the complex, additions of lithium salt solutions in CD3CN (2-3 x m ~ l - d m - ~were ) made at 298 K. (d) Isolation of New Metal Ion-Coronand Salts. The solid complexes of cyclic polyethers were prepared by dissolving
J. Phys. Chem., Vol. 98, No. 45, 1994 11797
stoichiometric amounts of the crown ether and the appropriate 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. (e) Enthalpy of Solution Measurements. These measurements were carried out at 298.15 & 0.01 K using the Tronac 450 calorimeter. The reliability of the equipment was tested by using the standard reaction suggested by Irving and W a d ~ o . ~ ~ For enthalpy of solution measurements, sealed ampules containing accurate amounts of the appropriate compound were loaded in the clamp. The volume of the solvent was 50.0 mL. The contents of the reaction vessel were kept in a constant temperature environment. After equilibration, the ampule was broken in the appropriate solvent and the thermogram was recorded. In all cases, electrical calibrations were carried out. Corrections were applied to account for the enthalpy of breaking of empty ampules in the solvent. For the determination of the enthalpy of solution of the coronate electrolytes, measurements were also carried out in the presence of different amounts of the ligand in order to ensure that no dissociation of the complex occurred during these measurements. (f) Conductance Measurements. Conductance measurements were carried out with a Wayne-Kerr Autobalance Universal Bridge, type B642. The cell constant was determined using aqueous potassium chloride solutions (0.1 m~ladm-~). A value of 0.1203 cm-' was calculated. A fresh solution (5 x m ~ l - d m - ~of) both the salt and the ligand (crown ethers) was made up for each experimental run. The conductance cell was cleaned with deionized water, dried, and weighed first. It was then filled with the solvent (- 100 mL) and left in the thermostated bath to reach thermal equilibrium while dry nitrogen was passed through the solvent. After thermal equilibrium was reached, the conductivity of the solvent was measured at 25 "C. The salt solution was added to the cell from a hypodermic syringe. The total amount of solution added was approximately 20 mL. Additions of crown solution were carried out using the same procedure.
Results and Discussion Thermodynamic Parameters of Complexation. Stability constants and derived Gibbs energies, enthalpies, and entropies for the complexation of lithium salts and crown ethers (15crown-5 and 1-aza-12-crown-4) at 298.15 K are reported in Table 1. The individual errors in K and A,H" expressed as twice the standard deviation (s) of the mean were calculated using s = [C(xi - x)2/(n - 1)]1'2, where n is the number of data (at least eight) for each calorimetric run. The results show that, within the experimental error, reasonable agreement is found among the thermodynamic parameters for the different lithium salts and these ligands in these solvents. For the process involving 15-crown-5 and lithium in these solvents some data are available in the literature. Thus, the thermodynamic parameters of complexation reported by Buschman26 for Li+ and 15C5 in propylene carbonate at 298.15 K (log Ks = 4.03; A,H" = -20.8 kJmol-'; A,S" = 7.0 J.K-'mol-') obtained calorimetrically are in good agreement with the data obtained by us for this system (see Table 1). As far as enthalpies are concerned, both values differ by about 4 kJmol-' from that reported by Smetana and PopovZ7(A,H" = -16.7 kJmol-') in propylene carbonate obtained by the same method. In the same paper, these authors reported the enthalpy of complexation for lithium and 15-crown-5 in acetonitrile (A,W = -21.0 kJmol-'), which again is not in close agreement with the data obtained by us. Smetana and Popov used LE104 for these measurements.
11798 J. Phys. Chem., Vol. 98, No. 45, 1994
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TABLE 1: Stability Constants and Derived Gibbs Energies, Enthalpies, and Entropies for the Complexation of Crown Ethers (15-Crown-5and 1-Aza-12-crown-4)and Lithium Salts in Acetonitrile and in Propylene Carbonate at 298.15 K electrolyte macrocycle log Ks AcG"(klmol-I) A,H" (kJmol-l) AcSo (JK'mol-') Acetonitrile 15C5 4.44 f 0.24 -25.34 f 1.36 Li+BF4-25.34 f 0.66 0.0 15C5 4.22 f 0.10 -24.09 f 0.56 -24.30 f 0.32 -0.7 Li+As&15C5 4.40 f 0.14 -25.12 f 0.80 Li'CF3SOs-24.15 f 1.74 3.0 Li+BF41-A-12C4 4.24 f 0.12 -24.20 f 0.44 -19.91 f 0.48 14.4 Li+AsFsl-A-12C4 4.23 f 0.34 -24.15 f 0.84 -18.84 f 1.70 17.8 LPCF3S031-A-12C4 4.23 f 0.86 -24.15 f 0.58 -18.69 f 1.06 18.3 Propylene Carbonate Li+BF415C5 4.21 f 0.84 -24.03 f 4.80 -20.44 f 0.78 12.0 Li+AsF615C5 3.97 f 0.46 -22.66 f 2.62 -21.27 f 0.28 4.7 Li+CF3S0315C5 4.23 f 0.28 -24.15 f 1.60 -23.11 f 0.62 3.5 Li+BF41-A-12C4 3.69 f 0.44 -21.06 f 2.50 -14.63 f 0.64 21.6 Li+AsF61-A-12C4 3.67 i 1.06 -20.95 f 6.04 -14.78 f 1.64 20.1 Li+CF3S031-A-12C4 3.87 f 0.84 -22.09 f 0.36 -15.08 f 0.66 23.5 As far as 1-aza-12-crown-4 is concerned, no data have been reported in the literature for these systems. Analysis of thermodynamic data in terms of the ligand, the cation, and the solvent shows that, as far as Gibbs energies are concerned, the results suggest that in propylene carbonate the interaction of 15-crown-5 and Li+ appears to be greater than that observed for the same cation and 1-aza-12-crown-4. However, the most interesting features of these data are shown in the enthalpy and entropy associated with the binding process. It is reasonable to assume that lithium undergoes greater desolvation when interacting with 1-aza-12-crown-4 than with 15-crown-5 since the circular array of electronegative atoms in the former is smaller than that in the latter. In addition, due to the size reduction of the hole and the presence of an NH group in 1-aza-12-crown-4, this ligand is more likely to undergo conformational changes than 15-crown-5. The contribution of any of these two processes would lead to a decrease in the enthalpic stability (AcHomore positive) and, consequently, an entropy increase (AcSomore positive) for 1-aza-12-crown-4 and the lithium cation relative to 15-crown-5 and the same ion. In fact, this is the pattern shown in the data listed in Table 1. Concerning the effect of cation-solvent interaction on the complexation process, it is important to consider the thermodynamic parameters for the transfer of lithium between these solvents, since these data provide information regarding the differences in solvation of this cation in these media. Thus, AtG0(Li+)pC-MeCN = 7.03 kJmol-' z8 (data based on the P b AsPbB c o n v e n t i ~ n )shows ~ ~ that lithium is better solvated in propylene carbonate than in acetonitrile (positive A,Go) and, therefore, this cation would be less readily available for interaction in the former than in the latter solvent. This is corroborated by the Gibbs energy data shown in Table 1, which clearly reflect that the stability of these complexes is slightly greater in acetonitrile (A,Go more negative) than in propylene carbonate. This is generally the case for the other alkali-metal cations and 15-crown-5 in these solvents as assessed from singleion-transfer Gibbs energiesz9 and complexation data (log Ks) reported in the literat~re.~O-~*
'H and 13C NMR Studies As mentioned earlier, the thermodynamics for the interaction between lithium salts and coronands (15-crown-5 and l-aza12-crown-4) in acetonitrile and in propylene carbonate reflect that these are referred exclusively to the complexation process since, within the experimental error, these data are independent of the anion. In order to verify it further 'H and 13C NMR studies for these systems in CD3CN at 298 K were carried out. It should be noted that' Li NMR studies of 15C5 and lithium at 298 K have been previously reported by Smetana and PopovZ9
TABLE 2: Ab Values for the Protons of 15-Crown-5 in the Presence of Lithium Salts in Acetonitrile at 298.15 Ku salt LiBF4 LiCF3SO3 LiAsFs
(PPm) 0.09 0.09 0.08
The chemical shift for the protons in 15-crown-5 in CD3CN is 3.57 PPm.
TABLE 3: Ab Values for the Carbons of 15-Crown-5 in the Presence of Lithium Salts in Acetonitrile at 298.15 Ka salt LBF4 LiCF3SO3 LiAsF6
Ad ( P P ~ ) -2.09 -2.08 -2.08
The proton decoupled I3C of 15-crown-5 in CD3CN at 298.15 K shows one chemical shift ( 6 ) at 71.20 ppm. using perchlorate as the anion component of the electrolyte. To our knowledge, NMR studies on lithium and 1-aza-12-crown-4 in acetonitrile have not been reported. In Table 2, the changes in the chemical shifts of the protons in 15C5 upon complex formation with lithium in CD3CN are shown. The deshielding effects observed may be attributed to the interaction of lithium with the oxygen atoms of the crown. As far as the proton decoupled I3C spectrum of 15-crown-5 in CD3CN at 298 K is concerned, the data reported in Table 3 (see footnote) for 15-crown-5 shows only one signal at 71.20 ppm. Although a deshielding effect would be expected for the carbon atoms adjacent to the oxygens when these interact with lithium, upfield shifts are found as a result of the predominant field effect of the cation. Both 13C and 'H NMR data clearly indicate that the changes observed in the chemical shifts upon complexation with lithium are independent of the anion component of the salt as suggested from thermodynamic data shown in Table 1. The 'H and 13C resonance values corresponding to l-aza12-crown-4 alone and in the presence of LiAsF6, LiBF4, and LiCSSO3 in CD3CN at 298 K are listed in Tables 4 and 5. For 1-aza-12-crown-4 (see footnote Table 5 ) , the chemical shifts were assigned on the basis of the I3C NMR spectrum of 12c r 0 w n - 4 ~and ~ nitrogen insertion increment.34 As far as the proton NMR studies are concerned, the larger proton shifts arising from complexation of the aza crown with the cation are those corresponding to the methylene hydrogens adjacent to the nitrogen. This could be taken as evidence in favor of a more marked interaction of the latter with the cation. However, the remaining protons also show a deshielding effect on complexation with lithium which indicates that to a lesser
Electrolytes Relevant to Lithium Battery Technology
J. Phys. Chem., Vol. 98,No. 45, I994 11799
TABLE 4: Ad Values for the Protons of 1-Am-12-crown-4 in the Presence of Lithium Salts in Acetonitrile at 298.15 K
4
A6 (ppmY proton
[LifBF4-]
[LifCF3S03-]
[Li'AsFs-]
H- 1 H-2 H-3 and H-4
0.18 0.09 0.12
0.18 0.09 0.12
0.18 0.09 0.11
L2 The chemical shifts (6) in ppm of the protons of I-ma-12-crown-4 in CD3CN are 6(H-1) = 2.63, 6(H-2) = 3.51, and 6(H-3)(H-4) = 3.58.
TABLE 5: Ad Values for the Carbons of 1-Aza-12-crown-4 in the Presence of Lithium Salts in Acetonitrile at 298.15 K
c;3 1
2
(,OS3 4
carbon
[Li+BF4-]
A6 (ppmY [Li+CF3SO3-]
c-1 c-2 c-3 c-4
-2.45 -1.69 -1.97 -3.02
-2.5 1 -1.69 -2.05 -2.97
[Li+AsF6-] -2.53
-1.74 -2.05 -2.93
a The chemical shifts (6) in ppm of the carbons of 1-ma-12-crown-4 in CD3CN at 298.15 K are 6(C-1) = 48.71,6(C-2) = 69.09,6(C-3) = 69.71, and 6(C-4) = 71.08.
TABLE 6: Microanalysis Data of Lithium Coronand Salts calculated salt
%C
%H
kilSCSlBF4 [LilSCS]CF3SO3 [Li 15C5]AsF6
38.25 35.11 28.86
6.42 5.36 4.84
[Li 1-A-12C4]BF4 [Li l-A-12C41CF3S03 [Li 1-A-12C41AsFs
35.72 32.64 25.89
6.37 5.17 4.62
experimental %N
5.21 4.23 3.77
%C
%H
37.97 34.72 28.90
6.55 5.40
35.32 32.65 25.97
6.48 5.29 4.74
%N
5.00 4.99 4.10 3.61
extent the oxygen atoms of the ring also interact with the cation. This seems to be supported by the 13Csignal shifts of the ligand in the presence of the salt. The comparatively larger influence shown by the resonance of C-1 might perhaps be the result of a conformational change required by the ligand to adopt the appropiate geometry involved which seems to affect the oxygen atom opposite to the nitrogen.
Enthalpies of Solution Since one of the main targets of this work is to select new electrolytes for use in lithium battery technology, based on the stability of these complexes, we proceeded with the isolation of new lithium coronand salts. Thus, Table 6 lists microanalysis data for six lithium compounds. Resonable agreement between calculated and experimental values is observed. At this stage it seems appropriate to emphasize that despite the current use of lithium salts containing highly polarizable anions (LiAsF6, LiCF3S03, and LiBF4) in batteries, the thermodynamic characterization of these electrolytes in nonaqueous media has hardly been considered. This is not surprising as far as solution Gibbs energy is concemed, since these salts are often highly solvated in these media and, therefore, this thermodynamic parameter cannot be accurately derived from solubility measurements.
However, such a limitation is not applicable in the determination of solution enthalpies. Therefore, we proceeded with the thermochemical characterization of these salts, the new lithium coronand electrolytes, and the ligands (15-crown-5 and l-aza12-crown-4) in acetonitrile and in propylene carbonate at the standard temperature in order to derive the enthalpies of solution of these compounds in these solvents and these are reported in Table 7. The concentration ranges used for these measurements are also included in this table. Within each range, at least five determinations were carried out for each compound in the appropriate solvent. For electrolytes which showed variations in the A&l with changes in their concentration (indicated in table), the standard enthalpy of solution reported is the value at c = 0 from a plot of A J I against c ~ ' where ~, c denotes the final concentration ( m ~ l d m - ~of) the salt in the calorimetric vessel. Most electrolytes did not show variations in AH with c1I2, and in these cases an average is given as A P . It should be noted that while the enthalpy of solution of 1-aza-12-crown-4 reflects the endothermic character of the dissolution process, the ASH value for 15-crown-5 is slightly exothermic in both solvents. Thus, the differences observed mainly reflect the energy input (endothermic process) associated with melting in the former (solid) relative to the latter (liquid). As far as the electrolytes are concerned, the standard enthalpies of solution result from the contribution of the solvation process (exothermic) and that of the crystal lattice (endothermic). As far as lattice enthalpies for ionic solids containing the same anion are concerned, it is expected that as the size of the cation increases (in going from lithium to lithium coronand), the amount of heat supplied to break up the solid decreases. On this basis and provided that no specific interactions occur between the ligand and the anion in the solid state, the heat associated with the crystal lattice process for the uncomplexed lithium salt would be considerably larger (more endothermic) than that for the lithium coronand salt. However, a distinctive feature is observed in the A P values listed in Table 7 since the dissolution of the uncomplexed electrolyte takes place with a release of energy (exothermic process) while the converse is true for the new electrolytes (endothermic process). These findings are striking and strongly suggest that the lithium coronand electrolytes are much less solvated in these solvents than common lithium salts; if so, this must have important implications on the conductivity of these electrolytes (see below), a relevant aspect to consider in the development of lithium batteries. In order to gain information regarding the differences in the enthalpy associated with solute-solvent interactions in these two solvents, the standard enthalpies of transfer, AtW, from propylene carbonate (reference solvent) to acetonitrile are calculated and these data are also shown in Table 7. In general terms, the crown ethers as well as the electrolytes are slightly more stable (in enthalpic terms) in acetonitrile than in propylene carbonate. Taking the AtW(Lif)28 from PC to MeCN based on the Ph+4sPW convention29and considering that the singleion AtHo values for the anion and cation (mainly those derived from the new coronands) constituents of these electrolytes have not been previously reported, these were calculated. Thus, AtWpC-MeCN) Values (kTmOl-') Of 1.04, -0.42, 1.98, -4.15, and -3.56 were obtained for BF4-, CF3S03-, As&-, Lif15C5, and Lif 1-A- 12C4, respectively. Single-ion-transfer enthalpies for the lithium coronand cations between these two solvents show that these are enthalpically more stable in acetonitrile than in propylene carbonate. Combination of these data leads to the calculated AtHo values for the electrolytes shown in the last
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TABLE 7: Standard Enthalpies of Solution of Crown Ethers (15-Crown-5 and l-Aza-12-crown-4),Lithium and Lithium Coronand Salts in Acetonitrile and in Propylene Carbonate at 298.15 K. Derived Enthalpies of Transfer from Propylene Carbonate to Acetonitrile in kJnol-1 acetonitrile propylene carbonate A,W(PC-AN) compound ASH0 concn range ( m ~ l d m - ~ ) ASH" concn range ( m ~ l d m - ~ ) obsd calcdb 15-crown-5 1-aza-12-crown-4 LiBF4 LiCF3SO3 LiAsF6 [Lil5C5]BF4 [Li15C5]CF3S03 [Li 15CS]AsF6 [LilA12C4]BF4 [Lil A 12C41CF3S03 [Lil A1 2C4]AsF6 a
-3.67 f 0.08 20.50 f 1.46 -14.57 f 1.64 -15.59 f 0.90 4.18 f 1.26 9.17 f 2.50 10.62 f 0.84 0.85 f 0.26 14.66 f 1.26" 3.72 f 1.14"
(3.01-6.24) x (0.96-2.04) x (2.05-15.4) x (2.96-13.80) x (1.87-6.83) (1.60-6.58) (1.48-5.93) (1.78-4.93) (1.34-4.39) (1.43-5.35)
x x x
x x x
-3.13 f 0.32 20.57 f 1.76 -15.55 f 1.78 -12.50 f 1.56 -15.14 f 0.44 1.29 f 1.62 13.34 f 1.30 13.19 f 0.94 3.73 f 0.24 18.74 f 1.52 4.84 f 0.36
(1.71-15.10) x (0.80-2.95) x (4.55-27.30) x (2.08-18.70) x (2.31-6.60) x (1.23-5.69) x (0.83-2.49) x (1.44-5.18) x (2.62-4.87) x (1.58-3.20) x (1.34-5.37) x
-0.54 -0.07 0.98 -3.09 -3.11 -4.17 -2.57 -2.88 -4.08 -1.12
-1.39 -2.85 -0.45 -3.11 -4.57 -2.17 -2.52 -3.98 - 1.58
Value at c = 0 frcim a plot of ASHvs c"*. From single ion-values (see text).
TABLE 8: Enthalpies of Coordination of Lithium Salts and Crown Ethers at 298.15 K in kJm01-~ [Lil5C5]BF4 [LilSCS]CF3S03 [Li 15C51AsFs [Li 1-A-12C4IBF4 [Lil-A-l2C4]CF3S03 [Li 1 -A-12C41AsF6
-47.76" (eq 2) -46.41b (eq 2) -52.87" (eq 3) -52.Ogb (eq 3) -52.73b (eq 4) -14.83" (eq 5) -13.34b (eq 5) -28.44" (eq 6) -25.45b (eq 6) -14.19b (eq 7)
av value = -47.08
av value = -52.47 av value = -14.08 av value = -26.94
"Derived from ASH" (Table 3) and AcW (Table 1) values in acetonitrile. Derived from A S P(Table 3) and A,H" (Table 1) values in propylene carbonate. column of Table 7, which are in close agreement with the observed AtHo data.
reported in these media,37 these can only be regarded as "apparent" values since in such a medium extensive ion-pair formation of the cation and cation-macrocycle containing electrolytes with the anion occurs and, therefore, these data are referred to an overall process where not only complexation but also ion-pair formation is involved. This is particularly relevant in the field of lithium batteries where low-permittivity media38 are often used and where the electrolytic conductance is a relevant factor to consider. In this context, quantitative knowledge about the factors which contribute to the behavior of macrocyclic electrolytes in such a medium are becoming increasingly important. Thus, the enthalpies of coordination of lithium coronand salts for the following processes LiBF,(sol) LiCF,SO,(sol)
-
+ 15C5(1) [Li5C5]BF4(sol) (2) + 15C5(1) - [Lil5C5]CF3SO3(sol) (3)
Enthalpies of Coordination On the basis of the statements made in the preceding section it is imperative to gain information regarding the enthalpy associated with the process referred to reactants and products in their pure physical states (solid or l i q ~ i d ) . , ~Thus, , ~ ~ the enthalpies of coordination, A,,,,,,dHO are calculated from enthalpies of solution, ASHo,for the uncomplexed (Li+X-) and complexed (Li+LX-) salts, for the ligands (L = 15C5 and 1-A12C4) in the appropriate solvent (s) (Table 7) using the following thermodynamic cycle (eq 1)
LiBF4(sol)
+ l-A-l2C4(sol)
-
LiCF,SO,(sol) 4-l-A-l2C4(sol)
[Lil-A-12C4]BF4(sol) (5)
-
[Li 1-A- 12C4]CF3SO,(sol) ( 6 ) LiAsF6(sol) 4- l-A-l2C4(sol)
-
[Lil-A-12C4]AsF6(sol)
(7)
The importance of deriving data for the coordination process is described as follows: (i) It provides information regarding the stability of the process in the absence of solvent. (ii) It must be stressed that for a given electrolyte salt and a given ligand (provided that no side reactions occur), the && value should be the same independent of the solvent from which this is derived. Therefore, eq 1 provides a suitable means to check the accuracy of the data. (iii) Provided that the cationmacrocyclic complex is stable enough to allow its isolation, coordination data combined with solution thermodynamic data for the host, the guest, and the resulting complex provide a quantitative basis to derive complexation data in low-dielectric media. It should be emphasized that although data have been
are reported in Table 8. Some interesting features are observed in the enthalpy of coordination data which reflect marked differences for 15crown-5 relative to I-aza-12-crown-4 and lithium salts. Thus, these results illustrate quantitatively that (i) Compounds involving 15-crown-5 (liquid) and lithium salts are enthalpically more stable than those involving 1-aza-12-crown-4 (solid). (ii) For 15-crown-5 and lithium salts, these data seem to be independent of the anion, and an enhanced stability (in enthalpic terms) is observed for the coordination process relative to the same process in solution. On the other hand for 1-aza-12-crown-4, the enthalpic stabilities observed in the solid state are very similar to those in solution (see AcHo values, Table 1) except for lithium trifluoromethanesulfonate (triflate). For this particular lithium compound, the enthalpy of coordination is greater (more negative) than the corresponding data for LiBF4 and LiAsF6. This unusual behavior is observed neither in the
J. Phys. Chem., Vol. 98, No. 45, 1994 11801
Electrolytes Relevant to Lithium Battery Technology
27
,
26
*,
25
2a
t
I
I
26
,
' i
33
.'
--
24 0
25
t
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0
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2
7
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a
c
c
(1 A- 12C 41/ [ELECT ROLYTEJ m LiBF4
0 LiCF3S03
LiBF4
LiCF3S03
Figure 1. Conductimetric titrations of 15-crown-5 and lithium salts in propylene carbonate at 298.15 K.
Figure 2. Conductimetric titrations of 1-ma-12-crown4 and lithium salts in propylene carbonate at 298.15 K.
coordination data for this salt and 15-crown-5 nor in the &&,TO values for cation triflate-macrobicyclic cryptand complexes.39 Since among these ligands I-aza-12-crown4 is the only macrocyclic containing an NH group and the additional stability (-13 kJmol-') observed is well within the range expected for hydrogen bond formation, it seems reasonable to suggest that the higher enthalpic stability of 1-aza-12-crown4 and LiCF3SO3 relative to other lithium salts (LiBF4 and LiAsF6) in the solid state must be attributed to an intermolecular hydrogen bond formation (between the anion and the NH group of the ligand) similar to that reported for a cation free lariat (N-@-tolylsul-
Final Remarks
From the results obtained in this work, it is concluded that additional researches are required and should focus upon (i) macrocyclic electrolytes of which little is known. As demonstrated in this paper these electrolytes show promise in lithium battery technology. (ii) the coordination process which offers a suitable mean to detect specific interactions between reactants in their pure physical state. In a more general context, the selection of macrocyclic electrolytes for use in lithium batteries would be greatly facilitated from fundamental thermodynamic studies on these electrolytes. Within this context, heat capacity fonyl)-4,10,16-triaza-1,7,13-trioxacyclooctadecane-9-7-dione) measurements are of particular relevance and these are now in where X-ray crystallographic studies@revealed hydrogen bondprogress. ing (intramolecular) between the S 0 group of the side arm and the NH group of the macrocyclic ring of this ligand. This Acknowledgment. The authors thank the US Army (Eurofinding illustrates further the importance of deriving coordination pean Office) for financial support. data and provides the f i s t example in macrocyclic chemistry of the suitability of this approach to detect specific interactions References and Notes in the solid state. Studies involving other monocyclic aza (1) Olsher, U.; Izatt, R. M.; Bradshaw, J. S.; Kent Dalley, N. Chem. crowns and oxygen-containing anion components of lithium salts Rev. 1991,91, 137. are now in progress.41 (2) Symposium Proceedings: Power Sourcesfor Biomedical Implant-
-
Conductance Studies Based on the conclusions derived from the thermochemical characterization of lithium and lithium coronand electrolytes regarding their solvation in these solvents, we proceeded with conductimetric titration studies of lithium salts (LiBF4 and LiCF3S03) with 15C5 and 1-A-12C4 in propylene carbonate at 298.15 K. Thus, Figures 1 and 2 clearly show that as the equilibrium position is shifted by the addition of these macrocycles as to favor the formation of the less solvated electrolyte (lithium coronand) a considerable increase in conductance is observed as predicted from solution thermochemical studies of these electrolytes in these solvents. Preliminary calculations in propylene carbonate for the single and complexed electrolytes show that the molar conductances at infinite dilution, A" are greater for the latter relative to the f0rmer.~1
able Applications and Ambient Temperature Lithium Batteries; Owens, B. B., Margalit, N., Eds.; Electrochem. Soc.: Princeton, NJ, 1980. (3) Symposium Proceedings: Electrode Materials and Processes for Energy Conversion and Storage; McIntyre, J. D. E.; Srinivasan, S.; Will, F. G., Eds.; Electrochem. Soc.: Princeton, NJ, 1977. (4) Krieger, J. C. Eng. News 1992,70, 17. (5) Scrosati, B. Electrochim. Acta. 1981,26 (1 1) 1559. (6) Tobishima, S.;Arakawa, M.; Yamaki, J. Electrochim. Acta 1990, 35,383. (7) Armand, M. In Materials for Advanced Batteries; Murphy, D. W., Broodhead, J., Steele, B. C. H., Eds.; Plenum Press: New York, 1980, p 145. (8) Scrosati, B. J. Electrochem. SOC. 1992,139 (lo), 2776. (9) Uribe, F. A.; Sammels, A. F. In Materials and Processes for Lithium Batteries; Abraham, K. M., Owens, B. B., Eds.; Proceedings Series; The Electrochem. Soc.: Pennington, NJ; 1989, p 201. (10) Nagaura, T.; Tazawa, K. Prog. Bafteries Sol. Cells 1990,9, 20. (11) Dudley, J. T.;Wilkinson, D. P.; Thomas,G.; LeVae, R.; Woo, S.; Blom, H.; Horvath, H.; Juzkow, M. W.; Denis, B.; Juric, P.; Aghakian, P.; Dahn, J. R. J. Power Sources. 1991,35,59. (12) Morita, M.; Hayashida, H.; Matsuda, Y .J. Electrochem. SOC.1987, 134, 2107.
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