J. Phys. Chem. 1994,98, 7923-7928
7923
'Li-NMR Determination of Stability Constants as a Function of Temperature for Lithium-Crown Ether Complexes in a Molten Salt Mixture Achim Gerhard,t**Daryl P. Cobranchi,? Ben A. Garland,t Aaron M. Highley,? Yo-Hsin Huang,? Gabor Konya,? Achim Zahl,* Rudi van Eldik,# Sergio Petrucci,s and Edward M. Eyring'J Department of Chemistry, University of Utah, Salt Lake City, Utah 841 12, Institute for Inorganic Chemistry, University of Witten/Herdecke, 58448 Witten, Germany, and Weber Research Institute, Polytechnic University, Farmingdale, New York 1 1 735 Received: January 21, 1994; In Final Form: April 28, 1994"
The stability constants of several crown ethers with lithium ion were determined by 'Li-NMR measurements. A room temperature, basic molten salt of the composition of 55/45 mol 3'% 1-methyl-3-ethyl-imidazolium chloride to aluminum(II1) chloride was used as solvent. On the basis of a 1:l complex formation the following order was found for the stability constants of the investigated crown ethers: 18-crown-6 C 12-crown-4 C benzo15-crown-5 C 15-crown-5. A temperature dependence study for 12-crown-4, benzo- 15-crown-5, and 15-crown-5 was undertaken for the range 5-84 OC. Values of AH and A S were calculated. At 5.5 OC the splitting of the single, fast exchange peak into two separate signals was observed for benzo- 15-crown-5, providing further evidence for the formation of the 1:l complex.
Introduction Since the discovery of crown ethers by Pedersen,lV2 macrocyclic chemistry has developed into a major field that is still expanding. It covers a wide range from phase transfer catalysis in organic chemistry to membrane transport phenomena in biochemistry. Spectral, electrochemical, structural, kinetic, and thermodynamic properties of macrocyclic complex formation have all been inve~tigated.~Izatt et al.4 have given a very detailed review of thermodynamic and kinetic data for macrocyclic complex formation with cations and anions through 1990. Molten salts constitute an unusual, interesting class of aprotic solvents for inorganic solutes and crown ethers. Room temperature chloroaluminate molten salts exhibit high electrical conductivity,moderate viscosity and density, a wide liquid range, and a broad electrochemical window all of which are favorable for use of these electrolytes in batteries.Characteristic properties of these systems are listed by V01kov.~ The disadvantage of molten salts is their extreme sensitivity to moisture. Rhinebarger et aLIO have determined stability constants for lithium ion-crown ether complexes in a 55/45 mol % Nbutylpyridinium chloride ((BP)Cl) to A1C13molten salt. In order to determine whether a change in the organic chloride would alter the complexation of crown ethers with lithium, we used 1-methyl-3-ethylimidazolium chloride (MEIC) and AlC13 as solvent. The MEIC/AIC13 molten salt exhibits a larger liquid range, which includes room temperature, and a larger electrochemical window11 than the (BP)Cl/AlC13 molten salts. Here, we report our findings for complex formation by the crown ethers 18-crown-6 (18C6), 15-crown-5 (1 5C5), benzo15-crown-5 (BlSCS), and 12-crown-4 (12C4) with lithium ion in the binary molten salt composed of 55/45 mol % l-methyl3-ethylimidazolium chloride to AlC13.
Experimental Section Materials. Both 15-crown-5(Aldrich) and 12-crown-4(Sigma) were distilled under reduced pressure (80-90 "C at ca. 0.6 Torr) and dried under vacuum. 18-Crown-6 (Aldrich) was recrystalt University of 1 University of
Utah. Witten/Herdecke. 8 Polytechnic University. 0 Abstract published in Advance ACS Abstracts, July 1, 1994.
0022-3654/94/2098-7923$04.50/0
lized from acetonitrile and dried under vacuum. Benzo-15crown-5 (Parish) was used as received. Anhydrous aluminum(II1)chloride (Fluka) was sublimed under vacuum in the presence of aluminum wire (Alfa) and NaCl (Mallinckrodt). This procedure is similar to that reported by Gale and Osteryoung.'* The LiCl (E M Science) was dried in an oven at 100 OC for 2 days. 1-Methylimidazole (Aldrich) was refluxed over barium(I1) oxide under vacuum for several days and then vacuum distilled. Acetonitrile (E M Science) was distilled repeatedly over P4010 and used immediately. 1-Methyl-3-ethylimidazolium chloride was prepared by mixing three equal volumes of acetonitrile and 1-methylimidazole with ethyl chloride (Matheson), which was condensed into the mixture through an acetoneldry ice condenser at a rate of approximately 1 drop per second. The clear, colorless solution was sealed with a butyl-rubber stopper. This solution was rapidly stirred with a Teflon-coated magnetic stirring bar for 5 days. The stopper was carefully removed, and the excess ethyl chloride was allowed to boil off through a drying tube. The product (MEIC) was crystallized by adding double the solution volume of ethylacetate (E M Science). The white crystals produced were dissolved in acetonitrile and recrystallized by addition of ethylacetate several times. The supernatant liquid was decanted under a dry nitrogen gas flow. The MEIC crystals were then vacuum-dried for 2 days. Preparation of the Melts. The MEIC/AlC13 melts were prepared by mixing (with a Teflon-coated, magnetic stirring bar) weighed (to iO.01 g) amounts of MEIC and AlC13in a Vacuum/ Atmosphere glovebox (Model MO-40-2H Dri-Train) under a dry nitrogen atmosphere. Since the reaction is highly exothermic,l3-15 initially small amounts of each component were added together to avoid thermal decomposition. Once a 3-5 mL volume of melt was prepared, larger amounts of each component were mixed until the desired mole ratio and melt volume were obtained. After stirring, each batch wasvacuum-filtered through a Pyrex Buchner funnel with medium porosity glass frit using the dryboxvacuum system. No attempts were made to further purify the melts, which are estimated to have millimolar amounts of protonic and oxygen impurities.l6 Preparationof Samples. A 1 mol % LiCl55/45 MEIC/AlCl3 basic molten salt batch was prepared in the drybox and then used to prepare subsequent samples containing varying amounts of the ligand. The 15-crown-5 and 12-crown-4 liquids dissolved readily and needed to be stirred for only 5 min. The solid benzo0 1994 American Chemical Society
7924
The Journal of Physical Chemistry, Vol. 98, No. 32, 1994
TABLE 1: Densities for a 55/45 mol % MEIC/AlC13 Molten Salt at 28.5 O C LiCl (mol %) density (g/mL) 1.o
0.0 0.0
1.27 18" 1.2716b 1.2721'
Determined by using the calibrated pycnometer. Determined from eq 2 of ref 7. Determined from eq 3 of ref 7. a
15-crown-5 and 18-crown-6 were less soluble and were stirred for 20-30 min until completely dissolved. Sample constituents were all measured by weight so that mole fractions were known, and the concentrations could be determined from the density of the solution. These samples were then transferred to NMR tubes (Wilmad 528-TR-7, 5 mm 0.d.) that had septum-sealed screw caps to prevent sample contact with air. To determine stability constants, solutions were prepared, in which the ligand concentration was varied while the substrate concentration was held almost constant. The addition of crown ethers to the melt caused a dilution effect up to 12%. The 15crown-5 melt used for variable temperature NMR measurements had only a 3% change in the total lithium ion concentration. In all cases, the exact lithium concentration was used in subsequent calculations. Density Measurements. Density measurements for the molten salts were carried out with a calibrated 5 mL pycnometer (SGA). The volume of the pycnometer was determined at 28.5 OC using triply distilled (conductivity less than 18.0 X 10-6 0'cm-I) degassed water as the standard. The density of the pure water was obtained from a fourth-order polynomial fit of density versus temperature data.17 From triplicate measurements, the 95% confidence interval for the pycnometer volume was calculated to be 5.0641 f 0.0018 mL. The change in volume was determined to be approximately 0.006%for a 4.8 OC decrease in temperature. Table 1 shows that our density measurement for a 55/45 MEIC/ AlC13melt with 1.Omol % LiCl compares quite well with densities determined by Fannin et aL7 for the same melt without LiCl. With this pycnometer, densities of molten salt samples were determined as a function of crown ether weight percent. In each case three data points were fitted with a least squares straight line. The equation for the best fit line describing the change in density was used to convert mole percent values into molarities. NMR Measurements. 7Li-NMR measurements were made on a Varian XL-300, variable temperature, Fourier transform spectrometer operating at a field strength of 7.05 T, a frequency of 116.6MHz, and a spin rate of 20 Hz. The data were recorded at a sweep width of 2054 Hz, an acquisition time of 7.3 s, and an average number of 40 collected transients. The chemical shifts for the lithium ion-crown ether complexes were externally referenced to a 1.47 X 10-1 M LiCl/D20 solution. "Externally" means that both sample and reference were housed in individual 5 mm 0.d. NMR tubes. Experimental parameters were obtained with theNMRspectrometer lockedontothedeuteriumfrequency of the reference sample. The XL-300 NMR spectrometer was very stable over the experimental time period (roughly 15 min) and allowed measurements to be made unlocked with virtually no field drift. The NMR spectrum of the sample was taken, and then spectral acqusition was resumed after the sample had been replaced with the reference. This allowed the reference spectrum to grow into the sample spectrum to yield the needed chemical shift differences. For weakly interacting complexes in dilute solution Chudeck et aI.l8Jg found a good agreement between values obtained from shift measurements using an external reference, as in our case, and those obtained from an internal reference. The use of an external reference in NMR experiments for determining stability constants has the consequence of lowering the accuracy of the shift measurements to fl Hz. The data reported here are comparable to those in the literature.20
Gerhard et al. The variable temperature data were obtained at the reported temperature with an uncertainty of *0.3 OC. The temperature control of the NMR spectrometer was carefully calibrated by using the methanol or ethylene glycol standard samples from Varian for the desired temperatures.
Results and Discussion Wilkes et al.8.21,22 described the following equilibria for chloroaluminate melts as having equilibrium constants much greater than unity MEIC'ClMEIC'AlCl,
+ AICl, MEIC'AlC1; + AlCl, * MEIC'Al,Cl,-
(2)
Both equations may be combined to obtain the overall chloroaluminate equilibrium 2MEIC'AlCl;
* MEIC'Cl- + MEIC'A12C1
0.5 the melt is acidic because it contains the Lewis acid A12C17-. In the neutral melt with AlC13 = 0.5, A1C14- is the only detectable anion. In preliminary experiments we prepared a molten salt batch that was 50/50 mol % ' in MEIC/AlCl3 with 1 mol % LiCI. In this neutral melt, heating and stirring were necessary to dissolve the LiCl. Since 15-crown-5 is liquid, we used it as a suitably easy to handle ligand. In freshly prepared NMR samples the LiCl began to precipitate after 15 min. Only in NMR samples containing more than 1 mol % 15-crown-5 was precipitation diminished. Of course, the NMR tubes containingno 15-crown-5 had the most precipitate. The 7Li-NMR measurements of these samples shbwed a very small lithium signal indicating that very little LiCl was actually dissolved. The stability constant for this series was not determined, since the actual concentrations of dissolved LiCl were not known. However, this preliminary work did indicate that higher LiCl solubility could be achieved in the presence of crown ethers. The ensuing studies reported were done in a basic melt (55/45 mol % MEIC/AlCls) in which the LiCl is much more soluble. Fannin et ala6state that roughly 10 mole % LiCl can be dissolved in basic melts of this type. The 7Li-NMR measurements of the complexes give a single resonance peak of time-averaged position of the complexed sites. In order to calculate the stability constant K ~ Iwe, assumed a 1:l complex formation model for the lithium ion-crown ether complexes. M+L+ML
(4)
Here M stands for the lithium, and L (ligand) denotes the crown ether. The concentration is the independent variable, and ML is the resulting Li+-macrocyclic complex. The 'Li-chemical shift data followed the two-site, fast exchange mode123 of eq 5 ' o b =fM6M +fML6ML
Here, d0b is defined as the experimental differences in chemical shifts obtained by subtracting the reference peak position from that of the sample. With this definition of doh, a downfield shift of the sample results in a positive change in bobs. The terms 6~ and ~ M refer L to the chemical shifts of the uncomplexed and complexed lithium; similarly,fM andfML refer to the fractional occupancies of each site.
7Li-NMR Determination of Stability Constants
The Journal of Physical Chemistry, Vol. 98, No. 32, 1994 7925
TABLE 2: Fitted Stability Constants, Chemical Shifts, and Their Respective Standard Deviations for the Studied Li+-Crown Ether Complexes in a 55/45 mol % MEIC/AICl3 Molten Salt 15C5
B15C5
12C4
18C6
T ("C) 4.7 5.7 22.0 25.1 26.6 39.1 57.5 81.6 84.3 22.0 30.4 39.7 59.0 65.3 78.3 5.5 22.0 28.8 37.6 54.8 78.2 22.0
KII 35449 f 40708 28915 f 20447 11417 f 2838 10943 i 5213 9810 f 2277 7524 f 1558 4027 f 731 2302 f 407 2402 f 201 7648 f 3021 3213 f 799 2363 f 545 1554 i 286 1344 f 219 1018 f 143 1302 f 383 1225 f 241 1251 f 258 1130 f 213 910 f 144 649 f 85 68f4
6~ ( P P ~ 1.15 f 0.07 1.01 f 0.03 1.04 f 0.02 1.05 f 0.03 0.99 f 0.02 1 . 1 1 f 0.01 1.23 f 0.02 1.30 f 0.02 1.26 f 0.02 1.12 f 0.03 1.06 f 0.03 1.08 f 0.03 1.16 f 0.03 1.18 f 0.03 1.23 f 0.02 0.90 f 0.03 1.01 f 0.03 1.00 f 0.02 1.04 f 0.02 1 . 1 1 f 0.02 1.22 f 0.02 1.05 f 0.01
~ M (L P P ~ ) -1.17 f 0.02 -1.07 f 0.03 -1.05 f 0.01 -1.00 f 0.03 -1.09 f 0.01 -0.98 f 0.02 -0.89 f 0.03 -0.88 f 0.03 -0.93 f 0.01 -0.70 f 0.02 -0.76 f 0.02 -0.72 f 0.02 -0.66 f 0.02 -0.64 f 0.02 -0.60 f 0.02 -0.40 f 0.03 -0.26 f 0.02 -0.30 f 0.02 -0.27 f 0.02 -0.22 f 0.02 -0.16 f 0.02 -0.59 f 0.02
The function we used to fit the experimental data to the 1:l complexation model can be derived from eq 4 and 5 as follows
1.0
7 2
0.5
Y
c 4J
{
0.0
0
3 a
-0.5
0
a
--1l.*5O
* 0.0
1.0
2.0
3.0
4.0
5.0
8.0
7.0
8.0
Crown/metal m o l e r a t i o
Figure 1. Plot of the chemical shift &bs versus the molar concentration ratio of crown ether to lithium in a 55/45 mol % MEIC to AlC13 molten salt at 22 OC. The crown ethers 15C5 (+), B15C5 (m), 12C4 (A), and 18C6 ( 0 )were used as ligands. The symbols are the experimental data points with the solid curve representing the nonlinear least squares fit.
TABLE 3 crown ether 12C4 15c5
Crown Ether Cavity Sizesa cavity radius (A) crown ether 0.60 0.85
cavity radius (A)
B15C5 18C6
0.85 1.30
OPedersen, C. J. J . Am. Chem. SOC.1970, 92, 386.
calculated using the Levenberg-Marquardt algorithm, which gave the most reasonable uncertainties within the 6 interval of fML = - f M confidence. Nevertheless, we are aware of the possible errors in the results as pointed out and discussed by Granot.26 CMdenotes the total lithium metal concentration, and brackets The 'Li-NMR results for the four crown ethers used in this denote equilibrium concentrations. Substitution of eq 6a and 6b study are shown in Figure 1. The crown ether 15-crown-5 spans into eq 5 gives a larger region of 6obs values with a sharper break at the concentration ratio [ 15C5]/ [LiCl] ofunitycompared to theother ligands. Thus, 15-crown-5 has a larger complex stability constant (7) and also shows good agreement with the 1:1 stoichiometric model used for the curve fitting. The sharp breaks a t unity for both the Using the mass law, the solution for [ML] is quadratic where the benzo-15-crown-5 and 12-crown-4 are also a good indication that negative root is retained as being physically meaningful: a 1:l complex has been formed. For benzo-15-crown-5 we find in our variable temperature study that the fast exchange regime turns into the slow exchange regime a t temperatures below 12 OC and crown ether to lithium ratios lower than unity. At 5.5 O C the splitted peak has a ratio of almost 1:l when 0.5 mol % benzo-15-crown-5 is added to the sample containing 1 mol % Therefore the chemical shift is given by lithium (Figure 4a). This provides further evidence for the assumption of a 1:l complex formation. 18-Crown-6 yields a very gentle curve without a definite breaking point. This behavior can be interpreted as a very weak interaction between the lithium ion and the oxygen donors of the 18-crown-6 ring. The stability constant increases in theorder 18-crown-6 < 12-crown-4 < benzo15-crown-5 < 15-crown-5 (see Table 2). The data in Table 2 are where 6 is the relative chemical shift of the species and CMand presented with all the digits necessary to avoid the propagation CLare the total concentrations of the lithium ion and the crown of a roundoff error in the weighted fitting that is discussed below. ether, respectively. K11 is the stability constant of the ML species, An enhanced stability constant is usually anticipated when the K1 I , C ~ M , and BML were the parameters fitted by the LevenbergMarquardt algorithm, simplex search, and Powell's m e t h ~ d . ~ ~ J s ratio of the cation radius to the cavity of the crown ether is near unity. The ionic radius of Li+ is 0.14 A, and the cavity radii for The deviation of the measured and fitted relative chemical shifts were minimized by least squares. The different optimization the crown ethers are listed in Table 3.27928 Considering these techniques, of which the first uses derivatives and the two others values, one could expect the Li+ ion to be bound more tightly by direct search, found the same sets of parameters using a variety the 12-crown-4 ligand, which has a cavity diameter close to that of reasonable initial guesses. The temperature dependence was of the Li+ ion, than by 15-crown-5 or benzo-15-crown-5, which first evalulated by the common linear relationship. These results have larger cavity diameters. However, the complex formation were the starting values for the more sophisticated nonlinear fit of 12-crown-4 with lithium was also studied29in methanol, and by the Levenberg-Marquardt algorithm. In this procedure the the authors concluded that even in this very common solvent uncertainties of the individual equilibrium constants were 12-crown-4 is not the ideal ligand for complexing an ion as small as lithium. The large difference in cavity size between the ligand introduced as weights. The reported data in Table 2 were
7926 The Journal of Physical Chemistry, Vol. 98, No. 32, 1994 1.5
5
10.0
1 I
-1 1.0
Gerhard et al.
-
0.5
-
Y
c 4
{
0.0
-
9.0
-
6.5
-
8.0
-
7.5
-
7.0
-
6.5
-
h
-
?4. c (
0 .r(
2
9.5
-0.5
-
c (
4
a I
0.0
1.0
2.0
3.0
4.0
5.0
8.0
7.0
,
I
6.0
Crown/metal mole ratio
2.60
2.90
3.00 3.10 3.20 l / T in 1000/K
3.30
3.40
Figure 2. Variable temperature plot of Sh versus the molar ratio of [15C5]/[Li+] in a 55/45 mol W MEIC to AlC13 molten salt at 4.7 OC (O), 22.0 OC (W), and 84.3 OC (+). The solid line is the fitted nonlinear least squares curve.
Figure 3. Van't Hoff plot of In K11 versus inverse temperature for the stability constants of 15C5 (W), B15C5 (e), and 12C4 ( 0 ) lithium complexes. The solid line was calculated using weighted nonlinear least squares for fitting of the K11 values at room temperature and above.
18-crown-6 and the diameter of the Li+ ion is consistent with our experimental finding that this complex is the weakest of the four studied in the MEIC/AlC13 melts. Our results parallel the findings of Rhinebarger et al.,l03where the same sequence for complex stability was determined in a basic 55/45 (BP)Cl/AlCI3 molten salt. Taulelle and Popov30 noted that the 'Li-chemical shift differed by 2.4 ppm between the basic and the acidic melts. They ascribed this difference to the existence of the lithium ion in two distinctly different chemical environments. Accordingly, they proposed that in the basic melt the lithium existed in a tetrahedral environment such as LiC12AlC142- ion. Furthermore, they suggested that the excess chloride ion present in the basic melt is responsible for the solubility of the LiCl producing the lithium dichloride ion LiC12-. This interpretation is also supported by our results which indicate a low solubility of LiCl in the neutral melt where, according to eqs 1-3, only extremely small equilibrium amounts of the Cl- ion were present. Considering the evidence, it is probable that the lithium ion complexing with the crown ethers is still bound to one or more CI- ions.31 Hence, lithium complexation with crown ethers in these molten salts cannot be simply described by the ratio of the cavity size to cation diameter. The discrepancy between the magnitude of the stability constant K11 for 12-crown-4 and those for 15-crown-5 and benzo-15-crown-5 can be partially explained by reasoning that the larger 15-crown-5 and benzo-15-crown-5 macrocycles provide a better fit for the cation, or possibly the lithium chloride species, than does the smaller 12-crown-4 ligand. The benzo-15-crown-5 complex is weaker than the 15-crown-5 complex because of its electron-withdrawing benzo group which renders the ether oxygens less effective for bonding. In addition, the benzo group introduces a steric hindrance to conformational changes of the ligand for adapting to the guest ion or molecule. Since the 12-crown-4 ligand may not be capable of expanding its cavity diameter sufficiently to accomodate the lithium chlorocomplex, this ligand probably forms a complex in which the lithium in the LiC12- species lies outside the crown ether oxygen plane. The out-of-plane structure can also be caused by steric repulsion between theether oxygens and thechlorine substituents. In either case, the experimental results show that the 12-crown-4 complex is much stronger than the 18-crown-6 complex where the cavity diameter is far too large for effective complex formation. A new batch of the 55/45 MEIC/AlCl3 melt containing approximatley 1 mol % LiCl was prepared for the variable temperature measurements with 15-crown-5. Later on, three further temperatures, namely 4.7,57.5, and 84.3 OC, were checked with the samples used at 26.6 OC. Figure 2 shows the shift in
TABLE 4 Temperature Dependence of the Li+-Crown Ether Complex Formation in a 55/45 mol % MEIC/AIC& Molten Salt ligand AH [kl/moll A S [J/(mol.K)l KII (at 22 "C) 12C4 -1Of 1 25 f 4 1.2 x 103 (-9 i 1 28 i 4)' B15C5 -20.7 i 0.6 -1f2 7.6 x 103 (-22 f 3 -6 f 9 ) 15C5 -22.4 A 0.8 2f2 1.1 x 104 (-22.6 f 0.8 1 3) 'The results in parentheses are based on the data over the full temperature range. Only five values of K11 were used for 12C4 and B15C5, whereas for 15C5 seven values were available. the lithium complexation curves of 15-crown-5 caused by the temperature change from 4.7 to 84.3 OC. Also, for 12-crown-4 and benzo- 15-crown-5 the same samples were measured for all temperatures, except the ones for 22 OC, representing a further batch of each crown ether-lithium complex. The stability constants were determined a t six different temperatures. The data indicate a noticeable dependence of the stability constant on temperature. As the temperature was increased, the strength of the Li+-crown ether complexes decreased, as expected. Table 2 lists the values obtained for the stability constants and the chemical shifts for both the complexed and uncomplexed lithium species a t each temperature studied. The average chemical shift difference between the complexed ( ~ M L )and the uncomplexed ( 8 ~ )species is 2.1 f 0.1, 1.8 f 0.1, and 1.3 f 0.1 ppm for 15crown-5, benzo- 15-crown-5, and 12-crown-4, respectively. These characteristic values of the individual crown ethers are constant with respect to the temperature. However, there is a downfield shift in the 6obvalues as the temperature is increased. This result suggests a change in the lithium environment resulting in a deshielding of the lithium by a faster exchange. Through the use of van't Hoffs equation, the values for AH and AS were calculated by weighted nonlinear regression analysis as described above. The linearity of the plotted data in Figure 3 for 15-crown-5, benzo-15-crown-5, and 12-crown-4 is quite good albeit only K I Ivalues down to room temperatures are used. Plots over the full temperature range deviated slightly from linearity. Despite this fact, both data sets were calculated, and the results are listed in Table 4. Below 22 OC the assumption of the fast exchange regime obviously no longer holds for benzo15-crown-5, thus making the deviation plausible (Figure 4). The entropy term of the 12-crown-4 complex adds considerably to the enthalpy term (Table 4). With the assumption that the complexation of the lithium ion is accompanied by a loss of chloride
’Li-NMR Determination of Stability Constants
The Journal of Physical Chemistry, Vol. 98, No. 32, 1994 1921
TABLE 5 Solvent Properties with Corresponding Enthalpy and Entropy Contributions to Li+-Crown Ether ComDlexeSa 15C5 12c4 18C6 SOIV.~ DNC AHd AS AHd ASe AHd M e CHjN02 CHoCN PCC CHoOH melts
2.7 14.1 15.1 25.7
-43.5 -21.3 -16.7 -11.3 -22.4
-36.8 >6.3 -16.3 >20.9 -10.9 -14.2 -3.3 -16.2 -10.3
27.2 >41.8
-12.6 -0.0 -15.9 11.7
>O.O
44.8 -2.5
6.7
All entries except the last line (melt) were taken from ref 34. All AS values shown here are referenced to a molarity concentration scale. Solvent. DN: Donor number; adapted from ref 32. AH [kJ/mol]. e AS [J/(mol.K)]. fpropylene carbonate. g MEIC/AlCI3 55/45 mol %.
densityvalues determined by the fitted lines for each of the crown ethers. The stability constants were calculated for each case, showing that K11 is essentially independent of the small changes in density.
Conclusions 3
2 1 0 -1 - 2 -3 chemical shift [ppm]
Figure 4. Variable temperature measurements of B15C5 lithium ion complex containing 1 mol % lithium and 0.5 mol % B15C5 as ligand in a 55/45 mol % MEIC to AlCl3 molten salt. The sample was measured at 5.1 OC (A), 11.5 OC (B), 14.5 OC (C), and 17.5 OC (D).
ions, a less ordered solvation state results, leading to an increase in the entropy. In contrast to that for 12-crown-4, the entropy term of the 15-crown-5 and benzo- 15-crown-5 complexes is almost zero, suggesting a complex in which lithium is still associated with one or more chloride ions. Therefore, the entropy does not change because the effective number of species is not changing. The enthalpic effect can be explained by the maximizing of eletrostatic interaction of thecationic charge with the donor atom lone pair electrons, which are delivering the driving force for the complex formation. The enthalpy stabilization of the 12-crown-4 complex is only half as big as those for the 15-crown-5 and benzo15-crown-5 complexes. Both of these latter complexes are strongly enthalpy-stabilized. In general, the “macrocyclic effectP3q4arises from enthalpy changes with the entropy term contributing to the stability in only a few cases, as is found here for the 12-crown-4 complex in the MEIC/AlCl3 molten salt. Rhinebarger20 concluded for the (BP)Cl/AlC13 molten salt that the monomerdimer equilibrium of LiC12- does not appreciably affect the stability constants for 15-crown-5 and benzo- 15-crown-5 lithium ion complexes. However, the stability constants for 12-crown-4 seem to be affected, giving a reasonable explanation for the slight deviation from linearity of the temperature dependence data below room temperature in our results. Gutmann donor number^^^,^^ can be used to describe the solvating ability of a solvent. Generally, as the solvent donor number goes up, the solvation of the cation increases, the complex stability constant decreases, and AHbecomes more positive. Table 5 summarizes some of the results of Smetana and P O P O Vwhich ,~~ show this general trend in A H for increasing solvent donor numbers. More importantly, our A H and AS values are of the same order of magnitude as the thermodynamic quantities for Li+-l5-crown-5 complexes in various solvents. Therefore, the 55/45 mol % MEIC/AlC13 molten salt solvent exhibits an overall solvation effect between that of nitromethane and acetonitrile. The dependence of the stability constants on the density of the molten salt was also investigated. A best fit linear equation describing the change in molten salt density as a function of crown ether weight percent was used to generate the needed molar concentrations. Molarities were also determined by using both the high-density value of 1.2718 g/mL (Table 1) and the lowest
Results obtained on the 50/50 MEIC/AlC13 molten salt mixture support the hypothesis of Taulelle and Popov3Oand Fannin et aL6 that the LiCl solubility is caused by the presence of excess C1-, which exists predominantly in basic melts to form the LiC12species. The partially soluble LiCl salt in the neutral MEIC/ AlC13 melt increases in solubility upon addition of a complexing agent such as 15-crown-5. In the basic MEIC/AlC13 melt the same sequence for lithium ion-crown ether complex stability is found as was determined by Rhinebarger et a1.10*20 in (BP)Cl/ AlCl3 melts. Comparing the stability constants in the two different melts, one can see that the Li+-crown ether complexes in the MEIC/AlCl3 molten salt are stronger. Variable temperature results on the 15-crown-5, benzo- 15-crown-5, and 12-crown-4 complexes suggest that the MEIC/AlC13 salt acts as a slightly poorer solvent for the lithium ion species than the (BP)Cl/AlC13 molten salt does. While it seems possible that the lithium ion in the crown ether complexes is associated with one or more chloride ions when it resides in the macrocyclic complexes, we are unable to say how many chloride ions are still bound to the complexed lithium ion. Referring to Table 4, the positive entropy value for the 12-crown-4 lithium complex adds to the enthalpy effect, whereas for 15-crown-5 and benzo-15-crown-5 the entropy is almost zero, leaving only the enthalpy for stabilizing the Li+crown ether complexes. Finally, the change in molten salt densities upon addition of crown ethers is small, resulting in a negligible change in the calculated apparent stability constants.
Acknowledgment. This work was supported by the Department of Energy, Office of Basic Energy Sciences, and by the Volkswagen Foundation. Informative discussions with Istviin Fiibiiin, Charles L. Hussey, Charles L. Mayne, Alexander I. Popov, and John S. Wilkes are gratefully acknowledged. References and Notes (1) Pedersen, C. J. J . Am. Chem. SOC.1967, 89, 2495. (2) Pedersen, C. J. J . Am. Chem. SOC.1967,89, 7017. (3) Lindoy, L. F. The Chemistry of Macrocyclic Ligand Complexes; Cambridge University Press: New York, 1989. (4) Izatt, R. M.;Pawlak, K.; Bradshaw, J. S.; Bruening, R. L. Chem. Rev. 1991, 91, 1721. ( 5 ) Lipsztajn, M.; Osteryoung, R. A. Inorg. Chem. 1984, 23, 1735. (6) Fannin, A. A., Jr.; King, L. A.; Levisky, J. A.; Wilkes, J. S. J . Phys. Chem. 1984,88, 2609. (7) Fannin, A. A.,Jr.; Floreani, D. A,;King, L. A.; Landers, J. S.;Piersma, B. J.; Stech, D. J.; Vaughn, R. L.; Wilkes, J. S.;Williams, J. L. J . Phys. Chem. 1984,88, 2614. (8) Dieter, K. M.;Dymek, C. J., Jr.; Heimer, N. E.; Rovang, J. W.; Wilkes, J. S. J. Am. Chem. SOC.1988, 110, 2722. (9) Volkov, S.V. Chem. SOC.Rev. 1990, 19, 21. (10) Rhinebarger, R.R.; Rovang, J. W.; Popov, A. I. Znorg. Chem. 1984, 23, 2557. (11) Hussey, C. L. Pure Appl. Chem. 1988, 60, 1763.
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