J . Phys. Chem. 1984, 88, 2609-2614 TABLE I: Self-AssociationEquilibrium Constants (in M-’) of 2-Azacyclotridecanone in CCl,‘
dielectric
IR
t, “C
Kz
K
K,
K
20 40
7.3
47.9 27.0 14.9
7.5 5.1 2.7
48.1 22.1
60
4.0 2.2
a AH, =-25 i: 8 kJ/mol, A H = - 2 4 . 7 (niobdeg), A S = -33 J/(mobdeg).
I1
14.3
kJ/mol, AS, = - 5 1 J/
of the C O N H group do not affect the dipole moment of trans amides. When calculating the dipole moments of multimers we assumed that the polarity of every hydrogen bond was 0.3 D. The molecular quantities, i.e., anisotropy coefficients and dipole moments, necessary for testing our theory may be, within the scope of such a formulated model, calculated without introducing any variable parameters. Eventually, the whole problem consists of finding such values of K2and K (at every temperature) to reproduce simultaneously the experimental results of polarization (eq 17) and nonlinear (eq 18) studies. The best obtained fitting (done by the least-squares method) is presented in Figure 2 (solid lines), and appropriate equilibrium constants are given in Table I. A good theoretical simulation of experimental results obtained for the proposed model indicates its adequacy. However, since the values of equilibrium constants obtained from dielectric data strongly depend on the geometry of multimers, it is worth it to compare these constants with those obtained from IR spectroscopy, which are not affected by geometric assumptions. We determined K2 and K (Table I) from the experimental dependence of molar extinction coefficients for the band of the free N-H group situated at 3454 cm-l (monomers and end groups in multimers) on amide concentration, using the same method as in the case of N M A solutions.’ The values of equilibrium constants obtained from IR spectroscopy agree, within the experimental accuracy, with those determined from dielectric data. This agreement provides strong support for the assumptions that we have made during the interpretation of the dielectric data.
2609
Analysis of the temperature dependence of equilibrium constants shows that the significant differences between dimerization and multimerization constants are connected with the entropy factor. This agrees with theoretical predictions of Sarolea-Mathot,22who showed that the change of entropy accompanying the formation of a dimer from two monomers is greater than in the case of i-mer formation from (i-1)-mer and monomer. The values of change in enthalpy and entropy shown in Table I were calculated by the van’t Hoff method, using the equilibrium constants expressed in mole fraction units.
Conclusions The volume of substituents on the peptide group, which determines the possibility of rotation of amide molecules around hydrogen bonds, is the main factor determining the geometry of chains. In the case of 2-ACT the ring consists of 11 methylene groups, which significantly reduce the possibility of this rotation. Consequently, the multimer structure should be very close to linear. Thus, the solution of 2-azacyclotridecanone in CCl, can be represented as a mixture of spherical solvent molecules and strongly anisotropic species of various lengths. The concentrations of these species depend on the nominal amide concentration, the temperature, and, after application of a strong electric field, also its strength. This simple model describes quite well the properties of a real system placed in an electric field-both in a weak field, when dielectric phenomena are proportional to the field strength, and in a strong field, when the changes in component concentrations take place. The conclusions resulting from dielectric studies have full spectroscopic support. These studies also give a positive test for the theory of Rivail et al. of nonlinear phenomena with its generalization for complex systems with many equilibria. Acknowledgment. This work was supported by the Polish Academy of Sciences within the framework of Project MR-1.9. Registry No. 2-Azacyclotridecanone, 947-04-6. (22) L. Sarolea-Mathot, Trans. Faraday SOC.,49, 8 (1953).
Properties of 1,3-DialkylimidazoIium Chloride-Aluminum Chloride Ionic Liquids. 1. Ion Interactions by Nuclear Magnetic Resonance Spectroscopy Armand A. Fannin, Jr., Lowell A. King, Joseph A. Levisky, and John S. Wilkes* The Frank J. Seiler Research Laboratory, US.Air Force Academy, Colorado Springs, Colorado 80840 (Received: February 18, 1983; In Final Form: November 14, 1983)
Mixtures of I-methyl-3-ethylimidazolium chloride with aluminum chloride are liquid at room temperature. The chemical shifts of protons on the cations of the melts are highly dependent on the proportions of aluminum chloride and organic chloride salt. The largest effect is in the basic melts, where there is an excess of organic chloride over aluminum chloride. The chemical shift behavior may be explained by assuming that the anions in the melt affect the chemical shift of the cation hydrogens and that the anions and cations interact in the fast-exchange regime on the NMR time scale. The model implies that the melts are relatively organized in the liquid state, especially basic compositions. The addition of LiCl or nonelectrolyte solvents to binary aluminum chloride-imidazolium chloride mixtures also affects chemical shifts of the cation protons.
Fused salts that are liquid at room temperature comprise an unconventional but interesting class of aprotic solvents for studying the chemistry of inorganic, organometallic, and organic solutes. In addition, these ionic liquids are potentially useful as electrolytes in batteries, photoelectrochemical cells, and electrodatina. In recent years some mixtures of aluminum chloride i n d I-ilkylpyridinium or 1,3-dialkylimidazolium chlorides3 have
been found to be liquid below room temperature. An understanding of the physical and chemical properties of the ionic liquids is vital to establishing the usefulness of the new materials. Especially important are the transport properties such as electric ering, (2) D, Chum, G.,Eds.; H. L.;Plenum Osteryoung, Press: R.New A. In York, “Ionic 1981; Liquids”; pp 407-23, Inman, D.; Lov-
(3) Wilkes, J. S.; Levisky, J. A,; Wilson, R. A,;Hussey, C . L. Inorg. Chem. (1) Hurley, F. H.; Wier, T. P. J . Electrochem. SOC.1951, 98, 203-6.
1982, 21, 1263-4.
This article not subject to U S . Copyright. Published 1 9 8 4 by the American Chemical Society
2610
The Journal of Physical Chemistry, Vol. 88, No. 12, 1984 N-CHa
e t h y l CHa
I
6
1
I
4
I
I
2
I
I
I
0
SHIFT (pprn)
Figure 1. Proton NMR spectrum of neat 1-methyl-3-ethylimidazolium chloride-aluminum chloride melt at 30 ' C (N = 0.667). Reference is Me2S0.
conductivity and viscosity. In turn, these physical properties are dependent on the detailed structure of the molten salt.4 In order to understand better the melt structure, we have studied the interactions between the anions and cations in the melts by nuclear magnetic resonance (NMR) spectroscopy. Proton N M R is an established technique for studying n-complexes5 and hydrogen-bonded systems! Since the chloroaluminate melts mentioned above have cations with nuclei that are readily observable by NMR, we have atttempted to deduce some aspects of the structure of the low melting ionic liquids from 'H N M R experiments. Similar experiments using 13CNMR have been done and have been reported separately.' The particular ionic liquid chosen for detailed study was the mixture of aluminum chloride chloride (MeEtImCl). That and 1-methyl-3-ethylimidazolium mixture is relatively easy to prepare and is liquid over a wide range or composition^.^ Experimental Section The 1-methyl-3-ethylimidazoliumchloride and chloroaluminate molten salts were prepared by methods described earlier., The 'H N M R spectra were recorded with a Varian T60A N M R spectrometer. The 'H chemical shifts were referenced to an external Me2S0 standard. Positive shifts indicate decreased shielding.
Fannin et al. TABLE I: Proton NMR Chemical Shifts of MeEtImC1-A1C13 chemical shift? ppm H-2 H-4, H-5 NCH, NCH? N CH? 0.33 0.34 0.35 0.38 0.40 0.41 0.42 0.44 0.46 0.48 0.49 0.50 0.52 0.54 0.56 0.58 0.60 0.61 0.62 0.64 0.66 0.67
7.63 7.64 7.57 7.37 7.22 7.18 7.03 6.87 6.61 6.37 6.19 6.05 6.06 6.07 6.07 6.05 6.03 6.03 6.03 5.99 5.97 5.97
(4) Fannin, A. A.; Floreani, D.; King, L. A.; Landers, J. S.; Piersma, B. J.; Stech, D.; Vaughn, R. L.; Wilkes, J. S.; Williams, J. L. J . Phys. Chem., following article in this issue. (5) Carper, W. R.; Buess, C. M.; Hipp, G. R. J. Phys. Chem. 1970, 74, 4229-34. (6) Hanna, M . W.; Ashbaugh, A. L. J . Phys. Chem. 1964, 68, 811-6. (7) Wilkes, J. S.; Reynolds, G. F.; Frye, J. S.Znorg. Chem. 1983, 22, 3870-2. (8) Caesar, F.; Overberger, C. G. J. Org. Chem. 1968, 33, 2971
1.99 1.99 2.00 1.97 1.93 1.96 1.96 1.93 1.92 1.92 1.87 1.90 1.90 1.92 1.94 1.94 1.94 1.96 1.95 1.96 1.96 1.95
1.67 1.67 1.67 1.63 1.62 1.63 1.63 1.62 1.59 1.63 1.55 1.58 1.60 1.60 1.61 1.61 1.62 1.62 1.63 1.63 1.63 1.64
-0.88 -0.88 -0.86 -0.87 -0.88 -0.85 -0.84 -0.85 -0.84 -0.83 -0.85 -0.83 -0.8 1 -0.81 -0.79 -0.78 -0.77 -0.77 -0.75 -0.73 -0.74 -0.75
"Numbers in column headings indicate imidazolium ring positions. NCH2 indicates methylene hydrogens of the ethyl group, CH3 the methyl hydrogens of the ethyl group, and NCH, the N-methyl hydrogens. Reference is Me,SO.
'.O
t
e e
k
6.0
v)
1
Results The 'H N M R spectrum of a MeEtImC1-A1C1, mixture and the assignments are shown in Figure 1. The assignments were made by comparison with similar dialkylimidazolium iodides reported by others.8 At 60 MHz the protons at the 4- and 5-positions could not be resolved, so chemical shifts for those protons are reported as the average of the unresolved multiplet. The spectra are of the neat liquids at slightly above room temperature and are very similar to that of MeEtImCl dissolved in a convention solvent. The signal to noise ratio is excellent, and the resolution of the multiplets is only slightly degraded from other solution spectra, probably due to the higher viscosity of the melts. We express melt compositions in terms of apparent mole fraction of AICl,, N , even though molecular AlCl, probably does not exist in the melts. The use of AICl, mole fraction is convenient because it is easily calculated from the amounts of materials used to prepare a melt sample, and the actual chloroaluminate species present in the melt may be calculated from it. The dependence of proton chemical shifts on AlCI, mole fraction in the melts is
5.66 5.66 5.61 5.49 5.41 5.40 5.35 5.28 5.18 5.12 5.04 5 -04 5.04 5.05 5.06 5.04 5.04 5.05 5.06 5.04 5.04 5.05
'
I 1
0 000
0 0
ooo
A
A
At,&
..1
2.0[ 1.6
000
1
I I 8 . 1
1 1
0 0 000 0
0
- 0.8 -1.2 0.30
AAA
0.40
AICI,
A
A
050
A
A
A
AAA
0.60
A
I
A
0.70
MOLE FRACTION
Figure 2. Proton chemical shifts in neat 1-methyl-3-ethylimidazolium chloride-aluminum chloride melts vs. N at 30 OC: ( 0 )H-2 proton, (A) H-4 and H-5 protons (average of unresolved multiplet), (m) N-CH, protons, (0) N-CH, protons, (A) CH, protons. Reference is Me,SO external standard.
shown in Table I and Figure 2. The chemical shifts of the ring protons are most sensitive to changes in N , and the side-chain protons show little composition dependence. Two things are immediately evident: there is a sharp change in behavior at N = 0.5, and chemical shifts in basic melts (N C 0.5) are more sensitive to composition changes than for acidic melts ( N > 0.5). The size of the alkyl side chains appears to have little effect on the composition dependence of chemical shifts. Figure 3 compares the shifts for the ring protons in 1-methyl-3-ethyl- and 1methyl-3-butylimidazolium melts. The addition of a third component to binary MeEtImchloroaluminate melts may also affect chemical shifts of protons
The Journal of Physical Chemistry, Vol. 88, No. 12, 1984 2611
1,3-Dialkylimidazolium Chloride-Aluminum Chloride 8.0
I
I
E
-
a a
& I c
7.0
!-
I-
Y
cn
6.0
-
5.0
-
A
a
2
I,I
I v)
1
I
0.30
I
I
I
0.50
0.40
AICI,
I
1
0.60
0
MOLE FRACTION
\
6 8 1
0
'
'
'
0 02
'
'
..
-\
0.06
0.04
-*
i
'
I
\
I
' 0.08
'
0 10
MOLE FRACTION OF THIRD COMPONENT
Figure 4. Proton chemical shifts of H-2 proton in neat ternary melts: (A) AICl,-MeEtImCl-BuPyCl, N = 0.40; (0) AlCl,-MeEtImCI-LiCl, N = 0.40; ( 0 )A1C13-MeEtImC1-LiC1, N = 0.35.
on the cation. The effects of LiCl and I-butylpyridinium chloride added to melts at constant N are shown in Figure 4. The effects of several nonelectrolytes on the binary MeEtIm melt are shown in Figure 5. For the cases shown in Figures 4 and 5 the addition of a third component results in upfield changes in chemical shifts, except for addition of 1-butylpyridinium chloride, where there is no effect. Discussion In order to understand the chemical shift behavior in terms of ion-ion interactions, the identity of the anions present at a given AlCl, mole fraction must be established. In chloroaluminate melts comprised of NaCl-AlCl, mixtures the following equilibria involving the anions have been shown to have equilibrium constants much greater than one:, M+Cl- + AlCl, M+A1Cl4-
+ AlCl,
F?
M+AlCl;
F?
M+A12C17-
(1) (2)
From eq 1 and 2 one may obtain the "chloroaluminate" equilibrium 2M+AlC14- ~t M+C1- + M'A12C17-
0.20
0.40
0.60
0.80
1.0
MOLE FRACTION OF THIRD COMPONENT
Figure 3. Proton chemical shifts in neat 1-methyl-3-ethylimidazolium chloride-aluminum chloride melts (open symbols) and l-methyl-3-b~tylimidazolium chloride-aluminum chloride melts (filled symbols) vs. mole fraction AlC1,: (0,0 ) H-2 proton, (A, A) H-4 and H-5 protons (average of unresolved multiplet).
'A
0
6.4
0.70
(3)
which has a value of in the MeEtImC1-AlC1, melts3 Equation 3 implies that in basic melts ( N < 0.5) the principal anions are C1- and AlC1,- and in acidic melts ( N > 0.5) the principal anions are AlC14- and Al2Cl,-. Some clues to the interpretation of the changes in shifts with changing proportion of AlCI, may be found in work done on some similar molten salts. Newman et al. have reported recently that
Figure 5. Proton chemical shifts of H-2 proton in neat l-methyl-3ethylimidazolium chloride-aluminum chloride-solvent melts vs. mole fraction solvent: (a) butyronitrile, (A) propionitrile, (M) acetonitriled3, (0) benzene-&
for some pyridinium and picolinium halides the chemical shifts of the N-H protons are more deshielded (shifted downfield to higher 6 values) in the order C1- > Br- > I-.1o This is in accordance with the relatively hdyrogen-bonding abilities of those halides and is evidence that the pyridinium chlorides and bromides exist as hydrogen-bonded complexes. The ring protons in the same melts exhibit the reverse order for deshielding by the halides, implying that simple electrostatic interactions do not occur between the anion and cation. The change in the shifts is in precisely the opposite direction in our experiments, where the C1- is being smoothly replaced by AlC14- as N increases in the basic melts. The ring protons in the 1-methyl-3-ethylimidazolium cations shift upfield as the Cl- is replaced with AlCl;, which should be a weaker inducer of negative charge. Angell and Shuppert have studied the proton N M R behavior of some pyridinium melts that are more closely related to 0urs.l' These workers added AlCl, or ZnClz to pyridinium chloride and observed changes in both N-H and ring protons. The N-H shifts moved upfield as AlC14- replaced C1- upon addition of AlC13,as expected from the results of Newman. The ring proton resonances on the pyridinium cations moved upfield with increasing AlCl, up to about 0.25 mole fraction A1Cl3, then moved downfield for some protons, and continued upfield for others above 0.25 mole fraction AlCl,. The explanations provided by Angell and by Newman for their data are not strictly applicable to our melts, since the existence of hydrogen bonding clearly has a profound effect on chemical shifts. Osteryoung and co-workers have done 'H and 13CN M R experiments on melts very similar to those we report here.12 While Osteryoung and co-workers did not study a large number of compositions, they did examine melts where bromine was substituted for chlorine (Le., 1-ethylpyridinium bromide-AlC1, and 1-ethylpyridinium bromideAlBr,). From the lack of differences in ring proton chemical shifts for the chloride- and bromidecontaining melts, they concluded that there was no pyridiniumhalide ion pairing in acidic melts. This is in accord with our observations (Figure 2), where we saw only small changes in chemical shifts in acidic melts. A major difference between our work and Osteryoung's is that we have concentrated our efforts on basic melts, where the C1- activity is high, rather than acidic melts, where C1- activity is very low. Popov and co-workers have studied the composition dependence for 'H and I3C shifts of the cations in the 1-butylpyridinium chloride-AlC1, system and obtained results similar to those described here for the l-methyl~~
~
~~~
(9) Fannin,A. A.,Jr.; King, L. A.; Seegmiller, D. W. J . Electrochem. Soc. 1972, 119, 801-7. ( 1 0 ) Newman, D. S.:Tillack, R. T.; Morgan, D. P.; Wan. W. J . Elecfrochem. SOC.1977, 124, 856-9. (11) Angell, C. A.; Shuppert, J. W. J . Phys. Chem. 1980, 84, 538-42. (12) Robinson, J.; Bugle, R. C.;Chum, H.L.;Koran, D.; Osteryoung, R. A. J . Am. Chem. SOC.1919, 101, 3776-9.
Fannin et al.
2612 The Journal of Physical Chemistry, Vola88, No. 12, 1984
0.8
0.6
0.4
0.2
0 0
0.1
0.2
AICI,
0.3
0.4
0.5
0.60.67
MOLE FRACTION
Figure 6. Calculated anion fractions vs. mole fraction A1C13 in chloroaluminate melts.
3-ethylimidazolium chloride-A1C13 system.13 In general terms, it appears that the less polarizable the anions are, the greater the effect on proton chemical shifts. The organic chloride salt-aluminum chloride melts that we have studied are in some ways simpler than Newman’s or Angell’s systems in terms of anion-cation interactions, since no hydrogen bonding is possible. They are more complex in the sense that three different anions need to be considered (Cl-, AlC14-, and Al2C1f), although never more than two will be present in significant quantities at any given composition. The observed direction of chemical shift changes agrees with the qualitative expectation that as greater negative charge is induced into the imidazolium cation (by C1- relative to AlC14-), the ring anisotropy is increased and the ring protons become more shielded. In order to interpret quantitatively the changes in chemical shifts with composition, we have made some assumptions about the anion-cation interctions and how they affect chemical shifts. The fundamental assumptions are that they can affect shifts by modifying the net charge of the cation and that the “complexes” formed are relatively weak so that fast exchange (on the N M R time scale) occurs at room temperature. The term “complex” means here an identifiable (but probably not isolable) structure or environment containing the dialkylimidazolium cation which may be formulated reasonably from the melt composition. The strength of the ion-ion interactions are probably stronger than the usual solute-solvent interactions and weaker than chargetransfer or hydrogen-bonded complexes. If the system is in the fast-exchange regime, the chemical shift of a proton on the cation is the population weighted average of the chemical shifts of the protons in the different complexes present, as expressed by eq 4,where 6 , and Xi are the chemical
0 5
(13) Taulelle, F.; Popov, A. I. Polyhedron 1983, 2, 889-94. (14) In the case of added nonelectrolytes, the mole fraction scale for the nonelectrolyte is defined as moles of nonelectrolyte/(moles of nonelectrolyte moles of melt). The moles of melts are calculated from the weight of the melt and the molecular weight of the melt, which is a function of N. The molecular weight of the melt ( M ) is calculated by
+
M = 146.64 + 133.34N/(1 - N) (15) Rodayne, .I.Williams, ; D. H. Annu. Rev.NMR Spectrosc. 1969, 2, 98. (16) Wilkes, J. S.; Levisky, J. A.; Pflug, J. L.; Hussey, C. L.; Scheffler, T. B. Anal. Chem. 1982, 54, 2378-9.
0.9
1 .o
MOLE FRACTION AICI,-
Figure 7. Proton chemical shifts of the H-2 proton in basic l-methyl3-ethylimidazoliumchloride-aluminum chloride melts vs. mole fraction AICI,. The solid lines are calculated as described in the text.
pairs should be present at a given melt composition and what their mole fractions are, it is instructive to calculate the fractions of the various anions as a function of composition (Figure 6) for the composition range 0 IN I0.67. The anion fractions (Y,)were calculated from eq 5-8. At N = 0, 0.5, and 0.67 only a single basic compositions Ycl-=
1 - 2N -= x1 1-N
(5)
N 1-N
Y*Ic14-=
-= x4
Y*,c4-=
2 - 3N -= x4
acidic compositions 1-N
(7)
anion is present in a significant amount: C1-, AlC14-, and Al,Clf, respectively. In basic melts varying proportions of C1- and AlC14are present, and in acidic melts varying proportions of A12C1f and AlCl; exist. For ion pairs there are three possible complexes CI-
AICIG-
I t
I
t
AlzYl7I
Im
I m
Irn
1
4
7
but a maximum of two will exist in any significant quantities at any given composition. (The notation for the complexes indicates the total number of chlorine atoms associated with the cation.) For model 1 eq 4 becomes 8obsd
shifts and mole fractions of the complexes, respectively. The simplest model (model 1) for the melt is one in which the ions exist as simple ion pairs; that is, each cation is coordinated with a single anion and vice versa. In order to determine which
0.8
0.7
06
=
61x1
-k 64x4
-k
’bx1
(9)
The observed shifts as a function of the X,’s may be least-squares fitted to eq 9, because the Xi for each species is equal to the anion mole fraction of the anion the species contains, which may be calculated from eq 5-8. The model can be tested well only for the basic compositions, because acidic compositions showed little change in shifts with AICl, mole fraction. Also, the model can be best tested by using the chemical shift data for the proton at the 2-position, because it shows the greatest change. Model 1 expressed as eq 9 is nonlinear on an AlC1, mole fraction axis but is linear on an AlC14mole fraction axis in the basic melts (X, = 0, X I = 1 - X4). 6 , is 6&d at N = 0.5 so the experimental value was used rather than a fitted value. The calculated straight line is plotted in Figure 7 along with the data for the basic melts. Since there appears to be systematic deviation from that straight line, one of the assumptions may be incorrect.
1,3-Dialkylimidazolium Chloride-Aluminum Chloride
The Journal of Physical Chemistry, Vol. 88, No. 12, 1984 2613
TABLE 11: Chemical Shift Parameters (ppm) for Eq 10
TABLE 111: Comrtarison of Fits to Proton Chemical Shift Data standard deviation, ppm N-CH2 model H-2 protons H-4, H-5 protons protons 1 0.053 0.018 0.014 2 0.020 0.014 0.014 2A 0.053 0.018 2B 0.049 0.019 0.014 3 0.020 0.014
proton H-2 H-4, H-5 N-CH,
82
85
ha
611
814a
8.25 6.56 2.13
8.14 5.54 1.98
6.05 5.03 1.89
6.07 5.06 1.97
5.97 5.04 1.95
Experimentally determined. All others were obtained from least-squares fits. That the anions and cations exist as ion pairs (Le., one nearest neighbor) is the weakest of the assumptions. It would be reasonable to assume that each cation may interact with two or more anions and vice versa. For two nearest neighbors, an analysis similar to the one above for ion pairs may be made. In this case (model 2) five different complexes are possible:
61I
t
I m
I m
2
5
Im
8
11
14
These species are not discrete, but we believe they represent portions of oligomeric chains of alternating cations and anions. For model 2 eq 4 becomes bbsd
=
62x2
+
+ s8x8 + 611xll + 61&14
The model where the coordination number of each ion was 3 (model 3) was tested by a procedure entirely analogous to that for the model 2 discussed above. The model involves seven possible complexes, and five of the 6,'s are adjustable parameters. With that many adjustable parameters almost any curve could be matched by the fit, but the fit was not improved relative to the two-coordinate model. The qualities of the fits, as indicated by standard deviations, for the models having different assumptions and coordination numbers are compared in Table 111. Since the number of adjustable parameters changes for the different models, the number of degrees of freedom ( q ) was included in the calculation of standard deviations.
(lo)
If the distribution of the five complexes is determined by a random distribution of the anions among the cations, then the X i s may be calculated from the anion mole fractions, similar to the way they were calculated for model 1. Thus, the X,'s for eq 10 are basic compositions
all compositions
acidic compositions
&, 88, and 614 should be the observed chemical shifts at N = 0, 0.5, and 0.67, respectively, so where possible experimental value$ were used rather than fitted values. No measurement was made at 0 mole fraction AICl3, since that composition is not a liquid at room temperature. The observed chemical shifts were leastsquares fitted to eq 10, and the values of the 6, coefficients are tabulated in Table 11. The calculated shifts from model 2 are plotted in Figure 7 on an AlC4- mole fraction axis, as was model 1. In this case, the equation for the model is quadratic in terms of X4 in the basic melts. In contrast to the earlier model, this one fits all of the experimental data very well. The assumption that the anions are randomly distributed among the cations could be false. One could envision that the anions are distributed such that there are chloride-rich, tetrachloroaluminate-rich, and heptachloroaluminate-rich regions, where complexes 2, 8, and 14 are preferred (model 2A). Alternatively, complexes 2, 8, and 14 may be selectively avoided, resulting in a melt structure that maximizes cations associated with different anions (model 2B). The complexes present would be the same as those presented earlier for model 2, but their mole fractions would be different. In order to test models 2A and 2B, we calappropriate to the culated the shifts according to eq 10 using x,'~ limiting nonrandom distributions (i.e., the preferred configurations exist to the limits imposed by stoichiometry). In both cases the fits to the chemical shift data for basic melts were poorer than for the random distribution model.
u is the standard deviation, x is the measured value, y is the calculated value, and n is the number of measured values. Clearly, model 2 fits the data well. In order for model 2 to be believkble, the values of the 6,'s obtained from the curve fitting must be reasonable physically. The values of 6,'s from model 2 for the H-2 proton are 82 = 8.25, 65 = 8.14, b8 = 6.05, 6,, = 6.07, and fi14 = 5.97 ppm. The 6,'s for the species present in acidic melts (8, 11, and 14) are all very similar, as expected due to the similarity of the chloroaluminate anions involved. Among the species prdent in basic melts (2,5, and 8), 2 is furthest downfield because it has two C1- nearest neighbors and 8 is furthest upfield because it has two AlC14- nearest neighbors. Species 5 has an intermediate shift, since it has one of each anion as nearest neighbors. Lithium chloride is soluble at room temperature in basic MeEtIm-chloroaluminate melts up to approximately 0.1 mole fraction LiCI. Since the chloride mole fraction in the melt has a great effect on the chemical shifts of the cation protons, the experiments with added LiCl were performed at constant AlCl, mole fraction and a constant sum of LiCl and MeEtImCl mole fractions. This ensures that the anion fractions of C1- and A1Cl4remain constant. The experiments may be viewed as a replacement of the imidazolium chloride with LiCl. Figure 4 shows that there is an apparently linear decrease in the chemical shift of the H-2 proton with increasing LiCl mole fraction. The behavior may be rationalized qualitatively by viewing Li+ as a better competitor for negative charge from the anions than the imidazolium cations. As the LiCl fraction increases, the fraction of Im+ ions whose nearest cation neighbor is Li+ increases, resulting in a higher proportion of Im+-containing complexes that have a diminished induction of negative charge. Li+ is such a good competitor for C1- relative to Im' that the discrete complex LiCl, probably forms in basic melts,I4 thus depriving the Im+ of C1- nearest neighbors. We have not developed a quantitative model that satisfactorily explains the LiCl ternary data. Ternary mixtures composed of MeEtImCl, AlC1, and an organic nonelectrolyte also show interesting changes in proton chemical shifts of the Im+ cation (Figure 5 ) . In the four nonelectrolytes studied the shifts move upfield with increasing mole fraction of n ~ n e l e c t r o l y t e .Again, ~~ we can interpret the changes in terms of the modification of the anion-cation interactions. We propose that the effect of added solvent is to break up the ionic interactions by solvation of the ions, thus reducing the electrostatic influence of the anion on the cation. Greater amounts of solvent would show a greater effect. Three of the solvents studied comprise the
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J . Phys. Chem. 1984,88, 2614-2621
homologous series: acetonitrile, propionitrile, and butyronitrile. Acetonitrile has the greatest effect, propionitrile a lesser effect, and butyronitrile the least. That order is the same as the order of dielectric constants for the solvents, and one would expect that the greater the dielectric constant the more effective the solvent would be in solvating the ions. Benzene presents an especially interesting case. Because it is quite nonpolar, it is somewhat surprising that benzene is soluble at all in the totally ionic liquids, although Osteryoung and coworkers have reported that it was soluble in the butylpyridinium chloroaluminates. They found that it modified the electric conductivity and it affected proton and chemical shifts of the cation.'* Benzene is not as soluble in the basic melts as the series of nitriles (which are miscible in all proportions), but it has a greater effect on the chemical shifts of the cation protons in the range where it is soluble. This may be due to a more specific interaction between the imidazolium cation and benzene, such as the complex between the two aromatic species suggested for the butylpyridinium melts.I2 This could result in aromatic solvent-induced shifts, where the ring current of the benzene shields the protons of the imidazolium ~ a t i 0 n s . l ~ The dependence of chemical shifts on N in basic melts implies that proton N M R measurements could be used as an analytical
method for composition determination. Such a method has been reported recently.I6 Conclusions
Proton N M R chemical shifts of the hydrogens located in the cations of organic chloride-aluminum chloride room-temperature molten salt mixtures are sensitive to composition. The changes in chemical shifts may be explained in terms of anion-cation interactions, where the imidazolium cation is distributed among five different anionetion complexes. Addition of lithium chloride or one of several organic solvents disrupts the complexes and changes the proton chemical shifts.
Acknowledgment. We thank Lloyd H u g for running the NMR spectra reported here and Dr. Charles Hussey for many helpful discussions. L.A.K. gratefully acknowledges the support of Universal Energy Systems, Inc., Dayton, OH, and of the Air Force Wright Aeronautical Laboratories, Wright-Patterson AFB, OH. Registry No. AlC13, 7446-70-0; LiC1, 7447-41-8; ImCl, 1467-16-9; MeEtImCl, 65039-09-0; 1-methyl-3-butylimidazolium chloride, 7991790-1; butyronitrile, 109-74-0;propionitrile, 107-12-0; acetonitrile, 7505-8; benzene, 71-43-2.
Properties of 1,3-Dialkylimldazollum Chloride-Aluminum Chloride Ionic Liquids. 2. Phase Transitions, Densities, Electrical Conductivities, and Viscosities Armand A. Fannin, Jr., Danilo A. Floreani, Lowell A. King, John S. Landers, Bernard J. Piersma, Daniel J. Stech, Robert L. Vaughn, John S. Wilkes,* and John L. Williams The Frank J . Seiler Research Laboratory, US.Air Force Academy, Colorado Springs, Colorado 80840 (Received: February 18, 1983; In Final Form: April 4, 1983)
Experimental values are reported for the densities and electrical conductivities of several 1,3-dialkylimidazoliurnchlorides and for the densities, electrical conductivities, and viscosities of representative compositions of binary mixtures of these salts with aluminum chloride. These data were collected over wide temperature and composition ranges. The transport properties are reported as specific and equivalent conductivities and as kinematic and absolute viscosities. These data and the densities are interpreted in terms of a model of the structure of the binary melts. The solid-liquid phase diagram for the 1methyl-3-ethylimidazolium-aluminumchloride system was determined, including glass transition temperatures.
The work described here is a part of a continuing study designed to develop new low-temperature electrolytes for battery applications. This study of the phase transitions, densities, electrical conductivities, and viscosities of dialkylimidazolium chloridealuminum chloride binary melts is the second in a series of papers on the properties of these melts.' In an earlier paperZ from this laboratory the densities, conductivities, and viscosities were reported for several 1-alkylpyridinium chlorides and their binary mixtures with AlCl,. The dialkylimidazolium chlorides used in this study are shown in Chart I. The 1-methyl-3-ethylimidazolium chloride (MeEtImCl)-AICl, system was chosen as a "base-line" melt, primarily because of its relatively favorable condpctivity and viscosity behavior and its wide liquid range. Only a few compositions were studied for the other binary systems, in order to rule out unexpected anomalies. Studies of the proton nuclear magnetic resonance (NMR) spec(1) Fannin, A. A., Jr.; King, L. A.; Levisky, J. A.; Wilkes, J. s.J . Phys. Chem., preceding article in this issue. (2) Carpio, R. A.; King, L. A.; Lindstrom, R. E.; Nardi, J. C.; Hussey, C. L. J . Electrochem. SOC.1979, 126, 1644-50.
CHART I.
R +C I R,/N+N\R3
compd MeMeImCl MeEtImCl MePrImCl MeBuImCl BuBuImCl
R1 methyl methyl methyl n-butyl n-butyl
R3
methyl ethyl n-propyl n-butyl n-butyl
troscopy of these systems were presented in the first paper in the present series.' Experimental Section Sample Preparation. The dialkylimidazolium salts and their binary mixtures with AlCl, were prepared as described elsewhere., All sample preparation and handling (except in sealed apparatus) were conducted in an argon-filled glovebox (Vacuum/Atmosphere Co. box and Model MO-40 DRI TRAIN), having a combined (3) Wilkes, J. S.; Levisky, J. A,; Wilson, R. A.; Hussey, C. L. Inorg. Chem. 1982, 21, 1263-4.
This article not subject to U.S. Copyright. Published 1984 by the American Chemical Society
