2420
The Journal of Physical Chemistry, Vol. 82, No. 22, 1978
J. Vaes, M. Chabanel, and M. L. Martin
Ionic Interactions in Lithium Thiocyanate Solutions. Nitrogen4 5 and Lithium-7 Nuclear Magnetic Resonance Studies J. Vaes, M. Chabanel," Laboratoire de Spectrochimie des Ions, Chemin de la Houssini&e, 44072 Nantes Cedex, France
and M. L. Martin Laboratoire de Chimie Organique Physique, CNRS-ERA 3 15, Chemin de la Houssini8re. 44072 Nantes Cedex, France (Received May 3, 1978) Publication costs assisted by the University of Nantes
The association of Li+ with SCN- has been studied in several aprotic solvents by 15Nand 'Li NMR. The chemical shift variations are interpreted by an ion pairing equilibrium (dimethylformamide) or by a dimerization equilibrium (dimethyl carbonate, ether). A decrease of B I S N with concentration is observed and this result shows with the formation of N-bonded species in both cases. However, in tetrahydrofuran, the slight increase of f N concentration indicates the formation of small amounts of S-bonded species. This result is confirmed by the appearance of a weak absorption band at a frequency (2087 cm-') which is much higher than for SCN- or for N-bonded species.
Introduction Owing to the recent facilities for a direct NMR observation of metal cations such as Li+, Na+, and Cs+, an increasing number of investigations is concerned with the problem of ionic interactions in sol~tions.l-~Although sometimes of qualitative nature, interesting information has been obtained about ion pairing in different solvents. The case of lithium thiocyanate is especially appropriate to a n NMR study because both anion and cation can be detected by means of the three convenient probes, nitrogen, lithium, and carbon. As the range of 7Li chemical shifts is known to be narrow, a relatively poor sensitivity of the shift to the mode of association is expected in this case. On the contrary the widespread range of nitrogen chemical shifts is quite favorable. Actually, due to the privileged role of the nitrogen lone pair 615Nis a remarkable criterion of electron distribution, and it is often more suitable than 613c for studying electron delocalization in conjugated structure^.^ Although the 14Nisotope benefits from good signal intensity performances, the lack of precision resulting form quadrupolar line broadening restricts the information to a somewhat qualitative nature. By contrast, very accurate chemical shift values can be obtained from 15N resonance. However, isotopically enriched samples are required to work on dilute solutions. In previous studies, carried out by infrared spectroscopy and by other techniques! we have identified the different species, ions, ion pairs, and dimers, which are formed when LiSCN is dissolved in different aprotic solvents. These data will remove the ambiguities which may arise from the uniqueness of the observed chemical shift. Experimental Section KSC15N was prepared according to the method given by Rykowsky and V a n d e r p l a ~ .A~ similar method was used for LiSC15N, with KOH being replaced by LiOH. The reaction product in methanol was evaporated under vacuum. The residue was dissolved in ether and the solution was evaporated. LiSCN was dried under vacuum a t 110 OC for 24 h. It was analyzed by Volhard's method. The four solvents were stored over molecular sieves. T H F and ether were first dried over sodium wire and then distilled on a 20-plate column; DMC was also distilled. 0022-3654/78/2082-2420$01 .OO/O
The purity of these solvents was checked by dielectric measurements. DMF (Merck, spectroscopic grade) did not need further purification. I5N spectra were obtained with a digital resolution of about 0.9 Hz by means of WP-60 and WH-90 Bruker spectrometers. The chemical shift reference was a solution of Na15N03 (30% enriched) in acidified D20. This solution, which was also used to lock the spectrometer, was contained in a cylindrical 2-mm tube coaxially inserted in the 10-mm sample tube. 'Li spectra were recorded on the same apparatus, with a resolution of about 0.3 Hz. The reference was a saturated solution of LiCl in D20. Positive chemical shifts correspond to low field effects vs. the reference. The measured chemical shifts should have been corrected for the magnetic susceptibility effects associated with the use of an external reference. In fact, owing t o imperfect shape factors, this correction is often inaccurate and was not performed. The variations in this contribution arising from changes in the salt concentration may be safely neglected and no influence upon the shapes of the curves has been suspected. However, it should be kept in mind that the absolute values of the chemical shifts contain a bulk susceptibility effect which may be of the order 0.5 ppm. This correction is nearly negligible as compared with the large range of nitrogen chemical shifts, but it is relatively important when lithium chemical shifts are concerned.
Results and Discussion Mode of Fixation of Lithium. LiSCN has been studied in four solvents which are, in order of decreasing ionizing ability N,N-dimethylformamide (DMF), tetrahydrofuran (THF), dimethyl carbonate (DMC), and diethyl ether (ether). The 15N chemical shift variations are given in Figure 1 as a function of the thiocyanate concentration, c. The 40-ppm range covered by these variations is large enough for an accurate study. However, this range is relatively narrow as compared with the variations observed when other MNCS bonds are formed (e.g., about 100 ppm for M = Zn or H). A similar behavior is observed in IR spectroscopy where small frequency shifts result from the fixation of an alkali ion on the SCN group. 0 1978 American Chemlcal Society
The Journal of Physical Chemistry, Vol. 82, No. 22, 1978 2421
Ionic Interactions in Lithium Thiocyanate Solutions
I
I
-16
i < 2 "7
U
-
17
-18
-191
, -
Ether (LiSCN)
-
-201
Ether
n
-
V - -
h
v
THF
1.5
0.5
c (3.11
L i SCN
c-.ntwion,M
Flgure 2. 'Li chemical shift of LiSCN solutions.
I
I
0.5
SCN
1 sas.mmion.M
I 15
where x is the fraction of thiocyanate which exists in the form of ion pairs. x is related to the equilibrium constant K by
Flgure 1. 15N chemical shift of LiSCN and KSCN solutions.
The lSN chemical shift of SCN- in DMF was obtained from a study of potassium thiocyanate solutions. According to IR spectra the apparent equilibrium constant for the formation of ion pairs is small (-0.5 M-9: and also K+ is characterized by a low polarizing ability. Therefore it is not surprising to find only a slight variation in the 61aN of KSCN when the concentration is increased (Figure 1). The extrapolated value GiN(DMF)= -159.5 ppm is close to the value (-166 ppm) obtained for SCN- in water through 14N r e s o n a n ~ e . ~ As shown in Figure 1 all nitrogen shifts are situated toward high fields with respect to the SCN- anion. According to the criterion established by Howarth et this behavior indicates that the associations preferentially involve the nitrogen atom (Li-N bonds) and are thus of an isothiocyanate type. In this respect, high field variations of 8laN with increasing concentration show the formation of increasing amounts of species associated with nitrogen. On the contrary formation of bonds with the sulfur atom introduces low field shifts. Consequently it will be seen later that the low field effects observed in THF solutions when the concentration is increased can be interpreted by the formation of small proportions of species where lithium is bound to sulfur. Free Ions-Ion Pair Equilibrium (DMF). For LiSCN solutions in DMF the curve &IaN = f(c) fairly extrapolates at the chemical shift of the SCN- anion (-160 ppm) and its shape is characteristic of a simple ion pairing equilibrium.' Owing to the accurate results obtained in this case, a determination of the apparent equilibrium constant, K, for ion pair formation is possible. The measured chemical shift 6 is the weighted mean of the chemical shifts in the ion, 6i, and in the ion pair, ,6, hence 6 =
6i
+ (6,
- 6i)X
x =
1
+ 2Kc - (1 + ~ K C ) ' ' * 2Kc
In this equation the activity correction has not been considered and K is an apparent formation constant. Therefore, it is assumed that the activity coefficients are nearly constant within the investigated concentration range. It should be emphasized that the value of K may be very different from the true equilibrium constant a t infinite dilution. As shown by the reported equations, K can be determined by searching for the value which yields the best correlation coefficient in a linear regression between 6 and x . The calculated value, K = 2.1 f 0.2 M-l, is in good agreement with the value K = 2 M-' deduced from IR measurements in the same concentration rangees This comparison is interesting because, in the latter case, we have separated absorption bands instead of a weighted mean. The nitrogen chemical shift in the ion pair has been estimated a t -182 ppm. The 13Cchemical shift was also measured, but it remains restricted to values between 150 and 152 ppm in the 0.1-1 M concentration range and exhibits no definite trend. Therefore the carbon nucleus does not appear to be a good probe to study thiocyanate associations. This behavior suggests that ion pair formation is accompanied by an electron redistribution between CN and CS bonds which leaves the chemical shift of the common carbon atom nearly unchanged. This interpretation is in agreement with the charge transfer from sulfur to nitrogen proposed on the basis of IR results.8 Although the range of variations in the lithium chemical shift is relatively narrow, a regular decrease of vs. concentration is observed (Figure 2). An upfield shift of about 0.8 ppm characterizes ion pair formation. Zon Puir-Dimer Equilibrium (DMC,Ether). A different kind of equilibrium certainly exists in DMC solutions since
2422
J. Vaes, M. Chabanei, and M. L. Martin
The Journal of Physical Chemistry, Vol. 82, No. 22, 1978
I
log
1
IC$
111111111111 11111111111 2050
cm-'
2100
2050
2050
2100
cm-l
cm-1
c
b (a) 2
a
Figure 3. Infrared spectra of LiSCN in tetrahydrofuran solutions: M LiSCN (film): (b) 0.013 M LiSCN ( e = 470 nm); (c) 0.013 M LiSCN 2 M LiC104 (e = 470 nm). p indicates ion pair; d, dimer; a, S-bonded aggregate: e , path length.
+
the curve 815N= f(c) is now entirely upfield with respect to that determined in DMF. Moreover the curve extrapolates to a value typical of the ion pair (Figure 1). This result is in agreement with an equilibrium between ion pairs and dimers as identified by IR spectroscopy.8 Owing to the rather narrow range of chemical shift variations no attempt was made to obtain K from the 6'6N curve. However, by using approximate values of the fractions of both species deduced from IR spectra, the shift's parameters could be determined: GPN(DR/IC)= -187 ppm for the ion pair, and 8dN(DMC)= -198 ppm for the dimer. It can be seen that the values of ~ 3 , in ~ DMC and in DMF (-182 ppm) are close. Similar results were obtained by lithium NMR (Figure 2) where an upfield shift of about 0.6 ppm is caused by the dimerization of ion pairs. In ether solutions, nitrogen and lithium chemical shifts are nearly constant (Figures 1 and 2) and the curve 6N = f(c) extrapolates to a value (-199 ppm) previously shown as typical of the dimer. Dimerization is accompanied with an increase in the screening constants of N and Li nuclei. The shift variations are of the same order of magnitude as for ion pair formation. This behavior indicates that dimerization involves an additional interaction a t the nitrogen and lithium sites, and this result confirms the quadrupole structure of the dimers8
.Li .. Nd!-S
S-C=N, .
*
'Li'
Equilibrium between Ion Pairs, Dimer, and S-Bonded Species (7°F). In dilute solutions only small changes are experienced by the nitrogen resonance and the extrapolated chemical shift value (-192 ppm) corresponds to the ion pair. This result is in agreement with the observation in dilute solutions (c 0.01 M) of a single IR band at a frequency (2064 cm-') characteristic of the ion pair. At higher concentrations (- 1 M) a noticeable low-field effect occurs. As previously mentioned, this behavior indicates that the ion pairs form higher aggregates bonded through the sulfur atom. This interpretation is corroborated by the IR spectra in concentrated solutions (Figure 3). The band of the ion pair is still the most important. A weak band, characteristic of the dimer, appears a t 2036 cm-'. On the other side a new band is observed a t 2087 cm-'. Relative to the free ion SCN- (uCN = 2050 cm-' in THF) the shift of this band is more than twice as large as for LiNCS. I t is well known that the shift of U C N is more important when the SCN group is S bonded or bridging
-
than when it is N bonded.'(' Such an effect is observed in the AgSCN ion pair." Furthermore, when LiC104is added to a dilute LiSCN solution in THF, the band at 2087 cm-' also appears (Figure 3). This behavior can be explained in the following manner. In the SCN ligand there are two basic sites, N and S. In the case of Li+ the first site to be used is the nitrogen atom, in accordance with the HSAB classification.12 However, when this site is occupied, a second Li+ cation can be fixed either on the same site (dimer formation), or on the sulfur atom as follows: Li
'
..
dimer
...Lit + Li ...NCS'
.
Li. .NCS. . .Li. .
chain polymer
The shape of the curves aN and gLi vs. c suggests that successive (LiSCN), chain species are formed when the concentration is increased. The fact that L i 4 interactions take place in T H F can be related to the higher value of its dielectric constant ( t = 7.6) with respect to the other solvents investigated ( t N 3-4). Polar species are then more favored, and it is also known13 that, when a competition is possible, S-bonded species are preferentially formed in a medium with a high dielectric constant. Actually the IR band at 2087 cm-' is also observed in very concentrated DMC solutions, but with a small relative intensity. Chemical Shift Variations. The kind of equilibrium observed here (ion pairing or dimerization) mainly depends on the dielectric constant and the basicity of the s01vent.l~ Good agreement between 15N and 'Li data is obtained despite the differences in the factors which govern the chemical shifts. As regards the small variations in the lithium screening constant both the diamagnetic and paramagnetic contributions may be involved and, in practice, the magnetic anisotopy effects of neighboring groups frequently predominate. As each step of the association process corresponds to the replacement of a solvent molecule by a thiocyanate ion in the first solvation shell of the lithium cation, either high or low field effects should be expected. In fact isothiocyanate formation is shown to introduce a high field effect. By contrast the variations in nitrogen chemical shifts are mainly governed by the paramagnetic term and the influence of electron excitation energies has been put forward to explain the behavior of 6 I r N in a lot of structurally different thiocyanates and isothio~yanates.~ A decrease in the charge localized a t nitrogen, brought about by association with lithium, would be accompanied by an increase in the
'%d
NMR Study of Cadmium(I1) Halide
Complexes
electron expansion term and therefore a low-field shift. Although the effect of electron redistribution in the thiocyanate anion may actually be rather complex it is likely that the large upfield variations following isothiocyanate bonding are mainly due to an increase of excitation energies. Such an increase arises from the bonding of the nitrogen lone pair with the lithium atom.
Acknowledgment. The infrar:d measurements were conducted by DaniBle Paoli to whom grateful appreciation is expressed. References and Notes (1) M. S. Greenberg, R. L. Bodner, and A. I. Popov, J . Phys. Chem., 77. 2449 (19731. (2) Y. 'M. Cahen, P.'R. Handy, E. T. Roach, and A. I. Popov, J . Phys. Chem., 79, 80 (1975).
The Journal of Physical Chemistry, Vol. 82, No. 22, 1978 2423
G. E. Maciel, J. K. k n & , L. F. Lafferty, P. A. Mueller, and K. M a k e r , Inorg. Chem., 5, 554 (1966). W. J. Dewitte, L. Liv, J. L. Dye, and A. I. Popov, J . Solution Chem., 6. 337 (1977). d. J. Martin, J. P. Gouesnard, J. Dorie, C. Rabiller, and M. L. Martin, J. Am. Chem. Soc., 99, 1381 (1977). M. Chabanel, C. MBnard, and G. GuihBneuf, C. R . Acad. S d . , Ser. C , 272, 253 (1971); D. Menard and M. Chabanel, J. Phys. Chem., 75, 1081 (1975). A. Rykowski and H. L. Vanderplas, Rec. J . Roy. Neth. Soc., 94, 204 (1975). D. Paoli, M. Lucon, and M. Chabanel, Spectrochim. Acta, in press. 0. W. Howarth, R. E. Richards, and L. M. Venanzi, J . Chem. SOC., 3335 (1964). A. H. Norbury, Adv. Coord. Chem., 17, 246 (1975). D. Paoli and M.Chabanel, C. R . Acad. Sci., Ser. C , 264, 95 (1977). R. G. Pearson, J . Chem. Educ., 45, 643 (1968). J. L. Burmeister, R. L. Massel, and R. J. Phelan, Chem. Commun., 679 (1970). C . MBnard. 8. Woitkowiak. and M. Chabanel, C. R . Acad. Sci., Ser. C , 278, 553 (19j4).
Cadmium-113 Nuclear Magnetic Resonance Study of Cadmium(I1) Halide Complexes in Water and Dimethyl Sulfoxide Torbjorn Drakenberg,' Nils-Olof Bjork,' and Roberto Portanovat Division of Physical Chemistry, and Division of Inorganic Chemistry, Chemical Center, University of Lund, Lund, Sweden, and Istituto di Impianti Nuclear!, Universiti di Palerrno, Palermo, Italy (Received March 9, 1978; Revised Manuscript Received June 12, 1978)
'H NMR of the MezSO protons, 13CNMR of the Me2S0 carbons, and l13Cd NMR have been used to study the formation of various cadmium(I1) halide complexes in ivIezSO solutions. The same complexes have also been studied in water solutions by l13Cd NMR. The l13Cd chemical shifts for the various complexes have been evaluated from both solvents. The differences in the chemical shifts of the individual complexes between water and Me2S0 solutions are in agreement with the fact that the change from octahedral to tetrahedral complexes take place at different complexation steps in the two solvents.
Introduction X-ray diffraction measurements carried out for aqueous solutions of cadmium(I1) perchlorate' have shown that the hydrated Cd2+ion is coordinated to six water molecules with an octahedral configuration. On the other hand, it has also been shown that the tetraiodo complex, CdId2-, is a regular tetrahedron in aqueous solution.' In water, therefore, a change of coordination certainly occurs, and the thermodynamic functions AC,",AHj",and ASj" for the cadmium(I1)-halide complex formation in aqueous solution2 are consistent with a change in coordination taking place with the formation of the third complex, so that the initially octahedral hydrated Cd2+ion is converted to a tetrahedral halide complex. More recently, X-ray diffraction investigations on solutions of cadmium(I1) perchlorate in MezSO have shown that the solvated metal ion is coordinated to six Me$O molecules3 and that the tetraiodo complex is tetrahedrally ~ o o r d i n a t e d .The ~ changes in AGjo, AHjo,and ASj" for cadmium halide complex formation have also been interpreted as due to a change in coordination from a purely Octahedral configuration to a tetrahedral 0118,~s~ similar to that taking place in aqueous solution. In Me2S0,however, this change seems to occur a t an earlier step than in ~ater.~,~ *Division of Physical Chemistry. 'Division of Inorganic Chemistry. t Instituto di Impianti Nucleari. 0022-3654/78/2082-2423$01 .OO/O
In an attempt to obtain more information about the conformational changes in the cadmium(I1) halide complexes, we decided to study these systems with various NMR methods and in the present work some results from the NMR studies on these systems will be presented. Information about the sensitivity of 'H, 13C, and l13Cd chemical shifts to variations of the cadmium(I1) species in solution is reported, and the advantage of l13Cd NMR compared to 'H and I3C is shown. Experimental Section (a) Measurements. The l13Cd NMR spectra were recorded on a modified Varian XL-100-15 spectrometer operating in the Fourier transform (FT)mode at 22.2 MHz and using 20-mm sample tubes. An external proton lock was used, and typical FT parameters were as follows: spectral width 8000 Hz, acquisition time 0.5 s, flip angle 20" (pulse width 30 p s ) and lo00 to 1OOo00 transients. The 13C NMR spectra were run on a Jeol F X 60 FT spectrometer a t 15 MHz using a 10-mm sample tube containing acetone-d, as lock signal surrounding a 5-mm tube containing the sample. Typical FT parameters were spectral width 500 Hz, acquisition time 16 s, and 60" flip angle (9 11s). The 'H NMR spectra were run at 100 MHz on a Jeol MH 100 spectrometer. All spectra were recorded at 30 f 2 "C. (b) Chemicals. The hexasolvate Cd(MezSO)6(C104)z was prepared and analyzed as described beforea7Ammonium chloride (BDH), ammonium bromide (B&A), and am@ 1978 American Chemical Society
