1830
J. Phys. Chem. 1983, 8 7 , 1830-1833
Cesium-I 33 Nuclear Magnetic Resonance Study of the Complexation of Cesium Salts by 18-Crown-6 in Methylamine and Liquid Ammonia. 2. 2:l Complex Formation Sadegh Kharaell, James L. Dye, and Alexander I. Popov' Department of Chemistry, Michigan State University, East Lansing, Michigan 48824 (Received: August 30, 1982)
The variation of the 133Cschemical shift (6) with the temperature and with the (l&crown-G)/(Cs+)mole ratio ( R )in methylamine and ammonia solutions indicated the formation of both 1:l and 2:l complexes. The mole ratio data for R > 1were analyzed according to the equilibrium CsC+.X- + C CsC2+.X-(KX,,Nox2), where X-, C, CsC+.X-,and CsC2+.X-are anion, ligand, ion-paired 1:l complex, and ion-paired 21 complex, respectively. The thermodynamic parameters were obtained for CsI and CsBPh4in methylamine solutions with corrections introduced for ion association of the salts and the 1:l complexes. The value of Kx,is larger for cesium tetraphenylborate than for the iodide, which indicates a competition between ion-pair formation and complex formation. The value of K,, for CsBPhl was much larger in ammonia than in methylamine, which reflects the difference in the degree of association of the complexes in these two solvents. In all cases, 2:l complex formation was enthalpy stabilized but entropy destabilized.
Introduction The complexation of the cesium cation by the crown ether 18-crown-6 (18C6) in aqueous and some nonaqueous solvents has been studied by potentiometric,' calorimetr i ~ , ~and b 133CsNMR technique^.^.^ In all cases the formation of both 1:l and 2:l complexes (18CS)/(Cs+))was detected. In our previous papers6*'we reported that the ion-pair formation constants of Cs+ salts and Cs+.18C6 complexes in methylamine are of comparable magnitudes. These results suggest that in methylamine solutions both the salts and 1:l complexes form mainly noncontact (Le., solventseparated or ligand-separated) ion pairs. We also measured free energy, enthalpy, and entropy changes for the formation of the 1:l complex between Cs+ and 18C6. In this paper we describe the formation of 21 complexes of cesium salts by 18-crown-6 in methylamine solutions as well as in liquid ammonia solutions. The effect of the anion and the solvent upon the complexation reaction is discussed. Experimental Section 1. Chemicals. The purification of the cesium salts and of 18C6 was described in previous publi~ations.~*' Methylamine (Matheson, anhydrous, 98%) was first dried over calcium hydride, distilled repeatedly onto an NaK mirror under high vacuum, and then distilled into a storage bottle.s Liquid ammonia (Matheson, anhydrous 99.99% ) was dried with NaK alloy in a similar way. The preliminary drying step over CaH, was omitted in this case because of the purity of the ammonia. Methylamine and ammonia solutions were prepared in 10-mm 0.d. precision NMR tubes (Wilmad) with wall thicknesses of 0.5 and 1.0 mm, respectively. All samples (1) Christensen, J. J.; Eatough, D. J.; Izatt, R. M. Chem. Reu. 1974, 74, 351. (2) Izatt, R. M.; Terry, R. E.; Haymore, B. L.; Hansen, L. D.; Dalley, N. K.; Avondet, A. G.; Christensen, J. J. J. Am. Chem. SOC.1976,98,7620. (3) Izatt, R. M.; Terry, R. E.; Nelson, D. P.; Chan. Y.;Eatough, D. J.; Bradshaw, J. S.; Hansen, L. D.; Christenwn, J. J. J.Am. Chem. SOC.1976, 98, 7626. (4) Mei, E.; Dye, J. L.; Popov, A. I. J.Am. Chem. SOC.1977,99, 5308. (5) Mei, E.; Popov, A. I.; Dye, J. L. J . Phys. Chem. 1977, 81, 1677. (6) Khazaeli, S.; Popov, A. I.; Dye, J. L. J. Phys. Chem. 1982,86,4238. (7) Khazaeli, S.; Popov, A. I.; Dye, J. L. J. Phys. Chem. 1982,86, 5018. (8) Dye, J. L. J . Phys. Chem. 1980,84, 1084.
were prepared under high vacuum ('.x-! PPm b 6 c s C , + . X - , P P deg-l ~ x2,
-
CsI in methylamine
CsBPh, in methylamine
CsBPh, in ammonia'
4.03 (5)b - 0 . 8 3 (1) -6.05 (8) -17.5 (3) -46.2 ( 6 )
22.8 ( 4 ) -1.85 ( 1 ) -7.4 (1) -18.4 ( 4 ) -49.4 (2)
700 ( 4 0 ) -3.88 ( 4 ) -4.9 (3) -3.5 (9) 40.1 ( 7 ) - 0 . 3 (1) 0.42
1.55 0.82 PPm a No correction was introduced for dissociation of the ion-paired 1:1 complex. "Thermodynamic" parameters are conditional. The standard deviation estimate of the last digit is given in parentheses; e.g., 4.03 ( 5 ) means 4.03 I 0.05. Standard deviation of the chemical shift. 06,c
to dMe+.X- due to the presence of small relative concentrations of (CsC+) was made according to (4) Lc+.x-= (1- Y ) b + . x - + Y b l c + in which y is the fraction of the 1:l complex present as CsC+. The value of y was computed from the equilibrium constants determined previ~usly.~ ( b ) Correction for (Cs+.X-). Since the formation of CsC+.X- from Cs+.X- and C is not complete nor of the same extent at various temperatures, the presence of Cs+-X-would affect the apparent chemical shift of the 1:l complex. The concentration of Cs+.X- depends on the concentration of the ligand and can be calculated at each concentration and temperature. Thus, a more exact "effective chemical shift" (~5'C~c+.~-)for CsC+.X- to be used with equilibrium 3c was calculated according to d'cse+.x-= (1- z)6csc+.x-f z6cs+.x(5) in which z is the fraction of CsC+.X- which is dissociated into Cs+.X- and C, and Bc8+c.x- is defined by eq 4. ( c ) Correction for the Total Concentration of the Salt. Since the samples were prepared by weighing the salt and complexant, the total salt concentrations were not exactly the same at various mole ratios. Therefore, the correct total concentration was used at each mole ratio (together with its proper standard deviation estimate) in the final calculation. The results with corrections a-c are given in Table I for the CsI and CsBPh, 2:l complexes in methylamine. All parameters are well determined with this simple model. The chemical shift of CsC2+.X-is essentially the same for both salts and agrees well with the chemical shift of the 2:l complex in other nonaqueous ~ o l v e n t s .The ~ value of Kx2 for CsBPh, is larger than that for CsI, which reflects differences in ion-association constants of their 1:l and 2:l complexes. Although the average standard deviation of the chemical shifts, s6,are higher than the estimated experimental errors, this is not unexpected since the complete scheme was not used to fit the data. In addition, other factors might contribute to the deviations: (1) 6csc2+.xmight be temperature dependent as was the case7 for Cs+.X- and CsC+.X-. This temperature dependence cannot be included since KA2cannot be calculated. (2) The mole ratio study of the complexation of CsBPh, by 18crown-6 in ammonia (section 2 below) indicated that in this case the solvent molecules do interact with the cesium cation in the 2:l complexes. Therefore, we cannot reject the possibility of similar interactions in methylamine although the larger size of the methylamine molecule makes this less likely. If such interactions occur, then a t least
I l r I l 0.8 1.6 2.4
*OOO
[Cs']
x
I@
(MI
Flguro 3. Concentration dependence of the I3%s chemical shift of CsI In the presence of a 6.0-fold excess of l&crown-6 in methylamine at various temperatures. Solid lines are calculated curves.
two kinds of ion pairs (ligand separated and solvent separated) may be present. The equilibrium between these two kinds of ion pairs might be temperature dependent and is not included in the present treatment of the data. (3) The chemical shift might depend on the concentration of the free ligand. If this effect exists in our system, it could not be determined from the mole ratio plots, since the variation of the chemical shift with mole ratio is large and masks minor effects. In addition to the mole ratio studies, the concentration dependence of the 133Cschemical shift for CsI in the presence of a 6.0-fold excess of 18-crown-6was studied in methylamine a t various temperatures. The results are shown in Figure 3. The calculated chemical shifts at each concentration and temperature were determined from the parameters given in Table I and are shown as the solid curves. The agreement is satisfactory in view of the problems discussed earlier. In summary, the ion-association equilibria involving the 2:l complex could not be completely described on the basis of the NMR technique. However, the fit of the data by a simple model permitted the determination of Kx2 and AHox2. Attempts to expand the model did not provide more information about the system because of the incapability of the NMR technique to separate different kinds of ion association. 2. Ammonia Solution. Cesium-133 chemical shifts of cesium tetraphenylborate in the presence of 18-crown-6 were measured in liquid ammonia solutions as a function of the mole ratio and temperature at a fixed salt concentration; the results are shown in Figure 4. In order to examine ion association of the salt, the chemical shift of cesium tetraphenylborate was also determined as a function of concentration a t 6.0 "C. The results are shown as the inset to Figure 4 and indicate that in liquid ammonia the cesium cation also forms ion pairs. Complete investigation of ion association and complexation in ammonia solutions would require an extensive study similar to that in methylamine solutions. However, it is expected that ion association in liquid ammonia would be considerably less extensive than in methylamine solutions because of the higher dielectric constant of the former solvent (D= 23 a t -33 "C). The mole ratio plots (Figure 4) show the formation of both 1:l and 2:l complexes. The data above R = 1 were analyzed according to equilibrium 3c, which assumes that the 1:l complex is completely formed and ion paired a t R = 1. This treatment ignores ion-pair dissociation and, therefore, gives only approximate values of the parameters. The chemical shift of CsC+.X- at each temperature was set equal to the observed chemical shift at R = 1. Since the limiting chemical shift of the 2:l complex in ammonia is very different from that in methylamine (Table I), it appears that the solvent can still
J. Phys. Chem. 1903, 87, 1833-1834
I I
40
I I I I
50
I
Flgwe 4. Cesium133 chemical shift vs. (18-crown-0)/(CsBPh,) mole ratio and temperature in ammonia; (CsBPh,) = 0.001 M. Solid lines are calculated curves. Dashed lines connect points at mole ratios of 0 and 1. Inset: Ceslum-133 chemical shift vs. concentration of CsBPh, in liquid ammonia at 6.0 OC.
interact with the cesium cation in the 2:l complex. Therefore, a linear temperature dependence was also assumed for 8cscz+.x-according to
(6cBcz+.x-)t = ( ~ c ~ c Z + . ~ - ) 2+5 "b(t c - 25OC)
(6)
Four parameters, Kx2 AHox2, (8cSc+ . x - ) ~ ~and ~ c ,b were adjusted. The results are given in $able I. The average standard deviation of the chemical shift (a, = 0.42 ppm) is about twice the experimental error. The mole ratio dependence of the chemical shift of CsBPhl in ammonia was also studied a t a total concentration of 0.0075 M of the salt a t 14.5 "C. The value of
1033
Kx2 obtained from these results was less than half the corresponding value given in Table I for (Cs+) = 0.001 M at the same temperature. These results indicate that ion association is important in liquid ammonia and would have to be included in a complete treatment. The "thermodynamic parameters" obtained for ammonia solutions are, therefore, conditional in nature. It appears that the formation constant for the 2:l complex in liquid ammonia is much larger than in methylamine. If ion association of the 1:l and 2:l complexes could be ignored, then one would expect the reverse order since ammonia is a better donor solvent than methylamine. The limiting chemical shift of the 2:l complex at high values of R in methylamine is almost temperature independent and practically equal to the values in other nonaqueous solvent^.^ Surprisingly, the limiting chemical shift in ammonia is more than 80 ppm downfield as compared to other solvents. This large downfield chemical shift could be due to the interaction of the cesium cation in the sandwhich complex with the solvent, since the chemical shift of the free cation in liquid ammonia is larger than 122 ppm. A more complete description of the formation of the 2:l complex in ammonia would require determination of the ion-association parameters of both the salt and the 1:l complex. Formation of the 2:l complex (equilibrium 3c) is enthalpy stabilized but entropy destabilized in both solvents. The entropy of formation of the 2:l complex seems to be nearly anion independent, but strongly solvent dependent. However, the enthalpy of the complexation reaction is both anion and solvent dependent. In methylamine solutions, differences in the stability of the 2:l complexes of cesium iodide and cesium tetraphenylborate with 18-crown-6 are mainly determined by the enthalpy contribution to the free energy of formation. The larger complexation constant for cesium tetraphenylborate compared to that for cesium iodide reflects the difference in the degree of ion association of their corresponding salts and complexes. The much more positive entropy of formation in liquid ammonia might be due to stronger solvation of the cesium cations by ammonia molecules than by methylamine molecules.
Acknowledgment. We gratefully acknowledge the support of this work by National Science Foundation Grants CHE-80-10808 (A.I.P.) and DMR 79-21979 (J.L.D.) Registry No. CsI, 7789-17-5; CsBPh,, 3087-82-9.
COMMENTS Decarboxylationof Acetate Ions
Sir: Electron spin resonance spectroscopic studies have revealed many details about radiation processes in recent year~.l-~In particular, the use of single crystals has shown (1) H. C. Box, 'Radiation Effects E.S.R. and ENDOR Analysis", Academic Press, New York, 1977. (2) M. C. R. Symons, Pure Appl. Chem., 53, 223 (1981). (3) S. Ya. Pshezhetakii, A. G. Kotov, V. K. Milinchuk, V. A. Rozinskii, and V. I. Tupikov, 'E.P.R. of Free Radicals in Radiation Chemistry", Wiley, New York, 1974. 0022-365418312087-1833$01.50/0
that neutral or ionic fragmenta are often trapped in specific sites with specific orientations, a general finding being that changes relative to the original geometry are often small. When a primary radical product undergoes unimolecular decomposition, the resulting fragments are usually expected to remain in their original sites. A clear example of this is the detection of intermolecular hyperfine coupling between the radical and nonradical fragments, as found, for example, for alkyl radical-halide ion adductsS4+ (4) E. D. Sprague and F. Williams, J . Chem. Phys., 54, 5425 (1971).
0 1983 American Chemical Society
