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Vapor phase osmometric studies of alkali metal salts in acetone and tetrahydrofuran solutions
.I. inorg, nucl. Chem., 1972, Vol. 34, pp. 3615-3622.
Pergamon Press.
Printed in Greal Britain
VAPOR PHASE OSMOMETRIC STUDIES OF ALKALI METAL SALTS IN ACETONE AND TETRAHYDROFURAN SOLUTIONS M I N G K E O N G W O N G and A L E X A N D E R I. POPOV Department of Chemistry, Michigan State University, East Lansing, Michigan 48823 (Received 24 February 1972) Abstract--Apparent mol. wts have been determined for a number of lithium and sodium salts in acetone and in tetrahydrofuran solutions in the 0.01-0.15 M concentration range. The degrees of association have been estimated from the data and are found to be strongly solvent and anion dependent. In acetone solutions, tetraphenylborates show nearly complete dissociation while thiocyanates, bromides and especially chlorides are strongly associated. In tetrahydrofuran solution all salts are strongly associated into ion pairs and for the chloride and the thiocyanate, to higher ionic aggregates. INTRODUCTION
recently shown that far i.r. spectra of alkali cations in acetone solutions are characterized by vibrational bands resulting from the motion of the cation in the solvation shell [1]. The frequencies were independent of the salt anion as long as large polyatomic anions were used, but the bands shifted to the lower frequencies when the anions were Cl- or Br-. These frequency shifts have been attributed to the penetration of the solvation shell by the halide ion and the formation of a contact ion pair. The situation is somewhat different in the tetrahydrofuran solutions where far i.r. bands of alkali metal salts show strong frequency dependence irrespective of the nature of the anion [2, 3]. It appears, therefore, that in this solvent the salts exist predominantly as contact ion pairs. Both acetone and tetrahydrofuran have medium donor strengths as given by Gutmann's donor numbers [4] of 17 and 20, respectively. However, acetone has a dielectric constant of 20.7615], compared to 7.39[6] for tetrahydrofuran. The difference in the dielectric constants may be reflected by the difference in dissociating power of the two solvents. A search of the literature showed that ion pair dissociation constants of electrolytes in acetone solutions have been determined conductometrically for a number of salts [7-11 ]. However, few such data W E HAVE
1. M. K. Wong, W. J. McKinney and A. I. Popov, J. phys. Chem. 75, 56 (1971). 2. (a) W. F. Edgell, A. T. Watts, J. Lyford and W. M. Risen, J. Am. chem. Soc. 88, 1815 (1966): (b) W. F. Edgell, J. Lyford, R. Wright, W. M. Risen and A. Watts, ibid. 92, 2240 (1970). 3. A.T. Tsatsas and W. M. Risen, ibid. 92, 1789 (1970). 4. V. Gutmann, Coordination Chemistry in Nonaqueous Solvents, p. 19. Springer-Verlag, Vienna (1968). 5. National Bureau of Standards Circular 514, U.S. Government Printing Office, Washington, D.C. (1951). 6. C. Carvajal, K. J. Tolle, J. Smid and M. Szwarc,J.Am. chem. Soc. 87, 5548 (1965). 7. P. G. Sears, E. D. Wilhoit and L. R. Dawson, J. phys. Chem. 60, 169 (1956). 8. D. F. Evans, C. Zawoyski and R. L. Kay, ibid. 69, 3878 (1965). 9. L. G. Savedoff, J. A m. chem. Soc. 88, 664 (1966). 10. W. A. Adams and K. J. Laidler, Can. J. Chem. 46, 1977, 1989, 2005 (1968). 11. G. S. Darbari and S. Petrucci, J. phys. Chem. 74, 268 (1970). 3615
3616
M . K . WONG and A. I. POPOV
are available for the tetrahydrofuran solutions [12, 13]. A comparison of the ion pair dissociation constants in the two solvents can be made for lithium chloride. The respective values are Kd = 3"3 × 10-6 in acetone and Kd ---- 1.2 × 10-1° in tetrahydrofuran indicating the influence of the low dielectric constant of the latter solvent. There are very few reports on vapor phase osmometric studies of electrolytes in acetone and tetrahydrofuran solutions. Peska, et al. [14] determined the apparent mol. wts, Maoo, of sodium perchlorate and potassium thiocyanate by thermoelectric method and found this value to be strongly concentration dependent. The authors calculated Mrea1 from Mao o by using the osmotic coefficients from vapor phase osmometric measurement and the degrees of association from conductance measurement. Sokolov and Lindberg[15] studied the effects of concentration and temperature on the M~oo of sodium iodide in acetone solutions by osmometric measurements. They also calculated the dissociation constant of the salt and found that both Mao o and Kd were concentration and temperature dependent. It was of interest to us to use vapor phase osmometry for the study of electrolyte solutions in acetone and tetrahydrofuran and to see if the information obtained by this technique can be correlated with the i.r. data. Both solvents have low boiling points and high vapor pressures, and thus are ideal for vapor phase osmometric measurements. EXPERIMENTAL
Reagents The purification of acetone and of the alkali metal salts has been described in a previous publication [1]. Reagent grade tetrahydrofuran from Matheson, Coleman & Bell was used without further purification. The water content of the solvent, as determined by the Karl-Fisher titration, was approximately 0.02%. Eastman Kodak reagent grade benzil was recrystallized twice from absolute ethanol and then vacuum dried at 50°C for 2 days [ 16]. Eastman Kodak reagent grade biphenyl was used without further treatment. Measurements All measurements were performed using a Mechrolab Inc., high temperature vapor pressure osmometer, Model 302. The method outlined in the operational manual[17] was used. The sample chamber was maintained at a constant temperature of 37°C, and at least 4 hr were allowed for equilibration. Benzil was used as calibration standard. The concentrations ranged from 0.02 to 0-2 M for benzil, and 0-02 to 0-16 M for other salts. Biphenyl was used as an additional check on the calibration constants obtained with benzyl. In all the measurements, the solvent and sample drops at the thermister beads were kept of approximately the same size. The resistance (AR) values were recorded 2 min (accurate to within 5 sec) after the sample drops were deposited on the thermister bead. Approximately 15-20 min prewarm periods were allowed for each solution. The solvent in the solvent cup (used to saturate the chamber) was replaced after every two or three series of runs. At least two readings of AR were taken for each solution and the averaged value was used for calculation. 12. 13. 14. 15. 16. 17.
W. Strohmeier, A. E. Mahgoub and F. Gernert, Z. Elektrochem. 65, 85 (1961). D. N. Bhattacharyya, C. L. Lee, J. Smid and M. Szwarc, J. phys. Chem. 69,608 (1965). J. Peska, J. Biros and M. Benes, Czech. Chem. Comm. 33, 1333 (1968). V.W. Sokoiov and J. J. Lindberg, Suomen Kemistilehti B43, 367 (1970). A. Adicoffand W. J. Murbach, Analyt. Chem. 39, 302 (1967). Mechrolab Inc., Model 302 vapor pressure o s m o m e t e r - Operating and Service Manual.
Studies of alkali metal salts RESULTS
AND
3617
DISCUSSION
Seven lithium and four sodium salts, listed in Table 1, were studied in both solvents. T w o calculation methods were used. The first was the limiting slope method [18] (or "constant K" method according to the Mechrolab Manual). The calibration constant K of the instrument was obtained by measuring AR for a series of benzil solutions at different concentrations and extrapolation AR/C,,, values to infinite dilution according to the equation
(AR/C,,,)c-o- K/M,, where AR = decastat reading of the osmometer, C,,, = concentration of benzil in moles lit -l, K = calibration constant of the instrument and M,, = mol. wt of benzil. Similar AR/C vs C plots were obtained for mol. wt determinations in electrolyte solutions. In this case, however, C was expressed in glit -1 so that the apparent mol. wt of the electrolyte was calculated from the equation
Maoo =
K -- (,~tR/C)c-o.
In the second method, the calibration curve was obtained by plotting AR/C,, vs hR. The AR value for each concentration of solution was entered into the calibration plot to find AR[Cm, from which a series of apparent Cm values could be calculated. The apparent mol. wt of the electrolyte at that concentration was obtained by dividing the weight concentration of the solution by the apparent Cm value. The AR/C,, vs Cm calibration curve for benzil in tetrahydrofuran is shown in Fig. 1. The calibration curve of AR/Cm vs AR is similar to Fig. 1, except that AR is used instead of C,~. The AR/Cg/~vs Cg/~ plots for some of the salts are shown in Table I. The apparent molecular weights of alkali metal salts in acetone. Calibration constant K , , - 424
Salt LiCl LiBr Li| LiNO:~ LiSCN LiCIO4 LiBPh~ Nal NaSCN NaCIOa NaBPh4 Biphenyl
(AR/CgIOC_o 7-60 5.40 4.46 6.00 6.80 5.30 2.24 3.70 5.65
4.59 2-06 2.68
Mol. wt (apparent)
Mol. wt (real)
55-8 78-5 95.1 70.7 62.4 80.0 189-3 114.6 75.0 92.4 205-8 158.2
42.4 86.9 133.8 68.9 65.0 106.4 326.2 149.9 81.1 122.4 342-2 154.2
18. R. U. Bonnar, M. Dimbat and F. H. Stross, N u m b e r - A v e r a g e N e w York (1958).
M~../Mreai Conc. range 1.32 0.90 0.71 1.03 0.96 0.75 /)-58 [).76 0.92 0-75 (/.60 1.03
Molecular
0-03-0.10 0.03-0.14 0.03-0.14 0.04-0.13 0.03-0.13 0.04-0.14 0.01-0.09 0.04-0-15 0.04-0.15 0.04-0.15 0.01-0-10 0.02-0-14
Weights.
[nterscience.
3618
M . K . W O N G and A. I. P O P O V
500
450
~
O0
0
0
.
0
0
°-°--%'--m°-~-°-°
--
o o---.
o
o_8_o__~
400
i
0 -04
i
I
0-08
0-12
i
0"16
I
0 -20
Cm
Fig. 1. The
AR/Cmvs Cm calibration plot. Benzil in tetrahydrofuran.
Fig. 2 for acetone and in Fig. 3 for tetrahydrofuran. The results obtained are given in Tables 1 and 2. It should be noted that in our solutions of electrolytes, we expect to have considerable amount of solute-solvent interaction as well as the formation of ion pairs and, possibly, higher ionic aggregates. In addition, with few exceptions, the nonlinearity of the AR/Cg/I VS Cg/1plots at concentrations lower than about 0.03 M for most the salts renders the apparent mol. wt obtained from extrapolation to infinite dilution subject to sizeable error. However, in the 0.03-0.14 M concentration range, a qualitative comparison of the apparent mol. wts of the various salts in a given solvent and that between the two solvents seems to be of significance. It is seen from Table 1 that for lithium salt in acetone solutions, the tetra-
4.5~..._o
q.....
4'01 4.4
@ ' " ~ ~"-"-"-~ ~
~)--~.£~ A
~ 0
4
8
12
16
20
Cgl(
Fig. 2. The
AR/Cg/Ivs Cg/l plots for (a) NaCIO4, (b) LiI, (c) NaI in acetone solutions.
Studies of alkali metal salts
3619
4.6 o
3.8
o o
°'°°~°~o--.qo ...o...4p~
3.0 4.2"~o o <11~ 3.6
o o0~..o~
'[
"~-ooa.
3-oJ: 4'~?
°'°'2"~ 00"---..~.Q
3"61" 3.01
B
0.....0.~.~"
/
| |
A
~@""0~-....~ o
,
,
o......o......°
,
4
.
8
,
12
C "7~ °'o---... ,
,
16
Cg/L Fig. 3. T h e
AR/Cg/I vs Cgll plots for (a) N a S C N ,
(b) LiBr, (c) LiC104 in T H F solutions.
Table 2. T h e apparent molecular weights of alkali metal salts in tetrahydrofuran. Calibration c o n s t a n t Km = 429
Salt
(AR/Cg/1)c=o
Mol. wt (apparent)
Mol. wt (real)
M~,,o/Mrea,
Conc. range
LiC1 LiBr Lil LiNO~ LiSCN LiC104 LiBPh4 NaI NaSCN NaC104 NaBPh4 Biphenyl
5.54 4-03 2-96 4.59 6.18 4.01 1.27 2.30 3.74 3.00 1.14 2.75
77.4 106.5 144.9 93.5 69.4 107.0 337.8 186.5 114.7 143.0 376.3 156.0
42.4 86.9 133-8 68.9 65.0 106.4 326.2 149.9 81.1 122.4 342.2 154.2
1.83 1.23 1.08 1.36 1.07 1.01 1.04 1.24 1.41 1.17 1.10 1.01
0.04-0-14 0.03-0.13 0.02-0.16 0.05-0-14 0.02-0.14 0-02-0.14 0.02-0.09 0.03-0.16 0.05-0-15 0.04-0-15 0.02-0.11 0.02-0.16
phenylborate is the most dissociated, the iodide and the perchlorate are dissociated to a lesser degree while no appreciable dissociation is observed for the nitrate, thiocyanate and the bromide. Apparently, higher aggregates are present in appreciable amounts in lithium chloride solutions. Sodium iodide, sodium perchlorate and sodium tetraphenylborate are also appreciably dissociated in this solvent, with greatest dissociation being observed
3620
M . K . WONG and A. I. POPOV
for the tetraphenylborate. No appreciable dissociation is observed, however, in the sodium thiocyanate solutions. From the greater dissociation of lithium tetraphenylborate, lithium iodide, lithium perchlorate and the sodium salts, we would expect that the above electrolytes exist in solutions mainly as free solvated ions or as solvent-separated ion pairs. The apparent mol. wts for lithium nitrate, lithium thiocyanate and lithium bromide indicate no appreciable dissociation in the concentration range studied. Even in these cases, solvent separated ion pairs cannot be ruled out especially for the salts with polyatomic anions. In the case of lithium chloride solutions in acetone it is reasonable to assume that, in addition to contact ion pairs, there exist higher ionic or molecular aggregates. In addition to the evidence from our spectroscopic studies [1] the fact that the chloride and the bromide ions do form strong bonds with lithium ion have also been shown by the low ion pair dissociation constants from conductance [9] and the low ion pair reactivity of lithium bromide from kinetic isotopic exchange measurements [ 19]. Thus the results from vapor phase osmometric measurements agree with the trends observed in the previous studies. With the exceptions of biphenyl and sodium tetraphenylborate, the apparent mol. wts of all the salts studied show some degree of concentration dependence in acetone solutions. The plots of the apparent mol. wts at each salt concentration vs molar concentration for some of the salts are shown in Fig. 4. In tetrahydrofuran solutions, from the apparent mol. wts obtained (Table 2), it
-oO-o (~ oo
°.--o-.o----o--
19
0 12(
-B
0 0. . . ~ 0 ~ " ~
g
100
•
.....~ _ . . ~ _ _ o o . _ _ _ ~
c ._~_.__D_ D
80
60 i
0'04
,
0"08
i
0.12
i
i
0-16
Cm Fig. 4. The apparent molecular weights (Mapp) vs molar concentration plots: (a) NaBPH4, (b) NaI, (c) NaCIO4, (d) NaSCN in acetone solutions. 19. P. Beronius, U. lsacsson and A. M. Nilsson,Acta chem. scand. 24, 189 (1970).
Studies of alkali metal salts
3621
is seen that in the concentration range studied, none of the salts show any evidence of dissociation. In fact, association to form higher aggregates is appreciable for most of the salts, especially at higher concentrations. This trend is obvious from the plots of the apparent mol. wt at each salt concentration vs concentration, some of which are shown in Fig. 5. The greatest degree of association was again exhibited by lithium chloride whose apparent mol. wt approaches that of a dimer.
,oJ 14o
A
--&---O--O------C,--
C ~ O - -
~.D--
19(3 B ~o
)6C
..L)._0,o__0___~.0_--00 -
13G
C
lOG .4~._.0dO__--.-O- O- O- o--O---O---&. D
7G
i
0.04
0-08
0.12
0.16
C,,,
Fig. 5. The apparent molecular weights (M~pp) vs molar concentration plots: (a) biphenyl, (b) Lil, (c) LiBr, (d) LiCI in tetrahydrofuran solutions.
The above observations are consistent with the far i.r. data of Edgell et al. [2] which indicate anion-cation association either in the form of contact ion pairs or as larger ionic aggregates. Similarly the conductance data of Edgell et al. [20] on solutions of sodium tetracarbonylcobaltate in tetrahydrofuran indicate formation of ion pairs even in dilute solution and the formation of ionic clusters, triplets and quadruplets, in the more concentrated solutions. A comparison of the Maoo/Mreal ratios of various salts in acetone and in tetrahydrofuran shows the greater dissociation power of the former solvent. Under identical conditions, solvent-separated ion pairs are more likely to dissociate than the contact ion pairs or the solvent-shared ion pairs. Thus, for the alkali metal salts studied, with few exceptions (notably the chloride and bromide salts of lithium), solvent-separated ion pairs should be the predominant species in acetone solutions, and the anion effect on the far i.r. cation-acetone vibrational frequencies is minimal. This is not the case for tetrahydrofuran solutions. The much weaker dissociating power of this solvent results in the predominant forma20. W. F. Edgell, M. T. Yang and N. Koizumi,J. Am. chem. Soc. 87, 2563 (1965).
3622
M . K . W O N G and A. I. P O P O V
tion of contact ion pairs. Thus the frequencies of the far i.r. vibrational bands are anion dependent. For many of the salts, the attraction between the cation and anion is so strong that at higher concentrations, aggregates higher than ion pairs will form to an appreciable degree. Acknowledgement- The authors gratefully acknowledge the support of this work by a grant from the National Science Foundation.
Pergamon Press.
Printed in Greal Britain
VAPOR PHASE OSMOMETRIC STUDIES OF ALKALI METAL SALTS IN ACETONE AND TETRAHYDROFURAN SOLUTIONS M I N G K E O N G W O N G and A L E X A N D E R I. POPOV Department of Chemistry, Michigan State University, East Lansing, Michigan 48823 (Received 24 February 1972) Abstract--Apparent mol. wts have been determined for a number of lithium and sodium salts in acetone and in tetrahydrofuran solutions in the 0.01-0.15 M concentration range. The degrees of association have been estimated from the data and are found to be strongly solvent and anion dependent. In acetone solutions, tetraphenylborates show nearly complete dissociation while thiocyanates, bromides and especially chlorides are strongly associated. In tetrahydrofuran solution all salts are strongly associated into ion pairs and for the chloride and the thiocyanate, to higher ionic aggregates. INTRODUCTION
recently shown that far i.r. spectra of alkali cations in acetone solutions are characterized by vibrational bands resulting from the motion of the cation in the solvation shell [1]. The frequencies were independent of the salt anion as long as large polyatomic anions were used, but the bands shifted to the lower frequencies when the anions were Cl- or Br-. These frequency shifts have been attributed to the penetration of the solvation shell by the halide ion and the formation of a contact ion pair. The situation is somewhat different in the tetrahydrofuran solutions where far i.r. bands of alkali metal salts show strong frequency dependence irrespective of the nature of the anion [2, 3]. It appears, therefore, that in this solvent the salts exist predominantly as contact ion pairs. Both acetone and tetrahydrofuran have medium donor strengths as given by Gutmann's donor numbers [4] of 17 and 20, respectively. However, acetone has a dielectric constant of 20.7615], compared to 7.39[6] for tetrahydrofuran. The difference in the dielectric constants may be reflected by the difference in dissociating power of the two solvents. A search of the literature showed that ion pair dissociation constants of electrolytes in acetone solutions have been determined conductometrically for a number of salts [7-11 ]. However, few such data W E HAVE
1. M. K. Wong, W. J. McKinney and A. I. Popov, J. phys. Chem. 75, 56 (1971). 2. (a) W. F. Edgell, A. T. Watts, J. Lyford and W. M. Risen, J. Am. chem. Soc. 88, 1815 (1966): (b) W. F. Edgell, J. Lyford, R. Wright, W. M. Risen and A. Watts, ibid. 92, 2240 (1970). 3. A.T. Tsatsas and W. M. Risen, ibid. 92, 1789 (1970). 4. V. Gutmann, Coordination Chemistry in Nonaqueous Solvents, p. 19. Springer-Verlag, Vienna (1968). 5. National Bureau of Standards Circular 514, U.S. Government Printing Office, Washington, D.C. (1951). 6. C. Carvajal, K. J. Tolle, J. Smid and M. Szwarc,J.Am. chem. Soc. 87, 5548 (1965). 7. P. G. Sears, E. D. Wilhoit and L. R. Dawson, J. phys. Chem. 60, 169 (1956). 8. D. F. Evans, C. Zawoyski and R. L. Kay, ibid. 69, 3878 (1965). 9. L. G. Savedoff, J. A m. chem. Soc. 88, 664 (1966). 10. W. A. Adams and K. J. Laidler, Can. J. Chem. 46, 1977, 1989, 2005 (1968). 11. G. S. Darbari and S. Petrucci, J. phys. Chem. 74, 268 (1970). 3615
3616
M . K . WONG and A. I. POPOV
are available for the tetrahydrofuran solutions [12, 13]. A comparison of the ion pair dissociation constants in the two solvents can be made for lithium chloride. The respective values are Kd = 3"3 × 10-6 in acetone and Kd ---- 1.2 × 10-1° in tetrahydrofuran indicating the influence of the low dielectric constant of the latter solvent. There are very few reports on vapor phase osmometric studies of electrolytes in acetone and tetrahydrofuran solutions. Peska, et al. [14] determined the apparent mol. wts, Maoo, of sodium perchlorate and potassium thiocyanate by thermoelectric method and found this value to be strongly concentration dependent. The authors calculated Mrea1 from Mao o by using the osmotic coefficients from vapor phase osmometric measurement and the degrees of association from conductance measurement. Sokolov and Lindberg[15] studied the effects of concentration and temperature on the M~oo of sodium iodide in acetone solutions by osmometric measurements. They also calculated the dissociation constant of the salt and found that both Mao o and Kd were concentration and temperature dependent. It was of interest to us to use vapor phase osmometry for the study of electrolyte solutions in acetone and tetrahydrofuran and to see if the information obtained by this technique can be correlated with the i.r. data. Both solvents have low boiling points and high vapor pressures, and thus are ideal for vapor phase osmometric measurements. EXPERIMENTAL
Reagents The purification of acetone and of the alkali metal salts has been described in a previous publication [1]. Reagent grade tetrahydrofuran from Matheson, Coleman & Bell was used without further purification. The water content of the solvent, as determined by the Karl-Fisher titration, was approximately 0.02%. Eastman Kodak reagent grade benzil was recrystallized twice from absolute ethanol and then vacuum dried at 50°C for 2 days [ 16]. Eastman Kodak reagent grade biphenyl was used without further treatment. Measurements All measurements were performed using a Mechrolab Inc., high temperature vapor pressure osmometer, Model 302. The method outlined in the operational manual[17] was used. The sample chamber was maintained at a constant temperature of 37°C, and at least 4 hr were allowed for equilibration. Benzil was used as calibration standard. The concentrations ranged from 0.02 to 0-2 M for benzil, and 0-02 to 0-16 M for other salts. Biphenyl was used as an additional check on the calibration constants obtained with benzyl. In all the measurements, the solvent and sample drops at the thermister beads were kept of approximately the same size. The resistance (AR) values were recorded 2 min (accurate to within 5 sec) after the sample drops were deposited on the thermister bead. Approximately 15-20 min prewarm periods were allowed for each solution. The solvent in the solvent cup (used to saturate the chamber) was replaced after every two or three series of runs. At least two readings of AR were taken for each solution and the averaged value was used for calculation. 12. 13. 14. 15. 16. 17.
W. Strohmeier, A. E. Mahgoub and F. Gernert, Z. Elektrochem. 65, 85 (1961). D. N. Bhattacharyya, C. L. Lee, J. Smid and M. Szwarc, J. phys. Chem. 69,608 (1965). J. Peska, J. Biros and M. Benes, Czech. Chem. Comm. 33, 1333 (1968). V.W. Sokoiov and J. J. Lindberg, Suomen Kemistilehti B43, 367 (1970). A. Adicoffand W. J. Murbach, Analyt. Chem. 39, 302 (1967). Mechrolab Inc., Model 302 vapor pressure o s m o m e t e r - Operating and Service Manual.
Studies of alkali metal salts RESULTS
AND
3617
DISCUSSION
Seven lithium and four sodium salts, listed in Table 1, were studied in both solvents. T w o calculation methods were used. The first was the limiting slope method [18] (or "constant K" method according to the Mechrolab Manual). The calibration constant K of the instrument was obtained by measuring AR for a series of benzil solutions at different concentrations and extrapolation AR/C,,, values to infinite dilution according to the equation
(AR/C,,,)c-o- K/M,, where AR = decastat reading of the osmometer, C,,, = concentration of benzil in moles lit -l, K = calibration constant of the instrument and M,, = mol. wt of benzil. Similar AR/C vs C plots were obtained for mol. wt determinations in electrolyte solutions. In this case, however, C was expressed in glit -1 so that the apparent mol. wt of the electrolyte was calculated from the equation
Maoo =
K -- (,~tR/C)c-o.
In the second method, the calibration curve was obtained by plotting AR/C,, vs hR. The AR value for each concentration of solution was entered into the calibration plot to find AR[Cm, from which a series of apparent Cm values could be calculated. The apparent mol. wt of the electrolyte at that concentration was obtained by dividing the weight concentration of the solution by the apparent Cm value. The AR/C,, vs Cm calibration curve for benzil in tetrahydrofuran is shown in Fig. 1. The calibration curve of AR/Cm vs AR is similar to Fig. 1, except that AR is used instead of C,~. The AR/Cg/~vs Cg/~ plots for some of the salts are shown in Table I. The apparent molecular weights of alkali metal salts in acetone. Calibration constant K , , - 424
Salt LiCl LiBr Li| LiNO:~ LiSCN LiCIO4 LiBPh~ Nal NaSCN NaCIOa NaBPh4 Biphenyl
(AR/CgIOC_o 7-60 5.40 4.46 6.00 6.80 5.30 2.24 3.70 5.65
4.59 2-06 2.68
Mol. wt (apparent)
Mol. wt (real)
55-8 78-5 95.1 70.7 62.4 80.0 189-3 114.6 75.0 92.4 205-8 158.2
42.4 86.9 133.8 68.9 65.0 106.4 326.2 149.9 81.1 122.4 342-2 154.2
18. R. U. Bonnar, M. Dimbat and F. H. Stross, N u m b e r - A v e r a g e N e w York (1958).
M~../Mreai Conc. range 1.32 0.90 0.71 1.03 0.96 0.75 /)-58 [).76 0.92 0-75 (/.60 1.03
Molecular
0-03-0.10 0.03-0.14 0.03-0.14 0.04-0.13 0.03-0.13 0.04-0.14 0.01-0.09 0.04-0-15 0.04-0.15 0.04-0.15 0.01-0-10 0.02-0-14
Weights.
[nterscience.
3618
M . K . W O N G and A. I. P O P O V
500
450
~
O0
0
0
.
0
0
°-°--%'--m°-~-°-°
--
o o---.
o
o_8_o__~
400
i
0 -04
i
I
0-08
0-12
i
0"16
I
0 -20
Cm
Fig. 1. The
AR/Cmvs Cm calibration plot. Benzil in tetrahydrofuran.
Fig. 2 for acetone and in Fig. 3 for tetrahydrofuran. The results obtained are given in Tables 1 and 2. It should be noted that in our solutions of electrolytes, we expect to have considerable amount of solute-solvent interaction as well as the formation of ion pairs and, possibly, higher ionic aggregates. In addition, with few exceptions, the nonlinearity of the AR/Cg/I VS Cg/1plots at concentrations lower than about 0.03 M for most the salts renders the apparent mol. wt obtained from extrapolation to infinite dilution subject to sizeable error. However, in the 0.03-0.14 M concentration range, a qualitative comparison of the apparent mol. wts of the various salts in a given solvent and that between the two solvents seems to be of significance. It is seen from Table 1 that for lithium salt in acetone solutions, the tetra-
4.5~..._o
q.....
4'01 4.4
@ ' " ~ ~"-"-"-~ ~
~)--~.£~ A
~ 0
4
8
12
16
20
Cgl(
Fig. 2. The
AR/Cg/Ivs Cg/l plots for (a) NaCIO4, (b) LiI, (c) NaI in acetone solutions.
Studies of alkali metal salts
3619
4.6 o
3.8
o o
°'°°~°~o--.qo ...o...4p~
3.0 4.2"~o o <11~ 3.6
o o0~..o~
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"~-ooa.
3-oJ: 4'~?
°'°'2"~ 00"---..~.Q
3"61" 3.01
B
0.....0.~.~"
/
| |
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,
,
o......o......°
,
4
.
8
,
12
C "7~ °'o---... ,
,
16
Cg/L Fig. 3. T h e
AR/Cg/I vs Cgll plots for (a) N a S C N ,
(b) LiBr, (c) LiC104 in T H F solutions.
Table 2. T h e apparent molecular weights of alkali metal salts in tetrahydrofuran. Calibration c o n s t a n t Km = 429
Salt
(AR/Cg/1)c=o
Mol. wt (apparent)
Mol. wt (real)
M~,,o/Mrea,
Conc. range
LiC1 LiBr Lil LiNO~ LiSCN LiC104 LiBPh4 NaI NaSCN NaC104 NaBPh4 Biphenyl
5.54 4-03 2-96 4.59 6.18 4.01 1.27 2.30 3.74 3.00 1.14 2.75
77.4 106.5 144.9 93.5 69.4 107.0 337.8 186.5 114.7 143.0 376.3 156.0
42.4 86.9 133-8 68.9 65.0 106.4 326.2 149.9 81.1 122.4 342.2 154.2
1.83 1.23 1.08 1.36 1.07 1.01 1.04 1.24 1.41 1.17 1.10 1.01
0.04-0-14 0.03-0.13 0.02-0.16 0.05-0-14 0.02-0.14 0-02-0.14 0.02-0.09 0.03-0.16 0.05-0-15 0.04-0-15 0.02-0.11 0.02-0.16
phenylborate is the most dissociated, the iodide and the perchlorate are dissociated to a lesser degree while no appreciable dissociation is observed for the nitrate, thiocyanate and the bromide. Apparently, higher aggregates are present in appreciable amounts in lithium chloride solutions. Sodium iodide, sodium perchlorate and sodium tetraphenylborate are also appreciably dissociated in this solvent, with greatest dissociation being observed
3620
M . K . WONG and A. I. POPOV
for the tetraphenylborate. No appreciable dissociation is observed, however, in the sodium thiocyanate solutions. From the greater dissociation of lithium tetraphenylborate, lithium iodide, lithium perchlorate and the sodium salts, we would expect that the above electrolytes exist in solutions mainly as free solvated ions or as solvent-separated ion pairs. The apparent mol. wts for lithium nitrate, lithium thiocyanate and lithium bromide indicate no appreciable dissociation in the concentration range studied. Even in these cases, solvent separated ion pairs cannot be ruled out especially for the salts with polyatomic anions. In the case of lithium chloride solutions in acetone it is reasonable to assume that, in addition to contact ion pairs, there exist higher ionic or molecular aggregates. In addition to the evidence from our spectroscopic studies [1] the fact that the chloride and the bromide ions do form strong bonds with lithium ion have also been shown by the low ion pair dissociation constants from conductance [9] and the low ion pair reactivity of lithium bromide from kinetic isotopic exchange measurements [ 19]. Thus the results from vapor phase osmometric measurements agree with the trends observed in the previous studies. With the exceptions of biphenyl and sodium tetraphenylborate, the apparent mol. wts of all the salts studied show some degree of concentration dependence in acetone solutions. The plots of the apparent mol. wts at each salt concentration vs molar concentration for some of the salts are shown in Fig. 4. In tetrahydrofuran solutions, from the apparent mol. wts obtained (Table 2), it
-oO-o (~ oo
°.--o-.o----o--
19
0 12(
-B
0 0. . . ~ 0 ~ " ~
g
100
•
.....~ _ . . ~ _ _ o o . _ _ _ ~
c ._~_.__D_ D
80
60 i
0'04
,
0"08
i
0.12
i
i
0-16
Cm Fig. 4. The apparent molecular weights (Mapp) vs molar concentration plots: (a) NaBPH4, (b) NaI, (c) NaCIO4, (d) NaSCN in acetone solutions. 19. P. Beronius, U. lsacsson and A. M. Nilsson,Acta chem. scand. 24, 189 (1970).
Studies of alkali metal salts
3621
is seen that in the concentration range studied, none of the salts show any evidence of dissociation. In fact, association to form higher aggregates is appreciable for most of the salts, especially at higher concentrations. This trend is obvious from the plots of the apparent mol. wt at each salt concentration vs concentration, some of which are shown in Fig. 5. The greatest degree of association was again exhibited by lithium chloride whose apparent mol. wt approaches that of a dimer.
,oJ 14o
A
--&---O--O------C,--
C ~ O - -
~.D--
19(3 B ~o
)6C
..L)._0,o__0___~.0_--00 -
13G
C
lOG .4~._.0dO__--.-O- O- O- o--O---O---&. D
7G
i
0.04
0-08
0.12
0.16
C,,,
Fig. 5. The apparent molecular weights (M~pp) vs molar concentration plots: (a) biphenyl, (b) Lil, (c) LiBr, (d) LiCI in tetrahydrofuran solutions.
The above observations are consistent with the far i.r. data of Edgell et al. [2] which indicate anion-cation association either in the form of contact ion pairs or as larger ionic aggregates. Similarly the conductance data of Edgell et al. [20] on solutions of sodium tetracarbonylcobaltate in tetrahydrofuran indicate formation of ion pairs even in dilute solution and the formation of ionic clusters, triplets and quadruplets, in the more concentrated solutions. A comparison of the Maoo/Mreal ratios of various salts in acetone and in tetrahydrofuran shows the greater dissociation power of the former solvent. Under identical conditions, solvent-separated ion pairs are more likely to dissociate than the contact ion pairs or the solvent-shared ion pairs. Thus, for the alkali metal salts studied, with few exceptions (notably the chloride and bromide salts of lithium), solvent-separated ion pairs should be the predominant species in acetone solutions, and the anion effect on the far i.r. cation-acetone vibrational frequencies is minimal. This is not the case for tetrahydrofuran solutions. The much weaker dissociating power of this solvent results in the predominant forma20. W. F. Edgell, M. T. Yang and N. Koizumi,J. Am. chem. Soc. 87, 2563 (1965).
3622
M . K . W O N G and A. I. P O P O V
tion of contact ion pairs. Thus the frequencies of the far i.r. vibrational bands are anion dependent. For many of the salts, the attraction between the cation and anion is so strong that at higher concentrations, aggregates higher than ion pairs will form to an appreciable degree. Acknowledgement- The authors gratefully acknowledge the support of this work by a grant from the National Science Foundation.