NOTES estimate, this number and the interim thermodynamic functions in Table I should be more reliable than those previously available, 2 - 5 , 9 , 26 in the absence of experimental data. It is interesting to note that even assuming equal force constants in As4 and Pd leads to an entropy for As4 about 1.7 cal/deg mol higher than that usually tabulated, supporting the conclusion that the commonly used entropy is too low. (26) K. K. Kelley and E. G. King, U. 8. Bureau of Mines Bulletin No. 592, U. S. Government Printing Office, Washington, D. C., 1961, p 16.
4329 I n order to calculate the concentrations in equivalents per liter the densities of the solutions were assumed to be the same as that of the solvent. All weights were corrected to vacuum. The following data for sulfolane at 30" were used in the calculations: density, 1.2618 g/ml;5 viscosity, 0.0987 P;2 dielectric constant, 43.3;6 conductivity, 2.0 X lo-* ohm-' cm-l. Results and Discussion The equivalent conductances, A, at various concentrations for tetraalkylammonium perchlorates are reported in Table I.
Solvation Numbers of Some Ions in Sulfolane
by Conductance Measurements
Table I : Conductance of Tetramethylammonium,
by Mario Della Monica and Ugo Lamanna
Tetraethylammonium, Tetrapropylammonium, and Tetrabutylammonium Perchlorates in Sulfolane a t 30'
Istituto d i Chimica, Uniaersitct d i B a r i , B a r i , I t a l y (Received M a y 28, 1068)
lo",
ohm-' cmt mol-1
M
M
Me4NClO4,
I n a previous paper,' we reported the limiting equivalent conductance of several ions in pure sulfolane solutions at 30". I n this paper the conductance measurements have been extended to some tetraalkylammonium perchlorate salts in order to provide more precise information concerning the solvation of all of these ions in sulfolane. Experimental Section Materials. Sulfolane (tetramethylene sulfone) supplied by Shell Italiana was purified according to Burwell and Langford.2 Tetraalkylammonium perchlorates were prepared by adding a silver perchlorate solution to tetraallkylammonium iodide solution in water; the addition of silver perchlorate solution was controlled potentiometrically until the equivalence point was obtained. After filtration the solution was evaporated. The perchlorate was recrystallized from conductivity water and dried in vacuo at 30" for 48 hr. Apparatus and Procedure. Resistance measurements were made with a Jones and Josephs bridge (Leeds and Northrup) at a frequency of 2500 cps. The three ,Jones and Bollinger3 type of cells employed had constants of 1.0672, 0.2733, and 0.05699 cm-'. The constant of the first cell was determined by using a 0.01 demal aqueous solution of KC1 at 30" and the Bremner and Thompson e q ~ a t i o n ;the ~ constants of the other two cells were determined by comparison with the first. Resistance ineasurements were carried out in an oilfilled thermostat at 30 =t0.01". The solutions of the perchlorates were obtained by dilution of a weighed amount of stock solution. The preparation of the solutions and the filling of the cells were performed in a drybox.
105~
Ao = 10.994
81.098 133.68 191.35 226.39 290.27 358.00 451.20 613.63 782.17 986.81
a
Et4NC104, = 10.632 131.12 10.132 255.07 9.946 350.30 9.832 439.66 9.744 599.45 9.615 712.97 9.538 926.37 9.406 A0
10,590 10.458 10.362 10.296 10.210 10.110 1o.oog 9.854 9.716 9.575
PrdNC104, Ao = 9.912 65.934 9.597 204.12 9.316 224.70 9.283 336.15 9.161 361.70 9.129 560.31 8.956 654.96 8.889
Aha ohm-' om2 mol-1
Bu4NC104, Ao = 9.486
78.373 160.98 233.55 328.01 489.81 711.77 934.34
9.121 8.964 8.861 8.751 8.601 8.440 8.316
The ohm used here is actually the international ohm.
The limiting conductances ho (Table I) were calculated by extrapolation to zero concentration7 of the function A' defined by the equation (1) M.Della Monica, U. Lamanna, and L. Senatore, J . P h y s . Chem.,
72, 2124 (1968).
(2) R. L. Burwell and C. H. Langford, J . A m e r . Chem. SOC.,8 1 , 3799
(1959). (3) G. Jones and G. M. Bollinger, ibid., 53, 411 (1931). (4) R. W. Bremner and T. G. Thompson, ibid., 59, 2372 (1937). (5) U. Lamanna, 0. Sciacovelli, and L. Jannelli, Gazz. Chim. Ital., 94, 567 (1964). (6) M. Della Monica, U. Lamanna, and L. Jannelli, ibid., 97, 367 (1967). (7) C. Treiner and R. M. Fuoss, J . P h y s . Chem., 69, 2576 (1965). Volume 79,Number 19 November 1968
4330 A’
NOTES
= Aobsd
+X
~ -C
EClog (6231’~)= A0
+ LC
(1) I n this equation X is the Onsager coefficient, E and El‘ are the Fuoss-Onsager coefficients, and L depends on the contact distance an8 The A’ us. c plots are straight lines in each case (see Figure 1); this means that the ionic association, if present, must be very small.’ On the other hand, a certain degree of ion-pair association seems to be very likely, since the values of contact distance a calculated from eq 1 are always smaller than the theoretical values. The values of the cation limiting conductance A0 (Table 11) were obtained by using the corresponding A0 values and the limiting conductance of C104- i0n.l These values increase from tetrabutyl- to tetramethyla,mmoniumion.
Ions in Sulfolane a t 30’“
xa
Me4N+
+
4.31 3.95 3.23 2.80
Et4N + Pr4N + Bu~N f
xo +cqo
rs
rD/TB
0.425 0.390 0.319 0.276
1.93 2.10 2.57 2.97
1.80 1.90 1.76 1.66
ra and rc in
X O + in (international ohm)-l cm* mol-’; qo in
8;
P.
As in the case of tetraalkylammonium salts in nitrob e n ~ e n ethe , ~ A0 increase from ethyl to methyl substituent is lower than the expected value. This small increase of the limiting conductance in sulfolane has been interpreted using the analogy of the behavior of nitrobenzene solutions as an interaction between the tetramethylammonium charge and the dipole of solvent molecules.
’05
2
1
3
4
5
P9,P
Figure 2. ro/rsvalues us. rs for tetraalkylammonium ions in sulfolane (curve l), acetonitrile” (curve 2), and waterlo (curve 3). Curves 2 and 3 are reported as examples of linear and nonlinear dependence of the function ro/rson rs.
On this basis, the xo values of unsolvated EtdnT+, Pr4K+, and B u ~ K +ions have been used to calculate the Stokes ionic radii r s according to the equation
Table I1 : Properties of Tetraalkylammonium
Ion
0
From these r s values, the corresponding Robinson and Stokes correction factors r c / r swere calculated by using the re (crystallographic radii) values reported by these authors.lo The r e / r s vs. r s linearized curve (Figure 2 ) drawn through these experimental points can be correctly utilized for the calculation of the correction factor of the alkali metal ions, since the corresponding rs values are between or near the tetraalkyl ammonium r Svalues. From this curve, the corrected radius of the solvated ions, rcor,is obtained and the volume of the solvation shell surrounding the ions is calculated from the relationship
v=
(4/3)74Tcor3
-
re3)
(3)
An average number of sulfolane molecules involved in the solvodynamic unit has been calculated, by considering the contraction of the solvent next to the ions because of electrostriction to be negligible. I n these coalculations a molecular volume of sulfolane (158 A3) has been assumed. The results are presented in Table 111.11,l2
t
(8) R. 11. Fuoss, L. Onsager, and J. F. Skinner, J . P h y s . Chem., 69, 2581 (1965). (9) E. G. Taylor and C. A. Kraus, J . A m e r . Chem. Soc., 69, 1731
(1947). I O
2
4
6
8
Figure 1. Plots of the A’ function us. normality of the tetraalkylammonium perchlorates. T h e Journal of Physical Chemistry
io
c 103
(10) R. A. Robinson and R. H. Stokes, “Electrolyte Solutions,” Butterworth and Co. Ltd., London, 1959, pp 124-125. (11) A. H. Harkness and H. ,M.Daggett, Can. J . Chem., 43, 1215
(1965). (12) F. Madaule-Aubry, Bull. SOC.Chim. France, 1456 (1966).
NOTES
4331
A comparison among the solvation numbers of alkali metal ions in this and other solvents13 shows that the smallest values for all of the cations are found in sulfolane and in nitrobenzene, solvents which have the greatest molecular volume.
Solubilization Behavior of a Polyoxyethylene Sulfate Type of Surfactant in Connection with the Micellar Charge
by Fumikatsu Tokiwa Table I11 : Solvation Numbers of Ions in Sulfolane a t 30""
Li + Na +
K+ Rb + c s+
NHr
+
c1BrI-
clod-
4 . ;33 3.61 4.05 4.16 4.27 4.97 9.30 8.92 7.22 6.685
1.92 2.30 2.05 2.00 1.95 1.67 0.89 0.93 1.15 1.24
3.74 4.23 3.92 3.85 3.78 3.36 1.98 2.06 2.47 2.64
0.60 0.95 1.33 1.48 1.69 1.48 1.81 1.95 2.16 2.40
1.4 2.0 1.5 1.4 1.3 0.9
... ... ... . ..
r in A. The a XO in (international ohm)-' cm2 mol-'; crystallographic radii ra are taken from ref 10, p 461. The crystallographic radius of the c l o d - ion was reported by MadauleAubry.12
As far as the anions are concerned the extrapolation of the curve of Figure 2 for the calculation of the rc/rs factors seems more arbitrary than in the case of the cations, owing to the greater distance of the values of rs for anions from the tetraalkylammonium ions rS values. Nevertheless, it must be pointed out that the error due t o extrapolation probably has very little influence on the calculated value of the corrected radius which expresses interaction between the anions and the solvent. For C1-, Br-, I-, and Clod- ions this interaction seems to be very small and they may be assumed as practically bare, since the radii obtained by the Robinson and Stokes method are close to the crystallographic ones (Table 111). The absence of solvation of the anions in sulfolane agrees with the views of Parker14 concerning the solvation of anions in dipolar aprotic solvents. I n addition, the association constants K A for LiC1, =~ 13,860 M-l; LiBr, and LiI in sulfolane's ( K A , L I C K A , L i B r = 278 M-'; K A , L < ~ I10 M-l) agree with the view that the anions are unsolvated; in fact, the interaction between Li+ and C1- can be explained, assuming that the anions are bare, in terms of the charge density decrease from (21- to I-.14
(13) R. Gopal and M. M. Husain, J . I n d i a n Chem. Soc., 40, 981 (1963). (14) A. J. Parker, Quart. Rev. (London), 16, 163 (1962). (15) R. Fernandeli-Prini and J. E. Prue, Trans. Faraday Soc., 62, 1257 (1966).
Research Lahoralories, K a o S o a p Company, Wakayama-shi, J a p a n (Received J u n e PI, 1968)
Although the phenomenon of solubilization of surfactants has frequently been discussed in connection with the structure of micelles, it has not been considered from the point of view of their electrical nature. The electrical nature of a micelle is probably one of the important factors which governs solubilization behavior, because the structure of the micelle depends highly on this nature. I n the previous work' studying the solubilization behavior of a polyoxyethylene sulfate type of surfactants, it has been suggested that the potential or charge on the surface of the micelle has an important effect on the extension of the polyoxyethylene chains in the micelle and the solubilization capacity of the micelle. I n the present note, the solubilization behavior of sodium dodecylpolyoxyethylene sulfate^^^^ (SDPS), C12H2j0(CH2CH20),SOs?;a, with different numbers of oxyethylene units ( p ) has been discussed in connection with the [ potential and the charge of the micelle in order to understand the solubilization mechanism of this type of surfactants. Solubilization results of SDPS with p from 0 to 10 in water and solutions of NaCl and CaClz are shown in Figure 1, in which the solubilizing power toward Yellow OB is plotted against the number of oxyethylene unit^.'^^^^ It has been shown from the spectral study of the solubilization of Yellow OB that, in the case of surfactants having a polyoxyethylene chain, solubilization occurs in both the hydrocarbon core of the micelle and the outer shell of the polyoxyethylene chains.' This accounts for the increase in the solubilizing power of SDPS with increasing oxyethylene content (shown in Figure 1). However, the effectiveness of the polyoxyethylene part for solubilization, i.e., the degree of solubilization per oxyethylene unit, decreases with increasing oxyethylene content, as shown in Figure 2 . (The degree of solubilization per oxyethylene unit was calculated by subtracting the solubilizing power of SDPS-0 (sodium dodecyl sulfate) (1) F. Tokiwa, J . P h y s . Chem., 72, 1214 (1968). (2) The samples were the same as those used in a previous work.3 (3) F. Tokiwa and K. Ohki, J . P h y s . Chem., 71, 1343 (1967). (4) Measurements of solubilizing Dowers and electroDhoretic mobilities of micelles were carried o u t using the procedures previously
reported.' 36 (5) F. Tokiwa and K. Ohki, Kollo.id-Z., 223, 38 (1968). V o l u m e 78, Number 18 November 1968
