DISSOCIATION OF TETRAETHYLAMMONIUM CHLORIDE IN ETHYLENE DICHLORIDE
87
Dissociation Constant and Degree of Dissociation for Tetraethylammonium Chloride in Ethylene Dichloride at 0, -15, and -30"
by David L. Lydy,l**b V. Alan Mode,lc and Jack G. Kay T h e W i l l i a m Albert Noyes Laboratory, University of Illinois, Urbana, Illinois
(Received M a y 16, 1964)
The conductivities of millimolar solutions of (CzHb)qNC1in 1,2-dichloroethane have been measured and the degree of dissociation of the solute has been calculated. The results indicate that for 6 X lop3 M solutions, the solute is approximately 27% dissociated a t O o , 30y0 at -15O, and 43% a t -30'.
I. Introduction In a series of papers, Kraus and eo-workers have reported the conductance of a number of electrolytes in ethylene dichloride a t 25" .z In conjunction with some isotopic exchange studies, we have needed to determine the free halide ion concentration for solutions of tetraethylammonium chloride in ethylene dichloride over a range of temperatures. We have obtained new conductivity data from which the degree of dissociation of (CzHJ4NC1 in ethylene dichloride at 0, -15, and -30" has been calculated. These data and calculations are reported herein.
11. Experimental Materials. Fisher Certified reagent grade ethylene dichloride (1,2-dichloroethane) (Karl Fischer water analysis 0.01%) was treated by repeated mixing with activated alumina, mixing with calcium hydride, and distilling from calcium hydride through a 1.2-m. column packed with glass helices. Samples were stored in stoppered Pyrex flasks in a desiccator containing anhydrous calcium chloride. Gas chromatographic analysis indicated that the water concentration was less than mole/l. Eastman White Label tetraethylammonium chloride was dried urlder vacuum with pzos, A typical analysis follows. Anal. Calcd. : C, 57.97; H, 12.16. Found: C, 57.63; H, 12.14. Apparatus cmd Proced.zLre. A conductivity cell designed - for low specific conductivity work and a Model RC-16B-2 Industrial Instruments Inc. conductivity bridge were used for the measurements. The cell was prepared by standard methods and cell
constants were determined a t 0 and 25" during the course of the study. Cell constants a t - 15 and -30" were obtained by a linear extrapolation of the specific conductance for a standard KCl solution with a correction for water conductance. The resulting cell constants were nearly unchanged over the teniperature range from 25 to - 30") confirming the coiiclusions of Robinson arid Stokes3for the type of cell used. The samples of tetraethylammonium chloride in ethylene dichloride were prepared by dilution of a stock solution, taking care to exclude moisture. Measurements a t each concentration were repeated until constant values were obtained for the resistance. Temperature control was *0.1".
111. Results The method of Fuoss and Shedlovsky4 was employed to obtain the limiting conductance (-io) and dissociation constant ( K ) . Following substitution of the expression for CY, the degree of dissociation, into the mass action equation, it was possible to obtain
(1) (a) This report is based on a portion of a thesis submitted by D. L. Lydy in partial fulfillment of the requirements for the degree of Doctor of Philosophy in the Department of Chemistry and Chemlcal Engineering a t the University of Illinois, June 1963, (b) Procter and Gamble Predoctoral Fellow, 1961-1962; University of Illinols Fellow, 1962-1963; (c) University of Illinois Fellow, 1962-1964. (2) W. E. Thompson and C. A. Kraus, J . Am. Chem. SOC.,69, 1016 (1947), and associated references. (3) R. A. Robinson and R. H. Stokes, "Electrolyte Solutions," 2nd E d . , Butterworth and Co., Ltd., London, 1959, pp. 97-99. (4) R . M . Fuoss and T. Shedlovsky, J . Am. Chem. S O C . ,71, 1496 (1949).
V o l u m e 69, Number 1
Januarit 196.5
D. L. LYDY,V. A. MODE,AND J. G. KAY
88
-Ao + CAS(z)f*' KAo2
- 1- - 1
AS(z)
I
where A is the equivalent conductance, C is the total amount of tetraethylammonium chloride per liter of solvent, and f* is the mean activity coefficient determined with the Debye-Huckel limiting law. In order to use tabulated values for the complex function S(z),5 it was necessary to determine the viscosity (v) and dielectric constant (E) of the pure solvent a t the temperatures of interest. Viscosity values a t 0, -15, and -30" were calculated using the equation log v
(i-> 532
=
I
I
I
I
I
IO
12
0.8
0.6
- 3.891
The dielectric constants, determined were 11.86, 13.13, and 14.39 for 0, -15, and -30°, respectively . Table I contains the values of A. and K for the temperatures of interest.
o'21 0.0 0
Table I : Limiting Conductance and Dissociation Constant for (C2H&NC1 in Ethylene Dichloride over the Temperature Range 25 to -30" Temp.,
1
IF
(1)
OC.
Limiting conductance, ohms-1 om.'
Diasociation constant, K x 10'
25" 0 - 15 -30
77.4 49.3 f 1 . 5 3 8 . 0 f 1.1 27.8 f 0 . 5
0.510 0.958 f 0.043 1.32 f 0.05 2.59 f 0.07
2
4 6 8 [(C~HS)~ NCl] X IO'
Figure 1. Degree of dissociation of (C2H6),rU'C1in ethylene dichloride over a range of solute concentrations.
developed an equation from which it is possible to a t which this minimum occurs. calculate the value of CYC As shown in Table 11, our observed minima differ
'See ref. 2.
Table I1 : Calculated and Observed Minima in the Degree of Dissociation of ( C2H6)4NC1in Ethylene Dichloride
By substitution of the mean activity coefficient in terms of CY into the mass action equation, a function was obtained which was solved for CY by successive approximations
where
A =
-8.395 X lo6 (eT)"S
7 -
Temp., OC.
0 - 15 - 30
ac x
10'
Obsd.
Calcd.
1.92 2.19 2.44
10.21 11.70 12.87
from the calculated minima by a factor of approximately 5.3. When compared with other work in this field, our results are in remarkably good agreement with the calculated values. It has been necessary in all calculations to use a
Values for CY are shown in Figure 1. IV. Discussion The appearance of a minimum in the dissociation curve is not unexpected for ionic salts in nonaqueous solvents of small dielectric strength. Daviesg has discussed the factors which produce a minimum and has The Journal of Physical Chemi8try
(5) H. .M.Daggett, Jr., J . A m . Chem. Soc., 73, 4977 (1951). ( 6 ) C. P. Smyth, R. W. Dornte, and E. B. Wilson, Jr., ibid., 53, 4242 (1931). (7) >I. Yasumi and M.Shirai, J . Chem. Phys., 20, 1325 (1950). (8) A. H. White and S.0. Morgan, ibid., 5 , 655 (1937). (9) C. W. Davies, "Ion Association," Butterworth. Inc., Washington, D . C., 1962, pp. 105-116.
ETHER-BORON HALIDEADDITIONCOMPOUNDS IN DICHLOROMETHANE
macroscopic dielectric constant. In a solvent such as ethylene dichloride, there appears to be a significant difference between the macroscopic dielectric constant and the effective LLmicroscopic”dielectric constant which exists near the ionic and molecular species in solution. Inami, Bodenseh, and Ramsey’o have indicated that the effective dielectric constant of ethylene dichloride may be as large as 12.4 in solutions of n(C4He)4NC104 a t 25” (eniscro = 10.232). Glueckauf’s’l corrections of the conventional DebyeHuckel expression t o include short-range changes in
89
the dielectric strength are insignificant a t the temperatures and concentrations used in this study. As both A and K in eq. 3 are functions of the dielectric constant, our reported degree of dissociation may be slightly in error. This could account for the variation between the calculated and observed minima in the degree of dissociation. Further work will be necessary to clarify this point. (10) Y.H.Inami, H. K. Bodenseh, and J. B. Ramsey, J . Am. Chem. SOC.,8 3 , 4745 (1961). (11) E.Glueckauf, Trans. Faraday Soc., 60, 776 (1964).
A Nuclear Magnetic Resonance Znvestigation of Ether-Boron Halide
Molecular Addition Compounds in Dichloromethane
by Ernest Gore and Steven S. Danyluk Department of Chemistry, Univereity of Toronto, Toronto 6 , Ontario, Canada
(Received M a y 87, 1964)
A study has been made of the stabilities of a number of ether-boron halide addition compounds in dichloromethane a t 23”. Equilibrium constants were determined for the reaction, RzO BX3 e RzO.BX8, by a least-squares analysis of the chemical shift-concentration curves for these systems. A shift to low field was observed for all of the ether protons on complexing with boron halide. The most marked deshielding (- 1.25 p.p.m.) was noted for the 1 : l diethyl ether-boron trichloride compound and has been attributed to the formation of ethyl ethoxychloroborate, CzHs+C2HSOBCl3-. Boron trichloride was found to be a stronger acceptor than boron trifluoride while the relative donor strengths of the ethers studied decreased in the order: (CzH&O 2 [(CH3)&H],0 > (CH3)zO. The disagreement between the present n.m.r. results and earlier infrared and gas-phase dissociation studies has been attributed to the influence of the solvent on the mean chemical shift of ether-boron halide solutions in dichloromethane.
+
Introduction 14olecular addition compounds formed with group 111-A acceptors such as boron and aluminum halides (A!&) have been widely investigated and recent reviews emphasize the scope and importance of these compounds as intermediates in many organic reactions. 1-5 A variety of techniques including gas-phase dissociation,6 c r y o s ~ o p y , electrical ~*~ conductivity,7~e~10 and infrared
spectroscopyll have been used to establish the structure and stabilities of group I11 addition compounds. Re( I ) W. Gerrard, “The Organic Chemistry of Boron,” Academic press, New N. y., (2) A. V. Topchiev S V. Zavgorodnii, and Ya. M .Pauskin, “Boron Fluoride,” Pergam;n London, 1959.
pres,,
(3) F, G, A, Stone, Chem, R e v , , 58, lol (1958), (4) R.s. Mulliken, J . p h y e . Chem., 5 6 , 801 (1952).
Volume 69, Number 1
January 1966
