1575
NOTES
Nov., 1957 ferential manometer filled with antimony entafluoride with an opposing pressure of dried nitrogen. $he nitrogen pressure in turn was determined by a mercury manometer that could be isolated readily from the system. Immediately after reading the mercury levels, the manometer was pumped out in order to prevent extensive reaction of the antimony pentafluoride with the mercury. Numerous vapor pressure measurements were made over the range 65 to 130' using a glass click gage. The amount of inert gas was determined by cooling the sample to a temperature where the antimony pentafluoride has a negligible vapor pressure. After applying the inert gas corrections, the static measurements agreed with those reported by Shair and Schurig.e
Results and Discussion The results of the density measurements are shown in Table I, together with their deviations from the formula p
= 3.193
- 3.67
X 10-*t
+3 X
10-'t2
(1)
which expresses the density of antimony pentafluoride in grams per cubic centimeter over the temperature range -10 t o 70" within f 0.3%. When converted, the two specific gravity values reported in the literature are in close agreement with values computed using eq. l . 0 s 7 TABLEI DENSITIES OF ANTIMONYPENTAFLUORIDE Density (g./cc.) (Ob.)
-9.1 -7.4 -6.8 -5.8 -3.1 6.5 11.0 15.3 26.3 46.0 56.0 59.5 62.7 71.8
Obsd.
Calcd.
3.230 3.224 3.220 3.216 3.206 3.172 3.155 3.140 3.102 3.040 3.004 2.992 2.980 2.952
3.227 3.221 3.218 3.215 3.205 3.169 3:153 3.137 3.099 3.031 2.997 2.985 2.975 2.945
Dev. X 108
3 3 2 1 1
3 2 3 3 9 7 7
Density (g./oc.) ("C.)
Obsd.
Dev. Calcd. X 10s
-8.1 -1.7 8.5 13.3 14.8 22.8 23.4 25.8 33.7 48.5 54.2 63.9
3.221 3.198 3.160 3.142 3.138 3.109 3.109 3.097 3.071 3.023 3.002 2.970
3.223 3.199 3.162 3.145 3.139 3.111 3.109 3.100 3.073 3.022 3.003 2.971
10.3 17.6 18.2 24.8 25.8
3.155 3.128 3.125 3.105 3.105
3.155 3.129 3.127 3.104 3.110
t
-2 -3 -1 -2 0 -3 -2 1 -1 -1
5
7
0 -1 -2 1 5
*
The melting point of 8.3 0.3" is slightly higher than the previously reported value of 7O.8 Experimental values of the vapor pressure from 9 to 50" are listed in Ta.ble I1 and plotted in Fig. 1 together with values previously reported.6 They can be represented by the equation loglo p(mm.) = 8.567
- (2364.1/T)
-.
-2 -1
(2)
which Shair and Schurig found valid over the temperature range 50 to 143". In the lower temperature range, eq. 2 is only precise t o f 15% due t o the difficultiesin measuring the low pressures. The high value, about 25, of Trouton's constant and the viscosity of the liquid are indicators of the high degree of association. (6) R. C. Shair and W. F. Schurig, Ind. Eng. Chem., 7 8 , 1624 (1951). (7) 0. Ruff and W. Plato, Ber., 39, 673 (1904). (8) 0. Ruff, J. Zedner, M. Knoch and H. Graf, Ber., 42, 4021 (1909).
Fig. 1.-Vapor
pressures of SbFh.
TABLE I1 VAPORPRESSURE MEASUREMENTS t ("C.)
Preseure (mm.) Obsd.
Calcd.
Dev.
48.5 40.0 29.7 19.9 9.6 29.9 19.6 19.6 9.8 9.6 8.7 44.0 34.9 34.3 25.6 24.8
14.0 10.0 5.8 3.2 1.5 5.5 3.4 3.3 1.6 1.5 2.0 12.5 7.3 7.0 5.0 4.6
16.5 10.4 5.8 3.2 1.6 5.8 3.1 3.1 1.6 1.6 1.5 12.9 7.8 7.6 4.5 4.3
-2.5 -0.4
-
.o .o .1 .3 .3 .2
.o
-
.1 .5
-
.4
.5 .6 .5 .3
THE DECOMPOSITION OF MALONIC ACID I N N-SUBSTITUTED ANILINE DERIVATIVES AND I N DIETHYLCYCLOHEXYLAMINE BYLOUIS WATTSCLARK Contribution from the Departvent of Chemistry, Saint Joseph CoEEege, Emmztsburg, Maruland Received Mag $1, 1967
Fundamental studies on the decayboxylation of malonic acid and its derivatives, the basis of the important malonic ester synthesis of substituted
1576
Vol. 61
NOTES
parable with that of aniline it has been shown that malonic acid does not ionize appreciably,6 therefore 1.6 in such media the undissociated acid must suffer cleavage. When stronger bases are present,, e.g., N-ethylpiperidine, the reaction is slowed downlm5 one carbonyl group is neut,ralized, and the acid malonate ion decarboxylates.* In still stronger 1.4 base, where both carboxyl hydrogens are neutralised, decarboxylation fails to take place.' 1.3A I n order to gain further insight into the effect of N structure and relative basicity on the reaction, kiI netic studies were carried out in this Laboratory on 1*2.5 the decomposition of malonic acid in N-methyl3 aniline, N,N-dimethylaniline, N,N-diethylaniline 1.1 and N,N-diethylcyclohexylamine. The results of this investigation are reported herein.
36 32 28
ea 24 v1 .la
g20 V u
916 12
1.0
Experimental Reagents.-(1) Malonic acid, 100.0% Assay; (2) Nmethylaniline, Rea ent Grade, b.p. 81-82' (14 mm.), 8 sp. gr. 0.989; (3) %,N-dimethylaniline, Reagent Grade, b.p. 193', sp. gr. 0.956; (4) N,N-diethylaniline, Reagent 4 o.8 Grade, b.p. 93-95' (15 mm.); (5) N,N-diethylcyclohexylamine, reagent grade, b.p. 73-75' (12 mm.). Apparatus and Technique.-The apparatus and technique have been described in a previous paper in this series.a I n these experiments 0.1857-g. samples of malonic acid 2 4 6 8 10 12 14 (the amount required to furnish 40.0 ml. of COa a t S.T.P. Minutes. on complete reaction) were introduced in the usual manner Fig. 1.-Decom osition of 0.1857 g. of malonic acid in 50 into 50 ml. of the amine in a ioo-ml. flask, the solvent being ml. of N,N-diethyfaniline a t 128.0' (cor.): I, ml. Of Con a t previously saturated with dry Coz. S.T.P.; 11, log ( U 2). Oa9
-
Results and Discussion
For the decarboxylation of malonic acid in N-substituted aniline derivatives as well as in N,N-diethylcyclohexylamine the maximum yield of COZ obtained in any experiment did not exceed 96% of the theoretical yield. This is in contrast with the decomposition of malonic acid in aniline which gave a 100% yield of COz in each e~periment.~
36 32
1
28
4
1.4 1.3 2 N
e 24 v1
I 1.2 5
0
cd
0" 20 u u 9 16
TABLE I FIRST-ORDER RATECONSTANTS FOR THE DECARBOXYLATION OF MALONIC ACID IN SEVERAL SECONDARY AND TERTIARY AMINESAT VARIOUS TEMPERATURES
bo
1.13
s 12
Medium
N-Methylaniline
-j1.0
I
I ' I
2
4
I
I
y:: I
1
6 8 1 0 1 2 Minutes. Fig. 2.-Decomposition of 0.1857 g. of malonic acid in 50 ml. of dieth lcyclohexylamine a t 123.5' (cor.): I, ml. of CO* a t s.T.~?;11, log ( a 2).
-
acetic acids, are prompted by the desirability of improving the synthetic possibilities of the reaction as well as throwing further light on the mechanism involved. Studies have been previously reported on the decomposition of malonic acid in a number of aromatic amines, as well as in other non-aqueous I n solvents whose basicities are com-
N, N-Dimethylaniline
N,N-Diethylaniline
N,N-Diethylcyclohexylamine
Temp, ("C.)
8 ecifio reaction ve?ooity oonstant (see. -9
121.7 125.5 131.5 138.1 121.2 128.0 135.8 138.8 123.3 128.05 136.2 115.52 120.2 123.5 129.45
0.000932 ,0013 .00218 ,00366 0.000808 .001453 .00278 .00354 0.001143 .001688 .00329 0.00111 .00167 .00221 .00364
(1) G. Fraenkel, R. L. Belford and P. E. Yankwioh, J . Am. Chsm. Soc., 76, 15 (1954).
L. W. Clark, THISJOURNAC, 60, 825 (1956). L. W. Clark, ibid., 60, 1150 (1956). L. W. Clark, ibid., 60, 1340 (1058). L. W. Clark, dbid., 60, 1583 (1956). (6) E. J. Corey, J . Am. Chem. Soc., 76, 1172 (1953). (7) G. A. Hall, Jr., ibid., 71, 2691 (1949). (2) (3) (4) (5)
h c
7
1577
NOTES
Nov., 1957
KINETIC DATAFOR
THE
TABLE I1 DECOMPOSITION OF MALONIC ACIDA N D OXALICACIDIN VARIOUSAMINES
System
+ + + + + +
(1) Malonic acid N-methylaniline (2) Malonic acid N,N-dimethylaniline (3) Malonic acid N,N-diethylaniline (4) Malonic acid N,N-diethylcyclohexylamine (5) Oxalic acid N-methylanilines (6) Oxalic acid N,N-dimethylanilines
E*
A
AS
AHS
106
28,800
486
26,200
-6.555
28,900
418
25,400
-8.24
28,800
452 870
27,400
26,600
27,000
8.35 X 10"
26,200
3.6
x
10"
k14Q'
-5.325
1 . 5 4 X 1012
-1)
x
(sea.
(e.J
(aea.
~ h d g (cal.)
(cal.)
(Gal.)
-1)
26,700
12.27 X 101l
25,900
-5.7
28,260
35,850
5.76 X 10"
35,500
+8.25
32,100
5.8
33,100
3 . 3 9 X loL3
32,400
+1.25
31,900
9.5
Data from typical results are shown graphically in Figs. 1and 2. T t will be seen from line I1 in Figs. 1 and 2 that plots of log (a- 2) os. time yield straight lines indicating the reaction to be first order in each case. Table I brings together the rate constants in each solvent for the various temperatures studied, calculated from the slopes of the logarithmic plots. Figure 3 shows the Arrhenius plot for three of the four solvents used. Table I1 lists the kinetic data calculated from the Arrhenius and Eyring equations, with data for oxalic acid included for comparison.8 Fraenkel and co-workers' obtained a value for k of 0.00014 sec.-l at 100.0" for the decomposition of malonic acid in dimethylaniline. Using the value of E* and A in line (2) of Table 11,and substituting in the Arrhenius equation, the rate constant at 100.0' is found by calculation to be 0.00015 sec.-I. a check of this sort obtained by two different investigators using different techniques and reagents from different sources inspires confidence in the validity of the results. The hypothesis of Fraenkel and co-workers' that malonic acid forms an unstable intermediate compound via the unshared pair of electrons on the nitrogen atom of the amine is further substantiated by the results shown in Table 11. Methylaniline would be expected to have a smaller steric effect on the formation of the activated complex than would dimethylaniline, consequently a higher entropy of activation. Further, since two methyl groups on nitrogen would have twice the positive inductive effect as one methyl group, dimethylaniline should be more basic than methylaniline, and the enthalpy of activation should be lower in the more basic solvent. Examination of the data in lines 1 and 2 of Table I1 reveals that these predictions are actually realized. One would expect a greater steric effect in the case of two ethyl groups on nitrogen than in that of two methyl groups. The solvent with the greater steric effect should have the lower entropy of activation. This expectation is also realized, as will be seen by comparing lines 2 and 3 of Table 11. The decrease in the enthalpy of activation of diethylaniline over dimethylaniline compensates for the dif(8) L. w.Clark, THISJOURNAL, 61, 099 (1957).
-3*2 -3.1
.
t 1 244
I
246
248 250 252 254 256 1/T X 106. Fig. 3.-Decomposition of malonic acid in various amines, Arrhenius plots: I, malonic acid in N,N-dimethylaniline; 11, malonic acid in N,N-diethylaniline; 111, malonic acid in N,N-diethylcyclohexylamine.
ference in entropy so the reaction takes place a t very nearly the same rate in both solvents. The enthalpy of activation in diethylcyclohexylamine is very nearly the same as in diethylaniline (lines 3 and 4 of Table 11). This would imply that the basicities of these solvents are nearly equal. Since no resonance obtains in diethylcyclohexylamine the C-N bond is greater than in diethylaniline, and in the former rotation about the C-N bond is possible. These two circumstances could account for the entropy of activation being lower in the aromatic than in the aliphatic amine. Because of the more favorable entropy factor malonic acid decomposes approximately twice as fast in diethylcyclohexylamine as in diethylaniline. The suggestion has been made that the decomposition of oxalic acid in non-aqueous solvents probably proceeds by a mechanism quite similar to
1578
NOTES
bhat of malonic acid.8 Evidence for the validity of this assumption is found by comparing lines 1, 2, 5 and 6 of Table 11. For the decomposition of oxalic acid in methylaniline and in dimethylaniline the differences in the entropy of activation and the enthalpy of activation exactly parallel those for the decomposition of malonic acid in these game solvents. Such a parallelism strongly suggests that the mechanism is the same for both acids. Acknowledgment.-The financial support of this research by the Raskob Foundation for Catholic Activities, Wilmington, Delaware, is gratefully acknowledged. SOLUBILITY O F LAURYL ALCOHOL I N AQUEOUS SOLUTIONS OF SODIUM LAURYL SULFATE BY M. B. EPSTEIN AND J. Ross Colgate Palmolive Company, Jersey City, New Jersey Received April 69,1967
Vol. 61
viously thicker. Below the observed break the bands are uniform and the films are clear and undistorted and the transition temperatures are presumed to correspond to homogeneous solutions unsaturated with respect to LOH. If this view is taken, then a solubility curve may be constructed. It will be a curve of solubility of LOH at maximum transition temperature versus concentration of SLS. The values above the CMC are essentially at one temperature, 33", but for the values below the CMC there is a different temperature for each concentration of SLS. I n Fig. 1, it is seen that above the CMC a line typical of solubility isotherms for similar systems is obtained. Below the CMC the slope is far less than that over the micellar region. An accurate value for the solubility of LOH in water is not available, but it must be exceedingly small. Sporck2 extrapolated solubilities for the lower alcohols and gave an estimate of 0.00019% mole/l.). Addison and Hutchinsona by a similar approach derived an equation for the solubility of the fatty alcohols in water which gives the value 0.000370for LOH. It is apparent from Fig. 1 that the solubility of LOH in water is markedly enhanced by the presence of SLS below as well as above the CMC.
It appears possible to obtain quantitative data for the solubility of lauryl alcohol (LOH) in aqueous solutions of sodium lauryl sulfate (SLS) from film drainage transition temperature data. When plotted as tt function of the amount of LOH in the sys(2) C. R. Sporck, J . Amer. Oil Chem. SOC.,$0, 190 (1953). tem a t any particular concentration of SLS, these (3) C. C. Addison and 8. K. Hutchinson, J . Chem. Boc., 3387 transition temperatures rise to a limit and then do (1949). not increase much upon addition of even large amounts of LOH. Below the critical micelle concentration (CMC) this effect is apparent from a NATURE OF THE METHANOL-METHYL direct plot of transition temperatures against the BORATE AZEOTROPE1 concentration of LOH at constant SLS. Above BY THOMAS J. TULLY~ AND PHOEBUS M. CHRIS TOP HER^ the CMC, it is preferable to use a function logarithmic in LOH. Contribution from the Departmen2 of Chemistry Newark College of Engineering, Newark 6, N . ' J . The observed break we interpret as being due to Received June 10, 1967 the limit of solubility of the LOH at the given concentration of SLS. Indeed it is observed that beMethanol and methyl borate form an azeotrope yond this break the films are visibly different. The that boils between 53 and 55". This is lower than interference colors are not always horizontal but the boiling point of any other mixture of the two often distorted by the superposition of local areas compounds. Due to the fact that the composition often showing a large number of bands and ob- of the azeotrope is very close to an equimolar ratio of the two components, it is generally believed that solutions of methyl borate in methanol contain a complex alkoxo acid formed by the reaction
The tetraalkoxo acid itself has not been isolated but stable salts of the acid have been prepared and characterized. 3-6 Mixtures of the two components give an acidic color with brom thymol green, showing that some reaction must occur. The acidity must be due to the formation of the methyloxonium ion, CH30H2+, since neither component by itself produces an acid color with the indicator. The present work comprises a study of the prop1.0 1.5 2.0 SLS, g./100 g. soh. of LOH in SLS from surface transition temperatures.
0.5
Fig. 1.-Solubility
( 1 ) M. B. Epstein, A. Wilson, C. W. Jakob. L. E. Conroy and J. Ross, THISJOURNAL, 68, 8GO (1954).
(1) Presented before the Physical Chemistry group of the North Jersey Section of the American Chemical Society Meeting-in-Miniature, January 30, 1956. (2) Department of Chemistry, Newark College of Engineering, Newark 2, N. J. (3) H.Cooaux. Comnt. rend.. 127. 719 (1898). (4j L. Cambi, Atti n'sndi.. 2S, I, 244 (1914): (5) H.Meerwein, Ann., 466, 227 (1927). (6) 11. Meerwein and T. Bersin, ibid., 476, 113 (1929).
