DECARBOXYLATION OF BENZYLMALONIC ACIDIN POLAR SOLVENTS
greater than 2. Part of the excess might be attributed to (FCOZ)CFz(FCO), but more likely it reflects an error in the CZF30I yields caused by a low calibration factor, as discussed earlier.
627
Acknowledgments. The author wishes to thank Mr. Ray Calloway for mass spectrometric analysis, Mr. Dennis Saunders for preparation of CzF4, and Mrs. Barbara Peer for assistance with the manuscript,
Further Studies on the Decarboxylation of Benzylmalonic Acid in Polar Solvents
by Louis Watts Clark Department of Chemistry, Weatern Carolina College, Cullowhee, North Carolina
(Received June 18, 1966)
The decarboxylation of benzylmalonic acid was studied in seven polar solvents: aniline, N-ethylaniline, N-sec-butylaniline, o-toluidine, N,N-dimethylaniline, quinoline, and 8methylquinoline. Rate constants and activation parameters were obtained and compared with results obtained previously for the reaction alone and in four additional solvents. An interesting parallelism between benzylmalonic acid and malonanilic acid was observed.
Kinetic data on the decarboxylation of benzylmalonic acid in the molten state,l in two of the fatty acids and two of the cresols,2have been reported. An enthalpyentropy of activation plot for the reaction series resulted in two parallel lines-one for the reaction in acids, the other for the reaction in the cresols-each having a slope of approximately 394°K or 121°C. Such a slope is known as the isokinetic temperature of the reaction series3since it corresponds to the temperature at which the rate constants of all the reactions conforming to the line are equal. (It is interesting to note that benzylmalonic acid melts at 121O.) Subsequent studies have shown4 that the isokinetic temperature for the decarboxylation of a large number of acids including malonic acid and many of its derivatives in all sorts of polar solvents is 422°K or 149°C. If two or more related reactions have the same isokinetic temperature, it is assumed that the different reactions all take place by the same me~hanism.~The fact that the decarboxylation of benzylmalonic acid in acids and cresols possesses a different isokinetic temperature from that of malonic acid suggested that, apparently, different mechanisms were involved in the two reactions. Subsequently, kinetic data on the decarboxylation of malonanilic acid in a large number of polar solvents
revealed that the mechanism of the malonanilic acid reactions was different from that of malonic acid.6 In order to try to obtain further insight into this aspect of the decarboxylation reaction, additional kinetic experiments were carried out in t,his laboratory on the decarboxylation of benzylmalonic acid in seven additional polar solvents, namely, aniline, N-ethylaniline, N-sec-butylaniline, o-toluidine, N,N-dimethylaniline, quinoline, and 8-methylquinoline. The results of this study are reported herein.
Experimental Section The benzylmalonic acid used in this research assayed 100.0% pure by titration with standard base using a Beckman Model H-2 glass electrode pH meter. The melting point of the benzylmalonic acid was 121” (cor). The solvents were reagent grade and were (1) L.W.Clark, J . Phya. Chem., 6 7 , 138 (1963). (2) L. W.Clark, ibid., 67, 1481 (1963). (3) 6 . L. Friess, E. S. Lewis, and A. Weissberger, Ed., “Technique of Organic Chemistry,” Vol. VIII, Part I, 2nd ed, Interscience Publishers, Inc., New York, N. Y., 1961, p 207. (4) L. W.Clark, J . Phys. Chem., 68, 3048 (1964). (5) J. E.Leffler, J. Org. Chem., 20, 1202 (1955). (6) L. W.Clark, J . Phya. Chem,, 68, 2150 (1964).
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distilled at atmospheric pressure immediately before use. The apparatus and technique have been described previously.' The course of the reaction was followed by measuring the volume of C 0 2 evolved at atmospheric pressure and at the temperature of a waterjacketed buret. The buret was calibrated by the U. S. Bureau of Standards at 20". Water maintained at 20.0 f 0.05" by means of a cooling coil and an electronic relay was pumped through the water jacket during the experiment. The temperature of the oil bath was controlled to within 0.005" using a completely transistorized temperature control unit equipped with a sensitive thermistor probe. A thermometer which also had been calibrated by the U. S. Bureau of Standards was used to read the temperature of the oil bath. The accuracy of the barometer used in this research was ensured by completely diassembling the barometer, replacing the barometer tube with a new clean tube of larger diameter, and refilling the tube with triply distilled mercury. In each decarboxylation experiment a 0.3489-g sample of benzylmalonic acid was introduced in the usual manner into the reaction flask. On complete reaction, this weight of acid will produce 40.0 ml of C02 at STP, calculated on the basis of the actual molar volume of C02 at STP, namely, 22,267 ml. About 60 g of solvent, saturated with dry C02 gas, was used in each experiment.
Table I : Apparent First-Order Rate Constants for the Decarboxylation of Benzylmalonic Acid in Several Solvents Solvent
Aniline
o-Toluidine
N-Ethylaniline
N-see-But ylaniline
N,N-Dimethylaniline
Quinoline
8-Methylquinoline
k
=
KT -,-AH*/RT h
eAS*/R
k X 104,
96.74 106.28 116.28 96.51 106.90 116,23 95.47 103.13 110.13 115.64 92.88 103.08 113.59 96.51 105.78 116.28 90.49 98.44 110.04 103.50 110.04 120.00
2.91 5.65 11.4 2.01 6.65 16.8 2.84 5.35 9.31 14.25 2.23 4.91 13.35 1.56 5.73 23.3 3.74 6.88 16.82 3.70 6.95 17.1
sec - 1
Table I1 : Activation Parameters for the Decarboxylation of Benzylmalonic Acid and Malonanilic Acid in Several Solvents
Results Two decarboxylation experiments were carried out in each solvent at each of three different temperatures over a 20" range. The decarboxylation of benzylmalonic acid gave smooth first-order kinetics in all of the solvents used in this research over the greater portion of the reaction in each case. I n each experiment the log ( V , - Vl) was a linear function of time over more than 70% of the reaction. Rate constants were calculated from the slopes of the logarithmic plots, reproducibility between duplicate experiments being generally 1-3%. Average rate constants thus obtained are shown in Table I. The parameters of the absolute reaction rate equations
Temp, (cor)
O C
Benzylmalonic ---acidSolvent
Ani1ine N-E thylaniline N-see-Butylaniline o-Toluidine N,N-Dimethylaniline Meltb n-Butyric acid" Decanoic acidc m-Cresol" p-Cresol' o-Cresol Quin o1ine 8-Methylquinoline 'See ref 6.
See ref 1.
Malonanilio --acida-
AH*,
AS*,
AH*,
AS/*,
kcal/ mole
eu/ mole
kcal/ mole
eu/ mole
19.8 21.9 26.56 29.9 38.4 29.4 23.0 26.9 22.0 27.1
-21.64 -15.8 -3.6 5,0 27.4 -2.6 -18.9 -9.0 -25.2 -14.2
27.6 31.9
-1.5 10.0
19.9 26.4
-19.94 -4.6
33.2 34.0 36.5 21.0 28.5
9.4 11.54 15.5 -17.5 0.4
See ref 2.
based upon the data in Table I are shown in Table 11, along with corresponding data for malonanilic acid where available.
decarboxylation of both benzylmalonic acid and malonanilic acid increases as the steric hindrance and
Discussion It will be observed in Table I1 that the enthalpy of activation as well as the entropy of activation for the
(7) L. W. Clark, J. Phys. Chem., 60, 1150 (1956). (8) 5. Glasstone, K. J. Laidler, and H. Eyring, "The Theory of Rate Processes," McGraw-Hill Book Co., Inc., New York, N. Y., 1941, p 14.
The Journal of Physical Chemistry
DECARBOXYLATION OF BENZYLMALONIC ACIDIN POLAR SOLVENTS
629
Table 111: Comparison of Activation Parameters for the Decarboxylation of Several Acids in Various Amines" Y alonic acidb Solvent
Aniline o-Toluidine Quinoline 8-Methylquinoline
AH*
AS*
26.9 -4.5 25.7 -7.0 26.7 - 2 . 4 24.4 -10.5
Cinnamalmalonic acidC AH*
AS*
23.8 21.9 23.5 21.6
-13.2 -17.5 -16.2 -21.8
Oxamic acide
Oxalic acidd AH*
AS*
38.9 15.8 37.7 13.7
Trichloroacetic acid* AH* AS*
Oxanilic acidi
AH*
AS*
AH*
AS*
59.7 53.7 47.0 36.0
68.0 57.1 37.5 12.2
49.8 47.8 38.6 35.6
46.3 39.9 16.0 10.0
'
24.5 23.8 24.0 22.3
-2.6 -6.8 -2.4 -8.4
8-Resorcyiic acidh AH*
AS*
34.5 5.95 2 2 . 9 -21.8
a Units: AH*, kcal/mole; AS*, eu/mole. L. W. Clark, J. Phys. Chem., 62, 79, 500 (1958). L. W. Clark, ibid., 66,836 (1962). L. W. Clark, ibid., 61,699 (1957); 62,633 (1958). L. W. Clark, ibid., 65, 180,659 (1961). L. W. Clark, ibid., 65,572,1460 (1961). L. W. Clark, ibid., 67, 2831 (1963). L. W. Clark, ibid., 63, 99 (1959).
nucleophilicity of the various solvents increase. This is exactly the reverse of the trend observed in the decarboxylation of the seven compounds listed in Table I11 which have been studied previously (malonic acid, cinnamalmalonic acid, oxalic acid, oxamic acid, oxanilic acid, trichloroacetic acid, and P-resorcylic acid). If we confine our attention to corresponding solvents in the nine different reactions in question, it will be observed that a methyl group introduced in a position near the nitrogen function in either aniline or quinoline has qualitatively the same effect on the activation parameters in the seven decarboxylation reactions shown in Table 111. However, in the decarboxylation of benzylmalonic acid and malonanilic acid (Table 11),the effect is large and in the opposite direction. These results indicate that, in the seven examples shown in Table 111, there is an effect due to the methyl group, possibly electronic but more probably steric. I11 the case of the decarboxylation of benzylmalonic acid and nialonanilic acid, on the other hand, the effect of the methyl group in aniline and quinoline is qualitatively different. It need not, however, be the same effect with opposite sign as one might at first be tempted to postulate. This reversal of the effect in the decarboxylation of benzylmalonic acid and malonanilic acid may be due in part to the fact that these two compounds (but none of the other seven) have aromatic substituents and a part structure Ph-X-C=C. A plot of enthalpy of activation vs. entropy of activation for the decarboxylation of benzylmalonic acid and malonanilic acid in the aniline series of amines, based upon the data in Table 11, is shown in Figure 1. Similar plots for the decarboxylation of these two acids in the cresols, and in quinoline and %methylquinoline, respectively, are shown in Figures 2 and 3. The white circles represent the data for benzylmalonic acid, the black circles those for nialonanilic acid. The slope of the line in Figure 1 (the isokinetic tem-
'
- 20
- 10
0
10
20
AS*, eu/mole.
Figure 1. Enthalpy of activation vs. entropy of activation plot for the decarboxylation of benzylmalonic acid and malonanilic acid in aniline and its derivatives based on data in Table 11. Reactants: 0, benzylmalonic acid; 0 , malonanilic acid. Slope of line = 378°K or 105°C.
perature of the reaction series) is 378°K or 105°C. The slope of each of the two parallel lines in Figure 2 is 393°K or 120°C. The slope of the line in Figure 3 is 423°K or 150°C. In the earlier studies on the decarboxylation of benzylmalonic acid in monocarboxylic acids and in cresols,2 it was suggested that there might be a correlation between the isokinetic temperature and the melting point of the reactant. This feeling was strengthened by the observation that the isokinetic temperature for the decarboxylation of oxanilic acid in a large variety of polar solvents (lZ0°)9 was also the same as its melting point. The present results indicate, however, that benzylmalonic acid does not have one but a multiplicity of isokinetic temperatures, one of (9) L. UT.Clark, J . Phys. Chem., 6 6 , 1543 (1962).
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LOUISWATTSCLARK
630
I
39
28 d
2 32
3 26
M
e
h
'
>
k
30
24
I-
&
2
28
26
22
20 -10
0
-5
5
10
15
AS*, eu/mole.
Figure 2. Enthalpy-entropy of activation plots for the decarboxylation of benzylmalonic acid in acids and cresols: line I, benzylmalonic acid in cresols; line 11, malonanilic acid in cresols. Slope = 393°K or 120°C.
which, by coincidence, is very nearly identical with its melting point. Petersen, et aZ.,1° have critically analyzed the problem of the validity of an observed linear enthalpyentropy of activation relationship. They have shown that such an observed relationship is probably invalid as a result of experimental error if the range of AH* values is less than twice the maximum possible error in AH*. Applying their mathematical interpretation to the data shown graphically in Figure 1, we find that the range of AH* values is nearly 20.0 kcal/mole, whereas the maximum possible error in AH* (assuming
The Journal of Physical Chemistry
-20
-10 AS*,eu/mole.
-15
-5
0
Figure 3. Enthalpy-entropy of activation plots for the decarboxylat,ion of benzylmalonic acid and malonanilic acid in quinoline and in 8-methylquinoline. Slope = 423°K or 150°C.
a maximum fractional error in the rate constants to be 0.05) turns out to be 1.4 kcal/mole. On this basis the range of AH* values is more than seven times as great as twice the maximum possible error in AH*. These results inspire considerable confidence in the validity of the relationship shown in Figure 1.
Acknowledgment. Acknowledgment is made to the donors of The Petroleum Research Fund, administered by the American Chemical Society, for support of this research. (10) R. C. Petersen, J. H. Markgraf, and 9. D. Ross, J . A m . Chem. Soc., 8 3 , 3819 (1961).
