NOTES
Sept., 1956
1341
TABLE I DECOMPOSITION OF MALONIC ACIDIN AROMATIC AMINES,KINETIC DATA E*
System
1. Malonic acid 2. Malonic acid plus quinoline 3. Malonic acid plus aniline 4. Malonic acid plus pyridine
W.)
A (sec. -1)
2.72
33,000
+ 4.5
31,100
0.00024
2.57 X 10'
1.78
17,800
-27.0
24,300
.00364
22,850
6.32 X lo9
2.01
22,800
-14.3
28,700
.0050
26,600
1.59 X lo1*
2.75
26,000
+ 7.2
23,000
.0126
34,500
1.8
18,500
x
1014
The above procedure was re eated at five different temperatures between 98-114' in e!t case of pyridine, and at six different temperatures between 126-144' in the case of aniline. The total volume of carbon dioxide obtained in each experiment was 100% of the theoretical yield within the limits of error of the experiment. Duplicate and triplicate runs at each temperature gave reproducible results.
Results and Discussion A plot of log (a - 2) us. t in the case of each amine at each temperature studied gave excellent straight lines, indicating that the decomposition of malonic acid in pyridine and in aniline is a first-order reaction. From the slopes of the lines thus obtained the specific reaction velocity constants for the reaction in the two solvents at the various temperatures studied were calculated. For the case of the decomposition of malonic acid in pyridine the temperatures studied and the corresponding specific reaction velocity constants obtained in sec. were as follows: 98.2', 0.000335; 102.3",0.000496; 107.0", 0.000764; 110.9", 0.00111; 113.5",0.00143. Results for the case of the decomposition of malonic acid in aniline were as follows: 126.6") 0.00184; 131.8') 0.002735; 135.7") 0.00358; 138.6")0.00437; 141.8", 0.00550; 144.1")0.00640. A straight line was obtained in each case when log /c was plotted against 1/T according to the Arrhenius equation. From the slopes of the lines thus obtained the activation energy and the frequency factor for the reaction in each solvent were calculated. On the basis of the Eyring equation the enthalpy of activation, entropy of activation and free energy of activation a t 140" for the reaction in each solvent were calculated. The kinetic data for the reactions studied are summarized in Table I which includes also for comparison data for molten malonic acid6and for malonic acid in quinoline.4 Some insight into the role of the solvent may be gained by considering the data in Table I. A comparison of the data in lines 1 and 2 shows that quinoline effects a very large lowering of the enthalpy of activation as well as a large decrease in the entropy of activation of the reaction with respect to molten malonic acid. This results in a lowering of the free energy of activation, so that, at this temperature (140°), malonic acid decomposes 15 times as fast in quinoline as it does alone. A comparison of lines 1 and 3 reveals that aniline lowers the enthalpy of activation by a large amount and also lowers the entropy of activation considerably, resulting in a small decrease in the free energy of activation a t 140' with respect to molten malonic acid. Malonic acid decomposes 21 times as fast in aniline a t (6) C. N. Hinabelwood, J . Chem. Soc., 111, 150 (1920).
140' as it does alone. A comparison of lines 1 and 4 reveals that pyridine lowers the enthalpy of activation less than does aniline and increases the entropy of activation, so that the probability of the formation of the activated state of malonic acid is greater in pyridine than in the molten state. The free energy of activation is lower in pyridine than in any other solvent shown. At 140' malonic acid decomposes 52 times as fast in pyridine as in the molten state. This order of increasing efficacy of the amines follows the order of increasing basicity as well as the order of decreasing molecular complexity of the solvent. Studies on the decarboxylation of acetonedicarboxylic acid in aniline' have been explained on the basis of the formation of an intermediate unstable compound between reactant and solvent molecules. The decomposition of malonic acid in nonaqueous solvents has been explained on the basis of decarboxylation of the undissociated form of the acid.6 Results reported herein strongly suggest the possibility that intermediate compound formation may be involved in the case of the decomposition of malonic acid in aromatic amines. The active methylene hydrogen atoms could attach to the unshared electrons on nitrogen. The greater the availability of the electrons the greater would be the probability of the formation of the activated state, other factors being equal. The electron structure of the molecule being thus distorted decarboxylation could occur a t either carboxyl group. Results obtained in non-nitrogenous, non-aquequs, basic type s o l ~ e n t are s ~ in ~ ~harmony with this mew. Further work on this problem is contemplated. (7) E. 0. Wiig, THISJOURNAL,S a , 961 (1928).
T H E PROTONATION OF N-METHYLACETAMIDE BY JAKEBELLO Contribution No. 8094 from the Coles and Crellin Laboralorjes of Chemistry, California Institute of Technolow, Paaadena, Calzfornza Received April 9, 1066
Mizushima, et u Z . , ~ have reported that N-methylacetamide (I) in aqueousohydrochloric acid has an absorption peak at 2695 A., which is not present in the absence of acid. This new absorpJion was attributed to the formation of the ion 11, smce (1) S. Miaushima, T. Simanouti, S. Nakagura. K. Kuratani, M. Tsuboi. H. Baba and 0. Fujioka, J . A m . Chem. Soc., 71, 3490 (1950); S. Mizushirna, Ada. Prot. Chem., 9, 299 (1954); 5. Mi~ushirna,"Structure of Moleciilcs and Internal Rotation," Academic Press, Inc.. New York, N. Y.,1954, 1 1 . 118.
4342
NOTES
I1 CHa-
8.
-NH&HP
it occurred near the 2650 8.band of acetone. On the assumption that the intensity of absorption of I1 and of acetone are similar, it was concluded that in a solution 2 M in I and 5 M in hydrochloric acid about 5% of I was converted to 11, t o give a concentration 0.1 M in 11. From such data there were also calculated an equilibrium constant for the dissociation of I1 and a value for the resonance energy of the "peptide" bond in I. We have had occasion t o investigate the spectra of solutions of I in acid preliminary to a study of protein-acid systems. We have found for 2 or 4 M I in 5 or 10 M hydrochloric acid no absorption a t 2695 A., or, indeed,oat any wave length above the beginning of 2050 A. band a t about 2400 A. According t o Mizushima, et al., a 2 M solution of I in 5 M hydrochloric acid has an optical density about equal to that of 0.1 M acetone. The optical density of 0.1 M acetone in water is about 1.6 and we could have detected readily an effect equal to 3% of the above. I n addition, no absorption a t 2695 A. was detected in concentrated sulfuric acid; that extensive solvolysis did not occur in this solvent was shown by dilution with water, resulting in an optical density a t 2050 A. equal (within 10%) to that of original I a t the same concentration. Acetic acid, which could be formed on solvolysis, absorbs weakly at 2050 A.2 I n addition, no absorption above 2400 A, was observed with 2 M acetamide or N-n-butylacetamide in 5 M hydrochloric acid. We conclude, therefore, that the reported absorption band a t 2695 8. does not exist, that there is no evidence for the formation of I1 and that the calculated resonance energy is not valid. We have COGfirmed the report2 that in acid solution the 2050 A. absorption peak of I is reduced. Experimental Freezing points are corrected. Materials.-Two samples of N-methylacetamide were used. One was White Label Grade purchased from the Eastman Kodak Co. and was purified by fractional distillation through a 12" Widmer column, b.p. 70-71' a t 2.5-3 mm., followed by two fractional crystallizations from the melt; yield 25%,f.p. 30.0°(reportedf. .28.80s; m.p. 28'1). The m.p. of the original amide is giveniy the seller as 27.529'. The second specimen was prepared by the dropwise addition of 2.2 moles of acetyl chloride to 6 moles of methylamine (Matheson, 96.5%) in 700 ml. of anhydrous ethyl ether cooled with Dry Ice. After filtration of the methylamine hydrochloride, the filtrate was distilled through the Widmer column. The fraction boiling at 79-81' a t 4.5-5 mm. was collected for use; yield 82%, f.p. 29.8". Anal. Calcd. for C3H,NO: N, 19.2. Found (Kjeldahl): N, 18.9. Both specimens gave identical infrared and ultraviolet spectra. Acetamide (Merck, Reagent) was used without additional purification; N-n-butylacetamide waR White Label (Eastinan Kodak Co.) grade nnd was distilled in vacuo through the Widmer column, b.p. 106" at 4 mm. (reported 103.5 a t 3.5 mm.z). Acet,one, Merck, Ileagent, "witable for spectrophotometric use'' was used. (2) A. R . Goldfarb, A. Mele and N. Gutatein, J . A m . Chem. SOC., 7'7, 6194 (1955). (3) I
