572
Vol. 6.5 in a hydrocarbon solvent, undecane; that result falls squarely on the line shown in Fig. 1. The pre-exponential factor obtained here is 1.1i X 1019. The corresponding entropy of activation, under the assumption that the reacting solution is ideal, is 28.2 e.u. While not unprecedented, these values are uncommonly large. They are entirely plausible, however, upon consideration that (1) they apply to the decomposition of an unstable compound, (2) resulting in an increased number of molecules, (3) including nitrogen, which has a high molar entropy.
-3
-4
T H E EFFECT OF ANILINE iiND ITS DERIVATIVES ON OXANILIC ACID
x'
BY LOUISWATTS CLARK^
W
0 _I
Department of Chemistry, Saznt M a r y of the Plains College, Dodge City, Kansas Received September 19, 1960
-5
-6
2.6
3.0
2.8
I T
3.2
3
x IO
Fig. 1.-The temperature dependence of the rate of thermal decomposition of dimethyl 2,2'-azo-bis-isobutyrate: 0,Lewis and Matheson; 0,Ziegler, Deparade and Meye; 0,this work.
rates of thermal decomposition. At temperatures of 53.6, 68.9, 84.5 and 95.3', the decomposition rates were 1.31 X 2.07 X 1.96 X 10-4 and 9.31 X set.-', respectively. These results are comparable to those that have been reported for similar azo compounds. The data were used to construct the Arrhenius plot shown in Fig. 1. The results of dilatometric rate determinations by earlier workersae4also are indicated. The following rate expression is obtained K(E;ec.-l) = 1.17 X 1019 e-37.3WIRT
The actival ion energy obtained here, 37.3 kcal., agrees within experimental error with that reported by Lewis and M a t h e ~ o n35.8 , ~ kcal. However, neither of these values agrees well with that reported by Ziegler, Deparade and Meye14 30.9 kcal. By reference t o Fig. 1 it can be seen that the discrepancy in activation energy might be regarded simply as a discrepancy in the rate of decomposition :it 80'. However, a preferable explanation is that it results from the influence of the solvent. The activation energies reported here, and by Lewis and Matheson, were determined using hydrocarbon solvents, while that low value was determined using nitrobenzene as solvent. Ziegler, Deparade and Meye give the result of a determination of the decomposition rate a t 80'
It has long been known that the first tm.0 members of the homologous series of dicarboxylic acids, oxalic acid and malonic acid, readily undergo decarboxylation when heated either alone or in certain solvents. Recently the mechanisms of both reactions have been established on the basis of kinetic studies.2 It appears highly probable that the rate-determining step of both reactions is essentially the same, namely, the formation, prior to cleavage, of a transition complex involving coordination between one of the polarized, electrophilic, carbonyl carbon atoms of the un-ionized diacid with an unshared pair of electrons on a nucleophilic atom of a solvent molecule. Since only one carboxyl group is involved in the rate-determining step it would be deduced that mono-derivatives of both acids also should be capable of undergoing decarboxylation in polar solvents according to the same mechanism. This deduction has been confirmed already in the case of the decarboxylation of oxamic acid in several aromatic amines.3 Preliminary experiments in this Laboratory recently revealed that oxanilic acid, another monoderivative of oxalic acid, likewise undergoes decarboxylation when warmed in polar solvents. Since this compound, apparently, has never been the subject of any kind of kinetic study, it was thought worthwhile to carry out an investigation of the kinetics of the reaction in order to ascertain whether or not it proceeded by the same mechanism as that of the other related acids. For this purpose kinetic studies where carried out in this Laboratory on the decarboxylation of oxanilic acid in aniline, o-toluidine and o-ethylaniline. The results of this investigation are reported herein. Experimental Reagents.-( 1) The oxanilic acid used in this research was 100.0% pure as revealed by the fact that in every decarboxylation experiment the theoretical volume of CO, was obtained. (2) The solvents were reagent grade. Each sample of each solvent was distilled a t atmospheric pressure directly into the dried reaction flask immediately before the beginning of each decarboxylation experiment.
~-
(1) Western Carolina College, Cullowhee, N C. (2) (a) G. Fraenkel, R. L. Belford and P. E. Yankwich. J . Am. Chem. SOC.,16, 15 (1954); (b) L. W. Clark, J . Phya. Chem., 61, 699 (1957). (3) L. W. Clark, %bid.,65, 180 (1961).
NOTES
March, 1961 Apparatus and Technique.-The kinetic experiments were conducted in a constant-temperature oil-bath (f0.01’) by measuring the volume of COZevolved a t constant pressure. Details are given in a previous paper.p Temperatures were determined by means of a thermometer calibrated by the U. S. Bureau of Standards. In each experiment a 0.2948-g. sample of oxanilic acid (the amount required to yield 40.0 ml. of COz a t STP on complete reaction) waB introduced in the usual manner into the reaction flask containing about 60 ml. of solvent saturated with dry COZgas. All atmospheric oxygen was flushed out of the apparatus with a stream of the dry COZgas, and the experiments were carried out in an atmosphere of COz in. order to prevent oxidation of the solvents during the kinetic experiments.
573
1.4
1.0
,
G 8
- 0.65
-\.
gJ
- 0.2d
c
F-
2 6 10 14 18 22 26 Results Time (minutes). In studies on the decarboxylation of oxamic acid 1.-Experimental data: decarboxylation of 0.2948 in aniline and in +toluidine3 some experimental g. Fig. of oxanilic acid in 130 g. of aniline at 160.58’ (cor.): difficulties were experienced due to the tendency of I, volume of COZ (ml.); 11, log ( I r m - IrJ. the reverse reaction to occur, COZ and formamide recombining to give oxamic acid. I n the present Discussion studies on the decarboxylation of oxanilic acid no Oxamic acid represents the substitution of an such difficulties were experienced, there being no amine group for one of the terminal hydroxyl observable tendency for the reverse reaction to groups of oxalic acid. Inasmuch as the +M effect occur. The theoretical volume of COZwas obtained of the amide group is greater than that of the in each decarboxylation experiment. No difference hydroxyl,6 the carbonyl carbon atom of oxamic acid in the rate of reaction a t a fixed temperature could which becomes involved in coordination with the be detected when the quantity of solvent used was nucleophilic atom of the solvent molecule will have varied from 50 to 135 g. I n the case of each of the a lower effective positive charge than will that of solvents used in this investigation the log ( V , - oxalic acid.3 On the basis of the principle that the Vt) was a linear function of time over practically activation energy or enthalpy decreases as the the eiitire course of the reaction. Results of a typi- attraction between two reagents increases16it would cal experiment are shown graphically in Fig. 1. be deduced that, in a given solvent, the AH* for the Duplicate experiments were performed in each decarboxylation of oxalic acid would be less than solvent at three or four temperatures over about a that of oxamic acid, in the event both reactions 20’ range. The average rate constants, calculated proceed by the same mechanism. The experimental in the usual manner from the slopes of the experi- data shown in lines 2 and 3 of Table I11 confirm mental logarithmic plots, are brought together in this deduction. Table I. The parameters of the Eyring equation, TABLE 111 based upon the data in Table I, are shown in Table 11, along with corresponding data previously ob- KINETICDATAFOR THE DECARBOXYLATION OF SEVERAL UNSTABLEACIDS IN ANILINE SOLUTION tained for the decarboxylation of oxamic acid.
TABLE I API~ARESTFIRST-ORDER RATE CONSTANTS FOR THE DECARBOXYLATION OF OXANILICACID IN THREE AROMATIC AMINES sol\ ent
Temp.,
k X 104,
OC.
880. -1
Av. dev.
Reactant
AH*, kcal./mole
AS*, e.u.
Malonic acid7 Oxalic acid2 Oxamic acid3 Oxanilic acid
26 9 40 3 59 7 49.8
- 4 45 $16 2 $68.0 +46.3
x lo4, 8ec. -1
k14Qo
50 0 1.5 0.25 1.5
Oxanilic acid represents the substitution of a phenyl group for one of the amide hydrogen atoms of oxamic acid. Since the phenyl group exerts a strong -I effect its presence will tend to reduce the o-Toluidine +E effect on the amide nitrogen atom; therefore, the effective positive charge on the antipodal, polarized, carbonyl carbon atom of oxanilic acid o-Ethylaniline will not be neutralized to as great an extent as it is in oxamic acid. This will mean that oxanilic acid will be a stronger acid than oxamic acid, and the effective positive charge on the carbonyl carbon atom involved in solvation will be higher in oxaiiilic TABLE I1 It would be predicted KINETICDATAFOR THE DECARBOXYLATION OF OXANILIC acid than in oxamic acid. ACID AND OXAMIC ACID IN SEVERAL AROMATICAMINES therefore, that, in a given solvent, the AH* would be less for the decarboxylation of oxanilic acid Oxanilic acid AH*, Oxamic acids than for that of oxamic acid, provided the two kcal./ AS*, AH*, AS*, Solvent mole e.u. kcal. e.u. reactions proceed by the same mechanism. ConAniline
Aniline o-Toluidine o-Ethylaniline
132 46 142 34 151.78 130.47 141.14 148 96 131 66 142 14 146 55 140.80
49.8 47.5 45.5
$46.3 $39.0 +34.3
0.51 2.16 8 44 0 59 3 20 8 23 0.75 3.18 5 76 9 10
59.7 53.7
(4) L. W. Clark, J . Phas. Chem.. 60, 1150 (1956).
f0.004 f .01 f .02 f .008 f .01 f .02 f .01 f .01 f .03 f .01
+68.0 $57.1
(5) A. E. Remlck. “Electronic Interpretations of Organic Chenm try,” John Wiley and Sone. Inc., New York, N. Y..2nd Ed., 1949 13 37 (6) K. J. Laidler, “Chemical Kinetics,” McGraw-11111 Book Co , Inc., New York. N. Y..1950,p. 138. (7) L. W. Clark, J . P k y s . Chem.. 63. 79 (1958).
574
NOTES
sidering also the greater bulk of the phenyl group as compared with the hydrogen atom, it mould be expected that, in the approach of the attacking molecule to the nitrogen atom of the amine, oxanilic acid should encounter more steric hindrance than would oxamic acid-in other words, for the reaction in a given solvent, the AS* for the reaction should be less for oxinilic acid than for oxamic acid. Both these predictioiis are seen to be confirmed by the data in Table 111. In the studies of the decarboxylation of oxamic acid in aniline and in o-toluidine3 it was shown that the presence of a methyl group ortho to the amino group has two effects, namely, (1) a +I effect which increases the electron density on the nitrogen atom of the amine giving rise to a decrease in AH*, and ( 2 ) an ortho or steric effectBwhich hinders the approach of the oxamic acid to the nitrogen atom, resulting in a decrease in AS*. Analogy suggests that oxanilic acid should behave in this respect in a manner similar to oxamic acid. The experimental data in Table I1 confirms this expectation. It is seen that, on passing from aniline to o-toluidine, a progressive decrease in both AH* and AS* takes place for both the oxamic acid and the oxanilic acid reactions. A, further decrease in the parameters takes place in the oxanilic acid reaction on going from o-toluidine to o-ethylaniline. For each solvent it will be noted, also, that both AH* and AS* are lower for the decarboxylation of oxanilic acid than for that of oxamic acid. This is due to the fact that, in the first place, the effective positive charge on the coordinating carbonyl carbon atom of oxaiiilic acid is greater than it is on that of oxamic acid, and, secondly, the molecule of the former is the more comp1e.u. These circumstances leave little doubt but that oxanilk acid decomposes in polar solvents by the same mechanism as does oxamic acid, oxalic acid and malonic acid. I n aniline solution, the AS* for the decarboxylation of oxanilic acid is 18.5 e.u. smaller than it is for that of oxamic acid (see Table 111, lines 3 and 4). This difference in AS* is commensurate with the increascld steric hindrance which would be expected from the substitution of a phenyl group for one of the amide hydrogen atoms of oxamic acid. Since oxaiiilic acid is considerably larger than oxalic acid the A S * for the reaction in aniline would be expected to be larger for oxalic acid than for that of oxanilk acid provided the same number of molecules of each acid were iiivolved in the solvation step. We find, however, that, in aniline, the AS* for the oxalic acid reaction is actually 30 e.u./mole less than it is for that of oxanilic acid. I t is well known that the dicarboxylic acids in solution associate through hydrogen bonding to form a cluster composed of at least 3-4 m~lecules.~Wheii one of the hydroxyl groups of a dicarboxylic acid is replaced by sorne other group such as the S-phenyl amide group association evidently can no longer take place past the dimer stage. In view of the great differences in the activation entropies of (8) L. 1’. Hamrnett, ”Physical Organic Chemistry,” 3IcGraw-Hill Book Co., Inc., Ncsw York, N. Y., 1940, p . 204 (9) \V. Huchel, “Tlieoretical Principles of Organic Chemistry,” Vol. 11. Elswier Pub Co , Kew York. N Y 1958. p. 329 el seq.
Vol. 65
these two reactions it may be deduced that, in the case of oxanilic acid, probably only a single molecule coordiiiates with the solvent, whereas, in the case of oxalic acid, a supermolecule consisting of an association cluster of 3-4 single molecules coordinates. h similar interpretation has been advanced for the oxamic acid data.3 In Table I11 it will be observed that, at 140°, oxalic acid and oxanilic acid decompose in aniline a t about the same rate. This evidently is due to the circumstance that, even though the AH* for the decarboxylation of oxalic acid is nearly 10 kca1.l mole less than it is for that of oxanilic acid, the improved entropy factor in the case of oxanilic acid exactly compensates for this disadvantage. Further work on this problem is contemplated. Acknowledgment.-The support of this research by the National Science Foundation, Washington, D. C., is gratefully acknowledged. HEAT OF NEUTRALIZATION OF STRONG ACIDS BY STRONG BASES IN MIXED WATER-DIOXANE SOLUTIONS’ BY HIDEHIKO KIDOAND W. CONARD FERNELIUS Department o j Chemistrv. The Pennsylvania Stale 1-niuersitg, Cniversitv Park, Pa. Received August 4, 1960
I n determining format,ion constants of complexes of organic ligands with met,al ions, it is common practice to use the mixed solvent wat’er-dioxane t,o secure the necessary h careful analysis of this mixed solvent has shown that measurements made with ordinary glass and calomel electrodes and a pH met,er can be used to calculate t’hermodynamic dissociation constants of weak electrolytes and formation constants of complexes.* This signifies that even a 75 volume dioxane-water mixt’ure (mole fraction of dioxane = 0.388) is functioning essent,ially as a “water” solvent. Indeed, the activity of water a t 25’ in a 757, dioxane solution differs by only 10% from that of pure water.9 Xevertheless, many are concerned that the state of hydration of ions in t’his (1) This investigation was carried out under contract AT(30-11-907 between The Pennsylvania State University and the U. S. Atomic Energy Commission. (2) M. Calvin and K. W. Wilson, J . Am. Ciiern. Sac., 67, 2003 (1915). (3) D. P. hIellor and L. E. hlaley, Nature, 159, 370 (1937); L. E. Jlaley and D. P. Mellor, Australian J . Sci. Research, 2 8 , 92 (1949). (4) H. Freieer, R. G. Charles and W. D. Johnston, J . Am. Chem. Sor., 74, 1383, 1386 (1952); W. D. Johnston and H. Freiser, A n d . Chem. Acta, 11, 201 (1954); T. R. Harkins and H. Freiscr, J . A m . Chem. Soc., 77, 1374 (1955); 78, 1143 (1956): 80, 1132 (1958); G . E. Cheney, 11. Freiser and Q. Fernando. ibid., 81, 2611 (1959). ( 5 ) F. Bnsolo, P. T. CIien and R. Kent Murniann, ibid., 76, 950 (1954). (6) L. G. Van Uitert. C. G. Ilaas, JV. C. Ferneliiis and B. E. Douglas, ibid.,75, 455 (1953); L. G. Van Uitert, W. C. Fernelius and B. E. Douglas, ibid., 75, 457, 2736, 2739 (1953); L. G. Van Uitert and W. C. Fernelius, ibid., 75, 3802 (1953); 76, 375 (1954); C. M. Callahan, W. C. Fernelius and B. P. Block, Anal. Chim. Acta, 16, 101 (1957); D. F. Martin and R‘. C. Fernelius, J . Am. Chem. Soc., 81, 1509 (1959); B. B. Martin and W. C. Fernelius, ibid., 81, 2342 (1959). (7) H. Irving and H. Rossotti, Acta Chem. Scand., 10, 87 (1956). (8) L. G. Van Uitert and C. G. Haas, J . Am. Chem. Soc., 75, 461 (1953); L. G. Van Uitert and W. C. Fernelius, ibid.. 76, 5887 (19.54). (9) P. Ilovorka, R. -4. Scliaefer and D. Dreisbach. J . A m . Chem. Sue., 58, 2264 (1930).
