July, 1960
DECARBOXYLATION O F MALONIC ACIDAND
However, the crystal samples were prepared by very different methods. The American Potash boron I was preparedg by a hot wire deposition process. It is believed that the U. S. Borax boron I1 was prepared by a process involving the recrystallization of a commercial “amorphous” boron a t high temperatures. If some of the amorphous boron still remains in boron 11, this would tend to explain the higher heat content. The high temperature data reported here may be combined with the low temperature data of Johnston, Hersh and Kerr’ to yield heat capacit,y equations for crystalline and amorphous borons of the form CpIR = D(OD/T) (9) D.
+ 2EdOdT) + E?(&/T)
K.Stern and L. Lynds, J. Electrochem. Soc., 106, 676 (1958).
TRICHLOROBCETATE
917
where the characteristic temperatures have the values 6~ = 875; 61 = 1075 and 62 = 4000 for crystalline boron and 6~ = 750, el = 1150 and 6:, = 4050 for amorphous boron. These equations yield the thermodynamic functions in Tables T’II and VIII. Acknowledgments.-The authors wish to acknowledge the financial support of this work by the Bureau of Aeronautics, U. S. Navy, through its contractor the Callery Chemical Company, and the du Pont Company. Also, we acknowledge the aid of Dr. M. V. Evans for spectrographic analyses. Samples of boron for study were furnished by the American Potash and Chemical Corporation, by the U. S. Borax Company and by Dr. Robert J. Thorn of the Argonne National Laboratory.
COMPARATIVE STUDIES ON THE DECARBOXYLATION OF MALONIC ACID AND THE TRICHLOROACETATE ION BY LOUISWATTSCLARK Depadment of Chemistry, Saint Mary of the Plains College, Dodge City,Kansas Received February 18, 1860
Kinetic data are reported for the decarboxylation of malonic acid in molten trichloroacetic acid and of the trichloroacetate ion in decanoic acid. The constants of the Eyring equation are evaluated. A comparison of these results with those previously obtained for the reaction of these two reagents with a wide variety of polar solvents indicates that the malonic acid tends to evoke favorable electromeric and inductive effects which produce a decrease in AH whereas the negative charge on the trichloroacetate ion tends to evoke unfavorable electromeric and inductive effects which cause an increase in AH An explanation is advanced for the fact that the aromatic amines are an exception to this pattern of behavior. Finally, a possible mechanism is suggested for the decarboxylation of molten trichloroacetic acid at high temperatures.
*,
*.
The mechanism of the decarboxylation of malonic acid in nucleophilic solvents has been well established.’ Evidence has been presented which strongly indicates that the trichloroacetate ion likewise decomposes in similar solvents by essentially the same mechanism as does malonic acid.2 The rate-determining step of both reactions appears to be the formation, prior to cleavage, of a transition complex, the polarized carbonyl carbon atom of the reactant coordinating with a pair of unshared electrons on the nucleophilic atom of a solvent molecule. If the cumulative inductive effects of three alpha halogens are considered, the effective positive charge on the carbonyl carbon atom of the trichloroacetate ion is evidently greater than that in the case of malonic acid. Therefore, since an increase in the attraction between two reagents lowers the AH+ of the reaction,a it would be predicted that, in identical solvents, the AH* would be less for the decomposition of the trichloroacetate ion than for malonic acid. Comparison of data on the two reactions in five aromatic amines reveals that this anticipated result is actually obtained.2 For example, in aniline, AH* (in kcal.) for the decarboxylation of malonic acid is 26.9, for the trichloroacetate ion, 24.5; in quinoline corre(1) G. Fraenkel, R. L. Belford and P. E. Yankwich, J. Am. Chum. Soc., 76, 15 (1954). (2) L. Clark, THIS JOURNAL, 68,99 (1959).
W.
(3) K. J. Laidler, “Chemical Kinetias.” McGraw-Hill Book Co., Inc., New York, N. Y.,1950. p. 138.
sponding figures are 26.7 and 24.0; in o-toluidine 25.7 and 23.8. However, when the reactions were carried out in several of the monocarboxylic acids and their derivatives, it was found that AH* was higher in each case for the decarboxylation of the trichloroacetate ion than for malonic acid.4 For example, in propionic acid, AH* (in kcal.) for the decomposition of malonic acid is 33.4, for the trichloroacetate ion, 35.3; in butyric acid, 32.3 and 35.1; in monochloroacetic acid, 31.7 and 48.9. These results suggest that a fundamental difference in electron activation obtains between reactions of malonic acid and the trichloroacetate ion in acid media. The source of this difference may be traced to the difference in the nature of the two reagents. I n malonic acid we are dealing with a neutral molecule containing a polarizable, electrophilic carbonyl group. The electrophilic carbonyl carbon atom may have two effects: (1) it may coordinate with an unshared pair of electrons on the nucleophilic atom of the solvent molecule, and (2) it may attract mobile electrons present in the solvent molecule-in other words it may evoke a favorable electromeric effect a t the moment of reaction. In the case of the trichloroacetate ion, however, the electrophilic carbonyl carbon atom is bound to two oxygen atoms which share, through resonance, a negative ionic charge. This negative charge will (4) L.
W. C‘srk, ibid., 63, 1760 (1959).
918
LOUISJT"V.ITTSCLARK
tend to repel any mobile electrons which may be conjugated with the nucleophilic center of the solvent molecule so that, a t the moment of reaction, the electron density on the nucleophilic atom of the solvent will be reduced, thus necessitating the acquisition of additional energy in order to form the nctivated complex. If these principles are applied to the case of the two reactions taking place in acid solvents, it will be deduced that the trichloroacetate ion will tend to repel electrons away from the point of attack, thus evoking a -E effect on the carbonyl oxygen atom of the solvent. Hence the effective positive charge on the carbonyl carbon atom will be increased, electrons will then be withdrawn from the hydroxyl group, and more energy will be needed to form the activated complex. In the case of the reaction with malonic acid, on the other hand, a -E effect will not be called into play a t the moment of reaction, hence the activation energy will be expected to be l o - ~ e r . ~ In order to further test the above poptulates, further kinetic studies have been carried out in this Laboratory, as: (1) the decarboxylation of malonic acid in molten trichloroacetic acid, and (2) the decarboxylation of the trichloroacetate ion in decanoic acid. Results of this investigation are reported herein. Experimental (a) Reagents.-The malonic acid, trichloroacetic acid and potassium trichloroacetate used in this research were Reagent Grade, 100.O~oassay. The decanoic acid was "Highest Purity" grade. The usual precaution mas taken of distilling this solvent a t atmospheric pressure directly into the reaction flask immediately before the beginning of each decarboxylation experiment. (b) Apparatus and Technique.-The details of the apparatus and technique have been described previously.6 Temperatures were controlled to within f0.01"and were determined by means of a thermometer calibrated by the U. S. Bureau of Standards. In studying the decomposition of malonic acid in molten trichloroacetic acid, 60 g. of trichloroacetic acid were placed in a 100-ml. reaction flask which was then immersed in the constant-temperature oilbath, and 0.1857 g. of malonic acid (the amount required to produce 40.0 ml. of COn a t STP on complete reaction) was introduced in the usual manner. In studying the decomposition of the trichloroacetate ion in decanoic acid, 330 mg. of potassium trichloroacetate (the amount required to furnish 40.0 ml. of COn a t STP on complete reaction) was added t o 50 ml. of solvent.
Results The Decarboxylation of Malonic Acid in Molten Trichloroacetic Acid.-It has been found that, when 1 mole (163.8 g.) of trichloroacetic acid is heated a t temperatures of about 156' and above, carbon dioxide is evolved at an appreciable rate.? Since the temperature coefficient of the reaction is about 5.4, the volume of COz evolved with time falls off rapidly as the temperature is decreased below about 156'. When smaller quantities (50-60 g.) of trichloroacetic acid are heated at temperatures considerably below 156' very little volume change occurs during the period of time required to study the decomposition of malonic acid therein. 1.
(5) A. E. Kemick, "Electronic Interpretations of Organic Chemistry," 2nd Ed., John Wiley and Sons, Inc , New York, N. Y . , pp. 6?. 53 and 54. (6) L. W. Clark, THIBJOURNAL, 60, 1150 (1956). ( 7 ) L, W,Clark, J, Am, Chsm, Boc., 7 7 , 3130 (1955).
Vol. 64
However, in determining the rate constants for the decarboxylation of malonic acid in trichloroacetic acid, allowance was made for the small increment of volume change due to this factor. Experiments were carried out, in duplicate, a t three different temperatures over a 16' range. The average rate constants, in sec. -I, calculated in the usual manner from slopes of the experimental logarithmic plots, at were as follows: a t 131.60°, 0.203 X a t 147.39', 1.065 X l P 4 . 138.14', 0.407 X Average deviations were less than 1% in each case. On the basis of these results, AH* for the reaction was found to be 34.9 kcal., and A S +5.5 e.u. 2. The Decarboxylation of the Trichloroacetate Ion in Decanoic Acid.-The decarboxylation of the trichloroacetate ion was first order in decanoic acid, and followed very nearly the same pattern of behavior as was found previously for the reaction in acetic acid, propionic acid and butyric acid.4 Experiments were carried out, in duplicate, a t three different temperatures over a 16' range. The average rate constants, in set.-', calculated from the slopes of the experimental logarithmic plots, were as follows: at 121.62', 1.19 X l o w 4 :at 132.13', 4.68 X at 137.70°, 9.58 X lo-'. -4verage deviations were not more than 1% in each case. The value of AH* for the reaction was found to be 41.4 kcal.; the ralue of A&'*, $27.7 e.u. Discussion The reasons for the selection of the particular systems for investigation reported in this research will become apparent from a study of Table I. TABLE I EXTI-IALPIES O F L4CTIV.1T10XFOR >fALONIC
THE ~ E C h R B O X Y L b T I O R 'Of
-4CID A V D O F THE TRICHLOROACETATE
10s IU
VARIOCSLIQVID~ AH
System
a
*,
kcal
Molten malonic acids 33.0 Malonic acid in propionic acidg 33 6 31 7 Malonic acid in monochloroacetic acidg 34 9 Malonic acid in trichloroacetic acid" 26.6 Malonic acid in decanoic acidlo 35.5 Trichloroacetate ion in acetic acid4 Trichloroacetate ion in propionic acid* 35.3 Trichloroacetate ion in monochloroacrtic acid4 48.9 62.0 Molten trichloroacetic acid' 41 4 Trichloroacetate ion in decanoic acid" Present research
On going from acetic acid to nionochloroacetic acid, AH* for the decomposition of the trichloroacetate ion increases from 35.5 to 48.9 kcal. (see lines 6 and 8 of Table 1)-a result which has been attributed to a -I effect of the halogen evoked by the ionic negative charge on the trichloroacetate ion.? On going from propionic acid to monochloroacetic acid, the AH* for the decomposition of malonic acid decreases from 33.6 to 31.7 kcal. (see lines 2 and 3 of Table I)-a result which has been attributed to a +E effect of the halogen evoked by the effective (8) C. N. Hinshelwood, J . Chon. Soc., 117, 156 (19801. (9) L. W. Clark, THISJ O U R A . ~ L , 64, 41 (1960). (10) f.ni. Clark, ibzd., 64, 692 (1960),
July, 1960
DECARBOXYLATIOS O F ~ I A L O ACID N I CA N D TRICHLOROACETATE
positive charge on the polarized carbonyl carbon atom of the neutral malonic acid.s It has been suggested that, in the decomposition of molten malonic acid, the transition complex is formed by the coordination of a polarized carbonyl carbon atom of one molecule of malonic acid with an unshared pair of electrons on the hydroxyl oxygen atom of r2 second molecule of malonic acid.1° h study of the data shon-n in lines 6, 8 and 9 of Table I sheds light on the probable mechanism of the decomposition of molten trichloroacetic acid. Iiiasmuch as un-ionized trichloroacetic acid is stable, and since, in molten trichloroacetic acid, especially at elevated temperatures near the boiling point, some trichloroacetic ions are undoubtedly present, it may be inferred that the electrophilic carbonyl carbon atom of the ion coordinates with one of the unshared pairs of electrons on the hydroxyl group of one of the un-ionized trichloroacetic acid molecules forming a transition complex similar to that in the case of monochloroacetic acid. The yery high activation energy for the reaction (62.0 kcal.) is in line with the strong electron withdrawing power of the three alpha halogen atoms, none of which can release its mobile electrons because of the electron repelling effect of the negative charge on the attacking ion. The increase in AH* for the reaction on going from monochloroacetic acid to trichloroacetic acid (13.1 kcal.) is almost exactly equal to the increase on going from acetic acid to monochloroacetic acid (13.4 kcal.). In other words, three halogen atoms on the same carbon atom have only about twice the inductive effect as does one halogen. This is in line with the fact that in cases of multiple substitution on the same atom the inductive effect is not cumulative but diminishes with successive substitution. l1 For the decarboxylation of malonic acid in molten trichloroacetic acid AH+ is only 3.2 kcal. greater than it is in monochloroacetic acid (lines 3 and 4 of Table I). This indicates that, in the reaction with trichloroacetic acid as in that with monochloroacetic acid, the malonic acid evokes on the chlorine atoms at the moment of reaction a +E effect, which tendi: partially to restore to the nucleophilic center those electrons which have been removed by the inductive effect of the halogens. In molten trichloroacetic acid the AH* for the decomposition of the trichloroacetate ion is nearly twice that for malonic acid (lines 4 and 9 of Table I). Again it appears strongly evident that, in the case of the trichloroacetate ion, the halogens are exerting a -I effect since the negative charge on the attacking ion prevents a +E effect from taking place. whereas in the case of malonic acid, the halogens manifest a +E effect, evoked by the effective positive charge on the carbonyl carbon atom of the attacking agent. The data for the reactions in propionic acid and in decanoic acid (lines 2. 5 , 7 and 10 of Table I) also are consistent with the postulate that the trichloroacetate ion evokes a -E effect on the carbonyl oxygen atom of the solvent, whereas no -E effect can take place in the presence of malonic acid. (11) E. A. Braude and F. C . Nachod, "Determination of Organic Structures by Physical Methods," Academic Press, Ino., New York, N. Y., 1955,p. 579.
919
The strong hydrogen bonding in the dimer of the lower fatty acids results in a withdrawal of electrons from the hydroxyl group somewhat analogous t o that which would be produced by a -E effect on the carbonyl oxygen atom of the monomer. Therefore, the electron repelling effect of the negative charge on the trichloroacetate ion is able to produce only a small additional decrease in the electron density on the nucleophilic atom of propionic acid. The electron density on the hydroxyl group of the undimerized decanoic acid is normally greater than that in the case of propionic acid since in the former no electrons are being called into play to form hydrogen bonds. Additional evidence in support of these deductions is furnished by a study of data obtained by Verhoek on the decomposition of sodium trichloroacetate ion in ethanol.12 According to his results, AH* for the reaction is 31.1 kcal., as compared with 27.2 kcal. for the decomposition of malonic acid in 1-butan01.l~ It appears evident from these figures that the AH* for the decomposition of the trichloroacetate ion in the aliphatic alcohols will be higher than will be that for malonic acid. In the reaction with malonic acid the alkyl groups of the alkanols can exert a +I effect, whereas, with the trichloroacetate ion, such a +I effect is prevented from taking place because of the electron repulsion of the negative ion. These results indicate that, in the attempted coordination between malonic acid and a polar molecule, the effective positive charge on the polarized, electrophilic, carbonyl carbon atom of the neutral malonic acid molecule will attract mobile electrons, in other words it will evoke favorable inductive and electromeric effects, in such a way as always to increase the electron density on the nucleophilic center at the moment of reaction, thus bringing about a decrease of AH*. In other words, the solvent molecule will always tend to respond to the advances of the attacking reagent. In the attempted coordination between the trichloroacetate ion and a polar molecule, on the other hand, the negative charge on the ion will tend to repel mobile electrons, in other words it will evoke unfavorable inductive and electromeric effects, in such a way as always to decrease the electron density on the nucleophilic center a t the moment of reaction, thus bringing about an increase of AH+. In other words, the solvent molecule will always tend to resist the advances of the attacking reagent. The only class of compounds so far investigated which does not follow the above pattern is the aromatic amines2 Evidently, in these liquids, the effective positive charge on the carbonyl carbon atom of the malonic acid is not sufficiently strong to evoke a -E effect on the nitrogen atom-in other words it is not able to attract back to the nitrogen those electrons which have been withdrawn from it by the pon-erful resonance effect of the aromatic nucleus. Since the negative charge on the trichloroacetate ion prevents a -E effect from taking place on the nitrogen atom the electron density on the nucleophilic center of these solvents will not be (12) F. H. Verhoek, J . Am. Chem. Soc., 66, 571 (1934). 64, 508 (1960). (13) L. W.Clark, THIBJOURNAL,
J. M. CREETH
920
affected appreciably by the field effects of the different electrophilic reagents. Further work on this problem is contemplated.
VOl. 64
Acknowledgments.-This research was supported hy the National Science Foundation, Washington,
D. C.
ACTIVITY COEFFICIENTS FOR THALLOUS SULFATE I N AQUEOUS SOLUTION AT 25’ BYJ. M. CREETII~ Contributionfrom the Department of Physical and Inorganic Chemistry of the Unii,ersity of Adelaide, South Australia Received February 98, 1960
Measurements of the e.m.f. of the cell Tl, Hg (2 phase)/T12SO~(m)/HgzS04(s), Hg have been made over the concentration range 0.005 to 0.1 molal, and by combination with the literature value for the appropriate standard potential have been used to calculate stoichiometric activity coefficients. In the lowest concentrations investigated (0.005-0 01 m ) anomalous effectswere observed, and failure of the HgtSO, electrode suspected. In general, the activity coefficients are considerably lower than for example, those of the alkali metal sulfates, in conformity with results obtained on other thallous salts. By applying the Bjerrum theory and a model equation for the activity coefficients, a series of values have been calculated for the apparent “dissociation constant” for the TISO,- ion: by this means a comparison is made possible with results previously obtained on this system. It is concluded that ion-association greater than is compatible with the Bjerrum theory occurs, in agreement with the previous findings. The limits of applirability of the model activity coefficient equation are discussed in the light of their effect on the value of the dissociation constant.
In connection with studies2 of the transport properties of the thallous sulfate-water system, information concerning the variation of the activity coefficient with concentration was required; accordingly a limited series of measurements of the activity coefficient was undertaken, the results of which are reported here as they have some general interest and applicability. For the concentration region of interest, a method based on e.m.f. measurements has advantages, and accordingly the cell T ~ - H ~ / T ~ ~ S O I / H ~Hg ~SOI, 2 phase m
was used: this has an e.m.f., E , given by E = EQ- 3RT __ In 4‘/arny, 2F
Here m is the molality of T12S04,y2t the stoichiometric mean activity coefficient on the molal scale and the coefficient of the logarithmic term has its usual significance. Measurements were made between 0.005 and 0.10 m, the lower limit being dictated bv the characteristics of the H d O d elec&ode3p4and the upper by the limited solugility Of T12SOa. Experimental Materials.-Commercial thallous sulfate (BDH, Ltd., England) was twice recrystallized from conductivity water and dried i n vacuo over phosphorus pentoxide. Metallic thallium was prepared by electrolysis, following the general procedure of Richards and Daniels,* and mixed immediately with the appropriate amount of mercury (twice redistilled) to give a composition of approximately 57% T1. The mixture thus prepared was clearly two-phase (the solubility of TI in Hg is approximately 44% by weights a t 25’); i t was stored under boiled conductivity water. Mercurous sulfate (1) Lister Institute of Preventive Medicine, London S.W. 1, England. (2) To be published shortly. (3) (3. N. Lewia and M. Randall, “Thermodynamics and the Free Energy of Chemical Substances,” McGraw-Hill Book Co., Inc., New York, N. Y., 1923, p. 356. (4) H. 6. Herned and R. D. Sturgia. J. Am. Cham. SOC.,47. 945 (1925). (5) T.
W. Richards
and F. Daniels, ibid., 41, 1732 (1919).
was prepared in the form of an intimate mixture of the salt and mercury by the method of Hattox and de Vries.6 It was stored in the dark after thorough drying. Solutions.-Three were prepared by weight, employing the appropriate air buoyancy corrections and density data. The Cell for E.M.F.Determination.-The design of this unit must allow a completely oxygen-free7solution of thallous sulfate (of precisely known composition) to be placed in contact with the thallium amalgam electrode: the latter rearts instantly with atmospheric oxygen and so is normally covered, if dry, by a black film of thallous oxide which.immediately dissolves on the addition of an aqueous solution. The cell shown in Fig. 1was therefore used; it is operated as follows. The dry amalgam is placed in the bottom of the ce!l compartment C through which R platinum wire E-1 1s sealed (a tube, not shown, leads this directly to the potentiometer terminals). The mercury-mercurous sulfate electrode E-2 is then assembled separately with dry materials4; about 1 ml. of the TlzSOcsolution to be used is then cautiously added to it and all air bubbles removed by alternately sucking and blowing with a capillary pipet. The electrode is then inserted into the cell arm F , connection to the potentiometer being made as before. The reservoir R (which holds about 25 ml.) is then filled with the TlSS04 solution; by opening the tap T-1 briefly, bubbles are removed from the tube joining R and C and simultaneously the surface of the amalgam in C is rinsed. Most of the solution in C is then removed, but enough is left to cover the surface of the amalgam. The remaining parts of the assembly are then put in place, and all taps are shut. In this state the cell is transferred to the thermostat (2‘ = 25.00 f 0.005”). and the solution in R deoxygenated at constant composition by the passage of purified hydrogen through a conventional presaturating system. During this time taps T-1 and T-4 aTe closed, T-2 and T-3 open: after about half an hour T-3 1s closed and T-4 opened so ensuring that the cell is thoroughly cleared of air As the junction of the tube I) with the cell occurs at a point just above the surface of the amalgam, most of the (slightly contaminated) solution left above the amalgam is displaced to D, from which it is removed. The amalgam is then rinsed four to six times with clean solution from the reservoir; this is done by closing T-2 and T-4 and opening T-1 and T-3, until the cell is half full, when T-1and T-3 are closed and T-2 and T-4 opened to displace the solution again to D. Finally the cell is filled completely by closing T-2 and T-4 and opening T-1 and T-3; if this is done cautiously, there is no disturbance of the layer of mercurous sulfate. (6) E. M. Hattox and T. de Vries. ibid., 68, 2126 (1936). (7) The sensitivity of the thallium amalgam electrode t o atmospheric oxygepheric oxygen has been discussed by G. N. Lewis and C. L. von Ende, J . Am. Chem. Soc., 78, 4520 (1956).
