DECARBOXYLATION OF OXALIC
ACIDIN POLAR SOLVENTS
A factor which may be of significance in the operation of this cell is the liquid-like layers purported to exist on an ice interface at temperatures below O 0 . I 3 These films of water on ice have often been invoked to explain the phenomena of regelation and would be expected to enhance the transport of hydrogen to every interfacial region. These same layers may be responsible in some
1597
measure for the spurious potentials observed when various electrodes are frozen into ice. The electrode described should be of use in the measurement of thermogalvanic potentials in ice. The results of this investigation are in accord with the position that protons are the only mobile carriers in ice. (13) W. A. Weyl, J. Colloid Sci., 6 , 389 (1951).
Further Studies on the Decarboxylation of Oxalic Acid in Polar Solvents
by Louis Watts Clark Department of Chemistry, Western Carolina College, Cullowhee, North Carolina (Received December 6 , 1966)
Rate constants and activation parameters are reported for the decarboxylation of oxalic acid in propylene glycol, 1,4-butanediol, and 2,3-butanediol. The results of this investigation are compared with previous data on the reaction in seven additional solvents, as well as in the vapor phase. Inductive and steric effects of the various solvents are discussed.
I n the past, kinetic studies have been carried out on the decarboxylation of oxalic acid in at least 15 nonaqueous solvents. These include dioxane,' g l y ~ e r o l dimethyl ,~ sulfoxide, triethyl phosphate, aniline, N-methylaniline, N,N-dimethylaniline, q ~ i n o l i n e , ~ 6-methylquinoline, 8-methylq~inoline,~o-cresol, mcresol, p-cresol, ethylene glycol, and 1,3-butanedi0l.~ Recently, Haleem and Yankwich have repeated the experiments using the solvent glyceroL6 Lapidus, Barton, and Yankwich have also studied the decarboxylation of oxalic acid in the vapor phase and have proposed a unimolecular mechanism for the reaction.' In an effort to gain a better understanding of the mechanism and energetics of this reaction, further experimentation has been carried out in this laboratory on the decarboxylation of oxalic acid in three additional polar solvents, propylene glycol, 1,4-butanediol1 and 2,3-butanediol. A comparison of the results of this investigation (reported herein) with previous data sheds light on the mechanism of the reaction and leads to a better understanding of other heterolytic reactions.
ExperimentaI Section Reagents. (1) Anhydrous oxalic acid, reagent grade, 100.0% assay, was used in this research. To ensure perfect dryness it was stored in a desiccator containing sulfuric acid. (2) The solvents were highest purity quality and were redistilled at atmospheric pressure directly into the dried reaction flask immediately before the beginning of each decarboxylation experiment. Apparatus and Technique. The details of the apparatus and technique used in this research have been described previously.* In these experiments, a sample (1) A. Dinglinger and E. Schroer, 2. Physik. Chem., A179, 401 (1937). (2) L. W. Clark, J. Am. Chem. Sac., 77, 6191 (1955). (3) L. W . Clark, J. Phys. Chem., 61, 699 (1957). (4) L. W . Clark, ibid., 62, 633 (1958). (5) L. W. Clark, ibid., 67, 1355 (1963). (6) M. A. Haleem and P. E. Yankwich, ibid., 69, 1729 (1965). (7) G . Lapidus, D. Barton, and P. E. Yankwich, ibid., 68, 1863 (1964).
Volume 70,Number 5 M a y 1966
LOUISWATTSCLARE
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of oxalic acid weighing 0.1618 g (corresponding to 40.0 ml of COz at STP on complete reaction, based upon the actual molar volume of C02 a t STP, 22,264 ml) was weighed into a fragile glass capsule weighing approximately 0.1 g and blown from 7-mm soft glass tubing. The course of the reaction was followed by measuring the volume of COz evolved at atmospheric pressure and at the temperature of a water-jacketed buret. The buret was calibrated by the National Bureau of Standards at 20.0’. Water maintained a t 20.0 k 0.05” by means of a cooling coil and an electronic relay was circulated through the water jacket during the experiment. The temperature of the oil bath was controlled to within ~ 0 . 0 0 5 ”using a completely transistorized temperature control unit equipped with a sensitive thermistor probe. A thermometer which also had been calibrated by the National Bureau of Standards was used to read the temperature of the oil bath. Appropriate corrections and calibrations were carefully applied to this thermometer in order to reduce errors in temperature readings to a minimum. About 60.0 g of solvent, saturated with dry C02 gas, was used in each experiment.
Results Decarboxylation experiments were carried out in each solvent at four different temperatures over about a 30” range. The experiments were repeated two or three times at each temperature. I n the case of each of the solvents the log ( V , - Vt) was a linear function of time over about the first 80% of the reaction. The average rate constants, calculated in the usual manner from the slopes of the experimental logarithmic plots, are brought together in Table I. The parameters of the absolute reaction rate equationg
k
=
KT -e-AH*/RTeAS*/R h
based upon the data in Table I, are shown in Table 11, along with previously published data which are pertinent to the discussion.
Discussion A comparison of the activation parameters for the decarboxylation of oxalic acid vapor and the decarboxylation of oxanilic acid melt (lines l and 2 of Table 11) reveals that the smaller molecule (oxalic acid) has the lower entropy of activation. Since an increase in molecular complexity is generally accompanied by a decrease in the entropy of activation,’O we would have expected the decarboxylation of oxalic acid to have a higher entropy of activation than that of oxanilic acid. This inversion, however, finds a ready The Journal of Physical Chemistry
Table I: Apparent First-Order Rate Constants for the Decarboxylation of Oxalic Acid in Several Solvents k X
Temp, O C (cor)
Solvent
Propylene glycol
2,3-Butanediol
1,4-Butanediol
104 (BY).
SeC-’
118.38 128.75 137.64 148.65 130.30 138.51 149.85 160.17 139.56 152.93 160.03 169.48
2.67 3.97 5.35 7.68 1.81 3.26 7.03 13.64 0.96 2.62 4.71 8.43
~
Table 11: Activation Parameters for the Decarboxylation of Oxalic Acid and Related Acids Alone and in Several Solvents
System
1. 2. 3. 4.
5. 6. 7. 8. 9. 10. 11. 12.
Oxanilic acid (melt)“ Oxalic acid (vapor)* Oxalic acid in glycerol’ Oxalic acid in glycerold Oxalic acid in ethylene glycol‘ Oxalic acid in propylene glycol’ Oxalic acid in quinoline’ Oxalic acid in l,.l-butanediol’ Oxalic acid in dimethyl sulfoxide’ Oxalic acid in 1,3-butanediole Oxalic acid in triethyl phosphate’ Oxalic acid in 2,3-butanediol’
AH*,
AS*,
kcal/ mole
mole
40.1 29.2 27.6 26.4 17.6 10.3 38.9 25.5 40.6 29.3 28.85 22.63
21.4 -6.6 -4.9 -7.8 -30.0 -49.1 15.75 -15.8 20.7 -4.9 -5.8 -20.2
4
See ref 6. a L. W. Clark, J. Phys. Chem., 66, 1543 (1962). Present research, based See ref 6. See ref 2. See ref 5. on Table I. See ref 3.
’
explanation in the fact that dicarboxylic acids tend to associate past the dimer stage.l’ The polymerization of oxalic acid could conceivably yield a larger entity for taking part in the rate determining step than would the dimerization of oxanilic acid. The higher enthalpy of activation in the case of the decarboxylation of oxanilic acid may be explained on the basis of the (8) L. W. Clark, J . Phgs. Chem., 60, 1150 (1956). (9) S. Glasstone, K. J. Laidler, and H. Eyring, “The Theory of Rate Processes,” McGraw-Hill Book Co., Inc., New York, N. Y., 1941, p 14. (10) L. P. Hammett, “Physical Organic Chemistry,” McGraw-Hill Book Co., New York, N. Y., 1940, p 126. (11) W. Huckel, “Theoretical Principles of Organic Chemistry,” Vol. 11, Elaevier Publishing Co., New York, N. Y., 1958, p 329.
DECARBOXYLATION OF OXALICACIDIN POLAR SOLVENTS
+E effect exerted by the unshared pair of electrons on the nitrogen atom which would tend to neutralize the effective positive charge on the carbonyl carbon atom of the anilide. The decarboxylations of oxalic acid in ethylene glycol and in propylene glycol (lines 4 and 5 of Table 11) are accompanied by very large negative entropies of activation which may no doubt be attributed to extensive polymerization of these two solvents. The decrease in AH* and in AS* on going from ethylene glycol to propylene glycol may be explained on the basis of the inductive and steric effects of the methyl group in propylene glycol. One would naturally expect that the presence of two methyl groups on opposite sides of the glycol (as in 2,3-butanediol) would further accentuate these effects and result in lowering still more the enthalpy and entropy of activation of the reaction. Actually, honever, it is found that both the enthalpy and entropy of activation for the reaction in 2,3-butanediol are highe. than they are in either ethylene glycol or propylene glycol (line 12 of Table 11). When the two hydroxyl groups are separated by one or two carbon atoms, as in 1,3-butanediol (line 10) and 1,4butanediol (line 8), the enthalpy as well as the entropy of activation of the oxalic acid reaction are larger than they are in those compounds in which the two hydroxyl groups are located on adjacent carbon atoms (ethylene glycol, propylene glycol, and 2,3-butanediol-lines 4, 5, and 12 of Table 11). These anomalous results find a ready explanation when it is observed that the various systems shown in Table I1 are actually representatives of four different reaction series: (1) the decarboxylation of oxalic acid (v:ipor) and oxanilic acid (melt) (a derivative of oxalic acid); (2) the decarboxylation of oxalic acid in the related solvents glycerol, ethylene glycol, and propylene glycol; (3) the decarboxylation of oxalic acid in quinoline and in 1,4-butanediol; (4) the decarboxylation of oxalic acid in dimethyl sulfoxide, triethyl phosphate, 133-butanediol,and 2,3-butanediol. Plots of enthalpy us. entropy of activation for these four reaction series are shown in Figure 1. Line I11 of Figure 1 combines the two reaction series: (1) the decarboxylation of oxalic acid (vapor) and oxanilic acid (melt), and (2) the decarboxylation of oxalic acid in the related solvents glycerol, ethylene glycol, and propylene glycol. The enthalpy of activation axis for the first series has been shifted purposely in such a manner as to superimpose it upon the second series. This device helps to point out the interesting fact that the two lines connecting the parameters have very nearly equal slopes. (It is often found in kinetic studies that a change in the structure of reactant or
1599
d 40
I2O I
I
-40
-30
I
I
-10 0 -20 AS*, ep/mole.
I
I
10
I
20
Figure 1. Enthalpy us. entropy of activation plots for the decarboxylation of oxalic acid and related compounds: line I, oxalic acid in 1,4-butanediol and in quinoline; line 11, oxalic acid in 1,3-butanediol, 2,3-butanediol, triethyl phosphate, and dimethyl sulfoxide; line I11 (open circles), oxalic acid in glycerol (two sets of data), in ethylene glycol, and in propylene glycol; (closed circles), oxalic acid (vapor) and oxanilic acid (melt).
solvent results in the formation of a second line parallel to the original line.'*) The slope of this line is 393°K or 120°C. This is the so-called isokinetic temperature of the reaction series,I3 i.e., the temperature at which the rate constants for all the reactions conforming to the line are equal. The intercept of the isokinetic temperature line on the zero entropy of activation axis yields A F " , the free energy of activation at the isokinetic temperature of all the reactions conforming to the line.12 Substitution of this value in the absolute reaction rate equationg enables the rate of reaction a t the isokinetic temperature to be calculated. For the decarboxylation of oxalic acid and its derivative in the absence of solvent AF" is found to be 31.8 kcal/ mole. For that of oxalic acid in glycerol and the two glycols AF" is 29.5 kcal/mole. The respective calculated rate constants a t 120" in sec-l for these two reaction series turn out to be 0.000017 and 0.000324. I n other words, at 120", oxalic acid mould react nineteen times as fast in glycerol and the glycols as it would in the absence of solvent. Glycerol and the glycols thus lower the free energy of activation at the isokinetic temperature for the decarboxylation of oxalic acid. The decrease in AS* for the decarboxylation of oxalic acid from 15.75 eu/mole in quinoline to - 15.8 eu/mole (12) J. Leffler, J. Org. Chem., 20, 1202 (1955). (13) S. L. Freiss, E. S. Lewis, and A. Weissberger, Ed., "Technique of Organic Chemistry," Vol. V I I I , Part I, "Investigations of Rates and Mechanisms of Reactions," 2nd ed, Interscience Publishers, Ino., New York, N. Y., 1961,p 207.
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LOUISWATTSCLARK
1600
in 1,4-butanediol (lines 7 and 8 in Table 11) points to the extensive linear polymerization of the diol. The AH* - AS* plot for the decarboxylation of oxalic acid in these two solvents has a slope of 423°K or 150°C (line I of Figure 1). The large positive entropy of activation for the decarboxylation of oxalic acid in dimethyl sulfoxide (line 9 of Table 11) may be attributed to the simplicity of the solvent rnolecule and its lack of association in the absence of hydrogen bonding. The negative entropy of activatiorl for the reaction in triethyl phosphate (line 11) may be attributed to the steric effect of the three methyl groups. Similarly, the decrease in the entropy of activation on going from l,&butanediol (line 10) to 2,3-butanediol (line 12) may be attributed to the greater steric hindrance to the reaction in the latter diol due to the presence of the two methyl groups flanking the nucleophilic center. A plot of enthalpy of activation us. entropy of activation for the decarboxylation of oxalic acid in these four solvents (line I1 of Figure 1)has a slope of 438°K or 165°C. The problem of the validity of an observed linear enthalpy-entropy of activation relationship has been critically analyzed by Petersen, et aZ.14 They have shown that (because of experimental error) the range of AH* values for a given reaction series must be much greater than twice the maximum possible error in AH* before any assurance as to the validity of the relationship can be entertained. The maximum possible error in AH* can be calculated by means of the equation
TIT T -T
6 = 2 R 7 - a
(1)
where 6 is the maximum possible error in AH* (in the positive direction), R is the gas constant, T' and T are the upper and lower temperature limits respectively, and a is the maximum fractional error in the rate constant. For purposes of discussion we may take as a typical system in the present research the decarboxylation of oxalic acid in 2,3-butanediol (Table I). If we assume a maximum fractional error in the rate constant of 0.05, 6 (according to eq 1) is found to be 1.14 kcal/mole and 26 is thus 2.28 kcal/mole. I n Table I11 are shown the range of AH* values for each of the four reaction series shown in Table 11, and the ratio dAH*/26 in each case. The fact that this ratio is much greater than unity for each series inspires considerable confidence in the validity of the observed linear relationships shown in Figure 1. It is interesting to note that several other reaction series have been reported previously having the same isokinetic temperatures as those found in this research The Journal of Phyakal Chemistry
Table 111: Validity Tests of Several Linear AH* Plots According to Petersen, et ~ 1 . 1 4
Reaction series
1. 2. 3. 4.
Oxalic acid (vapor) Oxalic acid glycerol Oxalic acid quinoline Oxalic acid dimethyl sulfoxide
+ + +
- AS*
dAH*, kcal/
ked/
dAH*/
mole
mole
26
Figure 1 ref
11.9 17.3 13.4 17.0
2.28 2.28 2.28 2.28
5.2 7.6 5.9 7.4
Line111 Line I11 Line I Line I1
28,
Table IV : Isokinetic Temperatures of Several Reaction Series-Decarboxylation Reactions Isokinetic temp System
Cinnamalmalonic acid in cresols" Oxalic acid in dimethyl sulfoxide, etc.* Malonic acid in nitro compoundsC Oxalic acid in quinoline, etc.* Benzylmalonic acid in quinoline, etcad Malonanilic acid in quinoline, etc.' Oxalic acid in various solventsf 8-Resorcylic acid in various solventsu Malonic acid in acidsh Picolinic acid in various solvents' Oxalic acid and oxanilic acid alongi Oxalic acid in glycerol, etc.b Benzylmalonic acid in acids, cresols" Malonanilic acid in acids, cresols" Benzylmalonic acid in aniline derivativesd Malonanilic acid in aniline derivatives"
OK
OC
468 453 438
195 180 165
423
150
408 393
135 120
378
105
L. W. Clark, J. Phys. Chem., 67, 1481 (1963). This research. ' L. W. Clark, J. Phys. Chem., 62,368 (1958). d L. W. Clark, ibid., 70, 627 (1966). * L. W. Clark, ibid., 68, 2150 L. W. Clark, (1964). L. W. Clark, ibid., 66, 1543 (1962). ibid., 67, 2831 (1963). L. W. Clark, ibid., 68, 3048 (1964). L. W. Clark, ibid.,69,2277 (1965). j See ref 7 and L. W. Clark, ibid., 66, 1543 (1962).
'
(see Table IV). The sixteen different reaction series listed in Table IV yield five isokinetic temperatures separated from one another by 15" intervals. In general, the higher isokinetic temperatures are associated with stronger mutual attractions between the electrophile-nucleophile pairs.
Acknowledgment. Acknowledgment is made to the Donors of the Petroleum Research Fund, administered by the American Chemical Society, for support of this research. (14) R. C. Petersen, J. H. Markgraf, and S. D. Ross, J . Am. C h m . SOC.,83, 3819 (1961).
