NOTES
1150
THE DECARBOXYLATION OF MALONIC ACID I N TRIETHYL PHOSPHATE BYLOUISWATTSCLARK Contribution from the Department of Chemistry, Saint Joseph College, Emmttsburg, MaTyland Received February 87, 1866
Kinetic studies on the thermal decarboxylation of malonic acid alone and in various solvents have been reported frequently in the literature.1-4 Hall2 has shown that the decomposition of malonic acid in aqueous solution can be explained by assuming two different mechanisms operating simultaneously: (1) decomposition of the unionized malonic acid, and (2) decomposition of the acid malonate ion. Reaction (1) is considerably faster than reaction (2). Despite the large amount of work done heretofore on this problem a great need still exists for more data, since every additional experiment throws further light on the reaction and aids kinetic theory. Preliminary experiments in this Laboratory revealed that malonic acid was smoothly decarboxylated in triethyl phosphate. The kinetics of the reaction were carefully investigated between 128150' and results are reported herein. Experimental Reagents.-Analytical reagent grade malonic acid was used. The triethyl phosphate (product of Tennessee Eastman Corporation) showed on the label the following analysis: specific gravity, 2Oo/4O, 1.065-1.070; chlorides, less than 5 p.p.m.; sulfates, none; acidity, not more than 0.03% calculated as HaP04;ester content, 97% minimum; boiling point, 216", refractive index, TP~D 1.404. Apparatus and Technique.-Considerable development in the apparatus and technique has occurred since the inception of these experiments.' For this reason a brief, general description a t this time appears warranted. The experiments are conducted in a well insulated, electrically heated, efficiently stirred constant temperature oilbath of 10-liters capacity. The temperature of the reaction is maintained thr2ugh the use of an oi1:bath thermostat regulated to h O . 1 The reaction is carried out in a 200ml. round-bottom Pyrex brand 2-neck (standard-taper) flask. Through the center neck is inserted an electrically driven, mercury seal stirrer controlled by a variac. In the other neck is inserted a reflux condenser. A short piece of glass tubing leads from the top of the condenser to the top of a 50-ml. water jacketed buret provided with a leveling bulb and an entraining liquid (20% by weight aqueous solution of sodium sulfate, 5% by volume sulfuric acid, plus a few drops of methyl orange). Tap water cvculates through the water jacket and condenser. Between the condenser and the buret is a small T tube, one end of which is sealed with a rubber policeman. A fine nichrome wire pierces the tip of the rubber policeman, forming a hermetic seal, and extends down the condenser tube. The lower end of the wire, near the lower neck of the condenser, terminates in a small loop bent a t right angles to its length. The upper end of the wire extends several inches beyond the tip of the rubber policeman. I n running an experiment the oil-bath is brought to the desired temperature, the solvent is added to the thoroughly dried flask, the tap water is turned on, and a stream of dry carbon dioxide gas is bubbled through the solvent for several minutes to ensure saturation. A suitable size sample of the compound being studied is weighed quantitatively into a thin glass capsule blown from 6 mm. soft glass tubing and
.
(1) C. N.Hinshelwood, J . Ckem. Sac., 117, 156 (1920). ( 2 ) G. A. Hall, Jr., J . A m . Chem. Soc., 71, 2691 (1949). (3) P. E.Yankwioh and R. Linn Belford, ibid., 76, 4178 (1953). (4) L.W.Clark, THEJOURNAL, 60,825(1956). (5) H. N. Barham and L. W. Clark, J . A m . Chem. Soc., 75, 4638 (1951).
Vol. 60
weighing approximately 0.2 g. The capsule is lodged on the loop inside the neck of the condenser and the condenser is fitted to the flask. The stirrer is started, and the tube leading from the top of the condenser is connected to the top of the buret through a short piece of rubber tubing, the entraining liquid in the leveling bulb and buret being set at zero volume. When the system reaches equilibrium, as shown by a steady state of the colored solution inside the buret, all variables-barometric pressure, water jacket temperature, volume, etc.-are recorded and the external end of the nichrome wire is twisted slightly to dislodge the capsule which falls into the rapidly stirred solvent initiating the reaction. The operator moves the leveling bulb by hand to maintain the inside and outside pressures equal a t all times, and records the volume, time and other variables at frequent intervals. That the experiments are carried out with a high degree of accuracy and precision is demonstrated by the fact that invariably the stoichiometric volume of carbon dioxjde is collected from every quantitative sample undergoing decarboxylation. Used in each experiment in the present study were 100 ml. of triethyl phosphate and a sample of malonic acid weighing 0.1856 g. (the amount required to furnish exactly 40.0 ml. of carbon dioxide at STP on complete reaction). The reaction was carried out a t seven different temperatures between 12S-15Oo. The experiment was often repeated two or three times a t the same temperature. A total of fifteen experiments were performed with excellent reproducibility.
Results and Discussion , The experimental data were converted to standard conditions and milliliters of evolved gas was plotted against time for each temperature. Values of x corresponding to different values of t were obtained from the resulting isotherms. Log (a - x) was then plotted against t (a is the theoretical stoichiometric volume of carbon dioxide, 40.0 ml.) . The points thus obtained for the middle 80% of the reaction fell on perfectly straight lines in every experiment. This fact indicates that the decomposition of malonic acid in triethyl phosphate is a first-order reaction. From the slopes of the lines thus obtained the specific reaction velocity constants for the decomposition of malonic acid in triethyl phosphate were calculated for the various temperatures on the basis of the equation for a first-order reaction. The temperatures studied, as well as the corresponding specific reaction velocity constants in set.-', were as follows: 128.3', 0.00203; 133.3", 0.00317; 139.1', 0.00506; 142.5', 0.0062; 144.6', 0.00816; 145.8", 0.00867; 149.4", 0.0107. When log k was plotted against 1/T according to the Arrhenius equation a straight line was produced. The energy of activation and the frequency factor, calculated from the slope of the line, are 26,900 cal. and 1.0 X 10l2, respectively. The temperature coefficient for the reaction is 2.5. Based on the Eyring equation, the enthalpy of activation is 26,100 cal.; the entropy of activation is -6.2 e.u.; the free energy of activation a t 140' is 28,600 cal. The specific reaction velocity constant a t 140' is 0.0052 set.-'. Some interesting conclusions can be drawn by comparing this reaction with data for the decomposition of molten malonic acid. According to the data of Hinshelwood,' the energy of activation for the thermal decomposition of molten malonic acid is 34,500 cal., the frequency factor is 1.8 X 1014, the temperature coefficient is 2.72, the enthalpy of activation is 33,000 cal., the entropy of activation is +2.5 e.u.; the free energy of activa-
Aug ., 1956
NOTES
tion a t 140' is 32,000 cal., and the rate constant a t 140" is 0.00024 sec.-l. Studies in this Laboratory are in good agreement with Hinshelwood's result^.^ It is thus seen that at 140" malonic acid decomposes in triethyl phosphate twenty-two times as fast as it does alone. We see that triethyl phosphate brings about a decrease in the entropy of activation, decreasing the probability of the formation of the activated state.6 However, this effect is compensated for by the lowering of the enthalpy of activation, so that less energy is required to bring about activation. The over-all effect is the lowering of the free energy of activation a t 140' from 32,000 to 28,600 cal., resulting in an increase in the rate of reaction. ( 6 ) S. Glasstone, K. J. Laidler and
H.Eyrinp, "The Theory of Rate
Processes," McGraw-Hill Book Co., New York, N. Y., 1941, p. 24.
POLAROGRAPHIC THIOUREAFORMALDEHYDE KINETIC STUDIES BYMARTING. CHASANOV AND CECILC. LYNCH Department of Chemzstry, University of Delaware, Newark, Del. Received February 69, 1966
The initial condensation of thiourea and formaldehyde in alkaline media was followed by polarographic analysis for formaldehyde. The reaction was found to be second order with respect to formaldehyde and thiourea. Values of the rate constant for the reaction were determined at 5, 15 and 25". The value for the activation energy for this reaction, determined from these rate constants, is 23.4 kcal./mole. Experimental The polarographic determinations were made with an E. H. Sargent and Company Model XI1 Polarograph. Electrodes of the original Heyrovsky-Shikata type' were employed. For the determination of the half-wave potential of formaldehyde versus the saturated calomel electrode (SCE) the H-cell of Lingane and Laitinen2was used. The formaldehyde employed was Merck and Company neutral reagent grade; analysis by the hydroxylamine methoda showed the formaldehyde content to be 36.5%; the solution was diluted to prepare a stock solution of molarity 0.01246 for use iii the polarographic measurements. The half-wave potential of the formaldehyde versus the mercury pool was -1.56 v.; versus the SCE, -1.66 v. The limiting current, ik, due to the formaldehyde was proportional to the formaldehyde concentration over the range studied. The thiourea was Eastman Kodak White Label. The thiourea crystals were dried a t 120" for two hours; volatile material was less than of 1%. The dried crystals were dissolved in water to make up a stock solution of molarity 0.02588. The supporting electrolyte was 0.05 M sodium hydroxide (Baker C.P.). The drop mass for the capillary was 1.48 mg./sec.; the drop time was 4.74 sec.; the mercury head was 34.6 em. For kinetic determinations solutions were prepared and thermostated prior to use. Formaldehyde and the supporting electrolyte were mixed f i s t ; the thiourea was then added. When half of the thiourea had been added the timer was started; the reaction mixture was agitated for 30 seconds, and then it was placed in the polarographic cell. The limiting current was measured as the difference between the upper and lower plateaus of the formaldehyde wave. (1) I. M. Kolthoff and J. J. Lingane, "Polarography," Interscience Publishers, N e w York,N. Y . , 1941. (2) J. J. Lingane and H. A . Laitinen, Ind. Eng. Chem., A n d . E d . , 11,
504 (1939). (3) L. E. Smythe, THISJOURNAL, 61, 369 (1947).
1151
The results of the kinetic determinations are shown in Table I. The data fitted a second-order rate plot for thiourea and formaldehyde. The values shown are the averages of three determinations for each concentration. The activation energy may be estimated from the Arrhenius relation d In k AEA
d(l/T)==-R
TABLE I THE REACTION OF FORMALDEHYDE
RATE CONSTANT FOR WITH THIOUREA I N STRONGLY ALKALINE MEDTAAT 5, 15 AND 25" Initial CH20 conon., M
Temp., OC.
25 25 15 15 5 5
415 208 312 208 415 375
x
Initial thiourea ooncn., M k(mo1e-1
431 X 132 205 273 547 656
10-6
.
sec.-1)
2.03 2.00 0.44 0.42 0.119 0.124
Discussion
Good yields of mono- and dimethylolthiourea have been produced by the reaction of thiourea and formaldehyde a t temperatures less than 5OoS4 Studies have shown that isomeric forms of the above exists. Pollak4 postulated the isomers of monomethylolthiourea as HzN-C-NH-CH20H
1
(acidic media)
A
HN=C-S-CHzOH
(alkaline media)
\
NHz
B
In alkaline media thiourea reacts to produce mercaptans, sulfur, and often HzS. The reaction to form B may be postulated as S
SH
H2N--Il-NH2
1
HN=C--NH2
(1)
ki + HzN-C=NH
(2)
tf
0
-NH2
+ H-C-HI1
1
S-CH~OH 0
SH
I
HN=C--NHZ 0
+ OH+ H-Il-H
kz
HzN-&-NH k-
(3)
0
0
H2N-LH2
+ HS-
HZN--Il-NHz
2
I
(4)
CHzOH
- -dCF = klCFCT + kzCFCIJ - k 2 C P dt
(a)
where CF is the formaldehyde concn. CT is the thiourea concn. CU Is the urea concn. Cp is the concn. of monomethylolurea
Crowe and Lynch5 have found a t 25" that ICz is 2.97 X mole-' see.-' and is 1.45 X (4) F. Pollak, Modern Plaslics, 16, No. 10, 45 (1939). (5) G. Crowe and C. C. Lynch, J . Am. Chem. Soc., 70, 3795 (1948).
