ISOTOPE EFFECT IN THE DECOMPOSITION OF OXALICACID
Oct., i952
897
THE ISOTOPE EFFECT IN THE DECOMPOSITION OF OXALIC BY ARTHURF R Y 3 AND MELVINCALVIN Radiation Laboratory and Department of Chemistry, University of Califarnia, Berkeley, Cdifmka Received March 6, is63
The C13 and CI4 isotope effects in the decomposition of oxalic acid have been studied simultaneously at two different temperatures. In each case the C14 effect is ap roximately double the ClBeffect, and both are in reasonable accord with theoretical calculations based on a simple postufated mechanism for the reaction. ing sulfuric to 150 cc. of ordinar concentrated sulfuric acid in a glass stop ered flask. $he solubon waa shaken Lindsay, McElcheran and Thode4 have studied vigorously and afowed to stand overnight. A sample of 1 the C1aisotope effect in the decomposition of oxalic cc., wt. = 1.8214 g., waa diluted with water and upon titraacid in concentrated sulfuric acid at 100'. The tion with standard 1 M sodium hydroxide required 36.92 cc. for neutralization, giving a molecular weight of 97.5 (99.4% results were reported as ratios of rate constants of theory). This same sulfuric acid stock waa used in all for the equations the decomposition experiments. Decomposition Experiments.-The decompositions were ki carried out using 99.4% sulfuric acid in the apparatus shown (C0OH)z +COz CO Hz0 (1) in Fig. 1. Approximately one-gram samples of oxalic acid were placed in the decomposition chamber, B. The aparatus was assembled, using the 99.4% sulfuric acid to ubricate the joints and stopcocks. Ten CC. of the 99.4% sulfuric acid was pipetted into the preheater, C, and dr carbon dioxide-free oxygen was allowed to sweep througg bOOH k3 the system for about one hour. The gas from the outlet of I -+ COZ HzO (3) the assembly shown in Fig. 1 was conducted,, in turn, COOH through a spiral trap cooled by a Dry Ice-isopropyl alcohol The ratio kn/ka, where C* indicates C13, was found mixture, a spiral bubbler filled with 2 M sodium hydroxide to remove the carbon dioxide, a copper oxide-packed furnace ka) was 1.034. to be 1.032, while kl/(kz at -700" to burn the carbon monoxide to carbon dioxide, Similar experiments have now been carried out and another sodium hydroxide filled bubbler to absorb this using oxalic acid-1,2-C214, and both the CIa and the carbon dioxide. When the oxygen had been sweeping C14 isotope effects have been determined on the through the apparatus for about an hour, the spiral bubblers Wed with 2 M sodium hydroxide, and the liquid in same samples. In addition, the decompositions were flask A was heated to reflux by means of a Glas-col mantle. have been carried out a t two different temperatures, The vapor of the boiling liquid served as a constant tem-
Introduction
+
+
P
+ 60 +
+
and the differences in heat and entropy of activation for decomposition by equations (2) and (3) have been calculated for both CL3 and C14.
WATER CONDENSER
Experimental Preparation of Oxalic A~id-l,Z-C:~.-The oxalic acidl,2-C1+ used was prepared by the reaction between carbon dioxide-CI4 and potassium on sand at 360O.6 The product from this reaction was purified by alternate calcium oxalate precipitations and continuous ether extractions. In the final step of the purification the ether extract was evaporated over a small amount of water. The solution was filtered and the water waa removed by lyophilization, leaving a pure white product of anhydrous oxalic acid-1 ,2-Ci4. Yields were from 45-55% with an additional 15-20% of the activity being recovered undiluted as carbon dioxide. The oxalic acid was checked for radioactive impurities by paper strip chromatography using a mixture of t-amyl alcohol and benzene saturated with 3 N hydrochloric acid as the solvent. No radioactive spots wwe found except for the oxalic acid spot, and an upper limit of 0.5% was set for the amount of radioactive impurity. An inactive sample prepared by the same procedure gave theoretical equivalent weight values when titrated with either base or permanganate. 'For use in the decomposition experiments, a sample of the oxalic acid-l,2-Ci4 was diluted with specially purified inactive oxalic acid by dissolving the two samples together, filtering and removing the water by lyophilization. Preparation of -100% Sulfuric Acid.-The -100% sulfuric acid used was prepared by adding 100 cc. of 30% fum-
x) COLD TRAP
THERMOMETER
-.
(1) The work described in this paper was sponsored by the U. 8. Atomic Energy Commission. (2) This paper was abstracted from part of the thesis submitted by Arthur Fry to the Graduate Division of the University of California in partial fulfillment of the requirements for the Ph.D. degree, June, 1951. (3) Department of Chemistry, University of Arkansas, Fayetteville, Arkansas. (4) J. G . Lindsay. D. E. McElcheran and H. G . Thode. J. Chem. Phg8.. 17,589 (1949). (5) F. A. Long, J. Am. Chern. Soc., 61, 570 (1939).
BOILING SOLVENT
Fig. 1.-Apparatus
for constant temperature decomposition of oxalic acid.
1
ARTHURFRYA N D MELVINCALVIN
898
perature bath to heat the sulfuric acid and oxalic acid to the desired reaction temperature. The boiling point of the liquid in flask A determined the temperature of decomposition. Dioxane and carbon tetrachloride were chosen to give temperatures of 103.0 and 80.1°, respectively. When the liquid in flask A had been refluxing vigorously for a proximately one-half hour, the stopcock was opened and t f e sulfuric acid was allowed to flow onto the oxalic acid. At 103.0° the reaction was very vigorous and care had to be exercised in adding the sulfuric acid to avoid frothing. In all caaes the sulfuric acid was added as rapidly as possible. The oxygen sweep served to stir the oxalic acid-sulfuric acid mixture. The sweep was continued until gas bubbles were no longer observed in the sulfuric acid solution, and then for an additional hour and one-half. The carbonate in the two bubblers was collected and weighed as barium carbonate. The yields of barium carbonate from the carbon dioxide and carbon monoxide were quantitative, and the blanks were negligible. Isotooic Comoosition Measurements.-The C1' activitv measuriments were made using an ionization chamber and vibrating reed electrometer connected to a Brown recorder. The ionization chambers were filled to a standard pressure with carbon dioxide generated in a vacuum system from the barium carbonate with concentrated sulfuric acid. A sample of this same carbon dioxide was stored in a bulb for later mass spectrometric analysis. The samples from the carbon monoxide and carbon dioxide from a given run were measured in immediate succession in the same ionization chamber and on the same instrument in order to reduce to B minimum-any variations in the procedure and instruments. Several independent measurements were made on each sample from each decomposition, starting with the barium carbonate in each case. The C1*measurementsewere made on a Consolidated Type 21-102 mass spectrometer, using a magnetic sweep to scan the peaks a t m/q = 44 and 45. Several independent scannings of each sample were made. I n several cases, the entire low m/q mass spectrum was taken using a voltage scan, and no peaks were found except those of the normal carbon dioxide pattern. Samples of normal carbon dioxide were also run duriqg the sample determinations to check the reproducibility of the mass spectrometer. The m / q = 45 peak height was corrected for the 0 1 ' content of the carbon dioxide.
Vol. 56
and the indicated errors are average deviations of these measurements. The C1* measurements were made at two different times, rum 1-4 and normal carbon dioxide 1 and 2 (carbon dioxide from Dry Ice) being measured at one time and runs 5-8 and normal carbon diomde 3 and 4 (same stock sample of carbon dioxide) being measured somewhat later. Since in the subsequent calculations, only the relative C*Q fractions are used, the variation in normal carbon dioxide value between the two times is not serious since the difference is reflected in the values from both the carbon monoxide and carbon dioxide. The C14 measurements were made using different instruments and a t dEerent times which probably accounts for the variations in specific activity from run to run. However, the carbon dioxide and carbon monoxide fractions from the same run were measured in immediate succession on the same instrument to reduce this variation aa much as ossible within a given run. I n the following w h a t i o n s only the monohbeled oxalic acid is considered although the oxalic acid actually used is called oxalic acid-l,2-Ci4, since by its method of preparation both carbon atoms would be labeled if ure C1402were used. Actually, the number of dilabeled mogcules was cdculated to be a very negligible fraction of the total. Referring to equations ( l ) , (2) and (3) and assuming that the reaction is first order with respect to oxalic acid, the specific activities or mole fractions of the carbon dioxide and carbon monoxide are given by the following equations, where Ox refers to oxalic acid containing only C12 and Ox* refers to oxalic acid containing C1a or C1*
c*o2 c*o*+ COI
=
and
c*o c*o + co = Combining the above equations gives equation ( 4 )
+
I
kz
z
+
(4) c*o2c*02 cot c * oc*oco = The C1402 specific activities and C1302mole fractions ob- I t should be noted that k2/k3 calculated in this manner is intained in the various runs are shown in Table I. I n all dependent of time, and hence of the amount of reaction, cases, the values given are averages of several measurements, while comparing the specific activity of the carbon dioxide or carbon monoxide to that of the oxalic acid is not. For ease of tabulation and convenience of reference, the TABLEI percentage isotope effect may be defined as 100 (kz/ks - 1). C1402 SPECIFICACTIVITIESAND C1802 MOLE FRACTIONSThe percentage isotope effect values calculated by equation (4)from the data in Table I are 'ven in Table 11. The FROM THE DECOMPOSITION OF OXALIC Ac1D-1,2-Ck4 data of Lindsay, McElcheran and qhode4 have been recalCl! culated in this manner, and show a much narrower spread Mo!e C'aC1a02" Speoifio aotivityb fraction ClzOz + C1302 Drift rate, volts/min. HzCnO4 than using their method of calculation, although the average
Results
Decomposition Temp., "C. Run No. 1 103.0 2 Normal COI 1
3 4 Normal Cor 2 Permanganate oxidation 5 80.1 6 ormal Con 3
7 8 Normal COa 4
x 106 x CO: C? coz Fraction Fraction Fraction 1089 & 6 1065 f 3 4585 f 40 1092 f 0 1058 f 2 4535 f 39
C? Fraction 4336 29 4277 f 37
*
1115 rt 1 1088 + 2 1059 4 4554 f 39 4328 65 1091 & 3 1062 f 5 4504 f 27 4286 f 29
*
*
1116 f 1 1074 f 2 4432 f 30 1114 f 1 1076 f 3 4507 rf2 21 4203 f 36 1110 f 2 1078 i 6 4636 f 43 4335 31
*
1137 f 3 1107 f 3 1073 f 3 4531 f 11 4247 f 20 1110 2 1075 f 3 4614 f 17 4352 f 24 1186 f 1
Corrected for presence of C12O1'016. activity = -30 dis./min./mg. BaCOs. 0
10s
b
Actual specific
(6) Thanks are due to Dr. A. Newton and Dr. L. Tohnan of this Laboratory for these measureiuents.
C13 AND
TABLE I1 C14 ISOTOPE EFFECTIN THE DECOMPOSITION OF OXALICACID
Decomposition Run No.
Temp., OC.
Per oent. isotope effect 100. C'*
(kz/ka
- 1)
Cld
1 103.0 2 . 3 5.7 3.2 6.0 2 2.7 5.2 3 2.7 5.1 4 Averugc 2.7 f 0 . 2 2 5 . 5 fO.30 5 80.1 3 . 5 7.2 3.0 6.9 6 3.2 6.7 7. Average 3.25 & 0 . 1 5 6 . 7 f .35 Lindsay, McElcheran and Thode" 100 2.9 1 3.5 2 3.2 Average Rccalculntcd froiii I h c origiiial dula using cyuetioii (4).
Oct., 1952
ISU~WIJN EFFNCT IN THE DXCOMPOSITION OF OXALICh c ~ u
899
molecule and either a symmetrical product of the ratedetermining step or products which are rapidly interconvertible compared to the irreversible folDiscussion lowing reactions, in which the isotope selection is The C I 3 isotope effect values shown in Table I1 made. If an equilibrium process occurs before the are thus in very satisfactory agreement with the re- rate-determining step, an isotope effect in the equisults obtained by Lindsay, McElcheran and T h ~ d e . ~librium could multiply an isotope effect in the rateThe C14 effect is approximately double the CI3 determining step, Such a multiplication of effects effect at tach temperature. could conceivably lead to either a larger or smaller The differences in heat and entropy of activa- overall effect, depending on whether.the equilibrium tion caused by the isotopic substitution can be cal- constant was greater or less than one. Several culated from the theory of absolute rates.' such pre-equilibria might be involved in some For equation (2), and temperature TI cases. In the present case a reasonable equilibrium kT I ( k 2 ) T I = -' e A S a / R e - A H : / R T prior to the rate determining step might be the h association of a hydrogen ion with one of the and for equation (3), and TI carbonyl groups of the acid as shown in equation (7) *
values are tho sttiiic. These rccalcultttod values are also shown in Table 11.
*
Combining
A simple calculation of the equilibrium constant, K , may be made using the equations derived by In a similar manner an equation is set up for T2, and Bigeleisen and Mayerg or by Urey.1° For lack of when this equation is combined with (5), equation complete data, we have assigned the frequency of ( 6 ) results 1750 ern.-' to the C=O vibration, 1100 em.-' to the normal C-0 vibration and 1500 cm.-' to vibrations which are hybridized the two C-OH by resonance. The isotopic shifts are then calFrom the data in Table 11, (AH? - AH?) can be culated using the approximation of simple harmonic calculated by equation (B), and this value can be oscillators. Using these values the equilibrium substituted back into equation (5), and (As? - constant at 100" is calculated to be 0.988 for C I 4 As?) can be calculated. The calculated values are substitution and 0.995 fDr C13 substitution. Thus this particular equilibrium decreases the calculated shown in Table 111. isotope effect, but it is conceivable that in other TABLE I11 cases the effect would be in the opposite direction. In calculating the ratio kz/ka, the above vibraHEATAND ENTROPY OF ACTIVATION DIFFERENCE BETWEEN tional frequencies are assigned to the normal EQUATIONS (2) AND (3) molecules (A) and (B). For the transition state - ( A H 2 - AH,), (AS*? - A s % ) , Isotope cal./mole e.u. we take a model in which water has been lost from the carbon originally attached to two OH groups C 18 61 -0.11 and the carbon-carbon bond is greatly weakened C'4 131 -0.24 leaving nearly free molecules of carbon monoxide The presence of the isotope effect measured here and WOOH. The carbon monoxide is assigned allows some rather specific conclusions to be drawn its accepted frequency of 2170 cm.-l and the concerning the mechanism of the decomposition.8 carbonyl and hydroxyl frequencies in eCOOH At complete reaction, all of the carbon-carbon are assumed to be unchanged. The force constant bonds, in all the molecules, are broken, so the iso- for the C-OH bond which is broken is set equal tope effect observed cannot be due to isotopic dis- to zero. The ratio kz/ks is then calculated using crimination in the rate of rupture of the carbon- equation (8), as derived by Bigeleisen." carbon bonds. Therefore, if the observed isotopic fractionation is due to an isotope effect in the ratedetermining step of the reaction, the ratedetermining step (8) must be the breaking of a carbon-oxygen bond, The function G(u) is defined by Bigeleisen and since the association of hydrogen with oxygen is Mayerg who have tabulated values of G(u) as a commonly considered to be rapid and reversible. However, the observed isotope effect might also function of u. The reaction coordinate reduced be due to some equilibrium step. An isotope ef- mass, p , is assumed to be the reduced mass of the fect is conceivable in rapid steps following the rate- atoms forming the bond being broken. The determining step only under very special circum- symbol =!= refers to the activated complex. The rate ratio, kz/ka, calculated in this manner, stances. These require a symmetrical starting when multiplied by the equilibrium constants (7) S. Glasstone, K. J. Laidler and H. Eyring, "The Theory of Rate calculated above, give over-all isotope effects of Processes," 1st ed., McGraw-Hill Book Co., Inc., New York, N. Y., 1941. (8) The authors wish to cxpress their thanks to Prof. Richard ~ i stLc i ~ decoiupuaitiun. i I'owcll for a fruitful discuaaioii uf the ~ n e c l ~ u ~ of
(9) J. Bigeleisen and M. G. Mayer, J . Chem. Phys., 15, 201 (1947). (10) 11. C . Urcy, 6.Chem. Soc., 562 (1047). (11) J. Bigeleisen, J , Chem. Phus., I T , G71, (1949).
ARTHURFRYAND MELVINCALVIN
900
6.0% for C14 substitution and 3.1% for Cl2 substitution a.t 100'. These compare to the values of 5.5% and 2.7% measured experimentally. The corresponding calculated values at 80" are 6.3 and 3.4% compared to the experimental values of 6.7 and 3.25%.12 The agreement in magnitude is quite satisfactory, but the calculated temperature coefficient is not large enough to explain that observed experimentally. Perhaps a better choice of frequencies and model would lead to a more satisfactory calculation. Ultimately the study of isotope effects in chemical reactions should enable us to elucidate ~ ~ of 1 1the intimate details of the mechanism of the reaction. When the proper equilibria and rate determining steps are chosen, the calculated isotope effect should coincide with that observed experimentally. The choice of a molecule which reacts a t one of two chemically Chart 1.-Mechanism of the Decomposition of Oxalic Acid in 100% Sulfuric Acid /OH
Vol. 56
identical groups, such as oxalic or malonic acid is a fortunate one, since net isotope effects are observed at complete reaction, thus making it unnecessary to study the isotopic composition of the products as a function of the amount of reaction. The presence of an isotope effect does not unequivocally establish the mechanism of the decomposition, but it does offer very strongevidence that the rupture of the carbon-carbon bond is not rate determining, and, consequently, that the ratedetermining step in this medium is the rupture of the carbon-oxygen bond. Actually, the rupture of the carbon-oxygen bond and of the carbon-arbon bond is probably all part of a concerted reaction with the loss of hydroxyl (or water) from one carboxyl or the other being the initiating step. The essential steps of what is considered to be the most likely mechanism for the decomposition are shown in Chart I. Path A is seen to be less likely than Path B, since we observe an enrichment of labelled atoms in the carbon dioxide.
\O
+ +
OH
H O '
K
@
H20
+
H20
Path A
Path B
+ >+. 0
H
.. C t 0 I
(12) We have recently learned, by private communication, of experiments in two other places (Research Chemistry Division, Atomic Energy Project, National Research Council, Chalk River, Canada: and Department of Chemistry, University of Illinois) in which the isotope effect in a chemical reaction was determined for both CIS and CI4 simultaneously on the same reaction. In both of these cases the reported C" effect is much more than twice the Cia effect. The reaults for the decomposition of meaitoic acid given by Stevens, Pepper and Lounsbury (Canada), J . Chrm. Phys., 20, 192 (1952). are as follows: k i d h i : 1.038 f 0.003; kit/kic 1.101 Et 0.005. Here the Cl4/CI' ratio is almost three. The reaults reported by Yankwich, Stivers and Nystrom (Illinois) at the Berkeley Meeting of the American Physical Society in December, 1951, and mentioned in the Abstract appearing iq the Bulletin of the American Physical Society, S I , No. 8, Paper B5, give values for both the CIS and Cl4 effects correaponding with those previously reported. Here the C" effect is about four timea the size of that for C". This latter report was for the deaarboxylation of malonic acid at its melting point. It is clear that from these three comparative studies, if results are accepted on their face value, that there are isotope factors involved in the determination of the rate of these decarboxylation reactions of which we are not yet cognizant. since it is diffiault to see how a change in mass from CIS to C14 could change the mechanism of the reaction, unless perhaps nuclear spin were involved in Borne way.
DISCUSSION J. BIGELEISEN (Brookhaven Nat. Lab.).-It is quite interesting to note that Fry and Calvin 6nd the quantity (&/IC8 1)C1*,C1*/(kp/ks - 1)ClZ,C**equal to 2.0 within the limits of exDerimenta1 error. This is in accord with the theoretical expeEtations aa discussed in my paper earlier in the symposium. I n footnote (12) it is su gested that nuclear spin effects may be responsible for the jeviation of this uantity in the work of Stevens, Pepper and Lounsbury an% Yankwich, Stivers and Nystrom. It has been pointed out several times previously in the present symposium that no effects due to nuclear spin can be observed in chemical reactions involving the isotopes of carbon and carried out a t temperatures above 10'K. It is improper to calculate an effect of isotopic substitution on the entropy of activation from equation (5). From equation (8)or e uation ( 5 ) of m paper ( J . Chem. Phys., 17,675 (1949)) oneffnds that the Jfference m entropies of activation should be calculated after the "reduced mass effect" has been factored out. E uation (5) in the paper by Fr.y and Calvin should have tge factor ( M ~ / M Z ) ' multiplied /* into the right-hand side. I am preparing a paper for publication which discusses the effect of isotopic subptitution on the entropy, enthalpy and heat capacity of equilibrium systems and the similar quantities in reaction rates. There will be extensive tables for the calculation of these quantities from molecular data. It is interesting to note, in connection with the calculation of k&, that the major contribution comes from the factor (Ms/&f2)'/*. Am I correct in assuming that the frequency shifts are calculated by assuming that Slater's theorem applies to all the vibrations? Slater has shown this to be approximately true for just one frequency, but not all. In fact, if it were assumed to be true for all frequencies, the product rule would be violated. A similar objection can be raised to the acid-base equilibrium calculation. The result in this case is close to what is predicted for acid-base isotopic exchan e equilibria (cf. ref. 9) and is quite insensitive to the a%solute value of ZG(u)Au for &her the acid or base pair.
-
A. Fm.-The fre uency shifts for those frequencies which are not assumed to%e unchanged are calculated using the
Oct., 1052
Cl4
ISOTOPE E F F E C T IN T H E DECARBOXYLATION
approximation that the bond may be considered as a simple harmonic oscillator between the two atomR forming the bond. This is probably no worse an a proximation than others involved, mch as the arbitrarily cRosen frequency assignments. The calculated and experimental values are both considcrably larger than the reduced mass term by itself. P. E. YANKWICH (University of Illinois).-I believe that we all agree that the statistical rate theory cannot permit a ratio of C14 and C1*intramolecular isotope effects which is sensibly greater than 2. It should be pointed out, however, that within this limitation we are still permitted considerable latitude in the choice of a detailed model, and, as Dr. Fry has shown, it is possible to construct reasonable models which are not in agreement with either the C18or C14experiments on some compounds. One of the compounds which has been studied by several groups of workers is unsubstituted malonic acid. There is apparent agreement on the values for the CIS isoto e effects Mr. and no agreement on the values obtained with E. C. Stivers and I have completed recently redeterminations of the intramolecular isotope effects in the decarboxylation of C1* and C14 labeled malonic and bromomalonic acids. The results for C14(expressed in the notation introduced by the Chairman) are: 100(k& - 1) = 9.9 f 0.6 (malonic) and 11.6 f 0.6 (bromomalonic); the corresponding results with C18 are 2.8 i 0.3 and 2.4 f 0.4. The stable isotope
814.
THE
OF
MALONIC ACIDS
901
results with bromomalonic acid leave much to be desired, but the very much greater than 2 experimental ratio of C14 to C'a isotope effects is shown clearly. In the matter of agreement among various workers on the magnitudes of these effects, it is interesting to note that H. G. Thode and his co-workers report C'a intramolecular isotope effects which differ by almost 30% for two different samples of commercial malonic acid. The sample which gave the higher effect was analyzed carbon-by-carbon and found to be isotopically homogeneous. With the permission of the Chairman and of Dr. Thode, I would like to ask Dr. Thode if one of his reported values is to be favored over the other, or if the difference between them can be attributed to experimental error.
H. G. THoDE.-The difference in the C1*intramolecular isotope effect for two commercial Sam les of malonic acid reported by us (Lindsay, Bourns and bhode) and referred to by Dr. Yankwich, we believe is real. As suggested in our paper, this difference could be explained by a depletion of only 0.5% in the C1* content of the methylene carbon of the British Drug House malonic acid. We would favor the result obtained with Eastman Kodak Co. malonic acid since the C18 content of the methylene carbon waa checked in this case and found to be very nearly the same as that for the COZ from the completely oxidized malonic acid.
ISOTOPE EFFECT IN THE DECARBOXYLATION OF CY-NAPHTHYLAND PHENYLMALONIC ACIDS'j2
C14
By ARTHUR FRY^ AND MELVINCALVIN Radiation Laboratory and Department of Chemistry, University of California, Berkeley, California Rsceived March 6 , 1968
The isotope effects in the decarboxylation of a-na hthylmalonic a~id-1-C'~ and phenylmalonic acid-l-CI4 have been measured, both in solution and on the liquid acids near tfeir melting points. The carboxyl group containing CIS is lost as carbon dioxide about 10% more fre uently than is the C'ccontaining carboxyl group. Some aspects of the theoretical calculations of intramolecular isotope e8ects are considered. Neither the models considered here nor the previously proposed models of Bigeleisen give agreement between theory and experiment for both the C'8 and 0 4 cases.
Introduction on malonic acid, both experimentally6-8 and In recent years the question of the effect of the theoretically.9-l The deaarboxylation of a malonic acid may be substitution of isotopic carbon on the rate of chemical reactions has received considerable attention. represented by the following equations where C* A considerable portion of this attention has been represents CI3or C14 directed toward determining the isotope effect in k.1 RCH(COOH)1+ COz + RCHzCOOH (1) the decarboxylation of malonic acids. Yankwich and Calvin4 studied the decarboxylation of malonic C*OOH kz RCH( acid-1-C14 and bromomalonic acid-l-CL4,and found +c*oz+ RCH~COOH (2) COOH rupture ratios for C1z-C12bonds to C14-C12 bonds C*OOH k3 of 1.12 f 0.03 and 1.41 f 0.08, respectively. The RCH< +coz RCH,C*OOH (3) authors mentioned that the bromomalonic acid case COOH was a single experiment on acid of doubtful purity.6 Since this time considerable work has been done For the labeled malonic acid molecule the rupture (1) The work described in this paper was sponsored by the U. 8. ratio C1z-C12/C1eC12is given by the ratio lcz/ks. Atomic Energy Commission. In the terminology of Lindsay, Bourns and Thode,' (2) This paper was abstracted from the thesis submitted by Arthur is an intramolecular isotope effect. this Fry to the Graduate Division of the University of California in partial
+
fulfillment of the requirements for the Ph.D. degree, June, 1951. (3) Department of Chemistry, University of Arkansan, Fayettoville, Arkansaa. (4) P. E. Yankwich and M. Calvin, J . Chum. Phys., 17. 109 (1949). (5) The bromomalonic acid case has recently been checked and the ratio found to be around 1.1 rather than 1.4; P. E. Yankwich. E. C. Stivers and R. F. Nystrom, J . C h s n . Phya., 20,344 (1952).
(6) J. Bigeleisen and L. Friedman, ibid., 17, 998 (1849). (7) J. G. Lindsay, A. N. Bourns and H. C. Thode, Con. J . C h s n , 28, 192 (1951). (8)A. Roe and M . Hellmanr~,J . Chum. Phya., 19, €360 (1951). (9) J. Bigeleisen. ibid., 17,426 (1949). (10) K. Pitrer, ibid., 17, 1341 (1949). (11) A. A, Bothner-By and J. Bieeleisen. ibid.. 19, 755 (19Sl),
