NOTES
1370
only important gaseous species and the evaporation coefficient was assumed to be unity., Akishin, et al.,3 have mass spectrometrically observed only Bl(g) a t 2100 OK., while Drowart and Verhaegen4bfound some Bz a t 2330°K. Their detectability limit for a given vapor species u as about 1%. The second law AH0298(s~bl) was found to be 133.4 f 5 kcal./mole while the third law calculatioii yielded AH0298(s~bl)= 136.9 f. 0.3 kcal./mole, in agreement with the second lam value. The equation for the sublimation pressure of B is 28,840 log Pat,(solid) = 7.239 - ___ (1781-2 152OK.) T The boiling point is estimated to be 4050 f 100°K. This estimate is obtained by use of the 136.9 kcal./mole for the third law heat of sublimation and extrapolated high temperature thermodynamic functions.10 The Laiigmuir sublimation pressure results are in close agreement with the results of Searcy and Myers, although their Knudsen data are slightly lower as might be expected if there was some reaction of the B samples with the Knudsen cells. The quaiititative results of Akishin, et uZ.,~ appear to have considerable scatter. Drowart and Verhaegendb feel that uncertainties in relative ionization cross sections and multiplier yields cannot account for the difference observed. They suggest that interactions between the Knudseii cell and the sample occur iii weight-loss experiments but do not explain how these same interactions are avoided in their work. Alcock and Grievesonll have also reported weight-loss data from graphite Knudsen cells in essential agreement with the work of Searcy and Myers.2 It appears that the mass spectrometer-Knudsen cell studies establish an upper limit for the sublimation pressure of boron and, further, that these Laiigmuir studies set a lower limit. The discrepancy could be explained if ae0,the Langmuir sublimation coefficient, is equal to 0.15 and if the ratios of hole size to sample area were somewhat larger in the work of Searcy and Myers2 and of R ~ b s o n . ~ The correction factor for channeling by the Ta-susceptor in the Laiigmuir experiments would also tend to lower the heat of sublimation, although it would appear that TaBz should form rapidly from the elements a t 1780-21OO'K. and eliminate the correction. Acknowledgments.-The authors are pleased to acknowledge the financial support of this work by the United States Atoniic Energy Commission and the Wisconsin Alumni Research Foundation. The sawples of boron were generously provided by Dr. Claude P. Talley of Texaco Experiment, Inc. (10) (a) S. R'ise, R. .4ltman, and J Margrave J . P h g s Chem., 64, 915 (1980); (b) R. McDonald and D. Stull, J . Chem. Eng. Data, 7, 84 (1982). (11) C . Alcock and P. Grieveson, in "Thermodynamics of Nuclear Xatelials," Intein. Atomic Energq Agency, Vienna, Austna, pp. 571-572.
THE THERXAL DECOMPOSITIOK OF CYCLOBCTANEI BY ROBERT W. CARR,J R . ,AND ~ W. D. WALTER~ Department of Chemistry of the Universitg of Rochester, Rochester, AT. Y . Recehed January 9,1968
Previous publication^^-^ have reported studies of the thermal decomposition of cyclobutane in the vapor
Vol. 67
phase over the temperature range 410-500° mainly a t pressures from 0,0005 to about 120 mni. At 449, 438, and 427' some work a t initial pressures as high as 740 nim. and slightly higher has been performed. The first-order rate constaiits were found to decrease noticeably a t pressures below about 20 mm. a t 449'. Very few experiments have been done a t pressures greater t,han 120 mm. where the first-order rate constants would be expected to be almost invariant with pressure if the decomposition is a quasi-unimolecular process. The present experiments at pressures up to 1500 mm. were performed in order to study more carefully the firstorder region of the reaction and to obtain values for the Arrhenius parameters a t pressures well above the onset of the fall-off. Moreover, if the rate constants should continue to increase as higher pressures are attained, the earlier conclusion that the decomposition is a simple unimolecular reaction (with no detectable free radical chain processes) would be opeii to question. On the other hand, if the rate constants should decrease with increasing pressure, a possibility which has been mentioned by Wilsonj8 the measurements inight provide interesting iiiformation about intramolecular energy transfer. Experimental The cyclobutane had been prepared earlier in this Laboratory by Genaux through the photolysis of cyclopentanone and had been purified in a Podbielniak distillation apparatus. In the present work gas chromatographic analysis with a 3-meter column cont,aining tetraisobutylene upon 60-80 mesh firebrick a t 28" showed the purity to be a t least 99.9%. The Phillips research grade ethylene used in this study was subjected to degassing, drying, trap-to-trap distillation, and a gas chromatographic check of purity. Kinetic experiments were performed in two different thickwalled spherical Pyrex vessels. One had a capacity of 27.2 ml. and the other, which was filled with thin-walled Pyrex tubes to give B 9-fold increase in S / V ratio, had a volume of 21.2 ml. The reaction vessel was placed in a spherical cavity inside a cylindrical aluminum block (7.5 cm. diam.) which just fit the iron core of an electrically heated furnace. Capillary tubing (1 mm. i.d.) connected the vessel to the rest of the vacuum system. Temperature measurements were made by the use of two standardized platinum, platinum-l3% rhodium thermocouple8 which lay tangent to the outside of the react'ion vessel. The temperature was controlled aut,omatically to about 1 0 . 1 " . For experiments a t pressures less than 1 at,m. the initial pressure was measured with a mercury manometer. Initial pressures greater than 1 atm. were calculated from the measured amounts of reactant introduced. I t was found that the undecomposed cyclobutane and the product ethylene could be separated satisfactorily on the basis of the difference in volatility a t - 130". From t,he amounts of ethylene and unreacted cyclobutane measured in a gas buret, the percentages of decomposition were obtained and first-order rate constants calculated. I n some experiments a portion of the reaction mixture was also analyzed mass spectrometrically and the results from the two methods of an.alysis were in good agreement.
Results Products.-To ascertain whether other products are formed a t the higher pressures, the reaction mixtures (1) This work was supported b y a grant from the National Science Foundation to the University of Rochester. (2) National Science Foundation Summer Fellon, 1960. (3) C . T.Genaux and W. D. Walters, J . Am. Chem. Soc., 73, 4497 (1951). (4) F. Kern and ?T. D. Walters, Proc. Natl. Acad. Sci. U . S., 38, 937 (1952). (5) C. T. Genaux, F. Kern, and W. D. Waltera, J . Am. Chem. Soc.. 76, 6196 (1953). ( 6 ) H. 0. Pritchard, R. G. Sowden, and A. F. Trotman-Dickeneon, Proc. Row. Soc. (London), A218, 416 (1953). (7) D. F. S>F-inehartand R. W. Vreeland, Paper presented at the 138th National Meeting of the American Cheniical Society, Division of Physical Chemistry, 1960; R. W. Vreeland, Ph.D. Thesis, University of Oregon, 1961. (8) D. J. Wilson, J . PICYE. Chem., 64, 323 (1960).
June, 1963 from several experiments a t about 1500 mm. (25-35% decomposition, 427-460') were separated into two portions, one of which was volatile a t -- 130' and the other non-volatile a t that temperature. Gas chromatographic analyses of the separated fractions showed the presence of only two substances, one with a retention time corresponding to ethylene and the other with the retention time of cyclobutane. The higher boiling fraction from a 1350 mm. experiment carried to 35% reaction at 460' was analyzed by mass spectrometry and infrared spectroscopy. The gas phase infrared spectrum was in good agreement with one found by Genauxg for pure cyclobutane. Comparison of the mass spectrum with that of pure cyclobutane showed that the cracking patterns were in good agreement, with all peaks accounted for except small ones a t m/e = 42, 43, and 44, and two in the Cs range. The substances contributing to these are probably no greater than 0.1 to 0.2% of the undecomposed cyclobutane. The separated fractions from a 732 mm. experiment after 25% decomposition a t 449' in the packed bulb were analyzed on the mass spectrometer. I n the fraction non-volatile a t -130' only cyclobutane was detected, but the fraction volatile a t -130' showed small peaks a t m/e = 40, 41, and 44. Although the substance giving the peaks was not identified, it probably did not exceed 0.1% of the ethylene. I n obher experiments near 1500 mm. (25% decomp., 420-449') measurements iii the gas buret showed that the initial amount of cyclobutane intxoduced into the vessel was equal, within experimental error, to the amount recovered plus one half of the amount of ethylene formed. A11 of the evidence indicates that the decomposition at two atmospheres proceeds essentially by the same reaction (CIHB + 2C2€I4)Jasthat observed earlier a t lower pressures. Kinetics.-The rate data for the decomposition of cyclobutane near 449' are presented in Table I. In the fourth column are the first-order rate constants corrected to 449O, the temperature dependence amounting to about 6% per 'C. An inspection of the values from 100 to 1500 mm. reveals that there is no noticeable trend with pressure, and in fact, the average deviation over the 50-fold pressure range investigated is only 2.3%. The present values are in reasonable agreement with those obtained earlier in this Laboratory and when plotted against pressure, they join quite smoothly to the pressure-dependent values in the region below 30 mm. Since the decomposition of cyclobutane a t 449' is somewhat, end other mi^,^ it is conceivable that the ternperature of the reacting gas a t the center of the vessel might be appreciably lower than near the malls. By the use of an equationlo taking into account only conduction of heat, it was calculated that for an initial pressure of 2 atm. at 449' the maximum temperature difference in the 27 ml. vessel might be -2.8' (unless lowered by another heat transfer process), but for a pressure of 100 mm. it mould not exceed 0.2'. As the data in Table I show, the rate constants from experiments in the packed vessel are in accord with those from the unpacked vessel. It was felt that the presence of packing would afford zi test for the occurrence of heterogeneous processes and would provide for a (9) C. T . Genaux, Master's Thesis, University of Rochester, 1953. (10) 5. W. Benson, "The Foundations of Chemical IGnetics," MaGrewHill Book Co., New York, N. Y.,1960, pp. 426-431.
1371
NOTES
1.38
-.,
1.40 108/T.
1.42
1.44
Fig. 1.-Temperature dependence of the first-order rate con733 mm., packed vessel; 0,1250stant: @, 650-790 mm.; 1500 mm.; 9,1490 mm., packed vessel; a,1320 mm., analysis on mass spectrometer.
TABLE I FIRST-ORDER RATECONSTARTS FOR THE THERMAL DECOMPOS[TION OF CYCLOBUTAEE NEAR449" Init. press., mm.
Temp., OC.
k X 101, sec. -1
k (cor.) X
104,O
see. -1
33.8 44.9.1 5.20 5.17 69.9 44-9.0 5.38 5.38 71.9 448.9 5.09 5.14 102 449.2 5.37 5.31 123" 449.0 5.24 5.24 126 449.0 5.40 5.40 136 449.1 5.34 5.31 243' 449.0 5.08 5.09 245 449.1 5.01 5.00 245 449.1 5.27 5.25 306 450.4 5.56 5.10 660 4510.1 5.17 4.85 673 44'3 . 4 5.09 4.97 733b 449.1 5.12 5.10 772 450.0 5.53 5.18 1260 450.0 5.35 5.04 1320' 449.1 5.32 5.29 14906 449.2 5.07 5.01 ' Percentage of decomposition determined by mass spectrome-. try. Experiments in the packed vessel. For the corrections, temperatures and rate constants with one more significant figure than those shown were used in the calculations and then rounded Off.
greater temperature uniformity throughout the vessel since the thermal conduct,ivity of glass is about lo2times that of cyclobutane. The fact that no observable effect on the rate was found in the packed vessel gives an indication that heterogeneous reactions and temperature gradients do not occur to a kinetically important extent under the conditions of this study. The temperature dependence of the reaction was studied over the temperature range 419-460' a t pres-
XOTES
1372
sures near one atmosphere and near two atmospheres. Since there appeared to be no significant change in the value of the rate constant in the present work resulting either from an increase in pressure from 650 to 1500 mm. or from packing the vessel, the resulta from all experiments in the packed and unpacked vesse!s over the pressure region 650-1500 mm. were used for determination of the activation energy. The experimental data, plotted as log IC us 1/T in Fig. 1,give a straight line, from the slope of which an activation energy of 62.5 kcal./mole was calculated. In addition, the activation energy was calculated by the method of least squares on an IBM 650 computer and found to be 62.5 f 0.4 kcal./mole. That the uncertainty in the value of the activation energy is probably in the neighborhood of 0.4-0.5 kcal./ mole is indicated not only from the least squares deviation but also from the fact that when the experiments shown in Fig. 1 were divided into groups on the basis of essentially constant initial pressure, the values of the activation energy (determined from the smaller number of experiments) ranged from 62.1 to 62.9 kcal./mole. For each experiment in Fig. 1 the pre-exponential factor A was calculated from the relationship k = A exp(-62,50O/RT). On the basis of the average value, which was 4.2 x l O I 5 sec.-l with an average deviation the rate expression of 0.1 X k = 4.2
f
Vol. 67
higher pressures, but more favorable conditions for testing the suggestion of Wilsoiis might be obtained by investigating the uniniolecular reaction of a more complex molecule in which the bond undergoing scission had to receive the needed energy from degrees of freedom some distance from it and only weakly coupled to it. This question has been considered also by Flowers and Frey,ll who studied the isomerizat'ion of 1,l-dimethylcyclopropane at pressures from 16 to 1596 mm. a t 460'. The first-order rate constants were invariant with pressure over this one hundredfold range which extended upward from the region where the high pressure limit is first approached. They concluded t,hat the high pressure rate constant for the unimolecular reaction of 1,l-dimethylcyclopropanedoes not go through a maximum. Acknowledgment.-The authors wish to thank Mr. Car: Whiteman, Jr., for his assistance in making the infrared absorption measurements and the least squares calculations. (11) >I. C. Flowers and H. 14. Frey, J . Phys. Chem., 66, 373 ( 1 9 6 1 ) .
FLUORIXE Y.M.R. SPECTROSCOPY. XIII. FURTHER COUPLING CONSTASTS FOR ISOTOPIC C3Fs BY GEORGEVAN DYKETIERS
0.1 X l O I 5 exp(-62,5OOR/T) set.-' (1)
will satisfactorily represent the data in the pressure region 650-1500 mm. Discussion Since the present study has shown that at pressures as high as 1500 mm. the pyrolysis of cyclobutane is a homogeneous reaction uncomplicated by side reactions and that the first-order constants a t 449' have not changed detectably over the pressure range 240-1500 mm., the rate expression given in equation 1 for the pressure region 650-1500 mm. should correspond rather closely to the rate expression for the high pressure limit of the unimolecular decomposition of cyclobutane (ICm). The values for k , and E , have been estimated earlier by Vreeland and Swinehart7 from their rate constants in the fall-off region (initial pressures, 20 nim. and below) by a sizable extrapolation of a l l k us. 1/P plot to a value of 0 for 1/P. From experiments a t 410-500' 35 exp(-61,800/RT) set.-' the relationship IC, = was obtained. They found, however, that a value k , 846 exp( -63,20O/RT) sec.-l would give Kassel fall-off curves in better agreement with their experimental data in the pressure dependent region. Either of these values of the activation energy could be regarded as in reasonable agreement with the value determined in this work. On the other hand it is noteworthy that for lc, a t 449' the former expression gives a value of 4.45 x 10W4see.-', but the latter expression gives 5.25 x lo-* sec.-l vhich is closer to the average value of k (5.12 x l o w 4sec.-l) for the experiments in the pressure region 700-1500 mm. a t 449' in Table I. The rate expression found earlier5 for experiments a t 120 mm. was k = 4.0 X 1015exp(-62,500/RT) sec.-l which is within 5% of that given by equation 1. I n the current investigation no evidence has been found that the first-order rate constants decrease with increasing initial pressure. S t would be of interest to measure the rate of decomposition of cyclobutane a t
Contribution X o . X69 f r o m the Central Research Department, Minnesota Mining and M f g . Company, St. Paul 19,Minnesota Received January 14, 195'3
I n a previous paper1 coupling constants were reported for the CF3 group in perfluoropropane. A t the time it was not possible to make analogous measurements upon the weaker and broader carbon-13 satellites of the CF2 peak; however, subsequent improvements in resolution and sensitivity have made such studies feasible. With the present paper, all but t,wo of the possible coupling constants for C3F8are nom reported, albeit no information on relative signs for the J-values is yet available. Experimental The sample, equipment and techniques have been described and shielding values have been reported.' Coupling constant's and isotope shifts, measured on the weak, partially resolved CI3 satellites to the CF2 peak, have been obtained with acceptable precision by averaging the results of twelve separate determinations on each peak. The results are given in Table I; error values are standard deviations for the averaged values. Firstorder spin-spin analysis was used throughout. TABLE I FLCORINE S.M.R. COUPLING CONSTANTS AND ISOTOPE EFFECTS FOR THE CF2 GROUPIN PERFLUOROPROPAXE Coupling system
J. e./aec.
Std. dev a
.
A+,
Std.
p.p.m.b
$0.125 f0.002 zkO.2 266.6 C13Fz $0.024 =t0.002 zk0.2 32.6 C'3CFz The isotope shift, A+, is defined a Std. dev. of the average. as +(C13 isomer) - +(ClZ isomer). A typographical error is present in footnote c, Table I1 of ref. 1.
Discussion The "direct" isotope shifts, A$(C13F2),are very similar to several previously reported for non-allylic CF3 groups,1,2but are substantially smaller than those for allylic CFB groups,3 for CF&l group^,^ and for CFClt (1) G. V. D. Tiers, J . Phys. Chem., 66, 945 (1962). (2) G. V. D. Tiers, J . Phys. Soc. J a p a n , 16, 354 (1960). (3) G , V. B. Tiers, J. Chem. Phya., 36, 2263 (1961).
