J. ERSKINE HAWKINS AND JAMES W. VOGH
902 11
VOl. 57
A first power law would be obtained by assuming the following reaction to take place H+RCONHa+
+ Hs0 +H+RCOOH + NHI+
The experimental rate constants, not considering medium effects, should then be given by 10 I
5 E
If Ka is small, kslpt1 is proportional to the first power of the hydrogen ion concentration, which is 2M the case up to about 2 N hydrochloric acid for the three heterocyclic amides. Any deviations can 2 then be ascribed to medium effects. k2 in equation 9 1 is a rate constant and K 3an equilibrium constant pertaining to the dissociation of the amide.2J A detailed discussion of mechanisms will be reserved until more data are available for dilute solutions, the main purpose of the present paper being to ascertain the dependence of the experimental 8 rate constants and the parameters of the Arrhen20 21 22 23 24 ius equation on hydrochloric acid in concentrated Energy of activation in kcal. Fig. 2.-Log A/[HCI] values as a function of correspond- solutions. The three heterocyclic amides do not show very ing energies of activation for picolinamide (A) and isonicopronounced differences amongst themselves. Thk tinamide (B). energies of activation for picolinamide are higher A mechanism for the hydrolysis of nicotinamide than those for the other two amides. It may also was presented in a previous paper.2 However, it be mentioned that the energies of activation for the is difficult to decide whether such a mechanism is three amides plotted against the logarithms of really operative because of medium effects. The their experimental rate constants for any one hyexperimental results indicate that a t low hydro- drochloric acid concentration do not give straight chloric acid concentrations the experimental rate lines with slopes of magnitude 2.303RT1 as was the constants are proportional to the first power of the case for other amides.l acid concentration and not to the second power as (7) 11. H. a. Jellinek and M. G. Wayne, THISJOURNAL, 111, 173 the mechanism, proposed previously, would require. (1951).
THE RATE OF THE THERMAL ISOMERIZATION OF 0-PINENE I N THE VAPOR PHASE1 BYJ. ERSKINE HAWKINS AND JAMES W. VOGH Contribution jrom the Department of Chemistry, University of Florida Received March 80. 1968
The rate of the thermal isomerization of &pinene in the vapor phase has been determined in the range 350-400". Sufficient variation in pressure, temperature and extent of decom osition was obtained to calculate the order of the reaction and the energies of activation for the formation of m cene and dimonene and for the decomposition of 8-pinene. The rate constants for the first-order reactions are q u a y t o 10".'e-"*w 'ET, 1016.5e-458"JO/RTand 1017.4 e-4*,800/RT min.-l, respectively. Two minor products of the reaction were found but have not been identified. A discussion of the mechanism of the reaction is included.
The rate of the thermal isomerization of p-pinene in the liquid phase has been reported by Hunt and Hawkins.2 It was shown that the presence of quinoline, hydroquinone and dipentene diluent had no effect upon the pyrolysis reactions. Myrcene polymers and I-limonene were the products. The activation energies for the formation of myrcene and E-limonene were found to be 47 and 50 kcal./mole, respectively, in the temperature range 220 to 235'. Other thermal isomerizations of @-pinenein the va(1) The material included in this paper is based upon a partial abstract of a dissertation presented to the Graduate Council of the University of Florida by James W. Vogh in partial fulfillment of the requirements for the degree of Doctor of Philosophy. June, 1953. (2) H. G. Hunt and J. E. Hawkins, J . A m . Chcm. Soc., 7 2 , 5618 (1950).
por phase have been carried out but no rates of the reactions were measured. The first work of this type was that of A r b u z o ~who , ~ stated, apparently erroneously, that P-pinene partially isomerized to alloocimene a t 345 to 350". Goldblatt and Palkin4J and Savich and Goldblatte have shown the pyrolysis products to be 1-limonene, myrcene and myrcene polymers and have discussed Arbuzov's work.8 The conditions they used were 5 seconds (3) B. A. Arbuzov, J . Om. Chem. (U.5.8.R.),6,297 (1936). (4) L. A. Goldblatt and 9. Palkin, J . A m . Chem. SOC.,63, 3517
(1941).
(5) L. A. Goldblatt and 8. Palkin, U. 8. Patent 2,420,131 (May 6, 1947). (6) T. R. Savich and L. A. Goldblatt, U. 8. Patent 2,507,546 ( M a y 16, 1950).
,
RATEOF THERMAL ISOMERIZATION OF PINENE
Dec., 1953
contact time at 40506 and 0.004 second contact time at 725 t o 750O.B At the higher temperature the yield was 85% myrcene and 0.8% polymer. Nothing else was reported. This virtual absence of polymer correlates with the small contact time, which does not permit the formation of the high boiling compounds, a-camphorene or other polymer~.~ Experimental Preparation of Materials.-Commercial I-8-pinene was distilled in a Lecky and Ewe11 column of about 60 plates under the conditions of operation. The @-pinene used in these studies was distilled from a single batch in order to ensure uniformity of the physical properties of the pinene and the pyrolysis products. The constants for the 8-pinene were: b.p. 59.5" (20 mm.), n26~1.4768, d% 0.8669 and -21.35. Myrcene was prepared according to the procedure of Goldblatt and Palkin.4 Approximately three liters of carefully distilled @-pinene was pyrolyzed. Myrcene was obtained from the product of distillation and the constants were: b.p. 65.5" (20 mm.), T Z ~ ~1.4682, D d% 0.7921, and [a126q 0. l-Limonene was repared by distillation of the same pyrolysis mixture an{ the constants found for it were: b.p. 71.5' (20 mm.), , 2 s ~ 1.4710, d26d 0.8385 and [cx].'~D -107.91. Separate redistillation of the fractions obtained prior to the myrcene and after the limonene produced two more components. Neither of these compounds was identified. Apparatus and Procedure.-All experimental runs reported in this paper were carried out in a pyrolysis train of the usual type. The pyrolysis tube consisted of four sections of 1.25-cm. Pyrex tubing, each 36 cm. in length. These were connected by narrow-bore U-bends, about 4mm. inside diameter. This design was found satisfactory since replacement of the upper half of the first section with a narrow tube had no effect on the calculated reaction rates. The volume determined for the pyrolysis tube was corrected for the calculated expansion at the operating temperature. The correction was approximately 0.7 cc. for the total volume of 200 cc. The tube was heated to the operating temperature in an electrically heRted thermostat which contained a fused mixture of LiNOa, 27.3 wt. %; NaNOs, 18.2 wt. % and KNOa, 54.5 W t . %. Three iron-constantan thermocouples were placed in the bath: one near the bottom, one near the top and one halfway down. Temperature control was attained by the use of a thermoregulator of the design given by Benedict.7 A nickel resistance thermometer was used to operate the thermoregulator. When in use the bath was found to have a maximum difference in temperature throughout the whole length of less than 0.2" at all times. The products of the reaction were discarded through a side arm until a steady rate was obtained. The drop rate into the receiver was noted in order to ascertain that the steady rate of production of pyrolysate was maintained. At the completion of the timed runs the samples were weighed and stored a t 0" under nitrogen. Analysis of Pyrolysate.-The determination of the rate of pyrolysis was carried out assuming that the amounts of materials other than 8-pinene, myrcene and I-limonene present in the product could be neglected. This was done since the complete decomposition of @-pineneunder the conditions of the experiment produced no more than 0.5 wt. % of material other than myrcene and Z-limonene. Density and optical rotation measurements were made for analysis of the product using the relations
+ [%I = [a.Pba + 1 = w@+ + WL
lid,
= wddg
Waa/dM
$. W d d L
f 0.0044 WSWM (1)
[(YLIWL
(2)
'WY (3) The symbols have the definitions: d = density, [a] = specific rotation a t 25O, w = weight fraction. The subscripts p, 8, L and M refer to the pyrolysis mixture, 6-pinene, I-limonene and myrcene, respectively. The relations were based upon the density and optical rotation values for
(7) M. Benedict, Rev. Sci. Instruments, 8. 252 (1937).
903
8-pinene-myrcene mixtures and Z-limonene-myrcene mixtures. In a supplementary test of the validity of this analytical method i t was found that the average discrepancy between the known and calculated values of weight fractions for the three components of three mixtures was 0.002 unit. These known mixtures were prepared to correspond to the compositions of typical pyrolysis products. There appeared to be no polymers in the pyrolysis roduct. Successive simple distillations of 30 cc. of fresf pyrolysis samples left less than 0.1 g. of residuein each case.
Results and Discussion The constants found for the component boiling at 78" (20 mm.) were: n Z 61.4836, ~ nZo 1.4862, ~ dZ54 0.8557, dZozo0.8612, [ a l Z 5-0.35. ~ Iskenderov* reported the following values for 1(7),8(9)p-menthadiene prepared by dehydration of Aat9dihydroperillic alcohol: b.p. 65-66' (1 1 mm.), dZo2,, 0.8735, nzoD 1.4870. When this b.p. is corrected to 20 mm. pressure by the method of Hassgit becomes 77.3-78.3". The mechanism (Fig. 1) for the thermal isomerization of 0-pinene permits the formation of l(7),8(9)-menthadiene. However, the infrared spectrum of the compound isolated in this work does not indicate its presence. lo Another unidentified component was obtained in a mixture with 0-pinene at 59-60" (20 mm.). Although only a partial separation could be made, the physical constants were estimated to be: b.p. 59-60' (20 mm.), n Z 5 D1.4661, 0.833, [Q1Iz5D0. Since the combined amount of the unidentified substances was less than 0.5% of the products, their structures were not investigated further. Table I contains the data of the several experimental runs. TABLE 1 CONDITIONS OF PYROLYSIS A N D RATECONSTANTS FOR FORMATION OF MYRCENE AND E-LIMONENE
Contact
Experiinent no.
Temp., OC.
Pressure, mm.
time,
min.
rain.-'
1 2 3 4 5 6 7 8 9 10 11 12 13
350.1 350.1 364.8 364.8 378.7 378.5 378.5 378.4 395.0 395.5 395.1 409.5 409.0
35.1 80.3 20.3 50.4 35.7 36.3 50.3 80.4 20.4 35.5 50.5 35.3 20.5
0.304 .482 .1519 .1732 .0617 .0457 .1444 .1143 .0736 .0508 .0581 .0397 .0438
1.72 1.62 4.17 4.08 9.72 9.69 9.55 9.34 24.58 25.59 24.56 53.0 48.7
+M,
min.
0.35 0.33 0.79 0.78 1.70 1.72 1.70 1.63 3.97 4.11 3.95 8.0 7.4
First-order reaction rate constants for the reactions were calculated according t o the equations kdeo
kM = kL
2'$3
=
-log W B
(4)
1
wM'wL +WM/WL
(5)
kdeo
kdeo
- kM
(6)
The symbols kdea, AM and AL are the rate constants for the decomposition of @pinene, the formation of myrcene and the formation of Llimonene, respec(8) M. A. Iskenderov, J . Gen. Chsm. ( U S S R ) , 1 , 1435 (1937). (9) H. B. Hass, J . Chem. Education. 13,490 (1936). (10) R. L. Webb and J. P. Bain, J . Am. Chem. SOC.,76.4279 (1953).
J. ERSKINE HAWKINS AND JAMES W. VOGH
904
tively. W M and WL ere the weight fractions of myrcene and I-limonene, respectively. The contact time is t, minutes, and the other symbols have meanings previously designated. Table I contains the rate constants for the formation of myrcene and 1-limonene. When log le's for the decomposition of @-pinene, for the formation of myrcene, and for the formation of E-limonene are plotted against 103/T, straight lines result. Activation energies for the decomposition of 0pinene and for the formation of myrcene and 1limonene were determined from the slope of these plots and were found to be 48,800,49,900 and 45,100 cal./mole, respectively. The dependence of the rate of these reactions may be expressed by kdec = 1017.4e-24600/2' min.-l kM = 1017.7e-26100/2' min.-l kL 1016.3e-Z27W/T d n . - l
(7)
VOl. 57
ous and first order and to be unaffected by the presence of propene or nitric oxide. This reaction is analogous to the breaking of the cyclobutane ring in p-pinene during the formation of myrcene. The formation of the menthadienes from 0-pinene is best explained by an intramolecular exchange of hydrogen as suggested by Burwell.11 The possibility that there is an exchange of hydrogen between the biradical and a stable molecule and a resulting complex mechanism is lessened by the fact that the reaction is first order and the rate constant is unaffected by a change of pressure within the range in which this work was carried out. On the basis of the proposed mechanism the following steps may be set up for the pyrolysis reactions where B, R, L and M are the symbols for ppinene, the biradical shown in Fig. 1, I-limonene, and myrcene, respectively. kl
B+R
The above relations indicate that the rate of the reaction is independent of pressure, contact time and amount of ,&pinene decomposed within the ranges covered. The preferred mechanism for this reaction in the vapor phase is the diradical scheme proposed by Burwellll which is shown in Fig. 1. This mechanism is in agreement with the findings of Hunt and Hawkins12 that the liquid phase pyrolysis of ppinene was not affected by the presence of quinoline, hydroquinone or dipentene.
6
l -B - Pinene
\1
kz
R+B
I
TI
/-Limonene
(9)
ka
R+L
(10)
k4
R-+M
(11)
The corresponding reactions have been observed in the pyrolysis of a-pinene by Fuguitt and Hawkins.14 The rate constant for the racemization of apinene was shown to be about one-tenth that of the sum of the rate constants for the formation of dipentene and alloocimene and the activation energy of the racemization to be 1.5 kcal. greater than that for alloocimene formation. If the same relation of rate constants and activation energies occurs in the pyrolytic reactions of 0-pinene the reaction R + B may be omitted without introducing a great error. Appropriate treatment of the rate equations and the energies of activation of the reactions in 8, 9,10 and 11 lead t o the conclusion that Edse =
I
(8)
where El is the activation energy for reaction 8. Since WM/WL is not independent of temperature, neither ELnor E M can be expected to be independent of temperature. However, temperature dependence could not be observed in the data available. EBand Ed,the energies of activation for the reactions in equations 10 and 11, could not be determined from the information available. However, the difference may be shown to be EM - E L = Ea
l(i%B@,-p-Menfhadien e
Myr ce n e
6
I-B-Pinene
Fig,
1.-The mechanism of the thermal isomerization of ppinene in the vapor phase.
The pyrolysis of cyclobutane to ethylene by Genaux and W a l t e r ~was ' ~ found to be homogene(11) R. L. Burwell, J . A m . Chem. Soc., 73,4461 (1951). (12) H.G. Hunt and J. E. Hawkins, %'bid.,72, 5618 (1950). (13) C. T. Genaux and W. D. Walters, ibid., 73, 4497 (1951).
(12)
E1
- E3
(13)
Reaction 8 represents dissociation of the 6-8 bond. The activation energy of this reaction can be taken as equivalent t o the dissociation energy of the 6-8 bond of the p-pinene with the resulting formation of the biradical. No direct comparison with reported dissociation energies can be made. Sehon and Szwarcl5 found a bond dissociation energy of 61.5 kcal. for the 3-4 bond of 1-butene. The difference between this and the experimentally (14) R. E. Fuguitt and J. E. Hawkins, ibid., 69, 319 (1947).
(15) A. H. Sehon and & Szwarc, 'I. Proc. Rog. Xoc. (London). A202, 263 (1950).
Dec., 1953
LIQUID-VAPOR EQUILIBRIUM IN
URANIUM HEXAFLUORIDE-HYDROGEN FLUORIDE 905 behavior and products observed is that of Hunt.I6 While the possibility of the occurrence of a chain reaction mechanism, particulady in the formation of the menthadienes, cannot be excluded on the basis of the information available, several possible chain mechanisms were analyzed but none fitted the experimental observations in a satisfactory manner.
determined activation of 48.8 kcal. may be attributed to the strained nature of the cyclobutene ring of the p-pinene, to the hyperconjugation effect due to the two methyl groups on carbon 8, to the fact that the distance between the unpaired electrons of the biradical is restricted by the molecular dimensions of the biradical, and to the fact that the two resonance forms shown are not symmetrical. The only mechanism other than that suggested by Burwell which completely agrees with the kinetic
(16) H. G . Hunt, “The Kinetics of the Thermal Isomerization of the Pinenes,” Ph.D. Dissertation, Department of Chemistry, University of Florida, June, 1950, p. 32.
LIQUID-VAPOR EQUILIBRIUM IN THE SYSTEM URANIUM HEXAFLUORIDE-HYDROGEN FLUORIDE1 BY ROGERL. JARRY,FREDD. ROSEN,CHARLEY F. HALEAND WALLACE DAVIS,JR. Contribution from the K-86 Laboratory Division, Carbide and Carbon Chemicals Company, Oak Ridge, Tennessee Received April 4, 1066
Liquid-vapor equilibrium in the system uranium hexatuoride-hydrogen fluoride has been determined over the whole range of compositions in the temperature interval 40 to 105 In this temperature range, and considerably beyond, the system is one of maximum pressure at constant temperature. Vapor-liquid separation factors, aetivity coefficient8 and other pertinent functions have been calculated with respect to a formula weight of hydrogen fluoride of 20.01. This procedure is convenient and practical, but it is not in agreement with molecular weights of hydrogen fluoride calculated from vapor density measurements. These weights vary from 60 to 100 in the temperature range 60 to 40” for saturated vapor in contact with solutions.
.
In a recent publication2 the nature of condensed phase equilibria of the binary system uranium hexafluoride-hydrogen fluoride was presented. The purpose of the present paper is to submit data on liquid-vapor equilibrium. Two factors should be noted concerning this system, as follows: first, a miscibility gap extends from 61.2 to 101 O; second, since the manner in which the molecular weight of hydrogen fluoride varies with temperature, pressure and concentration is not known, the system cannot be expressed in terms of molar quantities. Formula quantities have been used in this paper for convenience, although such units introduce very extensive distortions in graphical representation of this binary system.
Experimental Equilibrium Still.-The equilibrium still used to obtain data for this paper has been described in a previous communication,S as has the vapor pump.4 Briefly, i t consisted of a pot, an overhead vapor column that could be isolated from the pot, and a pump to force vapor through the liquid. This whole unit, with its pressure transmitter, was contained in one side of a two-compartment air thermostat; the other compartment contained the sampling manifold. One portion of this unit not previously described in adequate manner is the liquid pipet. It was formed by major modeling of two Hoke M 342 valves which were welded to the bottom of the pot, as shown in Fig. 1. The still was controlled at one of .eight temperatures by mercury regulator-electronic relay circuits. Temperatures were measured by means of a calibrated co er-constantan thermocouple and a Leeds and Northrup Fortable Potentiometer. Vapor and liquid temperatures were maintained equal within 0.1 to 0.2”. (1) This work was performed for the U. 8.Atomic Energy Commission by Union Carbide and Carbon Corporation at Oak Ridge, Tennessee. (2) G. P. Rutledge, R. L. Jarry and W. Davis, Jr., THISJOURNAL, 57,541 (1953). (3) R. L. Jarry and W. Davis, Jr., ibid., 57, 600 (1953). (4) F. D. Rosen, Reo. Sci. Instruments, 24, in press (1953).
Pressures were measured by use of a Booth-Cromer pressure transmitter6 and one of two Bourdon gages, each of which was calibrated against a dead weight gage.’ One, used in most of the experiments, had a 7-in. diameter scale of 0 to 160 p.s.i.g. in 2 pound increments; the second had a 4-in. diameter scale of 0 to 300 p.s.i.g. in 1 pound increments. Gas Density Flask.-This flask was made of spun nickel, about 10 mils thick, and had a volume of about 535 cc. With its light weight, right angle, metal valve (monel metal body, Z nickel stem, modified from Hoke type 431) this unit weighed less than 200 g. It waa connected to the sampling manifold through a I / l in. compression fitting in which the copper gasket was replaced by one made of Fluorothene. Materials.-Uranium hexafluoride used in these experiments contained less than 0.015 wt. % impurity,a the impurity being principally hydrogen fluoride. Hydrogen fluoride, whose purification has been described p r e v i ~ u s l y , ~ was about 99.74 mole yo pure. Procedure.-The pot (A, Fig. 1, ref. 3) was charged with a mixture of hydrogen fluoride and uranium hexafluoride, and then heated to one of the 8 different tem eratures used in this work. Vapor was circulated throughP the solution while the materials were heating and until constant temperature and pressure had been obtained. The pump was then stopped and the vapor and liquid materials valved o f f from each other. A sample of the vapor was obtained by cracking a valve between the vapor volume and the sampling manifold. The pressure and temperature of this sample as expanded into the gas density flask was taken, the flask closed, and then, along with a counterpoise, removed from the thermostat for weighing. A sample of the solution was obtained by use of the solution pipet of Fig. 1 . After collecting liquid between valves V-1 and V-2, valve V-1 was closed and valve V-2 rapidly and completely opened. By this operation all the liquid in the pipet was completely vaporized, without fractionation, into a part of the gas sampling manifold (Fig. 1, ref. 3) which included a re-evacuated and retared vapor density flask. It should be noted that the counterpoise was taken through the same heating and cooling cycles as W M the gas density flask. In three experiments the complete vapor sample was condensed from the isolated vapor volume into a weighing flask (5) 8. Cromer, “The Eleotronio Pressure Transmitter and SelfBalancing Relay,” MDDC-803 (1947). (6) W. B. Kay, J . A m . Chsm. SOC.,69, 1273 (1947).
