PHOTOLYSIS OF DIISOPROPYL KETONE
Sept., 1956
1315
ISOPROPYL RADICAL REACTIONS. I. PHOTOLYSIS OF DIISOPROPYL KETONE BY
c. A. HELLER.4ND ALVINs. GORDON
Chemistry Division, U.S . Naval Ordnance Test Station, China Lake, Calif. Received March 50, 1066
Diisopropyl ketone has been photolyzed using a medium pressure mercury lamp over the temperature range 100-400". Pressure has been varied from 40-100 mm. and light intensity varied about tenfold. The kinetics are explained by a primary break of the ketone to give isopropyl radicals. These radicals can dis roportionate, combine, or abstract hydrogen C3H6 (1); iso-C!H7 iso-C8H7+C6H14(2); iso-CaH7. from the parent ketone: iso-C3H7 iso-C~H7+C3H8 RH C3H8 R (3). The activation energy difference between disproportionation (b.eaction 1) and combination (reaction 2) is less than one kcal. per mole and the ratio of rate constants k l / k Z = 0.6 a t 200 . The activation energy difference between abstraction and combination E3 ' / ~ E= Z 8.5 f 0.1 kcal. per mole (error as standard deviation). Above 200' the ketonyl radicals decompose measurably into C & , , CO and isopropyl giving a chain reaction. A t 350" and above the isopropyl radicals decompose measurably into C3H6 and H atoms.
+
+
+
+
+
-
Introduction Diisopropyl ketone was chosen as a convenient source of isopropyl radicals. The primary step of the photolysis has been previously shown to be a break into radicals with no hydrogen shift such as occurs when a hydrogen on a carbon gamma to the carbonyl group is present.'-3 The CO quantum yield is unity from 50 to 150" and over a wide range of wave lengths.3 Isopropyl radicals generated from isobutyraldehyde, azoisopropane6 and diisopropylmercury have been studied previously. The thermal stability of diisopropyl ketone recommended it for work over a wide temperature range. Experimental The diisopropyl ketone (The Matheson Company, Inc., Pract. Grade) was stored over drierite and then distilled through a 30-plate column under 700 mm. of nitrogen. A small portion coming off at 122" was retained. This was vacuum distilled into a bulb a t -78" and subjected to evacuation a t this temperature. A small methanol impurity was removed by evacuation a t -78". This sample was stored in the darkat -78" since decomposition was shown to occur a t room temperature in diffuse sunlight. The original light source was a Hanovia 2537 lam but this produced slight photolysis of the pro ylene. @herefore, a Hanovia quartz Alpine sun burner, g-100, was used in conjunction with a 2 mm. thick Corning 7910 filter. Light intensity was varied by varying the distance from the lamp to reaction vessel. The cylindrical reaction vessel was 56 mm. i.d. and 66 mm. long with 2.75 mm. thick windows. The reaction vessel was set into an aluminum block furnace with double quartz windows so t,hat the temperature gradient along the vessel was less than ' / t o . A t low temperatures and high light intensity the oven changed temperature by a maximum of 1.5' during a run. Temperatures were measured with a chromel-alumel thermocouple. The vessel was evacuated by an oil diffusion pump. Gold leaf traps further protected the reaction vesse from mercury vapor. Pressure was measured by a spoon type Bourdon gage which was operated against a mercury manometer and measured by a Gaertner micrometer telescope. The glass tubing to the oven was electrically heated by nichrome ribbon. The products of reaction were Toepler pumped through two Ward-Leroy stills into a sample flask. Four thermal cuts were made with the two still8 a t -155 and -175", -78 arid - 130°,0 and -45", and 0 and -20", respectively. A . J. C. Nicholson, Trans. Faradag Soc., 60, 1067 (1954). ( 2 ) C. R. Masson, J . A m . Chem. Soc., 74, 4731 (1952). (3) S. G. Whiteway and C. R. Masson, ibid., 77, 1508 (1955). ( 4 ) F. E. Blacet and J. C. Calvert. ibid.,73, 661 (1951). (5) R . W. Durham and E. W. R. Steacie, Can. J . Chem., 31, 377 (1953). (6) H. T. J. Chilton and R. G . Gowenlock, T m n n . Fnradng S o c . , 49, 1451 (1053). (1)
Gases non-condensable a t - 175' were measured by a McLeod gage before they were pumped into a sample flask. The volumes of all the fractions were similarly measured, and then analyzed by mass spectrometer (Consolidated 21-103). Each fraction contained contaminants from the others and analysis of each fraction was necessary.' Surprisingly, hydrogen was present in the second fraction and occasionally in the third fraction. The first fract,ion included Hz, CO, CHI and CZ hydrocarbons. The second, C3 and C4 hydrocarbons, plus small amounts of Cz and Hz. The third fraction contained the parent ketone, c6, Cg and C3 hydrocarbons and isobutyraldehyde. Mass spectra of the cg indicated that it was 2,3-dimethylbutane (diisopropyl) to within 5% for all runs. It was specially checked for 2-methylpentane. Attempts were made to identify peaks due to large molecules other than diisopropyl ketone. A small peak a t 156 was found in one case and is attributed to propyl hexyl ketone. A peak a t 112 from high temperature runs is probably propyl propenyl ketone, and one a t 100, while doubtful, due to mercury peaks interference, may be propyl ethyl ketone. All these higher boiling products were dissolved in the large excess of the parent ketone and even qualitative separation for identification was difficult.
Products.-The following major products were measured: CO, C3Hs, C3H6 and C6H14. In addition, Hz, CH4, CZH4 and CzH6 were. present in easily measured quantities in the high temperature runs, and C2Hz,ethyl propyl ketone, propyl propenyl ketone, butanes and butenes were present in small quantities in the low temperature runs. A peak a t 72 which occurred a t all temperatures was assigned to isobutyraldehyde, although it might have been due to pentanes. Primary Reaction.-The only primary photochemical reaction which occurs in measurable amounts is the split ipto radicals.
+
C ~ H ~ C O C J H photon ~
+CaH7CO + C3H7
(i)
The isobutyryl radical appears to be stable enough to form trace quantities of isobutyraldehyde, but most of it decomposes into CO and isopropyl radicals. Secondary Reactions.-The following secondary reactions account for most of the products found in the analysis C3H7 .CaH7 --+ C3Hs CaHe (1) .C3H7 +Ce"4 .CaH7 (2)
+
.GH7 .CaH7
+
+
+ C ~ H ~ C O C ~+ H I C3H.5 + C~HTCOC.(C;H~)?
+ CaH7COCaH7 + CaH8 + C,H7COCH(CHa)CHz C3H7CO. --+ .C3H7 + CO 2C3H7COC3H6.---+ diketone
(34 (3b) (4) (5)
C. A. HELLER AND ALVINS. GORDON
1316
VOl. 60
sion permitted no measurement of activation energy difference. The abstraction reaction, which is first order in isopropyl radical concentration can be compared with the combination which is second order in isopropyl radical concentration by the ratio of rate constants ka/kz'/:
.CaH7
+ CaH7COCH(CI&)CHe. + Ca& + CaH7COC(C&) = CHz
.
.
(14)
Low Temperature Mechanism.-The most important reactions at 100' are combination and disproportionation. The ratio of the rate constants for these two reactions is given by h / k r = CoHs/Cd%r
(1)
As can be seen in Fig. 1, these reactions have almost equal activation energies. If the activation energy of the combination reaction E2 = 0 then El = 900 cal./mole.
\
t
0.la 1.4
0
RUNS AT 0.1LIGHT INTENSITY
t
t
00%
100*C,
250%
1.6
1.8
2.0
a
2.2
1000 T
2.4
2.6
2.8
Fig. 1.-Arrhenius plot of the ratio of rate constants Cor some of the products of the photolysis of diisopropyl
ketone.
The k l / k z ratios found in this work disagree with those found using azois~propane.~I n the azoisopropane work the ratio was 0.5 a t 30" and 0.37 a t 121". Thus there are unexplained differences in both the values and in the slope of the Arrhenius plot. The values of the ratios from isobutyraldehyde4 are scattered and cover the sets of values of both of the above data. In the pyrolysis of diisopropyl mercury' between 228.5 and 441 " the ratios were scattered between 0.2 and 0.6 and poor preci(7) A. Bywater and E. W. R. Steaoie, J . Chsm. Phys., 19, 319 (1951).
(Ca&
- CaHs)/CsH,,'/z[RH]
(11)
Lowering the light intensity favors the first-order reaction and permits a more precise measure of the propane formed by abstraction. The present data do not tell which hydrogen is abstracted from the parent ketone. If reaction8 3a and 3b have ditrerent activation energies the observed. activation energy E3, should increase with temperature. To check this two least squares calculations were made: one from 100 to 251' and one from 192 t o 398". The former gave E3 = 8.1 f kcal./mole, and the latter 9.1 f 0.2 kcal./mole. These two values indicate a change in the relative importance of 3a and 3b but are not the actual values of either 3a or 3b. It is of interest t o calculate the ratio of steric factors of abstraction and combination to see how this compares with that for methyl and ethyl radicals. Using the over-all value of 8.5 kcal./mole gave P~~U~/€'Z%TZ as 0.78 X 10-lo at 250". Collision {iameters were chosen as u3 = 3.0 8. and u2 = 2.1 A. which give P3/P2'/*= 1.8 X This is similar to the values found for other radicals.* The ketonyl radicals which are formed by the abstraction of hydrogen from the isopropyl ketone appear to be stable at low temperatures. The quantum yield of one for CO yield3 is one indication of stability. Presumably these ketoriyl radicals disappear by combination reactions such as reactions 5 and 6. The production of hydrogen and the lighter hydrocarbons in small amounts at low temperatures is not explained by the mechanism. It is possible that they are due to heterogeneous decomposition of the radicals. The 72 peak found in trace amounts a t low temperatures may be due t o isobutyraldehyde which is formed from isobutyryl radicals abstracting hydrogen before they decompose via reaction 4. Eigh Temperature Mechanism.-As the temperature is raised above 100" other reactions become more important although there is always measurable contribution from the combination and disproportionation. Above 200" there is a rapid increase in the propylene/dimethylbutane ratio and a parallel increase in the CO rate at constant light intensity. This is explained by the decomposition of the ketony1 radical (reaction 7) which gives both propylene and CO as products and which gives an isopropyl radical as a chain carrier. The propylene/dimethylbutane ratio increases at lower light intensity since the dimethylbutane is formed by a secondorder reaction while the ketonyl formation and decomposition are both first order in radical concentration. (8) E. W . R. Steacie, "Atom and Free Radical Reactions." Second Ed., ACS Monograph No. 125, Reinhold Publ. Corp., New York. N. Y.,1954.
PHOTOLYSIS OF DIISOPROPYL KETONE
Sept., 1956 m@a
a m
0 0 3
0 0
g;?
E2
. .
1317
e Propane added, 0.00567 mole 1.-' Calculated as iaobutene. c Rate of disappearance of ketone calculated on basis of hydrocarbon appearance. AC3 = CsHa-CaHs: Le., ACI from calculated CaHs, Le., propane from abstraction. C3Ha -N 0.6 CeHir.
Since there are two kinds of kctonyl radicals possible it would be of interest to see whether they both can decompose via reaction 7 . For the radical formed in reaction 3a to give propylene there would be required a hydrogen shift during the decomposition, while the ketonyl from reaction 3b decomposes by breaking one carbon-carbon bond with simultaneous formation of a double bond. The propylene formed by ketonyl decomposition complicates equation I1 since the quantity to be subtracted from the total propane in equation I1 is the propane formed by the disproportionation reaction (1). However, the disproportionation propylene can be fairly accurately measured by extrapolating the low temperature range of propylene/diinethylbutane ratios to high temperatures. As may be noted in Fig. 1, the disproportionation propane is only a small proportion of the total propane a t high temperature. When the propane is corrected, thc Arrhenius plot of kn/k2'/*is a straight line over the temperature range studied. The isopropyl radical apparently is quite stable up t o 299" but begins to decompose measurably before 353". The production of hydrogen gas is presumably via reactions 8a and 10. On this basis one can compare reactions 8a and 3a by the following relation k d k s , = H2[DIKl/C&
(111)
This does not permit an accurate measurement of the activation energy difference as can be seen in Fig. 1. The data indicate that the rate of formation of ethylene and methane are about equal, and that the rate increases with temperature and with intensity of radiation. It is possible that ethylene and a methyl radical results from the decompoqition of an isopropyl radical, since McNesby, et U Z . , ~ have shown that hydrogen atom transfer along the carbon skeleton can occur a t high temperatures. Another possible path is reaction 9 followed by reaction 8b. However, it is difficult to see why the relatively low concentration of product propane should be efficiently converted to n-propyl radicals when a much easier hydrogen source is available to the isopropyl radicals from the parent ketone. When propane was added t o the ketone prior to reaction (run 68) the rate of formation of hydrogen dropped, while rate of formation of methane and ethylene were almost the same as when no propane was present a t the start of the reaction. There is little possibility that the methane and ethylene could resuit from an impurity which contains n-propyl groups. High resolution mass spectrometer analysis of the diisopropyl ketone and of the (Y, a'-diisopropyl-d2 ketone showed that not over 0.6% of the total propyl content was n-propyl. The four-carbon hydrocarbons may be formed by reaction 13 and by addition of methyl t o propylene. The 72 peak a t high temperature may be due to pentanes formed by addition of propyl t o ethylene. (9) J.
R. McNesby and A.
9. Gordon, J . Chem. Phus., in press.
1318
CONSTANTINE NEUGEBAIJER AND JOHNL. MARGRAVE
Acknowledgment.-The authors wish to acknowledge the mass spectrometer analyses of A. V.
Vol. 60
Jensen and the help of S. R. Smith with the interpretation of mass spectrometer records.
THE HEATS OF FORMATION OF TETRAFLUOROETHYLENE, TETRAFLUOROMETHANE AND I,1-DIFLUOROETHYLENE BY CONSTANTINE A. NEUGEBAUER AND JOHNL. MARGRAVE Contributionf r o m the Department of Chemistry of the University of Wisconsin,Madison, Wisconsin Received April 18, 1966
The standard heats of formation of gaseous tetrafluoroeth lene and tetrafluoromethane have been determined by the hydrogenation and decomposition of tetrafluoroethylene in a t o m b calorimeter. They are -151.7 i 1.1 and -217.1 f 1.2 kcal./mole, respectively. The heat of formation of gaseous 1,l-difluoroethylene was found to be -77.5 i0.8 kcal./mole from combustion experiments,
Introduction The heat of formation of CF, by the direct interaction of the elements was reported by v. Wartenberg' to be -161 kcal. Later, v. Wartenberg2 studied the displacement reaction between CF4 and K to give K F and carbon, and found -231 kcal. for the heat of formation of CF4. Kirkbride and Davidson13using v. Wartenberg's later technique, obtained a value of -218 kcal. for the heat of formation of CF4. Jessup, McCosky and Nelson4 report -220.4 kcal. for the heat of formation of CF4, using the fluorination of methane t o give CF4 and gaseous HF. Later, Scott, Good m d Waddington,s by a method involving the heat of combustion of Teflon polymer in the presence of hydrocarbon oils, obtained -218.3 kcal. for CF4. Finally, Duus6 reports -212.7 kcal. for CF4, using the hydrogenation of CZF4 to give carbon and gaseous HF, and the decomposition reaction of C2F4 to give C and CF4. The final state of the carbon mas assumed to be graphite in both cases. Corrections for the final state of the HF were estimated. Furthermore, no corrections were made for methane produced in the hydrogenation. By the method described above, Duus also obtained the heat of formation of gaseous C2F4 as -151.3 kcal. On the other hand, both v. Wartenberg,7 and Kirkbride and Davidson,* report - 162 kcal. for the heat of formation of CzF4. I n general, the literature values for the heats of formation of both compounds show variations and uncertainties large enough to warrant new determinations. The reactions used to determine the heat of formation of C2F4 and CF4 are similar to those used by Duus. When gaseous C2F4 reacts with hydrogen in a calorimetric bomb that initially contains (1) H. v . Wartenberg and R. Schiitte, Z. anorg. Chem., 811, 222 (1933). (2) H.v . Wartenberg and G. Riteris. ibid., 868, 356 (1949). (3) F. W.Kirkbride and F. G. Davidson, Nature, 174, 79 (1954). (4) R. S. Jessup, R. E. RIeCosky and R. A . Nelson, J . Am. Chem. Soc.. 7 7 , 244 (1955). (5) D. W. Scott, W. D . Good, and G u y Waddington, J . Am. Chem. SOC., ibid., 77, 245 (1955). (6) H. C. Duus, Ind. E n g . Chem., 47, 1445 (1955). (7) H.v . Wartenberg, 2. anory. Chem., 1'78, 320 (1955). (8) F. W. ICirkhride and F. G. Davidson, ref. 3.
water, the net reaction may be represented by the equation CzFa(g) 2Hz(g) = 4HF(aq) 2C(amorphous) (I) From this reaction, the heat of formation of CzFa can be calculated, Tetrafluoroethylene decomposes when ignited according to the equation
+
+
CZFd(g) = CFa(g)
+ C(amorphous)
(11)
This reaction, when coupled with the one above, makes possible the determination of the heat of formation of CF4. The heat of formation of the amorphous carbon formed in the reactions can be determined by separate combustion experiments. The heat of formation of 1,l-difluoroethylene has been determined by measuring the heat of the reaction CHzCFz(g)
+ 202(g) = 2COz(g) f 2HF(aq)
(111)
Experimental Apparatus and Method.-A conventional isothermal calorimeter of the type described by Dickinsong was employed. The temperature of the water jacket was kept constant a t 24.80 f 0.002' with a vibrating mercury regulator in connection with an electronic relay. The calorimeter cup, of 5-1. capacity, was held in place by three Lucite wedges placed along the upper rim of the calorimeter cover and held firm by the lid of the water jacket. The calorimeter stirrer was driven by a small electric motor a t 600 r. .m. A double valve Parr illium bomb having a volume o f 3 6 0 ml. was used in all experiments. The temperature rise in the calorimeter was measured by means of a platinum resistance thermometer, calibrated by the National Bureau of Standards, in connection with a G-2 Mueller resistance bridge, also previously calibrated, and a high sensitivity galvanometer, all supplied by the Leeds and Northrup Co. A change of O.OOO1 ohm (0.001'). in the resistance of the thermometer caused a 3 mm. shift in the reflection from the galvanometer mirror on the scale. The galvanometer was used a's a null-point instrument, the time a t which a predetermined resistance was reached being read from an electric timer. The resistance values were converted t o temperature readings by using the method of Werner and Fraeer.10 The corrected temperature rise (in degrees) was calculated by the Dickinsonll method, assuming Newton's law of cooling to hold. The initial temperature was so chosen that the final temperature was about 0.05' above jacket temperature. For convenience in obtaining the appropriate initial tempera(9) H.C.Dickinson, Bull. Natl. Bur. Standards, 11, 230 (1914). (10) F. D . Werner and A. C. Frszer, Rev. Sci. Instr., 83, 103 (1952). (11) H.C.Diekinaon, ref. 9.
