2683
THERMAL DISSOCIATION OF CHLORINE TRIFLUORIDE also estimate the thermodynamic quantities for the process HC1 (gas) +HC1 (ice) and find that AH0 = -21.68 kcal/mol, AGO = -7.88 kcal/mol, and Ai30 = -48.2 cal/(mol deg). These results are obtained using wellknown thermodynamic data for the process HC1 (gas) -t HC1 (water).20 The equilibrium constant, K = (concentration of HC1 in ice)/(activity of HC1 in water), is 6.39 X 10-6, 1.78 X 10+, and 2.14 X at temperatures of -4, -11, and -18”, respectively. The mode of incorporation of chloride ion in ice and the path of diffusion is suggested by the large hexagonally shaped channels that appear in the ice struc-
ture. These channels define the C axis, and this is the direction of diffusion. It is proposed that diffusion proceeds through and dissolution occurs in these hexagonal channels. On the basis of the above model we may predict that diffusion perpendicular t o the C axis should be much slower than diffusion along the C axis. The technique described in this paper should be applicable to solubility studies in other solids that have low melting points, as long as they can be prepared in the form of large single crystals. I n cases where the solubility is appreciable, the requirement of large single crystals may be obviated.
The Thermal Dissociation of Chlorine Trifluoride behind Incident Shock Waves by J. A. Blauer, H. G. McMath, and F. C. Jaye Air Force Rocket Propulsion Laboratory, Air Force Systems Command, Edwards, California 93623 (Received January 8, 1969)
The thermal dissociation of CIFa behind incident shock waves has been studied in the temperature range of 800-1300°K. The course of the dissociation was followed by means of ultraviolet absorption spectroscopy exp - [(28,200 f 1700)/ centered a t 2200 =k 250 A. Initial slope measurements gave a value of kl = 101a.5*‘J-4 R T ] cc/mol sec for the rate constant of the reaction CIFa M = C1F2 F M. The collision efficiency of ClF relative to argon was found to be 35 =k 5 to 1. Evidence is presented which favors a heat of formation for CIFz of -19 =t2 kcal/mol a t OOK.
+
Introduction It is likely that the thermal decomposition of C1F3 proceeds initially by simple unimolecular bond rupture, i.e. C1F3 21 = ClFz F M (1)
+
+ +
It can be shown from the elementary theory1V2 of unimolecular reactions that the rate of this reaction should have fallen to 50% of its high-pressure limit a t approximately 30 atm if the reaction temperature is 1000°K. Furthermore, at pressures below 5 atm, the reaction should be bimolecular. The nature of the reactions suocessive to reaction 1 can be inferred by analogy from the mechanism given by Schumacher3 for the photochemical reaction of ClF3 and Fzto form C1F5. The reaction path is
+
+ +
ClFz M = C1F F M ClF2 F = CIF Fz ClF3 C1F = 2C1Fz F F M = Fz Ail:
+ + + +
+
+
(2) (3)
(4) (5)
+ +
I n addition to these reactions, one other possibility will be considered here, i.e.
F
+ ClFi = ClF2 + F2
(6) At present, there are no published kinetic results concerning the decomposition of CIFa; however, thermal data4are available. The choice of a technique based upon uv absorption for following the course of the dissociation was made possible by the very large extinction coefficient for C1F3 relative to those for C1F and FBunder the conditions of (1) 8. W. Benson, “The Foundations of Chemical Kinetics,” McGraw-Hill Book Co., Inc., New York, N. Y., 1960, p 234. (2) Values for the Kassel integral were furnished by Dr. T. A. Jacobs of Aerospace Cow., El Segundo, Calif. For this calculation we assume that the molecule has six effective oscillators and that the collisional deactivation efficiency has a value of about 0.1. The dissociation energy was assumed to be 37 kcal/mol with Y = 1014 sec-1. (3) R. L. Krieger, R. Gatti, and H. J. Schumacher, 2. Phys. Chem., 51, 240 (1966). (4) H. Sohmitz and H. J. Schumacher, 2. Naturforsch., Za, 362 (1947).
Volume 73, Number 8 August 1969
2684
J. A. BLAUER,H. G. MCMATH,AND F. C. JAYE
measurement. Although these ratios varied somewhat with temperature, at 900°K they were found to be represented approximately by 1:0.04:0.02. Any optical interference due to the intermediate species CIFz was ignored in the analysis, it being assumed that its concentration remained low during the course of the reaction and that its extinction coefficient is not large.
Experimental Section The shock tube, of Amos design, is of stainless steel and has an inside diameter of 3.75 em. The overall length of the test section is 7.5 m, and the entire inside surface is finished to a grad& smoothnesa. The observation port is equipped with sapphire windows held in compression by close-tolerance brass collets. Window-shock tube sealing is effected with indium wire gaskets. The driver, having an overall length of 1.7 m, was separated from the downstream section by means of scribed diaphragms of cold-rolled steel. The downstream section was in turn separated from a 220-1. dump tank by means of a thin sheet of Mylar." Shock detection was by means of moderate response (CQ. 7 rsec) piezoelectric detectors' having a spatial resolution of 2 mm and placed at intervals of 76.2 em along the entire length of the downstream section. After amplification, the outputs of these detectors were displayed on a Tektronix Model 535 oscilloscope which was equipped with a rastor sweep and a Radionics Model TWM crystal-driven timing generator. The course of the dissociation was followed by means of a once-through, single-lighbpath, ultraviolet absorp tion spectrometer.8 The source was a Beckman deuterium arc lamp. Spectral isolation was by means of a Baird-Atomic interference filter centered a t 2200 f 250 A. Detection was by means of a Type 1P28 photomultiplier tube. The instrument had a spatial resolution of 2 mm and an overall relaxation time of 3 psec. Argon having a purity of 99.998'% was purchased from Matheson and used without further purification. Gaseous FZ with a minimum purity of 98.2% was purchased from Allied Chemical Corp. A mass analysis revealed the presence of 0.7% Oz and 0.2% H F as the only significant impurities. After passage through a column of NaF pellets, the gas was used without further purification. Gaseous ClF, having a purity of 98.0%, was purchased from Oaark-Mahoning Co., and was further purified by trapto-trap distillation using a Freon-12 liquid slurry at -140". Gaseous CIFa, having a purity of approximately 98%, was purchased from Matheson and was further purified by forming the K F complex, KCIF,, a t ambient temperatures after which the ClFa was recovered by vacuum distillation a t 200". Infrared and mass spectral analyses of the purified ClFa and CIF samples showed no indication of H F or oxygen-containingcompounds. Mixtures of Ar, CIF,, F2, and CIF which contained O-l.SS% ClFa, 0-2.80% Fz, and 0-5% C1F were The J m d of P h y d C h m a W
Absorption trace ior test no. 6: 0.1-V ordinate divisions, ZO-rsec fast sweep and 200-ryec slow sweep abscissa divisions, O.94Y0CIF, in Ar, 912"K,5.7 atm incident shock pressure. Figure 1.
prepared and stored in stainless steel tanks, which were pre-passivated with CIFa and F2. Compositions were determined by differences in direct weighing for CIFa and by pressure differences for all other gases. Heise gauges, whose scales could be read to 0.1% of full scale, were used in all mixing operations. Immediately before using a gaseous mixture in a test, its CIFn content was determined by reading its optical density at 2200 A. Beer's law gave an excellent descrip tion of the results. No change in optical density was observed upon allowing the mixtures to stand for several days. The initial gas pressure within the shock tube was measured by means of a Wallace and Tiernan 0-60 psia gauge. Tests were conducted using either helium or nitrogen as the driver gas.
Results and Discussion With the assumption of Beer's law, the optical density of the reacting mix can be described by the relationship
A = K(C1Fa)
+ N(CIF) + L(Fz)
(7)
Here K, N,and L represent the absorption coefficients for CIF, ClF, and F,, respectively. The initial conditions behind the shock wave were obtained by means of a solution of the Rankine-Hugoniot equations. This permitted an evaluation of the absorption coefficient for ClFa by simple linear extrapolation of the oscillogram to the origin of no reaction; see Figure 1. The temperature dependence of these coefficients is illustrated in Figure 2. The absorption coefficient for fluorine was found by testing binary mixes of F2and argon. The absorption coefficient for C1F was taken (6) AVCO Corp., Wilmington. Mass. (6) Trsds name. (7) Xistle Instrument Corporation, Model 601. (8) Furnished by Rocketdyne. Inc., Cmoss Park. Calif., under
Contract No. AF 04(611)-5963.
2685
THERMAL DISSOCIATION OF CHLORINE TRIFLUORIDE Table I : Compositions and Shock Parameters for Individual Tests Test
mol/oo
(C1Fa)o X 108, mol/co
1 2 3 4 5 6 7 8 9 10
0.0406 0.0898 0.0396 0.0765 0.0937 0.0750 0.0863 0.173 0.151 0.146 0.145 0.128 0.208 0.0742 0.333 0.478 0.326 0.202 0.0889 0.0598
0.778 0.764 0.758 0.726 0 * 890 0.711 0.818 0.815 0.711 0.691 0.686 0.604 0.978 0.0351 0.801 1.12 0.697 0.528 0.889 0.637
(Arlo X IOa,
11
12 13 14 15 16 17 18
19 20 a
(Fzh X 1W, mol/cto
0.0 0.0 0.0 0.0 0.0 0.0 0.0 0.0 0.0 0.0 0.0 0.0 0.0 0.0
..700
P,
OK0
QKb
atm
0.0 0.0
1272 1011 703 856 1168 912 856 876 916 792 844 856 1159 910 996 805 946 1086 960 1086
1144 923 660 830 1109 874 820 858 897 775 826 837 1128 890 985 797 935 1060 919 1032
4.3 3.4 2.3 5.4 9.1 5.7 6.1 12.5 10.8 9.6 10.1 9.0 19.8 5.6 27.3 31.6 26.1 19.4 7.0 5.7
0.0 0.0 0.0 0.0 0.0 0.0 0.0 0.0 0.0 0.0 0.0 0.0 0.0 0.0 0.0
0.0 0.0 9.89 5.98 0.0 0.0
Conditions immediately behind the incident shock wave.
Tt 8
mol/oc
(C1F)o X 109,
0.0 4.44 3.18
Ti,
900
1000
1100
1200
x
10-1,
880 -1
20.9 1.2 N.R. 0.27 12.3 0.8 0.16 0.65 0.55 0.06 0.15 0.30 23.4 0.70 0.9 0.17 2.4 13.6 1.5 7.7
' Conditions at equilibrium.
D.0041 0.0024 0.0020
800
kl
A A
1300
T"k
e-
c
Figure 2. Molar absorption coefficients for ClFa illustrated as a function of temperature.
from the conditions prevailing a t equilibrium; see Figure 1. The values of N and L were found to vary only slightly with temperature. At 900°K the ratio K :N :L = 0.55 :0.02:0.01 was found to apply. Resort was made to the method of initial slopes for an estimation of the rate constant of reaction 1. The results are tabulated in Table I. A close examination of these results reveals that the estimated unimolecular rate constant is nearly a linear function of the concentration of possible collision partners. Consequently, the data were assumed to refer to a unimolecular reaction in its low-pressure region. The resulting bimolecular rate constant has the form
The results are graphically illustrated in Figure 3.
Figure 3. Temperature dependence illustrated for apparent second-order reaction rate constants. Solid triangles represent mixes containing 2.8% fluorine. Solid squares represent tests containing 5.0% CIF.
The experimental activation energy of a unimolecular reaction in its bimolecular region is related to the dissociation energy1 of the molecule by means of the relationship
(9) If 1000°K is taken as the average temperature at which data were taken, and the number of internal degrees of freedom, s, is assumed to be 6, eq 9 gives a value of 37 f 2 kcal/mol for the dissociation energy. Although this calculation is uncertain, due to the uncertainty in s, the result will be substantiated at a later point in the paper. Volume 73,Number 8 August 1969
J. A. BLAUER,H. G. MCMATH,AND F. C. JAYE
2686 When used in conjunction with the best thermal data presently a ~ a i l a b l e ,this ~ - ~translates ~~~ into a heat of formation for ClFx at 0°K of -19 I 2 kcal/mol. In subsequent analysis, more complete thermal data were required for the assumed intermediate, C1F2. These were generated by assuming for it a structure similar to that of HzO with vibrational frequencies of 750, 700, and 360 cm-l. The C1-F bond distance was estimated at 1.65 A. In an attempt to obtain further information from the data, the entire suggested mechanism consisting of eq 1 through 5 was subjected to numerical integration by means of a nonequilibrium computer program1I which soives the conservation equations simultaneously with the reaction rate equations. The overall reduction of rate data proceeded by matching observed and calculated reaction profiles and repeating the process with a new estimated set of rate constants. Only one rate constant was varied at a time until the best overall fit of the data was obtained. The initial estimate for kl was taken as eq 8. The initial estimate for kz was taken as k, = 1014.0e-Z8,OoO/RT cc /mol sec (10) The activation energy for step 2 was estimated from the best available thermal datal0 used in connection with our suggested value for the heat of formation of C1F2 and eq 9 with s = 3. Reaction 3 was assumed to have a near-zero activation energy and to have a steric factor of approximately 0.01, The resulting rate constant was assigned a value of
k3 = 1 X l0l2ccjmol sec
(11)
The effect of reaction 4 was obtained by difference between those tests having added ClF and those to which it was not added. As a consequence, it initially was deleted from consideration in those tests not containing added C1F. The rate constant for reaction 5 was taken from the work of Diesen.lz All reverse rates were computed from thermal data and the assumption of detailed balancing. An illustration of the comparison between the experimental data and the corresponding computed reaction profile for one test is shown in Figure 4. Here normalized optical density is plotted against laboratory time. Although the use of eq 8 is sufficient to force the computed profile to pass through the initial portion of the data, the bulk of the computed profile exhibits a far slower overall reaction rate than was experimentally observed. The reason for this behavior lies in the rapid temperature drop brought on by reaction 1 coupled to the slow recombination of fluorine atoms. No reasonable combination of reaction rates for the four reactions thus far considered was sufficient to describe the entire reaction profile. Increases in IC:! or ka by an order of magnitude were found to have only The Journal of Physicat Chemistry
Figure 4. Comparison of computed and observed reaction profiles for test no. I which was taken at 1272°K. The reaction model used excluded consideration of steps 4 and 6.
a minor effect upon the overall reaction rate, whereas changing the value of kl by a factor of 2 changed the computed reaction rate by approximately the same factor. Since the effect of large amounts of C1F upon the initial reaction rate is not gross, see Figure 3, it was decided to test the effect of the addition of reaction 6 to the model, If a heat of formation of -19 kcal/mol is assumed for C1F2, this reaction will be thermoneutral and should, as a consequence, exhibit only a small activation energy. The initial estimate for ke was
ke = 0.5 X 10” cc/mol sec
(12)
This value was based upon the near steady-state concentration given for F atoms by the above described attempt to compute the reaction profile coupled to the increase in reaction rate required to describe the lat’er stages of the reaction; see Figure 1. After adjusting the value of k~ until the best fit for all of the experimental data was obtained, the following expression resulted
ks
= 0.75
X 1011e-26m’RT cc/mol sec
(13)
During the course of these calculations, it was found necessary to make a slight adjustment in the expression for kl. The final result was
kl
= 0.7
X 1014e-301000’Rrcc/mol sec
(14)
The results of these calculations are illustrated in Figure 5 for several tests conducted over a wide range of experimental conditions. Increases in the values of LZ (9) H. H. Claassen, B. Weinstock, and J. G. Malm, J. Chem. Phys., 28, 285 (1958).
(10) “JANAF Thermochemical Tables,” Dow Chemical Co., Midland, Mioh., July 1968. (11) Furnished by Dr. T. A. Jacobs, Aerospace Corp., El Segundo, Calif. (12) R.W. Diesen, J. Chew. Phys., 44, 3662 (1966).
2687
THERMAL DISSOCIATION OF CHLORINE TRIFLUORIDE SYMBOL RUH
0 0 A
1 5 15
l a O K P ~ l i t m l BClf3 %F2 9.1 8.1 21.3.
1212 1168 898
0.50
0 A
0.0 0.0 0.0
0.14 0.24
18 20
846 1066
7.0 5.1
1.0 1.0
5.0 6.0
35
0.8
1 0.6
L AO
0.4
0.1
1
I
IOU
0
200
300
400
L A B O R h l O R Y TIME IPS1Cl -
,
I
10
0
l
l 30
20
