HEATSOF FORMATION OF TRIFLUOROMETHYL CHLORIDE AND BROMIDE
droxyl group. However, the AH" values (kcal mole-') are consistently more endothermic than those for the sugars (ie. cytidine, 10.7; uridine, 10.9; adenosine, 9.7; and ribose, 8.l), while the Ah'" values are consistently less negative than for adenosine and the monosaccharides. These differences are probably due to the slightly different environments of the dissociating proton. Thymidine contains a deoxyribose rather than a ribose sugar group. This probably accounts for the
2705
slightly higher pK (12.85)value for thymidine relative to that of uridine (12.59). This same effect (although to a greater degree) was observed2 with adenosine and 2'-deoxyadenosine, indicating that the deoxy derivatives of purines and pyrimidines are less acidic than the ribose derivatives. Acknowledgments. The authors wish to acknowledge the assistance of Mr. Wayne Allgaier and Mr. Joseph Richards in the determination of the pK values for several of the substances used in this study.
The Heats of Formation of Trifluoromethyl Chloride and Bromide
by Allan Lord, C. A. Goy, and H. 0. Pritchard Centre for Reaearch i n E x p i m e r t d Space Science, York University, Toronto, Canada (Received February $0, 1067)
+
+
+
+
The gaseous equilibria CF3Cl I2S CF31 IC1 and CF3Br Iz =: CF31 IBr have been studied over the temperature range 620-740°K. The heats (AH"298) for these reactions are calculated to be 17.27 f 0.13 and 9.55 f 0.03 kcal/mole, respectively. Using AHr"z@(CFaH,gas) = -165.1 kcal/mole as a key heat of formation, AHro29,(CF3C1,gas) = -167.5 kcal/mole and AHtozg8(CFJ3r,gas) = -154.1 kcal/mole.
We have previously noted that the presence of traces of CFaCl in commercial fluoroform accelerates the rate of iodination of fluoroform.' We have therefore studied the iodination of CF3Cl itself and have found that it is possible to establish the equilibrium CFaCl
+
CFaI
12
+ IC1 equilibrium constant = KT
at temperatures around 700°K. The probable mechanism of this reaction is
I
+ C F a C l z CF3 + IC1 IC1 I + c1 c1,
c1
+ c1
(1)
+ J_ CFaI + I CF3 + IC1 CF3I + C1 CF3
(3) (4) The over-all kinetics of this process will be complicated, but could be sorted out once data became available on reactions 2 and 4. Similarly, the equilibrium CF31 IBr CF3Br Iz can be established in the same temperature range, presumably by an analogous mechanism. 12
+
+
Experimental Section The experimental procedure used was similar to that used previously.ls2 Known amounts* of CFaCl (1) C. A. Goy, A. Lord, and H. 0. Pritchard, J . Phys. Chem., 71,1086 (1967). (2) C.A. Goy and H. 0. Pritchard, ibid., 69, 3040 (1965). (3) Full details of concentration and analysis conditions can be found in A. Lord, Ph.D. Thesis, York University, Toronto, 1967.
Volume 71, Numbe-r 8 July 1867
A. LORD,C. A. GOY,AND H. 0. PRITCHARD
2706
or CF3Br (5-10 cm pressure) and IZ (0.25-0.5 g) were sealed in a 300-ml Pyrex vessel and heated until equilibrium was established (1-350 hr depending on the temperature). The reaction was terminated by pumping the contents of the vessel into a liquid nitrogen trap and the ratio CF31/CF3C1 or CF31/CF3Br was determined gas chr~matographically.~The equilibrium constants were calculated as KT = [CSII2/ [CF3X][I2](i.e., assuming that [IX] = [CFd]) since it was not possible to make an accurate analysis for IC1 or IBr in the presence of a vast excess of 12. No significant disproportionation of IX to 1 2 and XZtakes place under these condition^.^ The trifluoromethyl chloride and trifluoromethyl bromide were obtained from Matheson of Canada Limited and were found by mass spectrometric and gas chromatographic analysis t o contain no important impurities. Iodine was Baker Analyzed Reagent grade and was purified by sublimation in vacuo.
Trifluoromethyl Chloride The results of 16 measurements on the equilibrium between CF3C1and I, are listed in Table I. The van’t Hoff plot of these equilibrium constants leads to a value of = 16.0 kcal/mole, somewhat lower than the more reliable value derived by third-law methods in Table I, i.e., AHozsa= 17.27 f 0.13 kcal/mole. (These standard heats were calculated using the thermodynamic data given in the JANAF table^,^ except that for CF31which was taken from the work of McGee, Cleveland, Meister, and Decker!) Combining the latter value with AHozs8= 17.10 i 0.17 kcal/mole for the equilibrium CF3H
+
12
CF3I
+ HI
which we derived previously, we have AH~ozga(CF3Cl, gas)
=
AHfozg8(CF3H,gas) -2.37
T, OK
Log K T
732 725 715 712 704 696 680 677 677 674 664 662 662 653 646 620
-4.6822 -4.7729 -4.7829 -4.8923 -4.7962 -5.0209 5.0631 -4.9899 -5.0951 -5.0314 -5.1642 5.2387 -5.1229 -5.2594 5.2723 -5.5471
-
-
+ Ia e CFJ + IC1
A[--(FOT - H 0 m ) / T ] , cal mole -1 AHOzos, deg-1 cal mole-’
2.335 2.332 2.329 2.328 2.325 2.321 2.315 2.313 2.313 2.312 2.308 2.307 2.307 2.304 2.300 2.291
17,388 17,525 17,321 17,597 17,075 17,597 17,325 17,014 17,345 17,081 17,215 17,402 17,033 17,225 17 ,066 17 , 158 Av 17,273 Standard deviation f.132
Table 11: The Equilibrium CFsBr
+ Iz E CF31 + IBr
OK
Log KT
A ( - ( F O T - HOaos)/T], cal mole-1 deg -1
739 736 729 726 711 696 686 667 65 1 637
-2.3833 -2.4011 -2.4288 -2.4414 -2.5193 -2.5888 - 2.6304 -2.7099 -2.8118 -2.8819
1.923 1.922 1.921 1.920 1.917 1.914 1,913 1,909 1.907 1.904
T,
AH’m, cal mole-’
9486 9504 9499 9508 9560 9571 9569 9551 9618 9617 Av 9548 Standard deviation *31
f 0.30
Thus, if D(CF3 - H) = 106.0 kcal/mole,1~6 D(CF3 C1) = 85.2 kcal/mole, independent of any uncertainties in the heat of formation of CF3H. Accepting, for the time being, AHyoZ98(CF3H, gas) = -165.1 kcal, we have AHfo2g3(CF&l,gas) = - 167.5 kcal/mole.
Trifluoromethyl Bromide The results of ten measurements on the equilibrium between CF3Br and IZ are listed in Table 11. The van’t Hoff plot of these equilibrium constants leads to a value of = 9.9 kcal/mole, compared with the more Of AH02g8 = 9*55 * kcal/mole derived in Table 11, again using thermodynamic data The Journal of Physical Chemistry
Table I: The Equilibrium CFaCl
from ref 4 and 5. Likewise, using our data on the iodination of fluoroform, we have AHfozgs(CF3Br, gas) = AHfoZs8(CF3H, gas)
+ 10.99 f 0.20
in fairly good agreement with the value of 11.5 f 0.3 derived by Coomber and Whittle from their experi(4) “JANAF Thermochemical Tables,” DOW Chemical Co., Midland, Mich., 1965, and supplement, 1966. (5) P.R. McGee, F. F. Cleveland, A. G. Meister, and C. E . Decker9 J . Chem. Phys., 21, 242 (1953). (6) J. W. Coomber and E . Whittle, Trans. Faraday SOC.,62, 2183 (1~56).
DISSOCIATION ENERGIES OF
THE
GASEOUS MONOXIDES OF
ments on the bromination of fluoroform? Our result leads to a value of D(CF3 - Br) = 69.6 kcal/mole, again independent of the heat of formation of CF3H, and AHrozss(CF3Br, gas) = - 154.1 kcal/mole. The values of A H 0 2 @ in Table I1 seem to show a systematic trend, in that those derived from the hightemperature runs are some 0.12 kcal lower than those derived from lower temperature experiments. The data of Coomber and Whittle' show precisely the same trend, both in magnitude and sign (independent of whether one uses the thermodynamic data of Gelles and Pitsers for CF3Br, as they do, or that taken from ref 5, as we do). Since such consistent trends were not observed in the iodination of CF3H or CF3C1, a possible explanation of this is that the entropy data6#* for CF3Br
Rare Earths.
IV.
THE
RAREEARTHS
2707
is incorrect. To try to test this hypothesis, we began astudyoftheequilibrium CFlCl
+ Brz
CF3Br
+ BrCl
to see if the same trend was apparent there, but we were unable to establish an equilibrium concentration of CF3Br in our temperature range. For the time being, therefore, we feel that the heat of formation and dissociation energy of CF3Br are subject to a somewhat larger uncertainty than is implied by the equilibrium measurements of this paper and of ref 7. (7) J. W. Coomber and E. Whittle, Trans. Faraday SOC.,63, 608 (1967). (8) E. Gelles and K. s. Pitzer, J . Am. Chem. S O C . , 7 5 , 5259 (19.53).
Dissociation Energies of the Gaseous
Monoxides of the Rare Earthsla
by L. L. Ames, P. N. Walsh, and David Whitelb Department of Chemistry, The Ohio State University, Co~umbun,Ohw
(Received March 9,1967)
The dissociation energy of the gaseous rare earth monoxides together with the closely related compounds ScO and YO have been determined from Knudsen effusion and mass spectrometric isomolecular oxygen-exchange reaction studies. It is shown that there is a double periodicity in the variation of the dissociation energies with atomic number and that these variations closely match those of the heats of sublimation of the rare earth it is shown that metals. Assuming a common bound state for the monoxides, M2+02-, the variations in the dissociation energies in the lanthanide series correspond to the magnitude of the 4f" --t 4fn-'5d transitions of the divalent ions.
I. Introduction In parts 12* and IIIZbof this series, it was shown that large variations occur in the dissociation energies of the gaseous monoxides of the rare earth metals. On the basis of some additional preliminary observations, it was suggested3 that these variations could be correlated with the filling of the 4f orbitals of the lanthanides. The largest dissociation energies corresponded
to the unfilled, half-filled, and filled 4f shell: between these points the dissociation energy decreased with (1) (a) This work was supported in part by the Office of Naval Research under Contract Nonr-495(12). (b) Address correspondence to this author a t the Department of Chemistry, University of Pennsylvania, Philadelphia, Pa. 19104. (2) (a) H. W. Goldstein, P. N. Walsh, and D. White, J . Phys. Chem., 65, 1400 (1961); (b) P. N. Walsh, D. F. Dever, and D. White, ibid., 65, 1410 (1961).
Volume 71, Number 8 July 1967
