NOTES
July, 1963 experimental data on the sublimation rate of MnFz had been reported in the literature. Brewer, Somayajulu, and Brackettl have estimated the heat of sublimation to be equal to 75 kcal./mole. Experimental Apparatus and Techniques The experiment was crzrried out with a microbalance built inside a vacuum system. The apparatus has been previously described in detail.2 It consists basically of a resistance furnace, a beam balance, and an electrical control circuit for the balance to measure the change of voltage with the change in weight. The MnFl single crystal was suspended from a holder made of tungsten wire inside a glass envelope contained in the furnace. The temperature was measured by means of a calibrated chromelalumel thermocouple placed inside the glass envelope. Several readings were taken of the rate of weight loss of the sample a t a given temperature and the time-weighted average of all of the weight loss readings was calculated and reported as one pressure point. These d.ata were used to derive a log P us. l / T equation with a Fortran program on a Control Data-1604 digital computer.
Discussion of Results MnFz(g) was considered the only important vapor species. The free energy functions for both solid and gaseous MnFp, were .taken from the data compiled by Brewer, Somayajulu, and Brackett.l The experimental results are summarized in Table I. A third law calculation from these data gives AH0298 = 76.4 =t 1 kcal./mole for the heat of sublimation while t'he second law least-square plot of log P vs. 1/T gives AHo = 73.5 :j= 0.5 kcal./mole at Tavg = 939.0'K. When corrected to 298'K. using the heat content values compiled by Mah3 and the molecular constants estimated by Brewer, Somayajulu, and Brackett,' the second law approach gives AHo298=; 76.0 1.0 kcal./ mole. TABLE I VAPORPRESSURE DATAON MnFz T , OK.
P
924.1 939.0 952.3 963.8 969.9 982.9 964.3 951.7 887.1
2.49 X 4.73 x 8.54 x 1.35 x 1.69 X 2.78 X 1.40 X 8.34 X '4.94 x
AH%,,, koal./mole
10+ 10-9 10-9 10-8
lo-* lo-*
IO-$ 10-10
76.46 76.46 76.38 76.39 76.43 76.44 76.36 76.38 76.34
-.
76.40 & 0 . 5 kcal./mole
Using the above sublimation and other available thermochemical data,4 one calculates the atomization energy of MnF2(g)to be 217 kcal./mole, which is in good agreement with the 220 kcal./mole estimate of Brewer, et a2.l This indicates 108.5 kcal./mole for the average bond energy in MnlF2 while Gayd.on5 has estimated DO(NnF) as 81 kcal./mole from molecular spectra via a linear Birge-Spooner extrapolation. (1) L. BreNer, G. Somayajulu, and E. Braokett, University of California Radiation Laboratory Report KO.9840, September, 1961; C h e m . Rez., 63, 111 (1963). (2) L. H. Dreger a.nd J. L . Margrave, J . P h y s . C h e m . , 6 4 , 1323 (1960); R. C. Paule, Ph.D. Thesis, University of Wisconsin, 1962. (3) A. D. Mah, U. S. BuTeau of Mines, Report of Investigation 5600 (1960). (4) D. R. Stull and G. C. Sinke. "Thermodynamic Properties of the Elements," Advan. in Chem. Series, ACS, 1956; F. D. Rossini, D. D. Wagman. W. H. Evans, 5.Levine, and I. Jaffe, National Bureau of Standards Circular 500 (1952); JANAF Interim Thermochemical Tables, edited by D. R. Stull, The Dow Chemical Company, Midland, Michigan, 1960. ( 5 ) A. G. Gaydon, "Dissociation Energies," Chapman and Hall, Lt,d., lY53.
1565
Acknowledgmen.ts,-The authors are pleased to acknowledge the financial support of the National Science Foundation and of the United States Atomic Energy Commission. ISOTOPIC FRACTIONATION I N T H E OH--HzO EXCHANGE REACTION BY MICHA.EL GREENAKD HESRYTAUBE' George Hevbert Jones Laboratory, T h e University o/ Chicago, Chicago, I l l i n o i s
Received X a r c h 16, 1863
It is import'ant to know with some certainty the equilibrium constant in the system HzO"iiq
+ 018H-aq2
HzO'*iiq
+ O1'HFaq
in order to interpret the fractionation of oxygen isotopes during the hydrolysis of certain complex ions. Values of this constant, K , a t 25' have been calculated. Hunt and Taube%have obtained a result of 1.035, from vibrational frequencies of H20 vapor and of OHin 10 M solution. More recently Thornton3has refin'ed the value to 1.0385 by using the parameters of liquid H20 and inc1udin.g librations. I n both calculations approximate meth.ods were employed in deriving the frequencies for HzO1*. We have therefore performed a direct experiment to measure this quantity, which we find to be 1.045 A: 0.003 a t 15'. When an allowance is made for temperature, the agreement with calculated values is quite goodl. The actual numerical value of K is an essenti.al feature in arguments a b w t the mechanisms of hydrolysis which we have developed el~ewhere.~ Experimental The isotopic composition of a stock quantity of redistilled water was determined by equilibration with carbon dioxide.& Freshly cut sodium was held under this water to make a solution approximately 3 M in sodium hydroxide. This solution was held a t approximately 15" A measured volume of liquid was drawn off a t a rate of 2-4 ml./hr. under reduced pressure as vapor which was condensed and equilibrated with carbon dioxide. The volume and molarity of the remaining solution were measured. The vapor was drawn off slowly and was therefore assumed to be in isot,opic equilibrium with the solution. A value of 1.009 was taken from the graph of Dostrovsky and Ravivo for 01, the distillation separation factor of H2016 relative to H2O1*. The relative fugacities of HzO'B and HzO18are unaffected by Na+ ions.? Values of 1.044 i 0.004 and 1.046 & 0.004 were obtained for K in two experiments where the respective mean concentrations of sodium hydroxide were 4.8 and 4.3 M . I
The method of Hunt and Taube3 gives a value ,of 1.044 for K a t Oo, which leads to 1.039 a t 15' by interpolation. However their procedure really refers to a m equilibrium between aqueous OH- and HzO vapor. When an allowance is made for this by multiplying by a, K becomes 1.048 at 15'. Inclusion of the libration of OH- reduces this quantity to 1.046, which agrees well with the values observed. If it is assumed that Thornton's valueS has the same temperature depend(1) Department of Chemistry, Stanford University, Stanford, Californ:a. (2) H. R. H u n t and H. Taube, J . Phys. Chem., 63, 124 (1959). (3) E. R. Thornton, J . A m . C h e m . Sac., 8 4 , 2474 (1962). (4) &I. Green and H. Taube, submitted for publication in I n o r g . C h e m . ( 5 ) C . A. Mills and H. C. Urey, J . Am. C h e m . Soc., 62, 1019 (1940). ( 6 ) "Proc. of International Symposium on Isotope Separation," NorthHolland Publ. Co., Amsterdam, 1958, p. 337, ed. by J. Kistemaker, J. Bigeleiaen, m d A . 0. C. Nier. (7) H. M. Feder and H. Taube, J. C h e n . Phys., 2 0 , 1335 (1952).
COMMUNICATIONS TO THE EDITOR
1566
ence as that of Hunt and Taube, his method would give a value of 1.0425 a t 15'. A value of 1.045 f 0.003 a t 15' seems to be compatible with both experimental and theoretical data. A similar sort of argument points to a value of 1.040 f 9.003 a t 25'.
Vol. 67
Acknowledgment.-This work was supported by the National Science Foundation under Grants G-5411 and G-17422. hl. G. wishes to express his gratitude to the United States Educational Commission in the United Kingdom for a travel grant.
COMMUNICATIONS TO THE EDITOR CATIONIC ;CIBBILITIES IK FUSED CESIUM KITRATE AND THALLOUS SITRATE
Sir : In the course of tlie past few years many papers dealing with the determination of transport numbers of ions in pure ionic melts have been published. Most of the experiments were made with porous plug cells1-4 in which the plugs act as reference frames for the ionic velocities. I n our present work we used the experimental technique of electrophoresis on thin layers for the determination of single ionic mobilities in molten salts. The layer consisted of fine alumina powder sprayed on a sintered non-porous alumina support strip (30 cm. X 1.5 cm. X 2 mm. thick). The thickness of the layer was about 10 mg./cm.2. It was impregnated with the pure salt under test and doped at one end with a small amount of radioactive cations. The cell terminals consisted of two platinum electrodes immersed in the melt contained in crucibles. The electric connection between the crucibles and the strip was achieved by asbestos paper bridges. The potential gradient along the alumina strip during the experiments was measured by two auxiliary platinum wire electrodes in contact with the strip at both ends. The mobility data reported are referred to this potential gradient. Under normal experimental conditions the field strength was 3-6 v./cm. and the running time was 1-2 hr. The current was 10-45 ma., which corresponds to a maximum Joule heat of 0.2 w./cm.2. This value is small enough to avoid temperature differences along the strip. When the experiment was completed, the strip was cooled to room temperature and the activity distribution scanned by a G.M. window counter in which the aperture was 0.5 mm. I n Table I we compare a few results obtained for alkali nitrate melts with those obtained by other authors who determined the porous plug transport numbers in the same systems. We also give original results for cesium nitrate a t 450' and thallous nitrate a t 250'. The mobilities in the third column are the results of runs carried out a t different potential gradients. (1) B. B. Owens and F. R. Duke, Ames Laboratory, Iowa State College, Unclassified Report USAEC ISC-992. ( 2 ) R. W. Laity and F. R. Duke, J . Electrochem. Sac., 205, 197 (1958); "Metals Reference Book," Butterworths Scientific Publications, London,
1955, pp. 614-627. (3) A. Klemm, Discussiuns Faradag Sac., 83, 203 (1961). (4) E. D. Wolf and F. R. Duke, Ames Laboratory, I o w a State University, Unclassified Report USAEC 19-334.
TABLE I CATIONIC MOBILITIES OF PUREFCSED SALTS
x 104, cm.2 v.-1 sec.-1 u
T, Salt
CSNO~
Tlxoa
oc.
450
250
1.56 1.69 1.64 1.08 0.99 1,04
Av.
Previous work
1 . 6 3 zt 0.07
...
1.05iO.06
...
1.11
NaNOB
350
3.86 3.90 3.84
3.87 zt 0.07
3.86 i 0.05a
KNOa
350
2.04 2.08 2.12
2.08 f 0.06
2.21 f 0.11"
AgN03
250
Reference I .
2.57 2.68 2.57 f 0.11 2.47 Reference 2.
2.87 f 0.1gb
We feel that electrophoresis on thin layers in fused salts is an accurate and useful method for determining electrical transport properties of ionic melts. DIPARTIMENTO MATERIALI CHIMICA ALTE TEMPERATURE C. C. R. EURATOM ISPRA-VABESE, ITALY RECEIVED JANUARY 29, 1963
S. FORCHERI C. MONFRINI
PHOTOCHEXICAL REACTION OF HYDROGES BROMIDE WITH OLEFINS AT LOW TEMPERATURE
Sir: A recent electron spin resonance study of the interaction between ultraviolet irradiated hydrogen bromide and various olefins1 has been interpreted in terms of radicals formed by addition of a bromine atom to the double bond in such a manner that bonding is effectively symmetrica1.l The alternative reaction considered was hydrogen atom addition, but this was excluded on the grounds that the spectra were not those expected for alkyl radicals, and, for reaction with cyclopentene and 2-butene, spectra were identical when deuterium bromide was used instead of hydrogen bromide. (1) P. I. Abell and L. H. Piette, J . A m . Chem. Sac., 84,916 (1962).
