NOTES
2703
a
with the second chlorine, which approaches as the asymmetry increases, should lie below 13 Mc/sec. The observed resonance, near 20 Mc/sec, indicates that in this salt the anion asymmetry is quite marked. If we make the usual assumDtion of 15% . - s character in the chlorine borlding Orbitals to hydrogen’ we may deduce an approximate net-charge density distribution -0.57 C1
+0.43
C1-H
Ci
-0.86
a
(1)
H-C1
aHCi
(111)
where the Cl-H in (I) and (11) is the purely covalent structure, not the hydrogen chloride molecule, and the contributions of (I) and (11) are equal. With the observed coupling constant of about -40 Mc/sec and 15% s character for the C1 bonding orbital, the approximate charge distribution is calculated to be -0.57 +0.14 -0.57 C1- - - - H - - - c1
-
A simple LCAO-MO treatment of this linear, symmetrical model following the Townes-Dailey procedures yielded the expression = eq
I;;[
Qatom
-
x2
where s2 = 0.15 is the s character of the C1 bonding atomic orbital @cl and X is a measure of the ionic character of the three-center bonding MO
4
= A h
Thermodynamic Properties of Plutonium Mononitride from Electromotive Force Measurements
by G. M. Campbell and J. A. Leary Los Alamos Scientific Laboratory, University of California, Los Alamos, New Mexico (Received February 18, 1966)
(11)
4-
Qmol
(8) E. Clementi and A. D. McLean, J . Chem. Phys., 36,745 (1962).
- - - - H . . . . . . c1
for this unsymmetrical model. The confirmation of this interpretation by the detection of another resonance in the 6-Mc/sec region is desirable particularly in view of the fact that one can also rationalize the observed resonance in terms of a symmetrical model. The valencebond picture for this considers the three important contributing hybrids to be
eq
such that for the present particular case, the tetramethylammonium hydrogen dichloride salt, we prefer the interpretation in terms of the unsymmetrical anion model.
+ @a’-
@Cl”
Here @H is the Is hydrogen atomic orbital and @ ~ 1 ’ and @ ~ 1 ” are the sp hybrid AO’s of the two chlorine atoms. The electron distribution deduced in this way was essentially the same as that obtained from the resonance hybrid picture given above. It is not an unreasonable one; the most sophisticated SCF-LCAO-MO treatment for the analogous anion, symmetrical [FHF]-,* yielded the result that the hydrogen atom is essentially neutral. However, the evidence from the vibrational datal is
The free energy of formation of PUN has not previously been reported. We have determined the emf of the cell P u ( s ) ] P ~ + ~ (LiCl-KCl~/LiCI-KCl, z~), Ag+(xz
=
O.O00568)]Ag (I)
over the temperature range 685 to 729°K and the composition range x1 = 0.000594 to 0.0215 mole fraction.’ I n addition, potentials were determined for four cells of the type Nz(g, p l ) , P U N ( S ) I P U + ~ ( ~ ~ ) L ~IC LiC1-KC1, ~-KC~\ Ag+(z2 = O.O00568)/Ag (11) over the temperature range 695 to 716°K and the composition range z1= 0.0177 to 0.0193. The cells behaved in a reversible manner; Le., the emf was independent of the direction of temperature approach and obeyed the Nernst equation with respect concentration, 21. to nitrogen pressure and the P u + ~ The emf was independent of time after an initial 24-hr cell equilibration period for cells of type 11.
Experimental Section Proper preparation of the electrolyte is one of the most important steps in the over-all procedure. In addition to the customary drying process for fused chlorides,2 it was necessary to perform repeated equi(1) G. hl. Camobell and J. A. Learv. “Thermodvnamic Prouerties of’Pu Compoun’ds from Emf Measurements I. Pu vs. Ag inLiC1KCl Eutectic,” LA-3399 (1965). (2) H. A. Laitinen, W. S. Ferguson, and R. R. Orteryoung, J . Electrochem. SOC.,104, 516 (1957).
Volume YO, Number 8 August I
N
NOTES
2704
librations with PUmetal, followed each time by filtrations of the fused salt, to obtain an electrolyte sufficiently free of metal ion impurities. Each cell consisted of a 24-in. long Pyrex container sealed at the top by a standard taper glass joint. Electrical contacts for electrodes and thermocouples were directed through Kovar seals at the top of the cell. This made it possible to evacuate the cell and to maintain an inert atmosphere over the electrolyte and electrodes. The electrolyte was contained in a Ta cup. Electrodes were suspended from a T a plate at the top of the cell, using Pyrex tubing as insulators. The Ag-AgC1 electrode was made by placing a polished Ag wire in a special Pyrex compartment containing electrolyte prepared specially for this purpose. Electrical contact between the electrolyte and reference electrode was maintained by sealing an asbestos fiber into the bottom of the special compartment. The Pu electrode was made by clamping a 0.25-in. diameter Pu rod to the end of a 0.125-in. diameter T a rod. The Pu metal had been electrorefined to purity of at least 99.98 wt % . 3 The Pun’ electrode was madeby clamping a donut-shaped sintered pellet of PUN between W washers at the end of a Ta rod. Nitrogen content of the P U S was 5.5 =t 0.1% by weight (5.53 theoret), X-Ray powder diffraction analysis indicated that the material was single-phase PUN, with a. = 4.9055 A. Spectroscopically pure IT2 was maintained at a carefully measured pressure in the cell when the PuX electrode was being used. A chromel-alumel thermocouple calibrated against the melting point of Xational Bureau of Standards pure Zn metal was used in all temperature measurements. The thermocouple was inserted into a 0.125-in. diameter T a tube which was closed at one end and immersed in the molten electrolyte. The cell potential was monitored continuously on a Sargent Model 1IR recorder which was periodically checked for calibration against a RIinneapolis Honeywell Rubicon potentiometer. Potentials could be read to + O . l mv on the recorder. Thermocouple potentials were measured exclusively on the Rubicon pot entiomet er. The P u electrode came to equilibrium with the electrolyte as soon as the temperature was stabilized. When there was no measurable drift in potential for 1 hr at any one temperature, this potential was accepted as being reliable. The PUNelectrode required up to 24 hr to reach equilibrium with the electrolyte. Potentials were accepted after there was no measurable drift in a 3-hr period. The Journal of Physical Chemistry
Results and Discussion Measurements of eight cells of type I gave the potential, E in volts, as E
=
EO
- RT/3F (In zl)
where eo = 2.190 - 0.001002‘, R is the gas constant, T is the temperature in OK, and F is the Faraday constant. Results from type I1 cells are summarized in Table I. Table I: Results from Type I1 Cells Pu V 8 . Nz PUN Cell no.
us. Ag,
Temp,
V
OK
21
pressure, atm
1 1 1 2
0.6924 0.7006 0.6890 0.6892 0.6951 0.6970 0.6900 0.6993 0,6947 0,6933 0.6920 0.6886
713.4 715.8 713.5 698.5 697.6 708.1 707.7 703.8 701.5 702.5 694.8 712.4
0.0193 0.0193 0.0193 0.0179 0.0179 0.0179 0.0179 0.0179 0.0177 0.0177 0.0177 0.0177
0.685 0.331 0.649 0.682 0.691 0.703 0.701 0.696 0.691 0.691 0.691 0.691
2 2
2 3 4 4 4 4
PuNat 1 atm of N2, V
0.866 0.863 0.870 0.884 0.879 0.868 0.875 0.869 0.876
0.877 0.885 0.873
Combining the results from each type of cell and correcting the emf data of cell I1 to a nitrogen pressure of 1 atm gave the free energy A G O T = -99 0.0552’kcal/mole of PUN for the reaction
+
Pu(s)
+ 0.5Nz(g, p = 1 atm) = PuN(s)
This corresponds to a free energy of formation of -60 1 kcal/mole of PUN at 700°K. At this temperature the free energy of formation of UN is -56 kcal/mole. The data for type I1 cells were obtained over a temperature range of about 21 O. Therefore, standard heat and entropy terms derived from these data would be unreliable in spite of the precision of measurement of free energy. Although no measurements have been made of the standard heat, entropy, or free energy of formation of PUN, it is of interest to compare these results with previous estimates. Brewer has estimated AHo2ss= -95 kcal/mole.6 (3) L. J. Mullins and J. A. Leary, Ind. Eng. Chem. Process Design Develop., 4, 394 (1965). (4)M. H. Rand and 0. Kubaschewski, “The Thermochemical Properties of Uranium Compounds,” Interscience Publishers, Inc., New York, N. Y.,1963
NOTES
2705
Olson and Mulford6 have studied the equilibrium 0.5 N2(g) in the PuN(s) = Pu(1, saturated with N) temperature range 2290-2770’K. It was not possible to obtain the standard thermodynamic properties from these results because the activity of saturated liquid plutonium is not known. They have estimated AH’298 I -64 kcal/mole, favoring a value of -70 kcal/mole which is equal to that reported for UN.? Rand and Kubaschewski have indicated? that for nitrides having the NaCl structure, generally S(MN) - S(M) = 1 f 1. This leads to ASozg8 = -22 f 1 for the standard reaction. Using 11.39 cal/mole deg for ST,,,, s Z 9 d for P u , ~ 3.02 cal/mole deg for STo0- Szgsfor 0.5N2,9and an average heat capacity of 12 cal/mole deg, one obtains AHo2g9= -76 kcal/mole from the emf results. The standard free energy of formation of PUN from these emf data indicates that AG’700 = -60 kcal/mole. If the suggested6 value of AHo2g8 = -70 kcal/mole is accepted, this leads to A S o z ~ 8= -13 cal/mole deg, which is not in agreement with the value of 1 1 for S(PuN) - S(Pu). Although it is not possible to reconcile these differences, the correct value for AH’298 is probably more negative than -70 kcal/mole.
+
*
Acknowledgment. This work was sponsored by the
U. S. Atomic Energy Commission. (5) L. Brewer, et ai., “The Transuranium Elements,” McGraw-Hill Book Co., Inc., New York, N.Y., 1949,p 863. (6) W. M. Olson and R. N. R. Mulford, J. Phys. Chem., 68, 1048 (1964). (7) M. H. Rand and 0. Kubaschewski, “The Thermochemical Properties of Uranium Compounds,” John Wiley and Sons, Inc., New York, N. Y., 1963,p 41. (8) R. Hultgren, et al., “Selected Values of Thermodynamic Properties of Metals and Alloys,” John Wiley and Sons, Inc., New York, N. Y., 1963,p 226. (9) K. K. Kelley, Li. S. Bureau of Mines Bulletin No. 584, U. S a Government Printing Office, Washington, D. C., 1960, p 132.
Deuterium and Tritium Isotope Effects in the Methoxide-Promoted Elimination Reaction of 2,2-Diphenylethyl Benzenesulfonatel
by A. V. Willi2 Chemistry Department, Brookhaven National Laboratory, Upton, New York, and Columbia University, College of Pharmacy, New York, New York 100BSa (Received April 19, 1966)
A kinetic study of primary deuterium and tritium isotope effects at different temperatures was under-
taken with special interest in the problem of tunneling. The reaction chosen for this purpose was the elimination of 2,2-diphenylethyl benzenesulfonate under the action of sodium methoxide in Methyl Cellosolve solution (H” = H, D, or T) (CeH6)zCHzCHzO02SCeHsf CH30- + PhZC=CHz
+ CH30HZ + CsH6S03-
Since previous work by Shiner and co-workers4 was concerned with the ethoxide-promoted elimination reaction of l-bromo-2-phenylpropane, it was decided to study the more acidic 2,2-diphenylethyl system. In this work, the reactions of the compounds with H” = H or D could easily be followed (in separate experiments) by observing the increase in the ultraviolet absorption at 255 and 260 mp. The rate of the tritiumlabeled compound was measured by following the decrease of radioactivity in the toluene-soluble fraction of the reacting solution.
Experimental Section Methods of preparation of starting materials are analogous to those reported by Shiner and Smith4”: 2,2-diphenylethanol was made from ethyl diphenylacetate by reduction with LiA1H4 and was allowed to react with benzenesulfonyl chloride in ether solution in the presence of pyridine. The crude sulfonic ester was purified by repeated crystallizations from ethyl acetate-n-hexane. For the preparation of the isotope-substituted materials, ethyl diphenylacetate had been exchanged with ethanol-d (99.0% isotopic purity, purchased from Merck Sharp and Dohme of Canada Ltd., Montreal) or with ethanol-t (prepared by hydrolysis of ethyl o-formate with tritium-labeled water) in the presence of catalytic amounts of sodium eth~xide.~”Three subsequent exchanges were carried out for the preparation of the D compound, and the isotopic purity at the 2 position (99.0% or better) was checked with the aid of nmr spectra. The activity of the tritium-labeled material was 0.5 mcurie/mmole. In a product analysis experiment, 1.3 g of 2,2diphenylethyl benzenesulfonate and 0.3 g of sodium methoxide in 50 ml of Methyl Cellosolve were heated to 71’ for 28 hr. The main product was 1,l-diphenylethylene as identified by ultraviolet (A, 250 mp) and nmr spectra of the ether-soluble fraction. There (1) Research performed under the auspices of the U. S. Atomic Energy Commission. (2) Visiting chemist to Brookhaven National Laboratory, 1964. (3) Where inquiries regarding this paper should be sent. (4) (a) V. J. Shiner and M. L. Smith, J . Am. Chem. Soc., 83, 593 (1961); (b) V. J. Shiner and B. Martin, Pure Appl. Chem., 8 , 371 (1964).
Volume 70, Number 8 August 1966
