NOTES
812
were read on a Wallace-Tiernan gage graduated in centimeters and having a recision of about =k1 mm. Procedure .-In t t e determination of vapor pressures, about 25 ml. of 2,2,3-trichloroheptafluorobutane was admitted to the evacuated still and a small amount of nitrogen added as a confining gas. The temperature waa then slowly raised until the condensed vapor was observed as (t Rteady drip through a fliiorothone observation window. After temperaturepressure equilibrium was established, the vapor pressure and temperature were recorded. An additional small amount of nitrogen (usually 1 to 5 cm.) was added to the still and the temperature raised as before until a second temperature-vapor pressure relationship WB8 obtained. This procedure was continued until the vapor The process of adding nipressure was measured to 173'. trogen was then reversed and small amounts of nitrogen were removed, with subsequent temperature decrease allowing vapor pressure readings to be made a t various points as the temperature was lowered.
Results.-The
vapor pressure equation
1131.324 log Pmm = 6.69356 - t 200
+
(4)
was fitted to the experimental data for the range 29.58 to 172.68'. Uncertainty of the equation, based on confidence limits a t the 0.95 level, is approximately 1.1%. The experimental vapor pressures of 2,2,3-trichloroheptafluorobutaneare presented in Table 11. The boiling point at 760 mm. pressure calculated from equation 4 is 96.7'. TABLE I1
*
VAPORPRESSURE OF 2,2,3-
TRICHLOROHEPTAFLUOROBUTANE Teomz,
29.58 31.25 42.15 55.01 62.66 73 * 33 78.40 94.99 98.23 98.91 103.47 106.27 115.82 121.37 123.30 138.29 140.39 145.53 148.48 166.66 170.02 172.68
Vapor pressure, mm. From eq. 4
Exptl.
58.5 65.0 104.4 178.1 240.5 355.0 424.5 714 809 825 908 1012 1293 1508 1545 2212 2317 3597 2822 4052 4354 4632
58.3 64.0 105.0 178.5 243.4 358.5 426.5 721.8 794.5 810.5 923.9 999.3 1292 1490 1564 2235 2344 2627 2800 4056 4326 4549
Dev. of expt. from eq. values
+ 0.2 + 1.0 - 0.6 - 0.4
- 2.9 - 3.5 - 2.0 - 7.8 +14.6 +14.5 -15.9 +12.7 + I +18 - 19 -23 -27 -30 +22 -4 +28 83
+
Heat of Vaporization.-The heat of vaporization of 2,2,3-trichloroheptafluorobutaneat various temperatures may be calculated from the exact form of the Clapeyron-Clausius equation gwen m equation 5 provided the vapor pressure,. liquid density and deviation of the vapor from ideality are known for the given temperature range. AH
dP tAV dT
2604.964P
n = ( t + 200)'
Since
v
=
v. - VI (7)
substitution of equations 6 and 7 into 5 yields 2604.964PT =( t f ) i
where
RT (P(*+Apj
- );
(8)
AHv = heat of vaporization per mole
P !i T R
nz
2
vapor pressure temperature, "C. temperature, O K . ideal gas constant molecular weight density = non-ideality coefficient (determined by Magnuson)s = = = = = =
Using equation 8, the heat of vaporization of 2,2,3trichloroheptafluorobutane was calculated for the temperatures shown in Table 111. TABLE I11 HEATOF VAPORIZATION
OF
2,2,3-
TRICHLOROHEPTAFLUOROBUTANE Temp., OC.
Heat of Vaporization, cal./mole
60 80 100 120 140
8207 7834 7487 7135 6728
Acknowledgments.-The authors wish to express appreciation to Mr. A. V. Faloon for purifying the 2,2,3-trichloroheptafluorobutaneused in these studies and to Dr. E. J. Barber for directing the work. (5) D. W. Magnuson. Carbide and Carbon Chemioala Company, K-25 Plant Report No. K-1130 (1954).
THE KINETICS OF THE REACTION BETWEEN Pu'~ AND P u O ~ + ~ ~ BY A. E. OGARDAND S. W. RABIDEAU Lou AEamoe Scionti& Laboratory, Undvsrsity of CaZifornia, Loa Alamoa, New Mexico Received December 6, 1066
In the course of the determination of the formal potential of the PuO$-PUO$~ couple in molar perchloric acid as a function of temperature,2 it was noted that near 3")following the addition of :P a to the plutonyl ion, approximately one minute was required for the establishment of a stable potential. It was of interest to determine whether a detectable slowness in the reaction Pu+S
+ PUOZ+*
kl
kz
Pu+4
+ PuOpf
(1)
(5)
could be observed spectrophotometrically. Connick8 has pointed out that this reaction usually can be considered to be in rapid reversible equilibrium. Heretofore no slowness has been detected in this reaction.
(6)
( 1 ) This work waa done under the auspices of the Atomic Energy Commission. (2) 8. W. Rabideau, J . A m . Chem. SOC.,78, June (1956). (3) R. E. Connick, ibid., 71, 1528 (1949).
Differentiation of equation 4 gives dP
Vol. 60
NOTES
June, 1956 Experimental The plutonium solutions were prepared by dissolving a weighed quantity of oxide-free metal in the re uired quantity of standardized J. T. Baker 71% perkloric acid. The plutonyl solutions were prepared by the prolonged ozonization of the Pd3 solutions. Water redistilled from alkaline potassium permanganate was used in the preparation of all solutions. The spectrophotometric results were obtained with the Cary Model 14 recording spectrophotometer. Water a t 5" was circulated through the walls of the cell compartment to minimize temperature changes of the solutions during the period of measurement. Dry helium was flushed through the cell, reference and phototube compartments to avoid condensation of moisture on the optical surfaces. A double-chambered spectrophotometric mixing cell was used. Known weights of solutions of P? and PuO:* were placed in their respective compartments. The cell was then immersed in an ice-bath for about an hour. At the end of this time the windows were dried, then the solutions were mixed in an atmosphere of helium which prevented the condensation of moisture on the cell windows. The first spectrophotometric readings were obtained approximately thirty seconds after mixing.
813
TABLE I1 REACTION RATECONSTANTS FOR THE REDUCTION OF PLUTONYL IONWITH Pu+*IN PERCHLORIC ACIDSOLUTIONS SPECIFIC
AT
+'
(PUOZ molea/l).i*
3"
1.488 X IO-* 1.250 X 5.583 X 10-3 4.703 X 4.099 X 8.087 X 3.838 X 8.037 X 5.108 x IO-' 3.749 x Equilibrium essentially mixing.
ki, 1. mole - 1 rnim-1
W+),
(Pu +a) i , moles/l.
molea/l.
10-
1.00 1.00 0.50 10-4 0.50 10-4 2.00 attained within
p
1 80 1 90 1 72 0.5 52 2 .. the time of
From a comparison of the results obtained.for k l at ionic strengths of 0.5, 1 and 2, it appears that an ionic strength dependence of kl is shown, and in
Results From qualitative spectrophotometric observa$3 +2 tions of mixed solutions of Pu and PuOz in molar perchloric acid at 3", the gradual disappearance of P i 3 (at 600 mp) and the appearance of P i 4 (at 652 mp) were noted which demonstrated a measurable slowness in reaction (1). The rate law for this reaction can be written -d(Pu+")/dt = k l ( P ~ +(PuO):*){I ~
- Keq/K*)
(2) +3
where K,, is the equilibrium quotient, (Pu (PUO;~)/(PU+~) (PuOt), and K* is this quotient a t any specified time, t. It has been found2that the value of K e , at 3" is 4.3. The value of ICl was obtained from the slopes of the straight line plots of d ( P ~ 3 ) ) / ( P ~ 3 versus )dt (PuO:~) ( 1 - Keq./K*J. The data for a typical experiment are given in Table I. The d(Pt3)/dt TABLE I THEREACTION OF Pu+3 A N D Pu02+*IN 1 M HClO, Time, inin.
(Pu+J) X IO4, moles/l.
0 1.0 1.5 2.0 3.0 4.0 5.0
4.70 3.13 2.69 2.34 1.84 1.59 1.45
(PuOB+*) d(Pu+*)/dt X IO', X 10+a, moles 1. - 1 moles/l. min. -1
5.58 5.43 5.38 5.35 5.30 5.27 5.26
... 1.04 0.78 .61 .32 .18
.os
AT
3"
...
(1
Keq./K*)
0.938 .879 ,809 .637 .503 .405
values were derived from measurements of slopes of +a plots of (Pu ) us. time. As a test of the mechanism of the reaction, the initial concentrations of each of the plutonium solutions were increased by a factor of approximately three. A value of k , was obtained in this experiment which agreed with the previously determined rate constant within 11%. I n another experiment the acidity was reduced to 0.50 M , but the ionic strength was maintained a t unity with added sodium perchlorate. Rate determinations were also made a t ionic strengths of 0.5 and 2. These results are summarized in Table 11.
0
O
i.0
I
3.0
I
I
4.0
5.0
(Pu02+")11 - K.,./K*l x 108. Fig. 1.-Evaluation of rate constant in molar perchloric acid at 3" for the reaction PuOz+e Pu+3 Pu+' Pu02+.
+
--f
+
agreement with the predicted rates for reaction (1) the value of ICl is greatest in solutions of highest ionic strength. However, since the variation in the values of klis rather large even a t constant ionic strength and acidity, it is difficult to distinguish with a high degree of 'certainty between ionic strength and acidity effects from these data.
THE REACTIONS OF AMMONIA AND HYDRAZINE WITH OXYGEN ATOMS AND HYDROGEN ATOMS IN ATOMIC FLAMES' BY GORDONE. MOORE,KURTE. SHULER,'~ SHIRLEIOH SILVERMAN A N D ROBERT HERMAN Contribution f r o m the Applied Physics Laboratory, The Johns I i o p k i n s University, Silver Spring, Maryland, and the Department of Physics, Uniueraity of Maryland, College P a r k , Maryland Received December 9 , 1966
There has been much work reported in recent A portion of this work waa supported by the Bureau of Ordnanco. Departiiient of the Navy, under Contract NOrd-738G and by Guggcnheiin Brothers, New Y o r k , under Contract GU-1. ( l a ) National Bureau of Standards, Washington, D. C. (1)
