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
SPECIFIC REACTION RATECONSTANTS FOR THE REDUCTION OF PLUTONYL IONWITH Pu+*IN PERCHLORIC ACIDSOLUTIONS AT 3" ki, 1. W+), mole - 1 (PUOZ (Pu +a) i , p rnim-1 moles/l. molea/l. molea/l).i*
+'
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.
10-
1.00 1.00 0.50 10-4 0.50 10-4 2.00 attained within
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+~) ( P u O t ) , 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)
NOTES
814
years on the oxidation, pyrolysis and photolysis of NH3 and N2H4. The results of these various studies have not as yet permitted the establishment of a consistent set of elementary mechanisms which adequately explains the experimental observations. There is thus a need to study in more detail the various elementary reactions which occur in these systems. I n order to obtain some information along these lines we have studied, both spectroscopically and by the chemical analysis of the trapped products, the reactions of hydrogen atoms and oxygen atoms with NH3 and NzH4in atomic flames. Experimental The atomic flame apparatus which was employed is shown schematically in Fig. 1. The reaction chamber was provided with a quartz window which permitted spectrographic observations to be made perpendicular to the flow of the reactants. Power for the dmharge was supplied from a 21/2 h a . transformer with a 5000 volt secondary operated from a Variac. The concentrations of a t o m issuing from the .discharge were determined semi-quantitatively by the isothermal filament technique.' The a a m e n t was withdrawn from the reaction chamber before the runs were started. The pressure in the reaction vessel was measured by a tilt-type McLeod gage. The pressures employed were in the range -0.5-1.0 mm. for most of the experimenta described in this paper.
Vol. 60
trapped products were carried out as follows: the total base was determined by titration with HCI to the methyl red endpoint; the hydrazine was then determined by titration in -5 HCI with KIOs solution using Brilliant Scarlet 3-R as an indicator.8 The difference between the total base and the hydrazine was taken to represent the amount of ammonia present.
Results and Discussion The chemical and spectroscopic data obtained in this study of the atomic flames of 0 or H atoms reacting with NH3 and with NzH4 are summarized in Table I and discussed below. TABLE I SUMMARYOF EXPERIMENTAL RESULTS Reac-
tants 0 4- NHs 0 N2H4
Emission band@ of N H OH NO .. vw B 8 8 s
NHa
Trap analysis 8 0 4 5 % of NHadestroyed No NHa found, all of the NZHI destroyed >95% of original NHa collected N2H4 mostly decomposed, 3055% of max. possible NHI collected
.. ..
+ H + NHi . . . . . . .. H + NaHk m w .. .. E,
strong; m, moderately strong; w, weak; vw, very
weak.
The band systems observed were the a-bands of NH2 in the region from -4000 A. to the long wave lzngth limits of the photographic plate a t ~ 6 5 0 0 A., the arI-3Z bands of N H at 3360 to 3370 A., the tZ-2rI bands of OH, especially the (0,O) band a t 3064 A., and the y-bands of NO which were strongest in the region from 2250-2850 An additional spectral feature observed in some of the atomic flames was the greenish-yellow air afterglow due to the recombination reaction4
A.
PRESSURE GAUGE
NO
+0
=
NO2
+ hv
+
Leads to Wheatstone
Y II
p".p*;l J
B
Fig. 1.-Schematic diagram of atomic flame apparatus showing: A, 1-1. reaction vessel; B, cold trap; C, ball and socket joint for inserting probes; D, discharge tube; E, tubular A1 electrodes; F, platinum filament; and G, H and J, capillaries to control flow rates of reactants. Pumplng speed -3 I./sec. All the spectrographic observations were made with a Bausch and Lomb medium uarts spectrograph on Eastman 103a-B plates. Slit widthsqfrom 100 to 1.0 mm. were employed in conjunction with exposure times ranging from 10 min. to about five hours. Commercial H,, 02 and NHa, all stated to exceed 99.6% purity, were used for these experiments without further purification. The hydrazine assayed -9774, the remainder being mostly water. Chemical analyses for ammonia and hydrazine in the (2) See e.g., -1. R . Dingle and D. J. LeRoy. J . Clrem. P h y r . . 18, 1632 (1950).
(1)
This continuum was found to be quite strong in cases where a large excess of atomic oxygen reacted with the nitrogen-containing f u e l ~ . ~ Reactions with Atomic Oxygen. (a) NH3 0. -The emission from the NH3 0 flame was very weak. Exposure times of about three hours with a slit of 100 p were required to record even the ttrong central &-branch of the N H band at 3360 A. The OH bands also appeared weakly in such :In exposure. However, a blank run with the same flow of oxygen and the same discharge voltage but with no ammonia present yielded the OH spectrum with even greater intensity. This indicates that the observed OH emission can be accounted for completely by the H20 impurity in the oxygen. The other bands prominent in the premixed atmospheric pressure flame of ammonia and oxygen6 (NH2,NO) were not observed under our experimental conditions. It should be noted, however, that bands of XH2 with intensity comparable to that found for N H could have been masked in these experiments by the overlapping emission of the air afterglow . Since the reaction was non-luminous, we were
+
(3) R . A. Penneman and L. F. Audrieth, Anal. Chrm., 40, 1058 (1948).
(4) A. G. Gaydon, "Spectroscopy and Combustion Theory," Chai)man and Hall, Ltd., London, 1948. p. 101. (5) The air afterglow is also obtained weakly in a discliarw tlirl)ll&!(h commercial tank oxygen owing to the inial1 aiiiount of riitrrigc.n impurities present. ( 0 ) B . B. F'ogarty and 11. C . Wolfhard, ,Vatrtrp. 168, 1122 (1951).
NOTES
June, 1956 unable to observe directly the size of the reaction zone. However, the fact that some ammonia was always found in the trap, even in the presence of a considerable excess of atomic oxygen, suggests that the reaction is a fairly slow one. (b) NzH4 0.-The observed emission spectrum of the NzH4 0 atomic flame consisted of bands of NH2, NH, OH and NO. These bands were all at least two orders of magnitude more intense than their counterparts in the NHI 0 atomic flame in the same apparatus and under the same experimental conditions. Ten minute exposures with a 100 p slit, width gave well exposed plates for all these species. The luminous zone of the NzH4 0 atomic flame is very small compared with that of other atomic flames a t the same pressure. The reaction volume is, for example, less than one tenth of the size of that for a comparable C2H2 0 atomic flame. This shows that the reaction between NzH4 and 0 is a very fast one. When a small amount of NZH4 is added to a large excess of atomic oxygen, the air afterglow appears very strongly below the mixing zone. This indicates that the oxidation of hydrazine by atomic oxygen involves the formation of NO, a fact which is also borne out by the emission of the 7-bands of NO in the reaction zone. The trap analysis showed that all the original hydrazine was destroyed in the reaction vessel. Since the thermal decomposition of hydrazine yields NH3'3 and since no NH3 was found in the trap, thermal decomposition of the N2H4 would seem to be ruled out. Qualitative analysis of the trap contents by the brown ring test indicated the presence of oxides of nitrogen. Reactions with Atomic Hydrogen. (a) NH3 H.-In agreement with Dixong no evidence was found which would indicate that NHI reacts with atomic hydrogen. No light emission was observed under our experimental conditions and with exposure times of -3 hr., and the ammonia was recovered quantitatively in the trap. Our techniques could not, however, be expected to provide any direct evidence for possible exchange reactions or stripping-reforming reactions which may proceed without the formation of electronically excited intermediates. It was found that the energy released by hydrogen atom recombination to a platinum filament inserted into the reaction chamber decreased with increasing partial pressure of ammonia under constant discharge conditions. This would seem to indicate that the presence of ammonia in the reaction system reduces the steady-state coricentration of atomic hydrogen. Our observations can readily be accounted for on the basis that the NHI reacts with the H atoms only by an exchange reaction of the type proposed by Farkasand Melville'o
+ +
+
+
+
+
NHgH'
+ H +NH3 + H'
(2)
(7) R. C. hlurry and A. R . Hall, Trans. Faraday Soc., 47, 742 (1951). ( 8 ) G . K. Adanis and G. W. Stocks, "Fourth Syiriposiiiin on Conibustion." The Willinnis and Wilkins C o . , Baitinlore, 1952. (9) .J. K . Dixon, J . Am. Chern. Sor., 54, 4 Z i 2 (1932). ( I O ) A. Barkas and 11. W. RIrirille, Pror. Rou. S o r . ( / , o n d o , , ) . 1578, 625 (193G).
815
and that the decrease in the steady state concentration of H atoms is due to a recombination with NH3 as the third body. (b) NzH4 H.-The reaction between hydrazine and atomic hydrogen took place with the emission of moderately strong NH2 bands and relatively weak N H bands. Trap analyses showed, in agreement with Dixon's resultslgthat ammonia was a major reaction product. The ratio p = (moles of NH3 formed/moles of N2H4 decomposed) was observed to depend on the partial pressure of the reactants and was found to vary from about 0.45 to about 1.14 f 0.05. In general a higher ratio of hydrazine to atomic hydrogen a t a given pressure favored the production of ammonia. The luminous reaction volume, although considerably larger than that for the reaction N2H4 0 at the same pressure, did not fill the reaction chamber. Order of magnitude calculations show that the observed reaction volume is compatible with the rate constant for the hydrazine-atomic hydrogen reaction obtained by Birse and Melville." The dependence of p on atomic hydrogen concentration suggests that the reaction between H atoms and N2H4 proceeds by at least two competing mechanisms of different order in H atoms.
+
+
(11) E. A. B. Birae and H. W. Melville, ibid., 175A, 1G4 (1940).
APPEARANCE POTENTIAL STUDIES. I' BY STEPHEN S. FRIED LAND^ A N D ROBERT E. STRAKNA Department of Physics, The Uniuersity of Connecticut, Storre, Connecticut Received December 88, 1966
The measurement of appearance potentials of molecular fragments leads to some understanding of molecular structure (for example, see ref. 3). I n addition, the magnitude of the appearance potential of polyatomic molecules, such as the esters and alcohols, may be used elsewhere. The proper combination of a polyatomic gas with an inert gas in a self-quenching Geiger-Muller counter is determined by the appearance potential of the polyatomic molecule. This is possible since the quenching process of the gas discharge in the counter depends upon the probability of electron transfer from the inert gas to the polyatomic m ~ l e c u l e . ~That is, the ionization potential of the inert gas has to be higher than the ionization potential of the polyatomic molecule or of the appearance potential of any of its fragments (if the fragment is t o participate in the quenching action) ; the closer the two values, the higher the probability of electron transfer and thus the more effective quenching action. Further, the value of the appearance potential of the fragments may be utilized to determine whether the fragments formed in the discharge5 will also act as quenching agents. (1) W o r k supported by the Research Corporation, New York City. (2) On leave with the Nuclear Development Corporation of America, White Plains, New York. (3) (a) H. D . Hagstrum, Rev. Modern Phys., 13, 185 (1951); ( b €1. B. Rosenstock, H. B. Wallenstein. A . I.. Wahvhaftig and, H. Eyrina, Pror. Nall. Acad. Sci., 38, (107 (1952). (4) S. A . Korff and R . D. Present, Phi/.*. Rei).. 65, 275 (1944). (5) R. S. Friedland, ibid., 74, 898 (1948).
