414
SHERMAN W. RABIDEAU
absorbancy index is attributed t o a decrease in concentration of light absorbing species with decreasing temperature. It is concluded that when NiCL dissolves in a LiCl-KCl mixture a t least two light a?.porbiiig species are formed, and that these species are in equilibrium such that those associated with bands I and I1 decrease in concentration with increasing temperature while those associated with bands 111, IV and V increase in concentration with increasing temperature. Small changes in the composition of the solvent salt were found to have a small but measurable effect on the absorption spectra. This effect is best shown in terms of the ratio of the absorbancy values of bands I1 and IV, that is, (&)II/(&)Iv. In Fig. 6 this ratio is plotted against mole % KC1 a t two temperatures, 364 and 398". These temperatures are near the freezing point of the eutectic so that the accessible composition range for LiC1KC1 mixtures was quite small. Values of the molar Concentration of NiClz are indicated for each point in Fig. 6. It will be noted that, to within experimental error, the effect of solvent composition is independent of the NiC12 concentration for the range of compositions measured. The absorbancy ratio for the 34.4 mole % KCl mixture was extrapolated from measurements a t 416" and above. This extrapolation was necessary because this mixture freezes a t the lower temperature while for mixtures near the eutectic at temperatures above 400°,band I1 is obliterated by overlap of band 111. At a constant temperature a change in the relative absorbancies of bands I1 and IV could be a
Vol. 62
measure of a change in the relative concentrations of the two species which give rise to them (ignoring the effect of band overlap), and these relative concentrations might be expected to shift with a change in the cation rakio of the solvent salt because of a corresponding shift either in the activity coefficients of the two species or in the equilibrium constant. On the other hand, chloronickel(I1) complex anions should be surrounded by Li+ and K+ cations which differ greatly in their ability to polarize neighboring ions. Consequently, a change in the average of the ratio of Li+ to K+ ions surrounding a chloronickel(I1) complex anion might well be expected t o have some effect on the width or even on the oscillator strength of absorption bands of the anion. The existing experimental data do not permit a separation of the contribution of the latter effects from that of a change in concentration. Absorption spectra measurements will be continued on this and other fused salt systems. I n addition, consideration is being given to measurements of paramagnetic susceptibility for fused salt systems. Such measurements, particularly in the case of nickel, should provide information regarding the electronic and spatial configurations of the complex species present. Acknowledgment.-The authors gratefully acknowledge the assistance of Mr. D. E. LaValle who prepared the compounds, Mr. D. F. Anthrop who made many numerical computations, and Dr. W. R. Laing and co-workers who did the Li and K analyses.
KINETICS OF OXIDATION-REDUCTION REACTIONS OF PLUTONIUM. THE REACTION BETWEEN PLUTONIUM(V1) AND VANADIUM(II1) IN PERCHLORATE SOLUTION1 BY SHERMAN W. RABIDEAU Contribution from the University of California, Los Alamos ScientiJc Laboratoru, Los Alamos, N . ?If. Received October t8#1967
The kinetics of the reaction between PUOZ +z and V f a has been studied by spectrophotometric measurements at 8304 A. The stoichiometry of the reaction is given by the equation PuOZf2 V+a HzO + PuOz+ VO +* 2H +. This secondorder reaction has been found to be dependent upon both an inverse first and an inverse second power of the hydrogen ion concentration. Values of k' and k " , the rate constants associated with the inverse first and the inverse second power of the hydrogen ion concentration, have been found to be 2.12 f 0.03 sec.-1 and 0.228 =!= 0.006 mole liter-' sec.-l, respectively, for perchlorate solutions of ionic strength two a t a temperature of 25'.
+
Introduction I n view of the observation that the rate of rea;can be followed tion between PUOZ+~and P u + ~ spectrophotometrically, interest has developed in the kinetics of other oxidation-reduction reactions previously thought to be too rapid to measure without the use of very rapid-mixing techniques. Also, @rice the absorption peak of plutonyl ion a t 8304 A. (molar absorptivity ca. 550 M-l cm.-l) (1) T h i s work was done under t h e auspices of the U. S. Atomic Energy Commission. (2) A. E. Ogard a n d S. W. R a b i d e a u , THIS JOURNAL, 60,812 (1956) : S. W. Rabideau a n d R. J. Kline, i b i d . , t o be published.
+
+
+
can be used to follow the progress of the reaction with spectrophotometric recording methods, the reactant concentrations can be as low as M to bring the reaction rates into a conveniently measurable range without the loss of precision. Experimental Preparation of Reagents.-Plutonium(V1) perchlorate solutions were prepared from oxide-free plutonium metal especially selected from a lot of high purity. A known weight of metal was dissolved in a weighed quantity of standardized 70oJ, perchloric acid. The oxidation to the plutonyl state was accomplished by prolonged ozonization, ca. 24 hours. An additional period of oxidation with ozone was performed on the diluted stock solution prior to
F
April, 1958
REACTION BETWEEN PLUTONIUM AND VANADIUM IN PERCHLORATE SOLUTION
use in the reaction with vanadium(II1). At the termination of the period of oxidation, the solution was flushed with helium to remove dissolved ozone. Baker and Adamson reagent grade 70% perchloric acid was boiled to remove traces of organic material; then filtered through a sintered glass disc prior to use. Standardization of this acid was performed on a weight aliquot basis with mercuric oxide used as the primary standard. All solutions were prepared with distilled water which had been redistilled from alkaline permanganate in an all-Pyrex apparatus. Sodium perchlorate was prepared by neutralizing C.P. sodium carbonate with perchloric acid, then recrystallizing the sodium perchlorate from water. Lithium perchlorate was similarly repared as the trihydrate and recrystallized from water. "his salt was analyzed for water and for perchlorate content before use. Mallinckrodt AR grade acetone and ammonium thiocyanate together with Baker and Adamson reagent grade stannous chloride were used in the spectrophotometric analysis of vanadium( 111). Vanadium( V) oxide was prepared by the ignition in air of Fisher purified grade ammonium metavanadate which had been recrystallized from water. Vanadium( 111) perchlorate was prepared by the electrolytic reduction of a 1 M perchloric acid suspension of vanadium(V) oxide with platinized platinum electrodes. The reduction apparatus consisted of a 50/50 Pyrex joint the male section of which was formed into a flat-bottomed container, was fitted with a wide tube for the entrance of the platinum cathode through a Teflon plug, and was provided with a stopcock through which the vanadium( 111) perchlorate was removed a t the termination of the electrolysis. The female section of the joint was equipped with a Dewar-sealed re-entrant tube terminating in a sintered glass disc of medium porosity which served to separate the anode and cathode compartments. A gas inlet tube was provided through which nitrogen purified with chromous chloride was admitted in the cathode compartment of the electrolysis cell to prevent the air oxidation of the vanadium(II1) perchlorate. The anode was supported in the re-entrant tube with a Teflon plug. Suitably located gas outlet tubes were provided. Analysis of Vanadium( 111) Solutions .-The stock solutions of vanadium( 111) were analyzed spectrophotometrically at 3960 A. as the thiocyanate complex using a minor modification of the procedure given by Crouthamel, et aL3 It was found that freshly prepared vanadium(II1) solutions when converted into the thiocyanate complex gave identical optical densities either with or without added stannous chloride. Thus, since it appears that the vanadium(II1) perchlorate is not significantly oxidized to vanadium(1V) under the conditions of this analysis, even in the absence of added stannous chloride, the vanadium( 111) solutions were analyzed without added reductant. Solutions of known total vanadium content were analyzed routinely as controls. Standard vanadium solutions were prepared by dissolving 99.7% pure fused vanadium metal obtained from A. D. McKay Co. in nitric acid solution; then removing the nitric acid by boiling with hydrochloric acid. The molar absorptivity e of the vanadium thiocyanate complex at a temperature of 25" was measured to be 12,250 f 93. I n good agreement with the results of Furman and the molar abso:ptivities of vanadium( 111) perchlorate at 4000 and 5900 A. at a temperature of 25" were found to be 8.42 and 5.57, respectively. Estimations of the vanadium(111) perchlorate stock solution concentrations were made by use of these molar absorptivities, and the final concentrations of the diluted stock solutions were obtained with the thiocyanate method. Spectrophotometric Measurement.-The progress of the reaction was followed by means of measurements a t 8304 A . with the Cary Model 14 recording spectrophotometer. The solutions of vanadium( 111) and plutonium( VI) perchlorgte were placed in the separate legs of a double-chambered 10 cm. spectrophotometric cell which had been flushed with nitrogen. The solutions were then brought to constant temperature within f0.1' by immersing the cell in a constant-temperature water-bath without mixing the solutions. At time zero, the solutions were mixed by removing the mixing cell from the constant temperature bath and tilting (3) C. E. Crouthamel, B. E. Hjelte and C. E. Johnson, Anal. Chem., 27, 507 (1955). (4) S. C. Furman and C. S. Garner, J . Am. Chem. Soc., 72, 1785 (1950).
415
the cell back and forth several times. The mixing cell was placed in the thermostated cell holder within 30 seconds of the time of mixing. In the spectrophotometric cell compartment, the cell was submerged in the water-bath surmounting a modifieds Cary cell holder which was maintained a t a temperature constant to within +0.2' by circulation of water through the heavy metal base of the modified cell holder. The walls of the cell compartment bath contained gasketed quartz windows for transmission of the light beam from the spectrophotometer. At temperatures below the dew point of the room, it was necessary to pass a current of dry helium across the faces of the windows to prevent fogging.
Results and Discussion Reaction Stoichiometry.-In a set of experiments designed to examine the stoichiometry of the reaction between plutonium(V1) and vanadium(III), the rate of disappearance of P u O ~ + was ~ followed a t 9512 8. (molar absorptivity = 24.8 M-I cm.-l) and tbe rate of appearance of Pu02+ was observed at 5690 A. (molar absorptivity = 17.1 M-I cm.-') a t a temperature of 2.4'. The initial reactant concentrations in the two experiments were essentially equal (ca. 3 X M ) . The rate of disappearance of P u O ~ +was ~ matched by a corresponding rate of increase of PuOz+ in accord with the assumption that the stoichiometry is given by the reaction P U O Z + ~ V+a He0 = PuOZ++ VO+*+ 2H+ (1) Although the usual procedure in this study has been to follow the disappearance of P u O ~ + through ~ the absorption a t 8304 8.,the progress of the reaction was followed by another stoichiometry check by observing the rate of formation of VO+Z from measurements at 7500 8. A rate constant of 0.24 liter mole-' sec.-l was obtained which is in good agreement with the values obtained at this temperature from measurements of the rate of disappearance of P U O Z + (vide ~ infra). It would appear that under the conditions of this study the rate of reaction between V+a and PuOz+ is rather small in comparison with the rate of re. at the plutonium conduction of P U O ~ + ~Also, centrations used, the disproportionation of PuOz+ appears to be negligibly small especially in the early phases of these experiments, because of the low plutonium(V) concentrations. Order of Reaction.-In Table I are given some of the results of experiments in which the second-order + ~ V+* character of the reaction between P U O ~ and was demonstrated. The apparent rate constant, k l , was calculated from the integrated equation for a bimolecular process
+
k -
+
2.303
- t[(PuOz+2)0 -
(V+3)01
(V+%i[(PUO2+2)0log (PuOz+*)o[(V+% -
51
(2)
21
in which the subscript zero refers to initial concentrations and x is the concentration of PuOz+ or VO+2formed a t time t. The linear relation between plots of the logaiithmic term in equation 2 versus time has been observed. Acidity and Ionic Strength Dependence.-The hydrogen ion concentration dependence of the reaction rate was obtained in a series of measurements ( 5 ) The modified cell holder was designed by T.W. Newton and F. Baker of this Laboratory.
SHERMAN W. RABIDEAU
416
TABLE I EVALUATION OF SPECIFIC RATECONSTANTS FOR Pu(V1)V(II1) REACTIONIN MOLARPERCHLORIC ACID AT 2.4' (Pu02+2)0 x 10-4,M ( V + % X 10-4, M k l , 1. mole-1 set.-' 1.911 2.201 0.23 1.736 6.818 .20 24.61 24.52 .24
a t a constant ionic strength of two a t a temperature of 25'. Sodium perchlorate was substituted for perchloric acid to maintain the ionic strength a t a constant value. The results are given in Table 11. TABLE I1 IONCONCENTRATION DEPENDENCE OF REACBETWEEN Pu(V1) A N D V(II1) IN PERCHLORATE SOLU-
HYDROQEN TION
TION
[PuOz +210
x
104,
M
1.939 1.718 2.119 2.123 2.122 1.666 1.704 1.910 2.085 1.731 1.734 2.049 2.081 2.003
[V' a h 104, M
[I€-], M
kl (obsd.). 1. mole-' set.-'
5.328 9.121 10.138 5.067 5.371 9.333 4.493 5.291 3.055 4.963 2.226 2.179 1.203 2.028
2.000 2.000 1.500 1.500 1.000 1.000 0.500 ,500 ,250 ,250 ,150 ,150 .IO0 .lo0
1.16 1.15 1.56 1.59 2.23 2.25 4.88 4.70 12.49 12.42 24.49 24.67 43.6 43.75
x
kl (ralcd.), 1. mole-' sec.' 1
1.12 1.51 2.34 4.80
VOl. 62
ionic strength and a t a temperature of 25'. The results of these experiments are given in Table 111. By the method of least squares the values of k' and k" for the perchlorate solutions of unit ionic strength and 25' have been computed to be 1.81 f 0.05 set.-' and 0.200 0.010 mole liter-' sec.-I, respectively. Thus, from a comparison of Tables I1 and 111it is seen that the effect of an increase in the ionic strength from one to two has been to increase both k' and k". The magnitude of the increase is about 10% for both k' and k".
*
TABLE I11 HYDROGEN IONCONCENTRATION DEPENDENCE OF REACTION BETWEEN Pu02+* AXD V+3 IN PERCHLORATE SOLUTION [V f a l o
[PuOt +*lo X 104, M
X 104, iM
[Htl, M
kl (obsd.)
1.620 1.584 1.347 1.387 1.170 1.245 1.953 1.669 1.586 2.289
3.161 1.583 1.271 1,244 1,186 1.227 1.882 1.918 1.819 2.390
1,000 1.000 0.500 ,500 .250 .250 .I50 ,150
2.08 2.08 4.20 4.20 10.26 10.35 22.4 21.4 37.3 37.5
.loo .loo
kl (calcd.)
2.01 4.42 10.44 20.96 38.1
In an experiment devised to illustrate the effect of the substitution of lithium perchlorate for sodium perchlorate as the added salt in maintain24.24 ing the ionic strength of the solution a t unity, the rate of reaction between P U O ~ and + ~ V+3 was meas43.96 ured in 0.500 M perchloric acid-0.500 M lithium perchlorate a t 25'. Values of 4.55 and 4.53 liters The values of Icl(oa~od) given in Table I1 were ob- mole -I see.-' were obtained in these solutions for lcl as compared to 4.20 liters mole-'sec.-' in solutained by use of the equation tions in which sodium perchlorate was the added kl = (k'/[H+] + k"/[H+]') (3) salt (see Table 111). Thus the substitution of in which k' and k" are the specific rate constants lithium perchlorate for sodium perchlorate occaassociated with an inverse first and an inverse sec- sions no large change in the specific rate constant ond power of the hydrogen ion concentration, for this readion. respectively. Values of k' and k" obtained by the Temperature Dependence.-The effect of the method of least squares are 2.12 f 0.03 set.-' and variation of temperature upon the rate of reac0.228 f 0.006 mole liter-' sec.-l, respectively, tion has been studied in molar perchloric acid for perchlorate solutions of ionic strength two solutions between 2.4 and 34.5'. In Table IV the and a t a temperature of 25'. The uncertainties variation of the apparent rate constant, kl, is given represent twice the standard deviation of the given as a function of temperature. From a linear average values of the intercept and slope. A plot of log k1 versus 1/T is computed a least squares possible mechanism which is in accord with the slopc of 3517 & 81 which corresponds to an exsecond-order nature of the reaction and with the perimental Rrrhenius activation energy of 16.1 & observed hydrogen ion concentration dependence is 0.4 kcal./mole. Since lcl = k' k" = A' e--E'/RT + V+3
PuOz+'
+ H20
+ VOH+' +
VOH++
(rapid)
+ VO+' + H +
PuOZ+
(610~)
TABLE IV TEMPERATURE DEPENDENCE OF THE RATEOF REACTION BETWEEK P U O Z + AND ~ V f 3 IN MOLARPERCHLORIC ACID
(5)
k"
+
+ VO+' + H20
PuOZ+
(slow)
+
(4)
k'
+
I ' U O ~ + ~ V(OH)2+
+ H+
12.11
(6)
It should be pointed out that alternative possibilities to reactions 5 and 6 consist of the reaction of Pu0201-1+ with Ti+++ and of Pu020H+ with VOH+2, respectively. Since these reactions are kinetically indistinguishable, both must be considered as possibilities. Measurements of the rate of reaction between P u 0 ~ +and ~ V+3 also were made as a function of aciditv in solutions of unit ~
Temp., OC.
k ~ 1., mole-1 sec.-1
log kl
34.5 25.0 15.2 2.4
4.57 2.08 0.74 .22
0.660 .318 - .I31 - .658
tJhe activation energy for the principal path is riot necessarily equal to the experimentally observed value of 16.1 kcal. However since the plot of li, versus 1/T is linear, the activation energies E' and E" are not greatly different from this
A"e-Eff/RT,
