that no ionic mechanism for the rate-determining step is indicated. The failure of stirring to alter the solution rate suggests that the process is not limited by any redox potentials existing in the immediate vicinity of the electrode. The kinetics for the initial dissolution process may be represented by the equation:
dD dt
-=
-hoc
where the rate constant, ho, depends only on the fading agent employed. I n the second stage the rate constant decreases with tkiickness and therefore with D - D2
dD
-= dt
-hc
where h is a function of (D- D2).I t follows from Equation 1 that ho is given by the first column of Table I. Integration of Equation 2 gives:
(3) Since D2 represents the contribution to the optical density not due to the film, the right-hand side of 3 is just the reciprocal of the rate constant integrated over the film density between the limits of 0 and the critical value of 0.54. This should depend only on the reactant and not its concentration. The constancy of c(t2 - t l ) follows a t once. The decrease in rate after the critical density is reached is subject to a t least two interpretations: 1. The film thickness is not uniform. Therefore the optical density decreases linearly until the regions that were initially thinnest have been stripped down to the substrate. Thereafter these areas do not contribute to the rate. As more and more of the substrate is exposed, the rate gradually falls to zero. 2. During the deposition, the initial coating is deposited on a crystalline tin oxide surface which influences
its structure and composition. After a suitable thickness is attained, the influence of the substrate can no longer be felt and the deposit has a fixed composition thereafter. During fading, the uniform part of the coating dissolves a t a constant rate. After the thickness has decreased to the critical value (at which D = D1),a progressively less reactive surface is exposed. The first mechanism seems unlikely for two reasons: I n the first place, the coated electrodes are shiny. Optical and electron microscopy failed to reveal any appreciable topographic irregularities. Secondly, the observed kinetics is inconsistent with the assumption made. If the reactivity of the film were the same throughout, it can be readily shown that a reciprocal relationship should exist between the linear rate, (Do - D i ) / t l , and the completion time (t2 - t l ) , which should be independent of the nature of the reactant. From the data of Table I we can see that the product (Do - D l ) / t l x ( t z - t l ) has the value 2.56 0.43 for Carbitol and 1.35 i 0.21 for hydroquinone. Thus the kinetics must change in the case of a t least one of these reactants when the optical density goes through the value D1. I t is most likely that the kinetics changes in both cases, but not to the same extent. The second mechanism is also consistent with the observation that inhibition of anodic deposition by hydroquinone is accompanied by a slight darkening of the anode which does not increase with time regardless of how long the voltage is applied.
*
Literature Cited
Criegee, R., Ber. 64B, 260 (1931). Davidson, A. W., Chappel, W., J . A m . Chem. SOC.55, 4624 (1933). Kuwana, T., Darlington, R. K., Leedy, D. W., Anal. Chem. 36, 2023 (1964). Mackenzie, J. D., J . A m . Ceram. SOC.47, 211 (1964). Tarbutton, G., Vosburgh, W. C., J . A m . Chem. SOC.54, 4537 (1932). RECEIVED for review July 8, 1968 ACCEPTED September 20, 1969
SSOLUTION OF PLUTONIUM IN DILUTE NITRIC ACID F .
J.
MINER,
J .
H .
N A I R N ' , AND
J .
W .
B E R R Y
Rocky Flats Division, The Dow Chemical Co., P . 0 . Box 888, Golden, Colo. 80401
THEreprocessing
of plutonium metal uza conventional aqueous chemical methods requires a conversion from the metal to a nitrate solution. Once in solution, the plutonium can be purified by ion exchange or solvent extraction, then precipitated as an oxalate, fluoride, or peroxide, and reduced back to the metal. In some methods used in the past, plutonium metal was converted to a nitrate solution by dissolving it in concentrated nitric Present address, Chemistry Department, Rice University, Houston, Tex. 77006 402
I & E C PRODUCT RESEARCH A N D DEVELOPMENT
acid containing trace quantities of fluoride. However, this dissolution was slow. I n addition, explosions sometimes occurred, caused perhaps by reaction products, which could include hydrogen. A method used more recently for converting the metal to a nitrate solution requires burning the metal to the oxide and then dissolving the oxide in a concentrated nitric-hydrofluoric acid solution (Molen, 1967; Stevenson and Paige, 1968). This method requires an extra step and, like the dissolution of the metal, is slow.
Rlwtes of dissolution were determined for plutonium metal as a function of nitric alcid concentration, hydrofluoric acid concentration, and temperature. The interactions of these parameters were determined, as well as the conditions that produced the maximum relative rate of dissolution. The maximum rate of dissolution occurred a t approximately 3M nitric acid, a hydrofluoric acid concentration of 0.13M (the highest investigated), and any temperature between 26' and 6 7 ' C . (the range investigated). The effect of temperature over this range was insignificant compared to the other two parameters investigated. The mechanism of the dissolution reaction is uncertain.
Recent work has shown that plutonium, contrary to expectations, dissolver; relatively rapidly in dilute nitric acid containing a low concentration of hydrofluoric acid. The rate of dissolution in dilute nitric acid was 10 to 20 times faster than in 10M nitric and 50 to 60 times faster than in 15M nitric acid containing comparable fluoride concentratioris ( H F 5 0.13M). The primary objective of the work described in this report was to obtain information on the functional relationship between the rate of plutonium metal dissolution and the three parameters of nitric acid concentration, hydrofluoric acid concentration, and temperature. Experimental
Equipment and Procedure. Unalloyed plutonium metal in the form of coupons was used in the experimental work. The coupons were cut from a cylindrical casting of plutonium and drilled in the center for fitting onto a tantalum rod. Before use, each coupon was soaked in 2-propanol to remove excess cutting oil, then washed in a nitric-hydrofluoric acid cleaning solution, and dried. Each coupon had a total exposed surface area of 20.4 sq. cm. and an average weight of 23.4 grams. The apparatus for the experimental dissolution work consisted of a 600-ml. stainless steel beaker in which the metal coupon was immersed in the test solution after being attached to a tantalum rod. A water bath was used to provide a constant (+1.0"C.) temperature throughout each run. The solution was stirred constantly using a Teflon-coated magnetic stirring bar. I n all runs, 200 ml. of test solution was used. The coupon was immersed in a test solution for 10 minutes, then removed from the solution, washed with distilled water, air-drLed, and weighed. The coupon was immediately replaced in the solution for another 10-minute interval and the procedure repeated. From the data obtained, the rate of dissolution in each test solution was calculated. Solutions with a low dissolution rate were initially light blue, indicating Pu(II1). This soon changed to brown, however, indicating oxidation of the Pu(II1) to Pu(1V). I n solutions with a higher dissolution rate, no initial Pu(II1) was observed. Gases formed during dissolution were analyzed by gas chromatography. Information on hydrogen in particular was sought because of the possibility of its generation in explosive quantities. The minimum detection limit for hydrogen was 0.05 volume 5.Gas samples were withdrawn from just above the surface of the solution at various times during dissolution. I n no case was hydrogen ever detected above approximately 0.3 volume 5 (4 volume '7 is the minimum explosive limit). The gas produced in the highest concentration was N?O.
There is a possibility of an electrochemical cell's being established by the plutonium and tantalum, which could influence the rate of dissolution. However, rates obtained when a plutonium coupon was submerged in solutions without contact with any other metal were comparable to rates when a tantalum rod was used to hold the plutonium. This would indicate that if an electrochemical cell were established, the current generated was small enough so that it did not significantly affect the rate of dissolution. Experimental Design
The region of experimentation was covered by a 33 factorial design. T o economize on the number of experiments, a second-order rotatable design was used (Box and Behnken, 1960; Phillips and Huber, 1967). This design efficiently estimates the coefficients of a second-degree polynomial by the method of least squares. I t requires the determination of only 13 of the 27 experimental points in the full 33 factorial design. Preliminary work had identified three parameters which have the greatest effect on the rate of dissolution of the metal. These parameters and the three equally spaced levels of experimentation which were selected for investigation are shown in Table I. Rates of dissolution were calculated from both the first and second 10-minute dissolution runs and used to calculate the coefficients ( p ) in the equation below. The data from the second 10-minute runs were more internally consistent than those from the first 10-minute runs. This internal consistency was determined by calculating the coefficients in the polynomial using rate data from both the first and second 10-minute runs. Then rates of dissolution at three points (not used in calculating the coefficients) were determined experimentally and compared with the calculated rates for these same points. The agreement between the experimentally determined and the calculated rates a t these three points was better for the second than for the first 10-minute runs. This was probably due to the greater uniformity of the surface of the metal coupons after the initial 10-minute runs. Rate data were obtained experimentally a t 13 points. The particular 13 points selected from the 27 available were determined by the statistical design being used (Box and Behnken, 1960). Employing these experimental rate
Table I. Parameters and levels Investigated
Parameter
Level 0
Leuel 1
Leuel 2
Xitric acid, M Hydrofluoric acid, M Temperature," C.
1.0 0.01 23
3.0 0.07 46
5.0 0.13 69
VOL. 8 N O . 4 D E C E M B E R 1 9 6 9
403
data, the p coefficients in the following polynomial were calculated using a least squares technique: 3
3
logR=do+
C pixL+ C t
=1
3
b,,XgX,
i = 1 , = 1
where R is the rate of dissolution (mg cm-' min ') and XI, X,,and X?are the H N 0 3 and HF concentrations ( M ) and the temperature (" C.), respectively. With these coefficients, and appropriate values of XI,X 2 , and X1, the rates of dissolution were calculated for the remaining 14 points in the full factorial design. Results and Discussion
Both experimental and calculated rates of dissolution are given in Table 11. A comparison of the calculated with the experimental rates a t the 13 points (where both are available) indicates the type of agreement obtained using the method of calculation described above. Triplicate experimental rates of dissolution were obtained a t two different experimental points; the relative standard deviations were 7.8 and 10%. These values represent the type of reproducibility obtained in the experimental work. The average effects of the parameters on the dissolution rate are shown in Figure 1 in the form of average effect curves. Data for these curves were obtained from the calculated values a t the 27 points in the full factorial, by averaging over all the combinations of levels of the two remaining parameters. The average effect curves show that the rate of dissolution reaches a maximum a t approximately 3M nitric acid. However, the rate is directly proportional to the concentra-
tion of hydrofluoric acid over the range investigated, indicating an optimum concentration of 0.13M or higher. (Higher concentrations of hydrofluoric acid were not investigated because of the problems they could cause in the subsequent processing of plutonium nitrate solutions.) The effect of temperature on the rate of dissolution is nearly negligible compared to that of the other two parameters. The largest interactions between parameters were found between the hydrofluoric acid concentration and temperature and between the hydrofluoric acid and nitric acid concentrations. Consequently, the average-effect curves alone do not completely summarize the effects of the parameters but can be used to give a general idea of the effect of a variable upon the dissolution rate. An alternative representation of the results, which is particularly useful in showing the interrelationship between the effects of hydrofluoric and nitric acid concentrations, is given in Figure 2. The rates of dissolution a t constant temperature (46"C.) and three different hydrofluoric acid concentrations were calculated and plotted as a function of nitric acid concentration. Since temperature has little influence on the rate of dissolution, analogous plots would
61
I
I
I
/'
/' ,/'
/ /' /'
*/'
Table II. Rate of Plutonium Dissolution
HNO,, M 1
Parameters HF, M 0.01 0.07 0.13
3
0.01
0.07 0.13
5
0.01 0.07 0.13
404
Temp., 23 46 69 23 46 69 23 46 69 23 46 69 23 46 69 23 46 69 23 46 69 23 46 69 23 46 69
O C .
Dissolution Rate, Mg. Cm. -'Min. Exptl. Calcd. 1.04 3.83 6.47 6.10
0.86 2.35 5.66 12.50 10.10 1.00
3.13 3.84 5.76
0.64 1.04 2.02 3.51 4.21 6.03 8.06 7.12 7.52
__
01
-
1
I
I 0.01 23"
3 0.07
1
1
5 M HNOq 0.13 M HF-
46'
69'
OC
Figure 1 . Average effect of parameters on rate of dissolution
0.94 1.40 2.52 5.11 5.66 7.47 11.67 9.52 9.28 0.62 0.86 1.42 3.36 3.43 4.19 7.64 5.75 5.18
I & E C PRODUCT RESEARCH A N D DEVELOPMENT
0 01
"03, M
Figure 2. Plutonium dissolution Interaction of HNOi and HF Temperature 46" C.
be obtained at the other temperatures. From Figure 2, it is apparent that the rate of dissolution increases with an increase in the hydrofluoric acid concentration and the maximum rate of dissolution is obtained a t approximately 3M nitric acid. The results as discussed show that fluoride is necessary for the dissolution of plutonium. However, the mechanism for the dissolution is not clear. Halide ions in general are known to accelerate the dissolution of metals (Kindlimann and Green, 1967; Kolotyrkin, 1961). I n investigation of the dissolution of zirconium in nitric acid containing hydrofluoric acid (DeCrescente et al., 1960; Meyer, 1965; Smith
