R. S. ONDREJCIN AND T. P. GARRETT,JR.
470
Vol. 65
THE THERMAL DECOMPOSITION OF ANHYDROUS URANYL NITRATE AND URANYL NITRATE DIHYDRATE1 BY R. S. ONDREJCIN AND T. P. GARRETT, JR. Savannah River Laboratory, E. I . du Pont de Nemours & Go., Aiken, South Carolina Received September 18, 1960
The thermal decomposition of anhydrous uranyl nitrate followed first-order kinetics in vacuo a t temperatures from 250 to 450”. Amorphous uranium trioxide was the only non-volatile product of decomposition. No intermediate com ounds were observed. The thermal decomposition of uranyl nitrate dihydrate occurred as two independent reactions fotowing first-order kinetics: dehydration to form anhydrous uranyl nitrate and denitration to form uranium trioxide. The decomposition was studied both in vacuo and at atmospheric pressure under nitrogen, at temperatures from 250 to 400”. Intermediate and final decomposition products were identified. The specific reaction rate constants were measured for the dehydration and denitration reactions. From these data, the heats and entropies of activation were calculated.
Introduction Normally, uranium from spent reactor fuels is processed to form uranyl nitrate hexahydrate, which is thermally decomposed in a complex series of reactions2 to yield uranium trioxide. The trioxide is reduced with hydrogen to uranium dioxide and then hydrofluorinated to uranium tetrafluoride. Conditions of thermal decomposition affect the reactivity of the trioxide and hence the efficiency of the reduction and hydrofluorination. Table I shows the products possible from the thermal decomposition of uranyl nitrate hexahydrate. TABLE I URANIUM TRIOXIDE AND ITS HYDR4TES LY-UOO Uo3(A) a a-UO3.2HzO P-uo~ U03.1/zH,0 fl-UO3.2Hzo r-UO, CY-UO~.H~O 6-UO3 P-UO3.HzO
f-UOa
Y-UO~.H~O 6-UO3.Hzo
Amorphous.
The nomenclature used for the allotropes of uranium trioxide was that suggested by Hoekstra and SiegeL3 0t)her notations have been used by various In Hoekstra and Siegel’s system, rU03 is Da\vson’s4 U03(I), and 6-U03 is Dawson’s cubic phase, UO3(I1). The cu-U03 is identical to U03(I) described by Katz and Rabinomitch.6 The a- and P-monohydrates correspond, respectively, to Dawson’s UO,. 0.8Hz0 and U03.H20. There is no variation in the nomenclature for y- and d-monohydrates. Infrared dat~aindicate hhat the dihydrates have only one true xvater of hydration, while the mono- and hemihydrah probably are not true hydrate^.^,^ (1) The information contained in this article was developed during the course of work under contract AT(07-2)-1 with the U. S. Atomic Energy Commission. Presented to the Boston Meeting of the American Chemical Society, April 1959. (2) B. A . J. Lister and R. J. Richardson, “The Preparation of Uranium Trioxide by Thermal Decomposition of Uranyl Nitrate.” Atomic Energy Research Establishment, Harwell, AERE C/R 1874, October 18, 1954. (3) H. R. Hoekstrn and S. Siegel, Proc. U. N . Intern. Conf. Peaceful Uses Atomic Energy, End, Geneua, 28, 231 (1958). P.t’l.548. (4) J. E(. Dawson, E. Wait, K. Alcock and D. R. Cliilton, J. Chem. boc., 3531 (1956). (5) J. J. Kat2 and E. Rabinowitch, “The Chemistry of Uranium,” Natl. Nuclear Energy Ser.. Div. VIII, Vol. 5 , IlcOrnm-Hill Book Co., New York, N. Y., 1951, pp. 277-285. ( 6 ) P. Perio, Bull. soc. chim. Fronce, 776 (1953). (7) J. J. Kats and D. M. Gruen, J . Am. Cham. Soel, 71, 9106 (1849).
These oxides are usually prepared by various treatments of uranyl nitrate or its hydrates, inferring that the mechanism of the thermal decomposition is extremely complex. The usual product of the thermal decomposition of UO,(XO,),. 6H20 is y-UOs, which has variable reduction characteristics with Hz.I0 A better understanding of the mechanism of decomposition could lead to a prediction of the process conditions necessary to produce U03 that would be reduced consistently under a given set of conditions.
Experimental Anhydrous uranyl nitrate was synthesized from uranium trioxide and nitrogen dioxide by the method of Gibson and Katz.11 The product was analyzed for nitrogen by the Devarda modification of the Kjeldahl procedure. The nitrogen content of the product varied from 6.4-6.7 weight % corresponding to 90-95 weight % anhydrous uranyl nitrate. The remainder of the material consisted of unreacted uranium trioxide. Uranyl nitrate dihydrate was prepared from reagentgrade uranyl nitrate hexahydrate that was placed in a vacuum desiccator over sulfuric acid for a minimum of 48 hours.12J3 Analysis of the dihydrate showed the salt contained 55.3 weight % ’ uranium which agreed with the theoretical value of 55.4 weight %. X-Ray diffraction showed only the dihydrate in the product. The thermal decompositions were carried out in an allglass system. Approximately 10-gram samples of each salt were heated in a reaction tube immersed in a molten metal bath. Two runs were made a t each of four temperature levels over the range of 250-400’. I n addition, one run with anhydrous uranyl nitrate was made at 450”. Gaseous products from the reactions were condensed in tr:qps cooled with liquid nitrogen. The condensed gaseous products taken periodically during the thermal decomposition were transferred to a storage bulb, weighed and the pressures measured a t room temperature with a high sensitivity Bourdon gauge.’* The samples were analyzed by infrared and mass spectroscopy. The liquid from hydrated samples remaining in the trap was weighed and titrated for acid content. The water of hydration was calculated from the liquid weight and acid content. Decomposition a t atmospheric pressure was carried out by flushing the system a t a rate of 30 ml./min. with dried nitrogen. Under theEe conditions the gaseous products were not ( 8 ) J. J. Kat2 and E. Rabinowitch. “Chemistry of Uranium,’ Natl. Nuclear Energy Ser. Div. VIII, Vol. 7, McGraw-Hill Book Co., New York, N.Y., 1951, p. 746. (9) R. E. DeMarco, National Lead Co., private communication. (IO) C. W. Kuhlman, Jr., and B. A. Swinehart, I n d . Eny. Chem., 50, 1774 (1958). (11) G. Gibson and J. J. Katz, J . A m . Chem. Soc., 73, 5436 (1951). (12) A. J. King, R. Pfeiffer and W. Zeek, “Thermal Stability of the Hydrates of Uranyl Nitrate,” Syracuse University. New York, N. Y., NYO 6313, August 1, 1957. (13) J. J. Katz and G . T. Seaborg, “The Chemistry of the Actinide Elements,” John Wiley and Sons, New York, N. Y., 1957, p . 193. (14) R. A. W. Hill and R. A. Hamilton, Research, 7 , 856 (1954).
THERMAL DECOMPOSITION OF ANHYDROUS URANYL XITRATE
March, 1961
47 1
removed from the reaction area immediately, so as to simulate the conditions normally found in the production of uranium trioxide. The pressure above the reaction remained within a few mm. of atmospheric pressure.
Results In the vacuum denitration of UOz(NO& the number of mol'es of this compound remaining a t time t was derived from analyses of the gaseous products its shown in Table 11. TABLE I1
-3 9
-8
i
TIIERXAL ~ ~ E C O M P O S I T I OO S F uOz( N 0 3 ) ~UNDER VACUUM
Temp., OC.
250
300
--Run Reaction time, min.
0 20 60 180 360
-I -log IUOz(NOa)2]=
1.85 1.96 2.12 2.52 3.26
-Run Reaction time, min.
0 30 90 150 210 330 0 10 20 30 60 122 0 6 16 30 48
11-
0 1.52 10 1.60 20 1.72 30 1.85 60 1.85 120 2.68 350 0 1.42 6 1.53 10 1.68 17 1.90 30 2.20 60 2.72 450 0 1.46 14 2.10 23 2.76 40 3.40 a Quantity of compound is expressed in moles.
-log [UOz-
(N0dzI0
1.55 1.77 1.95 2.14 2.42 2.64 1.48 1.52 1.62 1.72 1.98 2.52 1.45 1.55 1.81 2.09 2.35
Uranyl nitrate dihydrate was thermally decomposed under vacuum and under nitrogen a t one atmosphere. Fieom the analyses of the condensed gaseous products, the number of moles of water and uracyl nitrate present a t any given time were calculated as shown in Table 111. Least squares analyses were made on each set of data in Tables 11-111. The slopes of the lines were used to calculate the specific reaction rate constants shown in Table IV. The denitration reactions of both uranyl nitrate and uranyl nitrate dihydrate follow apparent first-order kinetics as does the dehyclratioii of the dihydrate. The over-all thermal clecomposition of uranium nitrate dihydrate proceeds by way of two fist-order mechanisms: the first, dehydration to the anhydrous nitrate; the second, denitration to the amorphous trioxide. The data in Table IV show that the reaction rates at atmospheric pressure are much greater than those under vacuum. Gaseous decomposition products are probably removed under vacuum as rapidly as produced. At atmospheric pressure some of the gaseous products are not removed immediately and may be acting as catalytic agents. Since the Arrhenius equation is an empirical relationship of the reaction rate constant-temperature dependence, the data in Table IV were plotted
-
0
U02(ND312 VACUUM DENITRATION
A
UO2ND3I1 2H20
3 !-.
L
-io S 1711
J
3214
- ATM
0 OZIS
PRESSURE DEHYDRATION
0 00 5
i T
0 31117
3 0613
3 0013
C 3013
OA-5
.
Fig. 1.-Thermal decomposition of uranyl nitrate and urnnyl nitrate dihydrate.
TABLE I11 THERMAL DECOMPOSITION OF UO2(N03)2,2HzO Pressure
Vacuum
-Run 1-7 -Run ReacReac-log tion tion Temp., time, -log iUOr- time, "C. min. [ H z O ](NOa)z]a ~ min.
250
-log
-log IHzO]"
0 1.25 20 1.80 60 2.82 .. 120 180 .. 240 .. 300 .. 0 1.32 15 1.84 30 2.34 45 .. 75 .. 0 1.34 10 2.16 20 3.03 30 .. 60 .. 0 1.23 5 1.52 12 2.37 18 .. Atmos0 1.40 pheric 10 1.81 16 2.02 21 2.77 0 1.18 3 1.32 6 1.85 15 2.60 350 0 1.38 1.62 0 1.44 3 2.36 1.99 3 1.89 6 2.82 2.43 6 2.38 400 0 .. .. 0 .. 3 1.51 3 1.70 1.71 5 2.16 4 2.11 1.89 7 2.42 6 2.82 2.38 0 Quantity of compound is expressed in molrs. 0 20 60 120 180 240 300 300 0 15 30 45 75 350 0 10 20 30 60 400 0 7 18 24 250 0 10 15 25 300 0 4 8
1.21 1.58 1.43 1.59 1 . 9 5 1.66 . . 1.72 . . 1.79 . . 1.86 . , 1.89 1.24 1.36 1.61 1.38 2.10 1.46 ,. 1.59 ,. 1.74 1.31 1.61 1 . 9 5 1.70 2.92 1.78 . . 1.96 . . 2.44 1.20 1.45 1 . 6 5 1.66 2.92 2.10 . . 2.34 1.12 1.42 1.87 1.73 2.31 2.00 3.16 2.52 1.18 1 . 6 1 1.55 1.80 2.12 2.60
11-
[TJOz(NOa)zlu
1.56 1.60 1.66 1.72 1.75 1.93 1.99 1.38 1.42 1.54 1.70 1.87 1.60 1.71 1.82 1.97 2.35 1.50 1.63 1.70 1.90 1.40 1.69 1.90 2.32 1.48 1.56 2.13 2.61 1.61 1.90 2.24
.. 1.69 2.25 2,70
according to absolute reaction rate theory on the basis of equation 1
R. S. OKDREJCIN AND T. P. GARRETT, Jit.
472
Vol. 65
nium. Uranium trioxide is 83.2 weight % uranium. Small amounts of occluded nitrates detectable by Plotting log k'/*T us. 1/T as shown in Fig. 1, Kjeldahl determinations caused the low experithe slopes of the lines are -AH*/2.30R, with the mental values. Alpha-UOa is the normal end product of the deintercepts equal to AS*/2.30R 4.48. The hydration of the hemihydrate of uranium trioxide. TABLEIV Since the hemihydrate is not formed above 250") a substantial portion of the denitration must have SPECIFIC REACTION RATESOF URANYL NITRATE occurred below this temperature. Heating at an k', min. 1 Temp., Vac., Vac., Atm., increased rate would minimize the initial formation OC. anhydrous dihydrate dihydrate of the hemihydrate. Denitration 250 0.0083 0.0028 0.11 Reactions accounting for the decomposition of 300 ,021 ,014 .22 anhydrous uranyl nitrate under vacuum are shown 350 .048 .032 .28 in Table VII. 400 ... .068 .55 k' = h e - ( A H * - T A S * ) / R T
(1)
+
Dehydration
450 250 300 350 400
.12
...
..
.044 .074 .I9 .22
.18
.25 .45 .69
equations of the lines in Fig. 1 were calculated by the least squares method and the heats and entropies of activation were calculated from them. These values are shown in Table V.
TABLEVI1 DECOMPOSITION REACTIONS ANHYDROUS URANYL NITRATE UNDER VACUUM 250-450" UOz(NO3)2 4 UOa(A) 2x02 '/zO? >5w0 U O ~ ( N O ~+=) ZU03(A) 2N02 '/zOz 3UOz(~o3)2 -+ U& 6N02 202
+ + + + + +
The vacuum denitration of uranyl nitrate dihydrate produced a compound with distinct X-ray diffraction and infrared absorption patterns not TABLE V identifiable from data in the literature. HEATSAND ENTROPIES OF ACT~VATION From previous experimental work and other AH*. data,'* the compound was assumed to be a uranyl Hydration kcal./mole AS*, e.u. Process state Vac. Atm. Vao. Atm. hydroxynitrate. To confirm this identification, Denitration Anhydrous 23.5 . . 7 . 0 ... uranyl hydroxynitrate was synthesized by adding Denitration Dihydrate 33.4 14.8 22.0 - 0 . 3 "active" amorphous uranium trioxide1' to an aqueDehydration Dihydrate 18.6 12.8 3.1 - 2 . 5 ous solution of uranyl nitrate. Vacuum dehydraThe least reliable data are those derived from the tion a t 30-35" of the solution produced a compound vacuum denitrations. The estimated accuracy of with the same X-ray diffraction pattern as the the values for AH* in Table V is =kl5%. The intermediate in denitration. The synthesized prodcalculated standard deviation of the plotted points uct had a rhombic structure. Analysis showed that titratable hydroxyl groups were present. affects AS* by 5 to 10%. Products and Intermediates.-The effects of Water of crystallization was determined by the temperature and pressure on the final products of Karl Fisher Reagent method and uranium by ignithe thermal decomposition of anhydrous uranyl tion; analysis showed 12 weight % water and 58.4 nitrate and uranyl nitrate dihydrate are summa- weight % U vs. 12 weight % water and 59.1 weight % U for a trihydrate of uranyl hydroxynitrate. rized in Table VI. Dehydration of this compound under vacuum a t room temperature produced a compound that was TABLE VI amorphous to X-ray diffraction, but showed an FINALDENITRATION PRODUCTS infrared pattern, with hydroxyl groupings, not preTyw., Salt C. Pressure hoduct viously reported in the literature. Hydration of UOd NOJ 2 250-450 Vac. uodA) the dehydrated compound a t room temperature 500 Vac. UOdA) u30, changed the diffraction pattern from that of the U02(N0&2H,O 250 Vac. UOs(A) + CY-UOS trihydrate. The water solubility of the compounds 300-400 Vac. UOa(A) formed proved the various allotropes of uranium U02(NO3)2.2H,0 250-450 Atm. yU03 trioxide were absent. This final compound was 500 km. yuo34-8-uo3 analyzed for water and uranium. The results agreed reasonably well with the values for uranyl The salts that were decomposed a t atmospheric hydroxynitrate tetrahydrate : uranium, 56.0 weight pressure were blanketed with nitrogen. The final yo vs. 56.5 weight %; wat8er, 16 weight yo us. products were identified from established X-ray dif- 17 weight yo. Listed in Table VI11 are the main fraction p a t t e r n ~ , ~ -infrared ~ ~ J ~ analyses, l7 and diffraction lines of these hydrates and their relative chemical analyses of the completely amorphous (A) intensities, plus the infrared absorption patterns for compounds. the hydrates and the anhydrous compound. Typical chemical analyses of the amorphous Differentiation is not made between the hydrates products were 83.0, 83.1 and 83.1 weight % ura- on the infrared pattern because the only variation is (15) ASTM X-Ray Powder Data, File Card No. 2-0278. a slight increase in the size of the 3390 and 1626 (16) J. R. Bridge, C. W. Melton. C. M. Schwartz and D. A.
+
Vaughan, Battelle Memorial Institute, BMI-1110, July 12, 1956. (17) J. W. NehleSavannah River Laboratory, Aiken, 9. C., private oommuniaation.
(18) R. H. Moore, "Factors Affecting the Reactivity of Uranium Trioxide " Interim Progress Report. IIanford Works IIW-31070, April 29. 1954.
March, 1961
DIFFUSION TO
A
PLANE WITH LANQMUIRIAN ADSORPTION
473
Thermal decomposition of uranyl nitrate dihyTABLEVI11 X-RAYDIFFRACTION AND INFRARED ABSORPTION PATTERNSdrate a t atmospheric pressure produced either r-UOa between 250 and 400' or y &UOa a t 500" as the OF URANYL HYDROXYNITRATES
+
Tetrahydrate d I/Io
7.15 6.52 6.30 5.19 4.70 4.49 4.33 3.83 3.56 3.24 2.16
70 60 85 35 45 40 100 35 35 70 50
Trihydrate d I/Io
6.33 5.68 5.58 5.28 5.20 5.01 4.81 4.08 3.75 3.52
55 55 100 80 75 70 70 70 70 75
-Hydrate? -AnhydrousWave Wave no., Characno., Characom.-' teristic ern.-' teristic
3390 1626 1613 1515 1381 1266 1026 943 845 803 749 742
1613 1515 1381 1266 1026 943 845 803 749 742
final product. From 525 to 550" the final product was p-U03. Amorphous anhydrous uranyl hydroxynitrate was found as the only intermediate. The equations in Table X summarize the reactions of uranyl nitrate dihydrate at atmospheric pressure.
TABLEX DECOMPOSITION REACTIONS URANYL NITRATE DIHYDRATE AT ATMOSPHERIC PRESSURE 250' UOz(NOs)z.ZHzO -* UOz(N0s)z 2Hz0 UOz(N0s)t -+ yUOo 2NOz '/zOz U02(NOs)z*2HzO-* UOz(0H)NOa HzO HNOs UOz(0H)NOS + ~ U O I HNOI 300-450' UOt(NOs)~.2HzO + UOz(NOJz 2H20 UOz(N0s)a -* r-UOa 2x02 '/zOt UOZ(NOS)Z*~HIO + UOs(y 8) 2N02 '/z500' 02 2Hz0 2N02 '/eOa 4525-550' UOz(NO~)z-2HzO 8-UOs 2Hz0
+ +
+
+
bands for the tetrahydrate as compared to the trihydrate. The trihvdrate, anhydrous uranyl hvdroxvnitrate and a-ura&um ikioxide monohydrate were identified by infrared and X-ray diffraction as intermediates in the den.itration of uranyl nitrate dihydrate under vacuum. Equations that account for these reactions are shown in Table IX.
+
+
-+
+
+
+ + + + + + +
The results presented show that by proper choice of conditions the thermal decomposition of uranyl TABLEIX nitrate dihydrate will produce a product that is DECOMPOSIT[ON REACTIONS-URANYL NITRATEDIHYDRATEprimarily @-UO,,7-UOa or amorphous uranium trioxide. The &phase is produced by the rapid UNDER VACUUM decomposition of the dihydrate at temperatures 250' Primary Reactions above 500" a t atmospheric pressure, 7-phase over 'I~O~(NC)~)Z*~HZO -* UOz(N08)2 2Hz0 the temperature range of 250-450". DecompositJ02(N08)o UOs(A 2N02 l/~Ot tion of the dihydrate in the temperature range of Secondary Reactions 300-450° under vacuum always produces amorphous uranium trioxide. Above 500" the product is UOz(NOs)z.ZHzO UOs(0H)NOa HzO contaminated with U308. Below 300' some aHNOu 'I~02(OE)NOa t 3Hz0 UOZ(OH)NOS.~HZO phase uranium trioxide is formed. UO~(OFI]NO~~3HpO + ~u-UOs.Hz0 2Hn0 Acknowledgment.-The authors are indebted to "02. Drs. C. H. Ice and R. C. Milham for technical ada-uO~*&o+ UOs(A) &O vice, Drs. W. R. Cornman and J. W. Nehls for Xray diffraction and infrared absorption patterns, 300400' 'I102(NOa)r*2HzO+ UOz(N0a)t 2He0 and Mrs. B. S. Russel' for chemical analyses. UOz(N0s)z + UOa(A) 2N02 '/zOz
-
+ +
+
(Y)
-+
+
-
+ +
+
+
+
+
+ +
DIFFUSION TO A PLANE WITH LANGMUIRIAN ADSORPTION BY W. H. REINMUTH Department of Chemistry of Columbia University, New York, N. Y. Received September 13, 1960
A theoretical treatment is given of semi-infinite linear diffusion to a stationary plane with Langmuirian adsorption at the boundary. The frmtion of the surface covered is a function of two variables C*/a and aZDt/I,Z where C* is the solution concentration of surfactant, a is the solution concentration which would correspond to half coverage, rmis the surface concentration of surfactant at full coverage, D, the diffusion coefficient of the surfactant and t, time. Comparisons of exact the0 with approximate treatments are given and discussed. A formal solution to the same problem at an expanding plane is inxded.
Many workers have concerned themselves with the effectsof surfaceactive agents on electrode procewes.' Delahay and Trachtenberg2 in particular Reilley, .*Re (1) For a review m: w, H. Reinmuth in c. cent Advancea in Analytical Chemistry and Inatrumentation." Interscience Publiehers, h a . , New York, N. Y.. 1980. (2) P. Delahay and I. Trachtenberg, J . Am. Chem. Soo., 79, 2365 (1957); 80, 2094 (1958).
emphasized the influence of the rate of diffusion of these species on the observed results. Delahay and Fikea later attempted to solve the differential equations describing semi-infinite linear diffusion to a plane boundary with Langmuirian adsorption by an unstated method with the aid of an electronic (3) P. Delahay and C. T. Eke, ibid., 80, 2628 (1958).
