I
G. I. CATHERS, M. R. BENNETT, and
R. L. JOLLEY
Oak Ridge National Laboratory, Oak Ridge, Tenn.
UF,-3NaF Complex Formation and Decomposition Sodium fluoride absorption of uranium hexafluoride is a convenient method of trapping, handling, and storing uranium hexafluoride LABoRAToRY-scale experiments were carried out on uranium hexafluoride absorption on and desorption from sodium fluoride, as used in a fluoridevolatility process for spent nuclear reactor fuels (7, 2, 4). Chemical interaction of UFO with NaF or K F was indicated first by Ruff and Heinzelmann (8). Martin, Albers, and Dust substantiated the earlier work, but Grosse indicated the presence of hydrogen fluoride was necessary in forming a ternary complex (3, 5 ) . T h e reaction is: UFe
+ 3NaF
-r
UF6.3NaF
(1)
the product being a yellow solid (5). This equation is assumed essentially correct as a basis for discussion, although in some tests the NaF/UF6 ratio was less than 3 to 1 .
Formation of UF8-NaF Complex T h e three types of NaF studied had distinctly different UFB absorption capacities (Table I) in the original state; fine reagent-grade NaF absorbed the largest amount. One valve, 1.9 grams of uranium per gram of NaF, was close to the 1.89 calculated for UF6.3NaF; 2.4 grams in three other runs, corresponds more closely to 2UF6.5NaF. Absorption by macrocrystalline and pelletized NaF, prior to grinding, was low. Pelletized NaF had a particle porosity of 4870 and therefore a higher initial capacity than the macrocrystalline material with a particle porosity of only 6%. The capacity of both materials was greatly increased by grinding, the greatest increase being for the macrocrystalline NaF, which originally had a low specific surface area. Only a qualitative correlation appears possible among particle size range, surface area, and UF6 absorption capacity. T h e experimental procedure consisted in conditioning 1 to 5 grams of NaF a t 100' C. under a vacuum of approximately 1 mm. of mercury in a nickel or copper tube, then admitting UF6 gas until the pressure exceeded 760 mm. The reaction was fairly complete in 1 to 2 minutes (as indicated by decreasing UFO pressure), although 5 minutes was usually allowed for equilibrium-i.e., for the pressure to become static. Excess UF6 gas was removed by evacuation, followed by filling and sweeping with dry nitrogen while the sample
tube was cooled to room temperature, Uranium absorption was determined by both increase in sample weight and colorimetric analysis of the material dissolved in 0.67M aluminum nitrate4M nitric acid (7). Results agreed closely. Small variations in experimental procedure did not affect capacity values significantly. For '/B-inch pellets, initial temperature of absorption was varied from 70' to 150' C. and U F O equilibrating pressure from 0.5 to 1.3 atm. Ground material was classified by dry-screening. Small weight loss on heating to approximately 300' C. indicated the presence of little hydrofluoric acid or water. A ternary complex as suggested by Grosse does not appear valid. T h e absorption complex was markedly hygroscopic. The complex x-ray diffraction pattern was different from that of UFe or NaF. The NaF pattern was abs,ent where saturation was complete.
Equilibrium Partial Pressure of UF6 over UF6-NaF Complex Formation of the UF6.3NaF complex (Equation 1) appears reversible, as UF6 gas is formed on heating. This reversibility does not result in appreciable UFOpressure below 200' C. T h e equilibrium partial pressure of UF6 over the complex was studied by transpiration
Table I. Reagent Grade NaF Absorbed Most UFO NaF Spec. Spec. U Surface Particle Capacity, Areaa, Mesh Size G U/G. NaF Sq.M./G. Reagent-grade NaF, Baker and Adamson As received 2.4 80% -275 -325 400 - 400
+
1.9 2.4 2.4
... ... 1.06
1.21
Macrocrystalline NaF, Allied Chemical -20
+ +
40 -100 120 - 140 -325 f 400
0.064 0.27 1.2 '1.5
0.25 0.77 0.75 1.16
Pelletized NaF, Harshaw Chemical '/*-inch pellets 0.61-0.90 1.28
+ 20 + 120
-12 -100 - 140 -325 (I
+ 400
0.93 1.2 1.6 2.1
1.51 1.49 1.43 2.10
Determined by nitrogen adsorption.
from SO' to 320' C. Partial pressure data were remarkably consistent; pressure varied as an exponential function of temperature over a wide range. Linearity of data indicated unit activity for the complex solid, and lack of phase or chemical transitions. Partial pressure data were fitted by: Log@ = 10.88
- (5.09
X lO3/T) (2)
where b, is partial pressure of UP6 (millimeters) and T is absolute temperature. Use of the Clausius-Clapeyron formula gave -23.2 kcal. per mole of UF6 for the enthalpy change of Equation 1. Decomposition temperature under atmospheric pressure was calculated as 363' C. The transpiration method consisted of passing a stream of dried nitrogen through a UF6-NaF complex bed, prepared by saturating 30 grams of NaF (12-20 mesh, Harshaw Chemical co.) with UF6 a t 100' c. in a vertical nickel reactor 1 inch in diameter. Nitrogen volume was measured with UFe vaporized in a wet-test meter. the nitrogen stream was determined by analysis of a 1M aluminum nitrate solution in which UFs was trapped. Nitrogen flow rates of 20, 50, and 100 ml. per minute (STP) were used in working a t different temperatures. Variation of UF6 partial pressure was made possible by the accuracy of uranium analytical procedures (&5%), over a wide range, and use of various time periods. At high temperatures, or high partial pressures of UF6, it was necessary to heat the line between reactor and hydrolysis trap.
Decomposition of UFB-NaFto UF6-NaF and UF4-NaF Complexes Martin, Albers, and Dust reported formation of two other sodium fluorouranate complexes when Ups. 3NaF complex was heated ( 5 ) : UF6.3NaF 4 UFr.NaF
+ 2NaF + 0.5F2
(3)
between 200' and 450' C. and UF5.NaF -+ UF4.NaF
+ 0.5F2
(4)
above 450' C. T h e UFs.NaF complex was colorless; the UFd. NaF complex was green. Data in this investigation agreed with theirs. HowFver, the fluorouranate decomposition product in Equation 3 may not be a l to l complex. VOL. 50, NO. I t
NOVEMBER 1958
1709
tion (Equation 4). Heating the white U(V) complex to 550" C. changed the color to pale green, with evolution of a gas that appeared to be fluorine.
THERMOCOUPLE
Application of NaF Absorption DRY
SUPPLY
This apparatus was used to determine eauilibrium partial pressure of UFs over UFe-Naf complex
Several exploratory tests of Equation 3 were made by heating UFs 3NaF complex in a nickel reactor to an elevated temperature for various periqds under UFs pressure. The excess UF6 pressure was required to prevent decomposition (Equation 1). The excess UFG. 3SaF complex present at the conclusion of each experiment was decomposed by flushing the system with dry nitrogen. In a typical case, heating to 400" C. produced a white residue with a U(V) content of 21%. The material was dissolved in hot 85% phosphoric acid, followed by determination of the b(1V) and (VI) formed by disproportionation of the U(V). The U(1V) was measured by oxidation with ferric chloride, the U(V1) by a thiocyanate colorimetric method (6, 7). The U(1V) and (VI) contents of the hydrolyzed material were always equal. The rate of UFG 3NaF decomposition to U(V) complex was studied from 245 " to 345' C. NaF samples (5 grams of 12-20 mesh in a I/ 2-inch nickel reactor) held at the desired temperature were subjected to a flowing stream of UFs gas (100 ml. per minute) at atmospheric pressure. 'This dynamic method prevented back-reaction by removing the fluorine formed, kept the material saturated, and avoided uncertainty in the time required to heat a UFe 3NaF complex bed already prepared. A fast UFG flow rate was used initially to saturate the NaF bed. The highest temperature (355' C.) was below estimated atmospheric decomposition temperature (363" C.). The V(V) content of the residue was assumed to be twice the U(1V) analysis, as excess T;F6 (present as UFG 3NaF) was not removed at the end of each run. The reaction was assumed to be firstorder, the specific reaction rates, k , being calculated and plotted against reciprocal absolute temperature (Table 11). The specific uranium capacity of 0.9 gram of uranium per gram of sodium fluoride was used in calculation of reaction rates. Although the data
1710
showed considerable scatter, the relation Log k = 6.09 - (5.22 x lO3/T) (5) was secured by statistical treatment (omitting the three lowest k values). An activation energy of 23.9 kcal. per mole was calculated, approximately the same as the enthalpy change for Equation 1, suggesting some relation between the two processes. The total V(V) content of the kinetic work was kept below 25%, the amount formed in pellet-type NaF. Repeatedly exposing reagent-grade NaF to UFS for 16 30-minute consecutive cycles (absorption at 100" C., decomposition at 400" c. under excess UFs pressure) gave awhite product with 51.6% uranium. A U(1V) analysis of 27.2% was secured after hydrolysis. corresponding to 54.4% U(V). As the calculated uranium content of U F j 3NaF is 51.9%, the reaction is probably UF8.3NaF (s) -+ U F d N a F (SI l / 2 F 1 (g) (6) The material after 16 cycles showed no affinity for UF6. The x-ray diffraction pattern was different from that of NaF or UFe 3NaF. The slight excess of U(1V) was probably due to decomposi-
+
Table 11. Specific Reaction Rate Constants for Decomposition of uFs.3NaF Complex to UF5.3NaF DeTemp.,
c.
245 246 281 288 306 309 314 317 322 327 335 335 345 355
INDUSTRIAL AND ENGINEERING CHEMISTRY
compn. Time, Mill
120 120 125 120 60 75 30 30 20 16 16 15 15 15
UF5 in
Product, Xt. % 3.82 2.46 7.72 3.09 8.09 15.28 3.50 18.77 5.52 4.37 8.47 18.89 9.37 13.34
Spec. Reaction Rate i?Iin.-1 x 104 2.54 1.60 4.93 2.06 10.7 16.2 9.3 16.6 22.0 21.7 42.0 29.0 49.6
70.5
I n laboratory tests the recovery of UFe from nitrogen or fluorine gas streams on NaF at 100" C. has generally been predictable from the equilibrium UFG partial pressure of the complex at this temperature. This value, 1.5 x mm. of mercury, is equivalent to the vapor pressure of UF6 at a cold trap temperature of -65" C. Extrapolation of the equation indicates that the UF6 vapor pressure over UFG.3NaF at 25" C. would be less than 10-6 mm. of mercury, much better than that practicable in cold trapping. UFGis almost completely desorbed by raising the temperature of the complex to 400" C. in the presence of fluorine or nitrogen. I n the presence of nitrogen decomposition to the U(V) complex results in a small retention of nonvolatile uranium. This is avoided by use of fluorine as a sweep gas; back-fluorination (Equation 6) apparently occurs. The mechanics of the desorption process are complex, owing to the problem of heat transfer through a granular bed. Much of the UFOis evolved close to the estimated decomposition temperature of 363" C. at atmospheric pressure. The rate of evolution is apparently controlled by the endothermic heat of 23 kcal. per mole of UFG for desorption. Acknowledgment
The authors thank J. F. B. Read and R. M. Duff for carrying out experimental procedures and G. W. Wilson, C. A. Pritchard, and R. L. Sherman for supervising analytical and x-ray crystallography work. Literature Cited (1) Cathers, G. I., Nuclear Sci. Eng. 2, 768
(1957).
(2)' Cathers, G. I., Leuze, R. E., "Reactor Operational Problems," vol. 11, p. 157, Pergamon Press, New York, 1957. (3) Grosse, A. V., "Chemical Properties of Uranium Hexafluoride," SAM Columbia 3, Rept. A-83 (1941). (4) Hyman, H. H., Vogel, R. C., Katz, J. J., Proc. of Intern. Conference on Peaceful Uses of .4tomic Energy, voi. 9,
p. 613, United Nations, New York, 1956.
(5) Martin, H., Aibers, A,, Dust, H. P., 2.anorg. u. allgem. Chem. 265, 128 (1951). (6) Pritchard, C. A , , Methods 1-21'3272
and 9-00719272, Oak Ridge Xatl. Lab. Master Analytical Manual, 1956.
(7) Rodden, C. J., "Analytical Chemistry of the Manhattan Project," NNES, vol. VIII-1, p. 104, McGraw-Hill, New
York, 1950.
(8) Ruff, O., Heinzelmann, A,, Z. anorg. u . allgem. Chem. 72, 63 (1911).
RECEIVED for review October I?, 1957 A C C E P T E D May 2, 1958 Divisions of Industrial and Engineering and Petroleum Chemistry, Symposium on Nuclear Technology in the Petroleum and Chemical Industries, 131st Meeting, ACS, Miami, Fla., April 1957.
'
