April, 1956
PHYSICAL PROPERTIES OF AQUEOUS URANYL SULFATE
413
PHYSICAL PROPERTIES OF AQUEOUS URANYL SULFATE SOLUTIONS FROM 20 TO 90’ BY EDWARD ORBAN,MARTINK. BARNETT, JANES. BOYLE, JOHNR. HEIKSAND LERROY V. JONES Mound Laboratory, Monsanto Chemical Company, Miamisburg, Ohio’ Received September 8, 1066
Measurements of the density, surface tension, viscosity and pH of aqueous solutions of uranyl sulfate up to as high aa 4 molal in concentration and in a temperature range between 20 and 90” are reported.
Introduction I n recent years, interest in the common physical properties of uranium salts has risen sharply, revealing the lack of adequate data. The measurement of the density, viscosity, surface tension and pH of a series of aqueous solutions of uranyl sulfate between the temperatures of 20 and 90” is described in the present paper. Some early work on the densities of solutions between 0.03 and 0.3 molar a t temperatures between 10 and 16” was reported by de Coninck.2 His results do not agree well with the more precise work done by Helmholtz and Friedlander3 on three concentrations between 0 and 90” and on eight other concentrations at 30 O. The surface tension of aqueous salt solutions has generally been found t o increase with increasing concentration. Of approximately 50 salt solutions covered in the International Critical Tables4 only two salts, ammonium chloride and magnesium acetate, violate this rule. However, Young and Coons6 found pronounced maxima and minima in the surface tension versus concentration curves. Grant,s et al., reported that uranyl acetate lowers the surface tension of water although a positive coefficient was found for uranyl nitrate. Values for uranyl sulfate have not been reported. Jones and Ray7 have noted minima for very low concentrations of KC1, KzS04 and Ce(N03)3,but Langmuir* has attributed this to an effective decrease in the radius of the capillary tube by an adhering film of solvent . Two groups of investigator^^.^ made measurements on the pH of uranyl sulfate solutions a t room temperature, but studies were not made with an excess of uranium trioxide. There have been no measurements of the viscosity of uranyl sulfate solutions reported in the literature. Equipment and Material The uranyl sulfate was prepared from uranium trioxide purchased from the Mallinckrodt Chemical Works. The (1) Mound Laboratory is operated by the Monsanto Chemical Company for the United States Atomic Energy Commission under Contract No. AT-33-1-GEN-53. (2) F. W. 0. de Coninck, Ann. chim. phys., 171 28, 5 (1903). (3) L. Helmholtz and G. Friedlander, Atomic Energy Commission Report MDDC-808, December 15, 1943. (4) “International Critical Tablea,” Vol. IV. McGraw-Hill Book Co., New York, N. Y., 1923, p. 463. (5) C. B. F. Young and K. N. Cooae. “Surface Active Agents,“ Chemical Publishing Company, New York. N. Y., 1945. (6) W. E. Grant, W. J . Darch, 8. T. Bowden and W. J. Jonea, THIS JOURNAL, 6 a , m 7 (1048). (7) G . Jones and W. A. Ray, J . Am. Chem. Soc., 67. 957 (1935); 6S, 187 (1937). (8) I. Langmuir, Science, 88, 430 (1938). (9) D. A. MacInnes and L. G . Longsworth, U. S. Atomic Energy Commission Report, MDDC-911, November 24, 1942.
oxide was first screened through a U. S. Standard Sieve #40 and then a #SO. It was washed with distilled water by stirring vigorously a t 80 to 90’ for six hours and then allowing to stand overnight at room temperature. Following decantation and filtration, the washing process was repeated until no nitrate could be detected in the wash water. The dried oxide was stirred with a stoichiometric amount of 2.6 M sulfuric acid. The mixture was heated with stirring to 80” until all of the oxide was dissolved. The solution was filtered and analyzed for uranium and sulfate giving results which permitted the adjustment of the solution by the addition of either oxide or sulfate to a molar ratio of l.W f 0.002. Some of this stock solution was diluted to prepare samples while another portion was recrystallized six times to give pure UOB04.3H20 crystals. No difference between the solutions prepared from this recrystallized material and the original stock solution could be detected in any of the measurements. The densities of the solutions were obtained by the hydrostatic weight method. A hollow bob, the weight of which had been adjusted by sealing a number of small steel balls inside, was suspended from the beam of an Ainsworth type DLB Chain-0-matic balance with a length of 36-gage platinum wire. Constant temperature was maintained around the test solution vessel by immersing it in an oil-bath which could be set within 3~0.05”of the test temperature. The viscosities were measured with Ostwald capillary viscometers of 5-ml. volume. A constant temperature waterbath maintained the desired temperature to within 10.05”. The surface tension measurements were made with the Pyrex capillarimeter similar in design to that employed by Young, Gross and Harkins.lo The large tube was about five centimeters in diameter. The capillary tubing was “Precision Grade” and had a nominal bore of 0.5 mm. The capillarimeter was immersed in water in a windowed thermostat that could also be controlled to =k0.05”. The heights of the menisci were observed with a Gaertner Micrometer-slide cathetometer. A Leeds and Northrup Universal pH poteptiometer assembly No. 7663-A-I was used to make pH measurements. Out of seven “L and N” glass electrodes available, two were selected which, when set with a buffer of p H equal to 4.01, agreed with the value of other buffers a t pH’s of 1.08, 2.08 and 6.86 a t 25”. The electrodes, buffers and test solutions were maintained a t equilibrium temperatures in a water-bath controlled to =k0.05”.
Procedure and Results Density.-The volume of the bob used in the density measurements was determined at each temperature by making weighings in doubly-distilled water. From the weight of the bob in air and the literature values11 of the densities of water at each temperature, the volume of the bob was calculated. The densities of 12, solutions were measured at 20, 30, 44.8, 59.8, 75 and 90 . The samples were stirred thoroughly. before each determination. Since evaporation occurred a t the higher temperatures, a small sample was taken for analysis, immediately after each measurement. Table I presents the density data obtained. Solutions were originally prepared a t 20” on a molar basis, and later converted to a molal basis by means of the meas(10) A. Weissberger, “Physical Methods of Organic Chemistry,” Vol. I, part 1, Interscience Publishers, Inc., New York, N. Y., 1949, p. 367. (11) C. D. Hodeman. “Handbook of Chemistry and Physics,” 30th Ed., Chemical Rubber Publishing Co., Cleveland, Ohio, 1947, p. 1695.
414
EDWARD ORBAN,M. K. BARNETT, J. S. BOYLE, J. R. HEIKSAND L. V. JONES TABLE I
DENSITYOF UEANYLSULFATE SOLUTIONS (GJCM.~) Molality 0.176 .177 .178 ,180 .452 .450 .454 .458 671 .669 .681 .864 .870 .868 .877 1.049 1.053 1.055 1.321 1.323 1.329 1.753 1.756 1.761 1.770 1.983 1. Q86 2.002 2.009 2.865 2.862 2.858 2.870 3.872 3.880 3.959 3.Q99
20.0 1.0539
30.0 1.0510
Temperature, OC. 44.8 1.0445
59.8
75.0
1.0293 1.0193 1.1384
1.1349
1.1282 1.1208 1.1113 1.1008
1.2050
1.2012
1.1939
1 1862 1.1761 1.1654
1.2630
1.2588
1.2513 1.2432 1.2334 1.2219
1.3141
1.3098
1.3023
1.2948 1.2845 1.2720
1.3911
1.3865
1.3779 1.3692
1.3458 1.3579
1.5089
1.5032
1.4956 1.4867 1.4735 1.4620
1.5672
1.5613
1.5516 1.5430 1.6328 1.6236
1.7797 1.7736 1 .7633 1,7533 2.0059
1.9983
1.9856 1.9752 1.9708 1.9646
ured densities. The statistical relation of the original molar concentration (M) to the density (d) in granla per cubic centimeter a t 20” was found to be M 3.18Qod - 3.1887 Upon extrapolating to M = 0. the density does not approach the density of water; to do so the line must assume a more gentle slope a t lower concentrations. This is in a eement with the work of Lee,” et al., who reported a smaKr slope at low than a t h g h concentrations. Viscosity.-The relative viscosities of the uranyl sulfate solutions were measured with an Ostwald viscometer according to standard procedures.18 The calculations were made using the literature values for the densities of water,” and the densities of the solutions as reported in this publication. The absolute viscosities listed in Table I1 were calculated from the absolute viscosities of water.“ Surface Tedsion.-For the measurement of surface tension, the solution was introduced into the ca illarimeter and the latter was sealed off. The space above t i e solution contained air at room temperature. The entire apparatus was immersed in a thermostat. Before each measurement the capillarimeter was tipped until the solution flowed up the capillary tube into the overflow tra thus giving a fresh surface. After leveling precisely anahlowing the capillary to come to equilibrium, the difference in the elevation of the large and small menisci were observed with the cathetometer. The bore of the capillary tube was determined by noting the capillary rise exhibited by Pure benzene at 20”. &om known coefficients of expansion it was determined that the change in size of the capillary became of the temperature change was beyond the limit of accuracy of the experiment. Enlarged photographs of the menisci showed that the con___-
(12) J. E. Lee, Jr.. R. Rowan, Jr., C. D. Susana and 0. Menis, United States Atomic Energy Commission Dwument ORNL 1332, July 28, 1952. (13) Joseph Reilly and William Rae, “PhysicdXmnioal Methods.” Vol. 1, D. Vnn Nostrand Co.. Ino., New York, N. Y.. 1843. p. 648. (14) E. C. Bingham and R. F. Jackaon. Bull. Bur. Standords, 14,75 (1918).
TABLE I1 ABSOLUTE VISCOSITYOF URANYLSULFATE SOLUTIONS (CENTIPOISE)
90.0
1.0343
Vol. 60
Molality
20.0
30.0
0.176 .452 .671 .864 1.049 1.321 1.531 1.753 1.978 2.865
1.12 1.37 1.58 1.80 2.04 2.50 2.91 3.51 4.18 8.53
0.881 1.06 1.23 1.42 1.58 1.93 2.22 2.64 3.15 6.21
Temperature, “C. 45.0
60.0
75.0
0.659 0.779 0.893 1.01 1.14 1.37 1.57 1.89 2.21 4.22
0.516 .610 .685 .779 .872 1.04 1.18 1.40 1.67 3.07
0.422 .494 .547 .608 .673 .836 .931
90.0
0.349 .403 .460 .501 .561 .656 .735 1.11 .865 1.26 1.01 2.30 1.82 tact angle was essentially zero. Also the precision of the measurements did not justify taking account of the density of the gas phase. Therefore, the equation used for the calculations reduces to the familiar form y = ((h r/3)dgr), where y is the surface tension, h is the capillary rise, r is the radius of the capillary, Q is the acceleration due to gravity and d is the density of the liquid. Table I11 summarizes the surface tensions of the uranyl sulfate solutions. TABLE I11 SURFACE TENSION OF AQUEOUSURANYL SULFATE SOLUTIONS (DYNEs/CM.)
+
Molality
20.0
Temperature, OC.
30.0
45.0
60.0
75.0
0.175 73.47 71.95 69.59 67.04 64.65 0.446 74.00 72.56 69.98 67.43 64.83 0.606 74.43 72.75 69.91 67.29 64.60 0.728 74.18 72.58 70.23 67.75 65.06 0.856 74.74 73.10 70.20 67.67 65.13 0.881 74.12 73.22 70.68 68.50 65.87 0.965 75.14 73.38 70.90 68.30 65.64 1.016 74.60 73.12 70..76 68.32 65.56 1.184 75.83 74.13 71.44 68.96 66.26 1.445 76.08 74.29 71.76 69.18 66.55 1.672 76.09 74.26 72.18 69.61 66.68 1.837 76.70 75.03 72.81 70.33 67.70 2.340 77.43 75.90 73.41 70.96 68.24 The data show that the average rate of increase in surface tension with increasing concentration over the concentration range embraced by the study and for five different temperatures is 2.02 dynes per centimeter for each unit increase in molality. Hence, dy/dm = 2.02 dynes per $m. er mole. Also, dr/dt = -0.168 dyne per cm. per C. keglecting the departure from linearity, the equation which described the influence of concentration and temperature on the surface tension becomes y = 2.02m - O.l68(t - 20) 72.75 where m is molality and t is degrees centigrade. JH Measurements.-Two glass electrodes, used for each p measurement, were standardized before each determination with a reference buffer solution of pH 4. The readings from the electrodes did not differ from each other by more than 0.03 pH unit in any of the measurements. The electrodes were cleaned with dilute hydrochloric acid after each series of four measurements except a t the higher concentrations of uranyl ion where i t was found necessary to clean the electrode after each measurement. At high concentrations a slow drift of pH was noticed indicating a reaction of the uranyl ion with the glass. The pH of each uranyl sulfate solution was measured a t 24.7,35.0,44.8and 59.8’, in the order named. After each measurement had been completed a t a particular temperature, the solution was immediately removed from the bath to avoid any change in concentration due to evaporation. A final measurement at the initial temperature agreed with the original measurement in each case. The measurements at the four temperatures over a concentration range from O.OOO6 to 3.85 m are given in Table IV.
+
PHYSICAL PROPERTIES OF AQUEOUS URANYL SULFATE
April, 1956
TABUIV TEEp H OF URANYL SULFATE SOLUTIONS Soh. no.
Molality
Temperature. "C.
24.7
35.0
44.8
0.00066 4.09 4.02 3.97 3.89 .0013 3.90 3.81 3.76 3.69 .0066 3.47 3.39 3.33 3.27 .0132 3.27 3.18 3.12 3.04 .0758 2.73 2.65 2.58 2.47 .157 2.48 2.38 2.29 2.18 .321 2.18 2.06 1.98 1.84 .824 1.72 1.61 1.52 1.38 1.731 1.25 1.13 1.01 0.89 3.117 0.68 0.52 0.41 0.30 11 3.850 0.33 0.18 0.04 -0.12" Obtained by extrapolating deflections to a zero reading. Similar measuremente were made with solutions that had excess UOs dissolved in them. The data are shown in Table V. The p H values of these solutions are considerably higher than the ure uranyl sulfate solutions as would be expected from the gasic nature of the oxide.
,.
OF
TABLEI V URANYLSULFA^ SOLUTIONS CONTAINING Ex-
Uranium, m
Sulfate,
1.700 1.904 0.460 .046 .0045 2.069 0.497 .049 .0049 2.234 0.536 .051 .0051 1.344 0.333 .033 .003
1.700 1.738 0.420 .042 .0041 1.730 0.416 .041 .0041 1.738 0.417 .039 .0039 0.954 0.236 .024 .0023
m
uo,
CESS Ratio U:SOd'
24.7
1.00 1.10 1.10 1.10 1.10 1.19 1.19 1.19 1.19 1.29 1.29 1.29 1.29 1.41 1.41 1.41 1.41
1.24 2.10 2.89 3.54 3.95 2.54 3.19 3.78 4.09 2.82 3.35 3.90 4.21 3.29 3.68 4.08 4.32
Temperature, O C . 35.0
44.8
1.18 1.96 2.76 3.38 3.83 2.38 3.08 3.59 3.96 2.68 3.18 3.72 4.08 3.12 3.54 3.93 4.18
0.98 1.72 2.42 3.17 3.68 2.09 2.71 3.46 3.88 2.35 2.91 3.60 3.89 2.89 3.42 3.82 4.12
proximately a straight line. These linear relations for the two properties may be combined to give a general e uation, b, where a and b are constants. Bo& 5 and lo q = oy b %ave been found to be h e a r functions of temperature so that the equation may be written as
+
59.8
1 2 3 4 5 6 7 8 9 10
THEpH
415
69.8
0.93 1.36 2.13 2.83 3.53 1.78 2.41 3.08 3.75 1.94 2.59 3.18 3.74 2.40 2.69 3.69 3.99
Discussion A plot at any one temperature of either surface tension ( 7 )
or the logarithm of the viscosity ( q ) against molality is ap-
+
+
(0.15 - 0.00046t)y (-11.32 0.046t) log 1) where t is the temperature in de ees centigrade. The hydrolysis of uranyl s x t e to form hydrogen ions may be expressed by a number of equations. some of which include the formation of complexes with the anion. The effect of the sulfate ion on the e uilibrium was studied by adding sodium sulfate to a urma sulfate solution. If the sulfate ion took part in the reaction, the equilibrium would have shifted,. changing the pH of the solution. However, the increase in pH was almost exactly that calculated from the sulfuric acid equilibrium expression assuming only the addition of the sulfate ion with no other reaction occurring. Thus, the sulfate ion ap arently did not enter into the equilibrium involving the gydrolysis of the uran 1 ion. A number of equations may be written to expyain the hydrolysia of the uranyl ion to produce hydrogen ions. Nine of these, in which the uranyl ion to water ratio was varied from 1 to 1 up to 6 to 1, were studied. The equilibrium constant for each of these equations was calculated at each concentration by neglecting the small but unknown liquid junction potenha1 pre&nt in the glass electrode measurements and calculating the activity of the hydrogen ion directly from the pH measurement. The activity coefficients used for the uranyl ion and the com lexes were the mean activity coe5cients measured by C. Secoyu for uranyl sulfate. Up to 0.025 m solutions of uranyl sulfate, constant values (8.8 X 104) for the equilibrium constant for the equation
%.
+
2U01++ B O
a
+
UOa*UOp++ 2H+
(1)
were obtained. From about 0.01 m up to the highest concentration the equation
4UOt++
+ &O
UO,.(UOa),'+
+ 2H+
(2)
gave near1 constant equilibrium constants of about 5.5 X 10-9. A &gher form of a uranyl complex would be expected in more concentrated solutions. Equation 1 is similar to the one found by MacInnes and Longsworth@to hold for dilute solutions of uranyl nitrate. It was assumed that by the addition of U08 to the solutions a reaction occurred with the simple uranyl ion to form the UOI complexes. When a small amount of UOa waa dissolved (10mole % in excess of the uranyl ion present) e UBtion 2 did, in fact, give ood constant values for the eq&brium constant, but witg larger amounts of UOa a marked deepenin of color occurred, and no equation could be found that woufd give constant values. Presumably, an ion, more complex than that shown in equation 2,was being formed. (15) C. H. Secoy, J . Am. Clrsm. BOG., 70, 3450 (1948).
