JOSEPH F. SPINNLER AND ANDREW PATTERSON, JR.
500
The Wien Effect in Uranyl Ion Solutions. I. and Perchlorate from 5 to 6 5 O .
Uranyl Nitrate
Negative Wien Effects'
by Joseph F. Spinnler and Andrew Patterson, Jr. Contribution No. 1 YO6 from the Sterling Chemistry Laboratory, Yale University, New Haven, Connecticut (Received August 24, 1964)
Experimental data on aqueous solutions of uranyl nitrate and uranyl perchlorate are presented for the temperatures 5, 15, 25, 50, and 65'. Plots of the pH us. log of concentration, of the equivalent conductance vs. square root of concentration, and of the high-field conductance as a function of field are presented. The range of concentration studied lies and loF3M. The high-field conductance results are unprecedented, since between under most circumstances application of the field decreases the conductance of the solutions, a phenomenon not heretofore observed with any other electrolytes.
Prompted by the unexpected observation of a decrease in the conductance of a solution of uranyl nitrate under the influence of a high electrical field,z we have studied the pH and low- and high-field conductance of a series of uranyl salts, including the flUOride,38sulfatelab nitrate, and as a function of concentration and temDerature. In this DaDer - are presented the results of n;easurementsof the p~ and conductance as a function Of for aaueous solutions of ,uranyl nitrate and perchlorate over a temperature range from 5 to 65'. In connection with this study we have computed the theoretical high-field conductances of these solutions using the theory of Onsager and Kim as programmed for machine computation. Extensive investigations have been made of the hydrolysis of the uranyl ion in the solutions of its various salts15-1g to which reference will be made in the Discussion. Of particular interest are the results of Baes and Meyer,l8 who have made acidity measurements on uranyl nitrate a t elevated temperatures up to 148'. They find a rapid increase of hydrolysis with increasing temperature and an increasing proportion of the monomeric hydrolysis product, UOzOH+, as predicted by Kraus6 and confirmed by the work of Hearne and White.15 The conductance of uranyl salts in solution has been investigated by numerous workers as well, but for the present purpose the measurements of Brown, The Journal of Phyeieal Chemistry
Bunger, Marshall, and SecoyZ0 and of West and Jonesz1 are relevant. The measurements of Brown, et al., were made a t sufficiently low concentrations (1) This paper is taken in part from a dissertation submitted by J. F. Spinnler to the Graduate School, Yale University, in partial fulfillment of the requirements for the degree of Doctor of Philosophy, May 1961. Fo; those who are interested, a limited number of copies of this dissertation are on hand and one can be provided on request. (2) F. E. Bailev. J. F. Soinnler. and A. Patterson. Jr.. J . A m . Chem. Sic., 83, 1761 i1961). (3) (a) J. F. Spinnler and A. Patterson, Jr., J. Phys. Chem., 69, 508 (1965); (b) ibid., 69, 513 (1965). (4) H. Freitag and A. Patterson, Jr., J . Electrochem. SOC.,108, 529 (1961). (5) J. Bjerrum, G. Schwarzenbach:, and L. S i l l h , "Stability Constants; Part 11, Inorganic Ligands, The Chemical Society, London, 1958, p. 9. (6) K. A. Kraus, Proc. Intern. Conf. Peaceful Uses At. Energy, Geneva, 7, 245 (1956). (7) J. Sutton, J . Chem. Soc., S275, S57 (1949). (8) J. Sutton, National Research Council of Canada, Atomic Energy Project Report CRC 325 (1947). (9) S. Ahrland, Acta Chem. Scand., 3, 374 (1949); 5 , 1151, 1271 (1951); 8, 1907 (1954). (10) B. Singh and G. Ahmad, J . Chem. Phys., 34, 351 (1937). (11) L. G. Longsworth and D . A. MacInnes, USAEC Report MDDC 911 (1947). (12) H . Guiter, Bull. S O C . chim. France. 64 (1947). (13) J. Paucherre, Compt. rend., 227, 200 (1948). (14) R. H . Betts and R. K. Michels, J . Chem. SOC.,5286, S58 (1949). (15) J. A. Hearne and A. G . White, ibid., 3168 (1957). (16) J. Sutton, Nature, 169, 235 (1952). (17) H . W. Crandall, USAEC Report MDDC 1294 (1947).
WIEN EFFECT IN URANYL IONSOLUTIONS
and with adequate precision for determination of the limiting equivalent or ionic conductances, but the nitrate and perchlorate salts were not studied. Reference 21 gives data for uranyl nitrate at 35, 50, and 65O, temperatures used in this work, but for concentrations higher than those of interest to us. Data for uranyl perchlorate are not available.
Experimental The experimental procedure was essentially that of Gledhill and Patterson.22 Two circuit changes contribute to increased precision of measurement, which is much to be desired, since the size of the effects observed is small: a vernier control of high voltage was provided, and small variable capacitors were shunted across the two primary windings of the bridge pulse transformer to compensate for small differences in the capacitance in the transfornier. The precision of ineasurement on one solution a t a given session of measurement was within 0.02 unit (absolute per cent) in AA/A(O); deviations of less than 0.05 unit were obtained in successive investigations a t different times on the same solution. The principal limitation on the precision of measurement is drift in the resistance of the solutions. These drifts are not the result of inadequate temperature control, this being held to within 0.002O a t all temperatures reported, but rather appear to be the result of passing high power pulses through the cells. The low-field conductances are measured to 0.1 ohm, but owing to these conductance drifts are not known to better than 0.05% during the entire course of an experiment. The uranyl nitrate was a sample provided some years ago by the Atomic Energy Coniniission for conductance measurements. The uranyl perchlorate was purchased from the A. 0. Mackay Co. Each salt was recrystallized froin conductivity water and air-dried. The airdry salts were added directly to the conductance cells until the desired cell resistance was obtained. The concentrations of the solutions whose Wien effects are reported were determined from the measured lowfield conductances and interpolation froin specific conductance-concentration graphs prepared in conjunction with the pH measurements referred to below. The concentrations are known to within 1 part in 500. This procedure, which leaves something to be desired in terms of precision, was chosen to avoid exposure of the cell contents to the atmosphere. The conductances reported in this paper were measured principally to aid in this determination of concentration and were made to four-figure precision only.
501
For the pH measurements, stock solutions of uranyl salts were prepared from the air-dry salts and conductivity water. The concentration was determined by the method of Frere,2ain which uranyl ion is precipitated with 8-hydroxyquinoline in an acetate buffer. The analyzed stock solut,ion was added to a conductance cell containing conductance water through a niicroburet whose tip projected into the upper part of the cell through a special cap which allowed passage of a constant stream of purified nitrogen through the upper part of the cell. The concentration of the solution was determined from the amount of stock solution added to the previously weighed conductivity water. Samples for the p H determination were withdrawn from the cell through the special cap after attainment of a stable conductance, indicating that' the solution was thoroughly mixed. The low-field conductance was measured on the differential pulse transformer (DPT) bridge to within 0.1 ohm and recorded for later use as noted in the preceding paragraph. The pH nieasurement.s were made with a Pye Catalog No. 11085 p H meter. The electrode compartnient, was suspended in the thermostat oil bath together with the conductance cells. The meter was calibrated against a suit'able reference buffer solution a t each temperature. All parts of the apparatus and the solutions were allowed to reach temperature equilibrium before a nieasurement was made. The pH electrodes were rinsed in a 100-ml. sample of solution withdrawn froiii the conductance cell and the p H measured on a separate sample taken from the cell shortly before use. The pH electrode compartment was kept under an atmosphere of purified nitrogen. The contents of the conductance and p H cells were magnetically stirred except while the p H measurement was actually being made. Additional stock solution was then added to the conductance cell and the measurement continued a t a higher concentration. The precision of measurement is 0.01 p H unit; the probable accuracy is estimated to be within 0.05 unit at' 25" and below, and 0.1 unit at 35Oand above. Stock solutions of the reference electrolytes were prepared and analyzed in accord with appropriate analytical and conductance techniques. The reference (18) N. P. Komar and Z. A. Tretyak, Zh. Analit. Khim., 10, 236 (1955). (19) C. F. Baes and N. J. Meyer, ORNL Reactor Chemistry Division Annual Progress Report, May 1961. (20) R. D. Brown, W. B. Bunger, W. L. Marshall, and C. H. Secoy. J . Am. Chem. Soc., 7 6 , 1532, 1580 (1954). (21) A. P. West and H. C . Jones, Am. Chem. J . , 44, 508 (1913). (22) J. A. Gledhill and A. Patterson, Jr., J . Phys. Chem., 5 6 , 999 (1952). (23) F. J. Frere, J . Am. Chem. Soc., 5 5 , 4362 (1933).
Volume 69, Number
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February 1966
JOSEPH F. SPINNLER AND ANDREWPATTERSON, JR.
502
electrolytes were prepared directly in the conductance cells by adding measured volumes of the stock solutions to known weights of conductivity water. In most measurements, the reference electrolytes for the uranyl salt, solutions were salt-acid mixtures; these were prepared by adding hydrochloric or nitric acid first and adjusting the pH to within 0.1 p H unit of that of the uranyl salt solution under study, followed by the addition of potassium chloride or potassium nitrate to establish the desired cell resistance. These reference electrolytes have been found to remain remarkably stable in conductance over extended periods of time in spite of the brutal electrical treatment they receive, and to obviate experimental difficulties due to differing polarization in the unknown and reference electrolytes. Because it is necessary to know some details of the behavior of these reference electrolytes in order to compute the Wien effects of these mixtures and of the uranyl salts, we have studied the high-field conductance of these acid-salt mixtures and report these results separately.24 The cell constants of the conductivity cells were determined using 0.01 demal solutions of recrystallized potassium chloride according to the method of Jones and B r a d ~ h a w . ~ ~
X
v
Results The observed p H values are plotted us. log c in Figures 1 through 3 for the temperatures 5 , 15, 25, 35, 50, and 65". All remaining data are presented for the same temperatures. The equivalent conductances, A, are plotted us. cl'' in Figures 4 and 5 . The highfield conductance quotients are plotted in Figures 6 and 7. In Figure 8, curves are plotted for the temperatures 15 and 50" with AA/A(O) data a t two different concentrations of uranyl nitrate represented a t each temperature. In Figure 9 is shown the effect on the high-field conductance quotients a t 15 and 65' of adding m a l l amounts of nitric acid to uranyl nitrate solutions. In Figure 10 are shown theoretical calculations of the high-field conductance quotient using two different models for the hydrolysis reactions.
Discussion In each of the plots just listed there appear distinctive differences between the behavior of uranyl nitrate and perchlorate solutions, both as a function of concentration and of temperature. Within the accuracy claimed for the pH data, uranyl perchlorate solutions exhibit a straight-line dependence of pH on log concentration, while uranyl nitrate exhibits an inflection point (in the range of concentration studied) at 5 and 1.5' which disappears a t higher temThe Journal of Physical Chemistry
-PIRCHLORITI
0-WITRATE
I
I
I 1 1 1 1 1 1
I
I
I
I I I I I I
I
I
I
I 1 l l 1
Figure 2. Observed p H values for uranyl nitrate ( 0 )and perchlorate ( X ) solutions. The pH values are plotted us. log c for temperatures of 25 and 35". See Figure 1 .
41 41 40
'"
-
I
I
I
I 1 1 1 1
I
I
I
I I
IIII
I
I
I l l l l
Figure 3. Observed pH values for uranyl nitrate (0)and perchlorate ( x )solutions. The p H values are plotted us. log c for temperatures of 50 and 65". See Figure 1.
peratures. With some exceptions displayed in the graphs themselves, the pH of the perchlorate solutions (24) J. F. Spinnler and A. Patterson, Jr.. J . Phys. Chern., 69, 658 (1965). (25) G. Jones and B. C. Bradshaw, J . Am. Chem. SOC.,5 5 , 1780 (1933).
WIEK EFFECT IN URANYL IONSOLUTIONS
503
.. 5' 1.1111 15' 13111
0.2
-- ~ - - - - . - 3 5 ' 0 , 1 8 l U
lo 0 x \ -02 x 4 -0k
25' lOlkY
- 0.8
50' 0.515Y
-0d
6 5 ' 0 kIOY
I 0
1
50
100
I
J
150
200
fl[uI. K U C U
Figure 4. The equivalent conductances of solutions of uranyl nitrate are plotted us. cl" for a range of temperatures.
Figure 7. The high-field conductance quotients of solutions of uranyl perchlorate are plotted as a function of field for a range of temperatures and concentrations. All concentrations given are to be multiplied by lo-'. URANYL YlIRAT€
t
1
I
50
0
IO0
FIELD.
I
I
150
200
XVICM
Figure 8. Effect of concentration on the high-field conductance quotient of uranyl nitrate: data are given a t two concentrations a t each of two temperatures, 15 and 50'. All concentrations given are to be multiplied by low4. Figure 5. The equivalent conductances of solutions of uranyl perchlorate are plotted us. c " ~for a range of temperatures.
Ok
0.d-
o Ok51M U0,1W0,12t0
-08 -10
I 8 2 Y UNO,
X 0 5 1 8 1 U02
IHtORmCAL RtACllOM
0
II
I
so IltLD
1 1
50
I
I
I
IW
I50
200
IlttO. I Y / E U
Figure 6. The high-field conductance quotients of solutions of uranyl nitrate are plotted as a function of field for a range of temperatures and concentrations. All concentrations given The lowest curve (long dashes) are to be multiplied by is derived from the curve immediately above (short dashes) by subtracting the high-field conductance quotient of the reference electrolyte
IO0 I Y I C I
I a0
8
150
Figure 9. Effect of added nitric acid on the high-field conductance quotient of uranyl nitrate. Data are given for two temperatures, 15 and 65". All concentrations given are t o be multiplied by
is lower, suggesting that hydrolysis is more extensive (See Figures 1 , 2 , and 3.) The conductance curves for the two electrolytes
in these solutions.
Volume 6.9, S u m b e r 2
F e b r u a r y i.96.5
JOSEPH F. SPINNLER AND ANDREWPATTERSON, JR.
504
t
Id0
t--=
-
,b
I IO0 IIELO, K V / C I
I
IS0
I 200
Figure 10. Theoretical calculations of the high-field conductance quotient of uranyl nitrate using two hydrolysis models; reaction 1 refers to eq. 1 (text) and reaction 2 to eq. 2. The concentrations used are those given in Figure 6 and corrections for the behavior of the reference electrolyte are included; i e . , the curves are relative to the same reference electrolyte employed in obtaining the experimental h t a of Figure 6.
are similar, but in every case the conductance of the nitrate salt is lower than that of the perchlorate salt. There is also a noticeable difference between the temperature-concentration data in the two sets of curves, the most significant one being the appreciably lower conductance of the nitrate salt at 25'. The nitrate data at 5, 15, and 25' are crowded together in a way quite dissimilar to the perchlorate data. As a result, the equivalent conductances of the two salts are more nearly similar a t 5 and 65', and least similar a t 25". It is of interest that the equivalent conductance of the nitrate ion is slightly higher than that of perchlorate ion and that the conductances of the two anions closely parallel each other as a function of temperature, while in these data the conductance of the nitrate salt is in every case lower than that of the perchlorate sali,. This observation, coupled with the lower p H of the uranyl perchlorate salts mentioned in the preceding paragraph, leads one to conclude that the nitrate and uranyl ions are associated to a significant degree and that the extent and variation of this association with temperature appear to be the most significant a t the lower temperatures. (See Figures 4 and 5.) The high-field conductance curves are distinctly unusual. In contrast with the typical curves shown in ref. 4, both uranyl nitrate and perchlorate high-field conductance quotients first exhibit an increase of conductance a t lower temperatures and fields and a decrease in conductance a t the higher temperatures and fields. These conductance changes, which one may appropriately denote negative Wien effects, are quite real and easily reproducible within the precision The Journal of Physical Chemistry
of measurements claimed in previous paragraphs. It must also be kept in mind that the curves plotted are for the conductance changes relative to those of a reference electrolyte, in these measurements always a salt-acid mixture to minimize polarization differentials between the solutions being compared. Since the reference electrolytes always exhibit positive Wien effects, the absolute Wien effect of the electrolyte under study will be shifted by an amount equal to the absolute Wien effect of the reference electrolyte, with appropriate regard to algebraic sign. The coniputation of the Wien effects of these salt-acid mixtures is discussed in another paper. 2 4 For comparison with the relative measurements, the 65" data for uranyl nitrate are corrected for the reference electrolyte Wien effect and replotted as the lowest curve in Figure 6. This correction removes the inflection in the curve a t low fields, which results from the sharp rise in reference electrolyte Wien effect toward an approximately constant value above fields of 50 kv./cm., and moves the curve downward to produce a larger absolute negative Wien effect. (See Figures 6 and 7.) These data are taken a t what is essentially constant specific conductance so the concentrations tend toward lower values a t the higher temperatures. The curves for the perchlorate salt show a consistent downward trend, corresponding to an increasing negative Wien effect, with increasing temperature. The curves for the nitrate salt are similar, except that a t the higher fields the 25 and 35' curves cross each other and are inverted in order. To indicate the effect of concentration on the magnitude of the high-field conductance quotient, data for two different concentrations have been plotted for the temperatures 15 and 50'; the concentrations of uranyl nitrate involved differ by the ratios 1.904/1.814 a t 15' and 0.792/0.709 at 50'. In each case, the higher the concentration, the lower is the conductance quotient. The inversion in the uranyl nitrate 25 and 35' curves cannot be accounted for as an effect of concentration, since the magnitude of the effect of concentration difference is insufficient to account for the difference in high-field conductance quotients observed. I t will be recalled that the nitrate conductance curves were oddly spaced, with the greatest difference occurring at 25', so the highfield conductance results reflect also the competition between uranyl and nitrate ion association and hydrolysis of the uranyl ion, the pattern of which changes with increasing temperature. (See Figure 8.) I n itself, the observation of a smaller Wien effect in a more concentrated solution is abnormal. Curves for potassium chloride (Figure 1, ref. 4) show increased conductance quotients at increased concentrations.
WIENEFFECT IN URANYL IONSOLUTIONS
Calculations made on a presumably normal 2-1 electrolyte, calcium chloride, show the same treQd; these calculations agree satisfactorily with experimental measurements. Since uranyl nitrate solutions are significantly hydrolyzed, and in anticipation of the probable role of hydrogen ion in the Wien effect results, measurements were made in which solutions of nitric acid were added to those of uranyl nitrate and the Wien effects of the mixtures observed, a t the temperatures 15 and 65’. (See Figure 9.) It is not immediately clear how correctly to interpret these results; we shall discuss them further below in connection with the negative Wien effect phenomenon. To perform theoretical calculations of the high-field conductance quotients, it is necessary to have the limiting equivalent conductances and, ultimately, the limiting ionic conductances of the ions involved. To obtain these from the available data, we have resorted to the use of the method of Kraus and Bray28and have employed Walden’s rule2’ to obtain the values desired. Since the method of Kraus and Bray involves extrapolation to zero concentration and is an approximation technique, it is assumed that the values obtained with its use are known to be no better than 10%. The effect of such an error on the theoretical Wien effect calculations is not serious, however, as will be pointed out. The limiting ionic conductances of the well-studied ions are known to much higher degrees of precision. The data we have used for the limiting ionic conductances of ions such as H+, S O 3 - , and C104- as a function of temperature were taken from Harned and Owen,28 Robinson and Stokes,29and other Limiting equivalent conduhances of uranyl nitrate were taken from ref. 21; since no data were available for uranyl perchlorate, other than the data we have presented here which indicate essential similarities between the behavior of the two salts, the behavior of the cations (UOz+2,H+) in these solutions was assumed to be the same as in the nitrate solutions. Following the procedure outlined in ref. 24, we have then calculated a combined valence and a combined limiting conductance for the cations in the mixture. This procedure overlooks any possible mixture effects or specific interactions. It is convenient to proceed in this way since it is unnecessary to know the limiting ionic conductance of the uranyl ion itself along with the individual conductances of the hydrolysis products, all of which are uncertain or unknown. AIoreover, it is simple to determine Ao+ by subtracting from the limiting equivalent conductance of the total solution the limiting ionic conductance of the anion, while the
505
net z+ is determined from the total concentration and p H measurements. We have assumed that either of two reactions
2U02+2
+ H2O
+
Uz0,+2 2H+
or
U02+2
+ Hz0
UO2OH+
+ H+
is responsible for the hydrolysis and the observed p H of the hydrolyzed solution and computed the combined valence accordingly in the absence of an applied field. This entire procedure allows one to compute a Wien effect for a mixture of ions, but it does not take into account any kind of ionic or other interactions under the influence of field beyond those which we might term the classical Onsagera2and Onsager-Kima3 types for weak and strong electrolytes. Theoretical calculations are plotted assuming both reactions 1 and 2 as the hydrolysis equilibrium, and using conductance data for uranyl nitrate. Siiiiilar results are obtained for uranyl perchlorate, so these calculations are not reproduced here; however, the uranyl perchlorate calculations are perfectly regular, and any inversions or other unusual features found in the uranyl nitrate results are absent. I n each case the conductance quotient obtained using reaction 2 is lower than when reaction 1 is assumed for the hydrolysis equilibrium. The quotients decrease with increasing temperature. The 25 and 15’ curves are inverted in order when reaction 1 is employed; calculations were made only a t 5 , 25, 50, and 65’ for reaction 2. At 6 5 O , reaction 2 gives an apparent negative Wien effect; this is in reality due to the slower approach of the reference electrolyte Wien effect to its maximum value than that of the uranyl nitrate. (See Figure 10.) The differences between the calculated and observed values of the conductance quotient exceed those which can be attributed to experiniental error, the choice of limiting conductances, errors in extrapolations or interpolations employed, or erroneous assumptions of the equilibrium status of the solutions a t low fields. (26) C. A. Kraus and W. C. Bray, J . Am. C h m . SOC.,35, 1315 (1913). (27) P. Walden, Z . physik. Chem., 5 5 , 207, 246 (1906). (28) H. S. Harned and B. B. Owen, “The Physical Chemistry of Electrolytic Solutions,” American Chemical Society Monograph Series, 3rd Ed., Reinhold Publishing Corp., New York, N. Y., 1958. (29) R . A. Robinson and R . H . Stokes, “Electrolyte Solutions,” Butterworths Scientific Publications, London, 1959. (30) A. N . Campbell and E. Bock, Can. J . Chem., 36, 330 (1958). (31) J. Johnston, J . Am. Chem. SOC.,31, 1010 (1909). (32) L. Onsager, J . Chem. Phys., 2 , 599 (1934). (33) L. Onsager and S. K. Kim, J . Phys. Chem., 61, 198 (1957).
Volume 69,hTumber8
February 1966
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Keeping in mind as an example that the experimental conductance quotient of uranyl nitrate a t 65' is -0.80% while the calculated value is 0.12%, we may estimate the influence of the factors just noted. An error in p H determination gives rise to an error in the computed valence factor fed into the computation. Comparing the results of the calculation a t 25' using the two different hydrolysis reactions, we find that a difference in the valence factor of (1.82 - 1.67) = 0.15 unit gives rise to a change in the conductance quotient of 0.123 unit a t 200 kv./cm. and corresponds to a change in hydrogen ion concentration of 0.27 X mole/l. Since the p H can be determined a t this temperature with sufficient accuracy that the hydrogen ion concentration is known within 0.05 X lop4 mole/l. , an error in p H or in the valence factor is not sufficient to account for the discrepancy between the experimental and calculated results. Using data for the 15' run where nitric acid has been added to the uranyl nitrate solution (Figure 9), there is a decrease of 0.19 in the cationic valence, a concentration change of 0.312 X mole/l., an increase of 33 ohm-' c m 2 equiv.-' in the cationic limiting ionic conductance, and a resulting decrease of 0.164 unit in A h / A ( O ) in the theoretical computation made with these data. The corresponding experimental change in A A / h ( O ) was 0.01 unit. Yariations in the majority of the quantities used in the calculation are considerably less important than is the valence factor. It thus appears that the processes in solution under the influence of high fields are sufficiently complex that they are not accounted for by the theoretical approach which we have used. It can be concluded that both uranyl nitrate and perchlorate are nearly completely dissociated over the temperature and concentration ranges covered, for otherwise the small Wien effects a t lower temperatures and the negative Wien effects a t higher temperatures would be overshadowed by the effect of field on association reactions giving rise to large positive Wien effects. The importance of association will be noted in papers describing measurements on uranyl fluoride3. and uranyl sulfate.3b In spite of the observation just made, there is involvement between the nitrate ion and the uranyl ion or its hydrolysis products, since the low- and high-field conductances and the pH measurements are distinctively different from those on the presumably unassociated uranyl perchlorate, and since the specific conductance of the uranyl nitrate is lower than that of uranyl perchlorate in spite of the fact that the nitrate ion has a higher ionic conductance than does perchlorate ion. It also appears that the limiting conductance of the The Journal of Physical Chemistry
JOSEPH F. SPINNLER AND ANDREWPATTERSON, JR.
uranyl ion in solutions of these two compounds is abnormally high. At 65' more than half the uranyl ions are in hydrolyzed form with UOzOH+ predominating.lg Using our conductance data for uranyl nitrate at the same temperature and making a hydrolysis correction by the method of Owen and G ~ r r y ~ ~ assuming hydrolysis equilibrium (2), we find that the ionic conductance due to uranyl ion or the various possible types of uranyl complexes is approximately For comparison, a t this 320 ohm-' cm.2 equiv.-'. temperature the limiting ionic conductance of potassium ion is 135 and that of calcium ion is 240. Thus the limiting ionic conductance observed is decidedly higher than that of a typical univalent ion, an ordinary divalent ion, or a combination of the two. It is the concensus of the references cited (ref. 5-19) that monoand divalent cations are the principal types involved in any of the proposed hydrolysis schemes; although nonionic and anionic species have been proposed by Sutton, these are not to be expected in the range of concentration involved here. One is thus led to suspect some unusual conduction niechanisni a t low fields, possibly operating through an oxygen bridging mechanism as proposed by Longsworth. l 1 To explain the inusual conductance behavior and the lack of agreement with theoretical calculations a t high fields, an influence of field upon the conduction niechanism must be invoked. Two possibilities are suggested by the observations of the previous paragraphs. If some bridging process is responsible for the abnormal conductance a t low fields, orientation times for the conducting uranyl species at high fields may be long compared to the pulse lengths involved. d decrease in conductance might result from a failure of the ordinary conduction mechanism within the period of the highfield pulse, corresponding to a bulk polarization, I t must be pointed out that there was no evidence of a relaxation process operating in the range of 4-psec. time intervals, so any time-dependent conduction mechanism must either respond in times shorter than 0.5 psec., the shortest time me can conveniently observe, in which case it would not be responsible for the experimental results, or require times larger than 10 psec., which we cannot observe a t higher fields where the phenomenon becomes significant. For this reason, and since the solutions are appreciably hydrolyzed into hydrogen ions and polymeric uraniuiii species, we prefer the explanation that the applied field causes increased collisions between the fast-moving hydrogen ions and the slower uranyl complexes, reversing the (34) B. B. Owen and R. W. Gurry, J . Am. C'hem. Soc., 6 0 , 3074 (1938).
WIEN EFFECTIN URANYL IONSOLUTIONS
hydrolysis equilibria (e.g., eq. 1 or 2). I n either casechange of conduction mechanism or reversal of hydrolysis equilibria-the number of ions available for conduction decreases and the conductance decreases with increasing field. The experimental data adduced are not adequate to answer which of these processes is operating. Additional confidence would be lent to the second suggestion if a model compound could be found which showed similar abnormal high-field conductance but had more nearly normal low-field conductance behavior. I n search for such a compound we have examined a number of aquo complex compounds, including aquopentaamminecobalt (111) perchlorate, the several chromium(111) aquo complexes, and the formally similar titanyl, vanadyl, and related compounds. None has been found to behave peculiarly in any way so far as they have been studied. We have not yet studied neptunium(V1) compounds, although we have hopes of so doing. As another approach, we have made three kinds of additions of reagents to the uranyl nitrate solutions, utilizing nitric acid, potassium nitrate, and potassium hydroxide. Addition of the first and third is an obvious way to change the pH, while addition of nitrate ion is a way to examine the significance of nitrate-uranyl ion complexing on the high-field conductance results. When it became apparent that the effects of adding hydroxyl and nitrate ion were much less significant than that of adding hydrogen ion, these experiments were not carried out in as much detail as would otherwise have been appropriate. The addition of hydroxide changed the shape of the high-field conductance curves, putting a dip in the middle and raising the low- and high-field ends of the curve above that of uranyl ion alone. The only experiments were done a t 25'; it was found that stable solutions could be prepared with appreciably higher pH than that of the hydrolyzed uranyl ion, but if the increase of p H went much further than about 0.4 pH unit, the solutions were prone to precipitate when the high-field measurements were made, although they would remain stable if no ' such measurements were performed. The net effect of the addition of such small amounts of hydroxide (0.5 X M final concentration) was to lower the high-field conductance curve in the direction of a greater negative Wien effect a t the given temperature. The effect of addition of nitrate was much smaller, not much larger than the precision of measurement in the case of an addition of nitrate to yield a 0.635 X M concentration of potassium nitrate. With M , the general twice this concentration, 1.210 ?
