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UNITED STATES ATOMIC ENERGY COMMISSION
THE MEASUREMENT AND INTERPRETATION OF pH AND CONDUCTANCE
VALUES OF AQUEOUS SOLUTIONS OF URANYL SALTS
D. A. Machnnes
L. G. Longsworth Rsr
Rockefeller Institute for Medical Research
This document consists of 10 pages.
Date of Manuscript: November 24, 1942
Date Declassified: March 5, 1947
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THE MEASUREMENT AND INTERPRETATION OF pH AND CONDUCTANCE
VALUES OF AQUEOUS SOLUTIONS OF URANYL SALTS
By D. A. MacInnes and L. G. Longsworth
The hydrolysis of aqueous solutions of the uranyl salts has introduced difficulties in manipulation
and in the interpretation of many of the experimental results pertaining to such solutions. It is our
purpose to report pH measurements on solutions of UOCL UO(3NO',, UO.SO,, and UO,(C,HsO2,)
together with pH, conductance, and solubility measurements of selected solutions of UO, in aqueous
HCI. With the aid of these data it has been possible to draw certain conclusions, which are included
in this report, regarding the nature of the canons in these solutions and the strength of UO0 as abase.
Since the corrosiveness of aqueous uranyl salt solutions is due largely to the free acid resulting
from hydrolysis, the data reported here should be of value in the prediction and control of that cor-
rosion. Moreover, we have found that the pH and conductance of a given dranyl salt solution may be
utilized for the rapid analysis of such solutions on a semimicro scale.
Standard solutions of stoichiometric uranyl salts were prepared by solution of a weighed sample
of UOs* (H,0), (previously assayed for uranium by ignition to UsO.) in the proper volume of a con-
centrated standard acid solution, followed by dilution to the desired final volume.
The samples of UO, *(H20)x used in this work were prepared by two different methods. One
sample was obtained by ignition of UO,(NOs), 6H.O, first at 290C and then at 390*C, until nitric
oxide fumes were no longer evolved, after which the cake was pulverized and repeatedly extracted with
water. The oxide was then dried in the oven at 150'C, thoroughly ground, again washed with water,
dried and heated at 390*C for six hours, and then stored in shallow dishes over a saturated solution of
Mg(NOs)* 6H.O. The material soon attained a weight sufficiently constant to permit direct weighing
and was then assayed for uranium.
The insolubility of uranium peroxide has also been used in the preparation of the oxide. The
following procedure has proved practical. One liter of a 10 per cent solution of UO,(NO,)2 .6H20 was
heated from 60 to 80C and to this was added slowly, and with continuous stirring, 250 ml of 3 per cent
aqueous H,-O2. The precipitate was allowed to settle and was then collected on a Buchner funnel. The
filtrate had a pH of 0.63 and contained 10 g UO2(NO,), -6H.O, thus indicating 90 per cent precipitation
of the peroxide. The filter cake was removed, dispersed in 500 ml H20 and refiltered. This was re-
peated four Limes and the successive filtrates had the following pF values: 1.52, 2.45, 3.39, and 4.38.
The peroxide was then dried, pulverized, and converted to the oxide by ignition in shallow dishes at
280' for three hours. The product thus obtained was hydrated by suspension in water, followed by dry-
ing overnight m an oven at 110C. The resulting hydrate was again ground and proved to be sufficiently
stable in air for direct weighing.
The pH measurements were made at a room temperature of 20 to 25C with the aid of a glass
electrode. The latter was calibrated with 0.05M potassium acid phthalate as pH 4.00.
Table 1. pH values of solutions of stoichiometric uranyl salts.
equivs liter UO2,C UO,(NOS,) UOSO, UO,(C2HO,)2 r
0.001 4.05 4.05 4.17 4.76 .99t
.002 3.85 3.85 3.99 4.69 .98t
.005 3.60 3.59 3.76 4.61 .97t
.01 3.41 3.41 3.59, 4.55 .96
.02 3.22 3.22 3.40 4.48 .94
.05 2.95 2.96 3.14 4.36 .92
.1 2.76 2.7B 2.91, 4.24 .89
.2 2.58 2.58, 2.67 4.07 .87
.5 2.25 2.25 2.30, .97
1.0 1.92 1.91 1.97 1.16
The pH values of stoichlometric solutions of UOC12, UO2(NOs),, U02SO,, and UO,(C2H,02)2 are
recorded in Table 1. Since the values for the chloride and nitrate are essentially identical at a
given concentration, etiher set of data may be taken as typical of the behavior of UO, with a com-
pletely ionized acid. The data for the sulfate, and to a much greater extent for the acetate, indicate
the behavior of UO, with incompletely dissociated acids.
The significance of the values for y* in Table 1, and also in Table 2, will be considered later in
The pH and conductance values of mixtures of aqueous HCI and UOs at constant ion concentrations
of 0.1, 0.5, and l.ON, respectively, are recorded in Table 2. At a given chloride ion concentration,
say 0.1N, the solutions for which the ratio of moles of base to acir' i.e., [UO] [HCI], is less than
0.5. These were prepared by mixing aliquots of 0.1N UO2C12 and 0.1N HCI. For values of IUOs] '
[HCI] > 0.5 weighed increments of UO, were dissolved in aliquots of 0.1N UO2C4l. The change in
normality resulting from this latter procedure is quite small and has been neglected in the com-
putations of Table 1.
THE IONIZATION OF UO,
Uranium trioxide, written with the formula U(OH),, is potentially a hexavalent base; therefore,
it is of interest to ascertain the nature of the ions that can exist in aqueous solution. Some evideAce
as to the nature of these ions is furnished by the data of this report. The pH values for the 0.1N a
solutions of Table 2 are plotted as circles in Figure I and show a single point of inflection when the
molecular ratio of base to acid is 1,2, that is, [UO,] HCI] = 0.5. The simplest ionization
mechanism that can be assumed is these equilibriums
U(OH), = U(OH)* OH-
U(OH)5 = U(OH)" + OH"
U(OH)I* U(OH)+++ + OH, etc.
As will be shown, both the pH and conductance data of Table 2 indicate that no appreciable con-
centrations of U(OH)++ exist in aqueous solution. Consequently, the constant, K., for the third
equilibrium must be essentially zero and we are left with the problem as to whether or not the uranyl
ion, U(OH)'* (that is, UO,*), is formed as indicated.
At all pH values at which UO, is soluble, the hydroxyl ion concentration is negligible in com-
parison with that of the other ion species. It is more convenient to write the equilibriums I and II as
the equivalent hydrolytic reactions
UO0OH' = UO, H* (1)
UO+ *+ HO0 = UO, *OH" + H (2)
Table 2. Electrometric and conductometric titration of U03.
[Cl-] = 1.0 [Cl-] = 0.5 [Cl-"]= 0.1
IUO.] / [HCI] pH Ko pH pH K25 y*
0.000 0.25 0.3326 0.50 0.1786 1.08 0.03899 0.83
.125 .30 .2550 .57 .1387 1.19 .03107 .84
.250 .40 .1818 .64 .1012 1.36 .02341 .85
.375 .58 .1145 .91 .06672 1.64 .01606 .87
.425 .75 .0897 1.11 .05214 1.85 .01319 .88
.450 .89 .0777 1.27 .04647 2.02 .01199 .88
.475 1.14 .0665 1.54 .04019 2.26 .01052 .88
.500 1.92 .0565 2.25 .03499 2.76 .009391 .89
.525 2.45 .05485 2.64 .03411 3.11 .009104 .89t
.550 2.61 .0542 2.81 .03378 3.26 .009011 .89j
.575 3.34 .008947 .89t
.600 2.73 .0530 2.97 .03321 3.43 .008895 .89t
.625 3.49 .008852 .89t
.650 2.86 .0517 3.08 .03256 3.54 .008822 .89t
.700 2.94 .0503 3.19 .03198
.750 3.00 .0490 3.27 .03136
.800 3.06 .0476 3.33 .03071
.850 3.12 .04615 3.37 .02999
.900 3.17 3.80
As a matter of fact, these equilibrums can be made to represent the titration data satisfactorily. The
dash curve of Figure I was computed with K, = 4.6 x 10" and K2 = 6.0 x 10-'. Below the equivalence
point, it is identical with the solid curve.
The theory on which these computations are based is this. If we let x = fUOnsi], y = [U020H'],
z = [UO,], h = [Hi], and a = [CI-], in which the brackets indicate gram ion or gram mole concentration,
then K,y = hz and K1x = hy. Electrical neutrality gives the relation, 2x + y + h = a, and the total
concentration of base, b, is x = y z = b. Elimination of x, y, and z between these four relations gives
b aK,K, + (aK K,K.)h 4 (a K,)h' hs
2h2 + hK,
an expression explicit in b, and hence, suitable for computation.
The simple mechanism represented by the reactions in equations I and 2 does not, however,
explain satisfactorily the hydrolysis data of stoichiometric uranyl chloride of Table 1. This is shown
in Figure 2, i which the pH is plotted against the logarithm of the salt concentration. In this plot,
the experimental values are indicated by the circles, while the broken line represents values com-
puted with the aid of the theory and constants suggested. Although these constants are apparent
constants and strictly valid only at a concentration of 0.IN, the discrepancy between the observed
and computed pH values shown in Figure 2 is probably too great to be attributed to a variation of K,
and K, with ionic strength.
Moreover, the simple theory outlined leads to improbable values for the concentrations of the
various ion species in solutions for which [uO,] .' [HCI] is greater than 0.5. Thus for [UO,]/ [HCI]
= 0.75, (UO,] = 0.04454, [UO, OH] = 0.01067 and [UO, undissociated] = 0.01967. It is difficult to
reconcile the value of 0.01967 for the concentration of dissolved, undissociated UOs with the very low
solubility of this material in pure water.
The hydrolysis reaction that accounts best for the great majority of our results is
2UO" + HO = UOSUO0 *. 2H* (3)
The full curves of Figures 1 and 2 were computed with the constant, K', for the reaction equal to
1.35 x 10 The theory is this. Letting w = [IO UO], then, K' x2 = h'w, 2x + 2w + h = a
(electrical neutrality) and x + 2w = b (total base). Elimination of x and w between these relations
leads to the quadratic in b, 2(a b h)2 K' = (2b a 4 h)h2. For the special case that a = 2b, i.e.,
stoichiometric uranyl chloride, this reduces to 2(b h)2 K' = h'. Since h is small in comparison with
b, 2b' K' = h' and 2 log b + log 2K' = 3 log h; -log h = pH = -2/3 log b 1/3 log 2K'. Hence, if the
pH values of the stoichiometric UO2C1, solutions are plotted against the logarithm of the concentration,
as in Figure 2, a straight line of slope -2/3 should result. As explained, this is approximately the
As may be seen in Figure 2, the observed pH values of stoicluometric uranyl chloride agree
satisfactorily with the computed values for all concentrations below IN. In Figure 1, the observed
and computed pH values are in agreement for all values of [UOJ] [HCI] less than 0.65. Above this,
the computed values increase much more rapidly than the observed pH values.
The equation 3 does not, of course, provide for a further reaction of the UO, UO* ion with
water, and hence, cannot account for the relatively low pH of solutions of basic uranyl chloride shown
by the difference between the solid and broken lines of Figure 1. Moreover, the reaction UO,-UO:'+ +
HO =- 2UO, + 2H+ is improbable since this leads to appreciable concentrations of dissolved, un-
dissociated UO, in solutions for which [UOn] / [HCI] >. 1. This step does not introduce this difficulty
UO1 + UOs. UO + H20 = (UO,) -UOU + 2H+. The (UOs)2 "UO+ complex will be recognized as the
ion obtained by the addition of two protons to uranosic oxide, U,08, and corresponds to the primary
MDDC 911 5
[U0,1/[HCI] HCl]- 0.1 Nomal
LO I I I I I
0 (11 0.2 03 0.4 0.5 0. 0.7 0.8 0.9 10
dissbociation of uranium oxide trimer (UO,), 2H* = (UOs)2 UO04 + H20. Cryoscopic measurements
on solutions of basic uranyl chloride should tell whether the principle ion in these solutions is UO, *OH4
or UO, UOI .
In the preceding computations of the various hydrolysis and titration curves, it has been necessary
to determine values of the hydrogen ion concentration, [H*], from the pH measurements, pH = -log
I[H] in which '* is a coefficient that includes both ion interaction effects and a small but unknown
liquid junction potential. The values of used in the computations are recorded in Tables 1 and 2 and
were obtained as follows. Taking pH = -log [HI] as a first.approximation, values of [H*] were computed
for the various solutions of UO,. These solutions were then duplicated with the unhydrolyzed Ba ion
substituting for the UO0* ion and their pH values determined and denoted pH'. Since the hydrogen ion
concentrations, [H]*, of these mixtures of HCI and BaCl2 are known accurately, / = antilog (-pH*)/
THE SOLUBILITY OF UO,
We have determined the pH and composition of a great many solutions of aqueous HCI that have
been shaken for several days with excess solid UO,. The results for a typical sample of oxide are
reported in Table 3 and plotted in Figure 3. These data probably do not indicate precisely the sol-
ubility of the oxide, however, since the attainment of equilibrium appears to be very slow and different
preparations behave somewhat differently. In no case has our saturating phase been homogeneous, nor
did it contain an integral number of moles of water of hydration. The water content of the hydrated
oxide that is the stable phase at room temperature has not been definitely established, but our results
thus far indicate that it is the monohydrate.
Experiments with solutions supersaturated with respect to UOs indicate, however, that the
solubilities given in Table 3 are low by not more than a few per cent. The basic chloride, UO,] ,.' [HCl]
= 1, is a crystalline solid that dissolves, without decomposition, in either a small or a large volume
of water. At intermediate concentrations, however, the solution of the basic chloride is accompanied
by a slow precipitation of the oxide and the concentration interval over which this hydrolytic decompo-
sition occurs agrees with that portion of the curve in Figure 3, from 0.016N to 2N, in which [UO,] / [HCI]
is less than unity. Supersaturation with respect to the oxide may also be achieved by dilution, with
water, of a fairly strong solution of HCI saturated with the oxide, and by the addition of NaOH to a
solution of the normal chloride. On standing, such solutions usually deposit UO,; analysis of the super-
natant has generally given results a few per cent above those of Table 3. In no instance do we have
any evidence that UOs was dispersed in a colloidal state.
Table 3. pH values and densities of solutions of HCI saturated (approximate) with UO,.
[HCI] [UOs] [UO,],' [HCI] pH d4
2.785 2.832 1.0169 2.51 1.7754
1.013 .958 .946 3.20 1.2640
.2979 .2718 .9124 3.56 1.0733
.09909 .0900 .908 3.80 1.0222
.03140 .02986 .951 4.07 1.0056
.01082 .01138 1.052 4.28 1.0004
.00298 .00343 1.15 4.53
.00108 .00150 1.39 4.775
MDDC 911 [ 7
EVIDENCE FROM CONDUCTANCE AND TRANSFERENCE DATA
CONCERNING THE IONIZATION OF UO,
The assumption was tacitly made in computing the curves of Figure 1 that no further dissociation
of the uranyl ion to form a trivalent ion of the type U(OH)4'" occurred. The close agreement between
the observed and computed values of pH for values of [UO,] ,' [HCl] between 0 and 0.5 justifies this
assumption. The conductance data are, however, somewhat more convincing on this point than the pH
measurements since the latter are not sufficiently sensitive in the strongly acid solutions in which
U(OH)" ions might be expected to exist.
The observed conductances of mixtures of HCI and UOCI, are recorded in the fourth column of
Table 4. Values computed with the aid of the assumption that the ion conductances in these mixtures
are additive. These are given in the third column. The differences, column 5, are small but slightly
greater than the values, column 6, that have been observed for similar mixtures of HCI and CaCI,.
These differences are explained qualitatively by the Onsager and Fuoss theory as due to the braking
action of the slow UOV* ion on the fast H* ion and leave no room for the rather large effects to be
expected in the replacement of H" ion by U(OH);+.
Table 4. Additivity of ion conductances in mixtures of HCI and UO,CI,.
1 2 3 4 5 6
CHCI CUOC1 Acomp. .Aobsd. A A(HCl-CaCl,)
0.1 0.000 389.9
.075 .025 314.6 310.7 3.9 2.8
.050 .050 239.3 234.1 5.2 4.2
.025 .075 164.0 160.6 3.4 3.3
.000 .100 88.7"
*Corrected for hydrolysis.
Since the acid present in the mixtures listed in Table 4 suppresses the hydrolysis of UOC1l, it
was necessary to use as the conductance of that material a value corrected. The equivalent oon-
ductance, A, of UOCL1 may be written
h 2K y
A = a Ah 4 x + y (4)
if the hydrolysis corresponds to reaction (2). Here x, y, h, and a represent, as before, gram ion
concentrations of UO% UOaOH+, and Cl- ions respectively. The equivalent ion conductances, A,
are proportional to the mobility of the corresponding ion. The mobility of UO, .OH', or UOs.UOr*,
is shown later in this report to be 18.5. We may also safely assume, as will be shown,that h and
Aa have the values, 60.86 and 305.2, respectively, as in the corresponding mixtures of CaCL, and HCI.
From the pH and )* values for 0.1N UOCI, given in Table 1, [H*] = 0.00195. Hence,A UO++ = 27.82 and
the equivalent conductance of 0.1N UO2C12, after correction for hydrolysis, is 88.68. If the UO,'UOr
ion is the hydrolysis product, the last term of equation 4 should be replaced by 2w/ a yw. However,
since the equivalent weight of uranium is the same in the UO, .OH* ion as in UO,. UOr", neither
transference nor conductance measurements can distinguish between the two.
The foregoing ion conductances for 0.IN UO2C 2 correspond to the following transference
numbers: TUO, = 0.289, TCl- = 0.648, and TH+ = 0.063. The first two figures are in agreement
with values determined experimentally by Kraus and his associates. These values are given in
the second line of Table 5, which also includes all of the transference data given in their report for
From the data of Table 2, the equivalent conductances of the solutions listed in Table 5 have been
interpolated and are given in the fourth column of that table. The chloride ton conductances,X Q
= TCI- A are given in column 5. The constancy of A Cl- with increasing values of [UO,1 / 1HCI` above
0.5 affords convincing evidence that the chloride ion is nut involved in any of the complexes in these
The values of TUO, given in Table 5 are based on an equivalent weight for uranium equal to half
of its atomic weight, and hence, represent twice the number of gram atoms of U that are transferred
per Faraday. The actual transference numbers fur the uranium ions, x = [U044] and y = [UO2OH+],
T, 2x (5)
T y y (6)
in which a is again the chloride ion concentration. Since the equivalent weight of U in y is twice its
value in x,
TUO, Tx + 2Ty (7)
Eliminating y \y between equations 4, 5, 6, and 7 and solving for x
a A 12 TUO) 2a \a 2h \h
Thus for the solution, Table 5, in which [UO,] [HCl] = 0.6565, .\ = 88.1, TUO = 0.372, and h = 0.00032.
Therefore, x = IUOa"] = 0.0349. Moreover, solution of the relations for v\ y and setting y = 0.)6565 -
0.03491 = 0.03074 gives\ y = 18.5.
A 0.IN solution for which jUO,] ,, [HCI] = 0.6565 contains, it will be recalled, 0.01565 moles of
dissolved UO, per liter of stoichiometric uranyl chloride, per 0.05 gram ions of UO+. Hence, if the
reaction, on dissolving UO, in aqueous UOCL,, is simply UO1 + UOs + HO = 2UOO'OH and goes
essentially to completion, that is, no dissolved, undissociated UO, is present in solution, the con-
centration of UOf* in the solution would be 0.05 0.01565 = 0.0344. The close agreement between
this value and that, 0.0349, obtained from a combination of conductance and transference data indicate
that the latter are consistent with the ionization mechanism suggested in this report.
If the reaction is
UO + UO, = UO UO'.
then [UOs UO,'*] = 0.0349 2 and k UOUO UO = OH+ = 18.5. As mentioned previously, conductance
and transference data cannot distinguish between these Iwo possibilities.
In connection with their transference measurements, Kraus and his associates mention that
although [UO,] / [HCl] remained constant in the middle compartment during electrolysis, this ratio
always increased in the anode compartment and decreased in the cathode compartment. As the con-
siderations will show, this is consistent with the values of k UO+ and A UO2 *OH* already derived.
Table 5. Ion conductances of solutions of UO in 0.1N UO,CI,.
1 2 3 4 5
[UO,] / [HCI] TUO, TCI A TCI A
0.490 0.286 0.627 98.2 61.6
.500 .298 .646 93.9 60.7
.503 .298 .654 92.9 60.7
.5495 .297 .679 90.1 61.2
.6565 .372 .692 88.1 61.0
.804 .429 .694
.977t .471 .655
T Precipitation of UO, occurred in this solution during
If we consider a cathode compartment of unit volume in which the initial value of [UO,] / [HCl] is
b/a and pass 1 Faraday of electricity, 1/2 Tx + Ty gram atoms of U will enter this volume, Ta
equivalents of Cl" will leave and I equivalent of Cl" will be formed at the Ag, AgCl cathode. The
increase in b, Ab, is 1/ 2 Tx + Ty and that in a, Aa, is I Ta. The final value of [UO,] / [HCl] Is thus,
(b + Ab) / (a + Aa) and this equals b/a only if Ab 'Aa = b/a. With the relations 2x + y + h = a and
x + y = b and equations (5) and (6)
Ab x- x yAy
Aa 2xAX y Ay+h^h
For solutions in which h is small, the term hAh may be neglected andAb/Aa = b/a = (x + y) / (2x + y)
only if Ay = X which, as we have seen, is not the case. With A > y, Ab/Aa < b/a and the value of
the ratio [UOJ]/ [HCI] decreases in the cathode compartment as electrolysis proceeds.
In this document we have reported pH values of aqueous solutions of UOCI,, UO(NOa),,
UO(C,HsO,),, and UOSO4, and pH and conductance measurements on solutions of UO0 in aqueous HCI.
Approximate values for the solubility of UO, in aqueous HCl are also included. Although only slightly
soluble in water, UO, dissolves in aqueous HCI to the extent of about 1 mole of base to 1 mole of acid.
Both electrometric and conductometric titrations of UOs with HCI yield, however, a sharp end-
point only at 2 moles of acid per mole of base, corresponding to the divalent uranyl ion UO++.
The additional solubility of UO, in aqueous UOCI may be due to either of these reactions:
UO,' + UO, HE0 = 2UO,OH*
UOr + UO, = UO, UO.
Conductance and transference measurements cannot distinguish between these two mechanisms.
Measurements for pH appear to favor the second process.
UNIVERSITY OF FLORIDA
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