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UNITED STATES ATOMIC ENERGY COMMISSION
A COMPARISON OF ANALYTICAL METHODS FOR THE DETERMINATION OF URANIUM
D. A. Maclnnes
L. G. Longsworth
Rockefeller Institute for Medical Research
This document consists of 5 pages.
Date of Manuscript: November 24, 1942
Date Declassified: March 4, 1947
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A COMPARISON OF ANALYTICAL METHODS FOR THE DETERMINATION OF URANIUM
By D. A. MacInnes and L. G. Longsworth
In the analysis for uranium, conversion to the oxide UsO, is frequently used. The conditions under
which this oxide exists are, however, still in some doubt. We have attempted to develop a precise in-
dependent analytical method by which the composition of the UsO, oxide might be checked after varying
ignition procedures. The most promising independent method is the standard oxidimetric titration in
which a uranyl solution is reduced somewhat below the quadrivalent stage in the Jones reductor and then
reoxidized with a standard oxidyzing agent, the titer between the two end-points, U+3-U+4 and U+4 -U+*
representing the uranium present. Successful results have been reported' for the adaptation of the method
of differential potentiometric titration to this system. Unfortunately, as will be discussed later in this
report, the technique fails when the dilute solutions and small titrating increments required for accurate
results are used. For our comparisons, we were obliged to use the standard volumetric procedure,2 in
which the over-reduced solution is oxidized to the first end-point by air, before the subsequent titration
with potassium permanganate or ceric sulfate. Samples of a preparation of hydrated UO% were analyzed
in this way; other portions were ignited in air to U,O, for prolonged periods at several temperatures
and finally reduced with hydrogen to UO2.
Ignition to U,06
Four 0.4 g samples of UO,.xHzO (prepared from U04) were weighed into small platinum crucibles
(0.9 g) on the microbalance, and transferred to an electrically heated furnace where the temperature
was gradually raised to 700 and held at that point for 7 hours. After removal from the furnace, they
were cooled in a desiccator for several hours and reweighed. They were then reheated three times at
the same temperature. Subsequent series of hearings and weighing were made at 8200 and 9300 and
finally again at 700. The average of the weights of the four portions of resulting lower oxide calculated
as the per cent of the initial weight of UO,-xH.0, and the average deviation of the four from the mean are
recorded with the corresponding temperature and length of heating in the first 11 lines of Table 1.
Following these ignitions, the samples were stored in the atmosphere above a saturated solution of
Mg(NO, -6H2O (which has a vapor pressure of water corresponding to an average humidity of about 50%)
to determine whether the uptake of water was sufficiently rapid to have taken place to a weighable extent
in the few minutes that the samples were exposed to the laboratory air during weighing. The absorption
of water under these conditions was evidently slight as is shown by a comparison of the weights in lines
11, 12 and 13 of Table 1.
If the weights recorded in Table I are plotted against the corresponding cumulative hours of heating
and the seemingly anomalous value of line 6 is ignored, it is seen that the oxide tends to approach a con-
stant weight at a given temperature. Furthermore, the loss or uptake of oxygen is apparently reversible,
as is shown by the asymptotic approach to constant weight at 700 from both sides. The true constant
weight is presumably somewhere between the value of line 4 and of line 11. In the results reported, the
value of line 4 is assumed to be correct and cannot differ by more than 0.013% from the true value. For
the purposes of comparison, this weight is arbitrarily assumed to represent stoichiometric U3, and the
per cent UO. in the original sample of UO3-xH20 calculated, using this assumption, is:
93.314 x 280.= 95.09% UO,
Reduction to UO%
The four crucibles, containing their portions of oxide, were transferred to a quartz tube heated by
a combustion furnace. Dry, oxygen-free hydrogen was passed in and the temperature raised to 900
and held at that point for I hour. The flow of hydrogen was maintained during the subsequent cooling.
When room temperature was reached, the samples were removed individually and weighed, the final
weight on the microbalance being taken 2 to 3 minutes after a sample was first exposed to air. The
weights before and after reduction are given in lines 13 and 14 of Table 1.
Table 1. Ignition of UO-xHO to U3O, and the reduction to UO,.
Line Treatment Oxide Hours Total Percentage of Avg deviation
treat- hours original wt (4 samples)
1 Ignition at '"U1O" 7 7 93.345 .003
2 700 21 28 93.329 .005
3 to 46 74 93.314 .004
4 49 123 93.314 .002
5 Ignition at "UIO 18 18 93.301 .003
6 820 '" 59 77 93.248 (?) .005
7 to 48 125 93.267 .002
8 Ignition at "UO"08 19 19 93.245 .003
9 930 44 63 93.238 .003
10 Ignition at "U3O6" 52 52 93.284 .005
11 700 92 144 93.288 .004
12 Atm of "UOs" 5 5 93.293 .005
13 15 20 93.294 .005
14 Reduction to UO2 1 1 89.798 .006
15 Atm of UO.xHO 16 16 89.814 .002
As with the U,O,, the samples of UO, were stored in the Mg(NOQ).6HO atmosphere to determine the
rate of water absorption. Comparison of lines 14 and 15 shows that the uptake of water was slight.
The reduction of uranium oxides to UO, has been criticized as an analytical method.2 However,
Biltz and Mullers found that UOz was produced by the reduction of UsO. with hydrogen very nearly in
the theoretical ratio. In the reduction described, we have followed their procedure. If line 14 is assumed
to represent pure UOa the per cent UO, in the original sample of UO,.xHO is:
89.798 x 2!6 -= 95.12% UO,
Titration with ceric sulfate
Ceric sulfate solution was standardized against well-dried Bureau of Standards sodium oxalate, by
adding an excess of the oxidizing agent from a weight buret to a sulfuric acid solution of the oxalate,
heating, cooling, and titrating the excess ceric sulfate with 0.01N ferrous ammonium sulfate to the
ferrous-o-phenanthroline end-point. Four determinations gave a weight normality (equivalents per 1000 g
solution in air) of 0.10506 with an average deviation of 0.013%
0.5 g samples of the same UOs.xH2O taken for the ignitions were dissolved in 50 ml dilute H2SO, (5:95)
and reduced in the Jones reductor. Air was bubbled through the resulting olive-green solution until the
color changed to clear blue-green. Five drops of 0.005N ferrous-o-phenanthroline indicator were added,
followed by the ceric sulfate in slight excess from a weight buret. The excess was titrated to the pink end-
point with 0.01N ferrous ammonium sulfate. Four blank determinations (indicator correction plus reagent
correction) were made, averaging 0.022 g of ceric sulfate, or, about 0.07% of the net titer.
The average of five uranium titers was 63.328 g ceric sulfate per gram UO.xHRO with an average
deviation of 0.017%. The per cent UO, in the UO3.xH.O is then:
63.328 x 0.10506 x 0.14304 x 100 = 95.17% UO,
A comparison of the final results of the ignitions, reduction, and titration are recorded in Table 2.
Table 2. Per cent UO, in UO,.xHO by three independent methods.
Ignition UpO,, const. wt 7000 95.09%
Reduction to UO, 95.12%
Titration U4 -U+e 95.17%
It will be seen, therefore, that the oxide assigned to formula U3,6 has that composition within the
experimental error of the analytical methods when brought to constant weight at 700.* This is shown
by an independent titration method and by direct reduction of the oxide to UGQ.
Application of the titration to small amounts of uranium
If the titration procedure described above is adapted to 5 mg of U03 the results generally agree
closely but fail to represent the exact quantity of uranium.
In one experiment, portions of a standard uranyl sulfate solution, each equivalent to 5 mg UO, were
dispensed from a weight buret into a micro Jones reductor, and the resulting 20 ml of over-reduced so-
lution aerated briefly to the usual color change. Standard .01N permanganate was added in slight excess
from a weight buret, and the excess determined by differential potentiometric titration with .001N ferrous
sulfate. Blank corrections were made. Four determinations yielded the following values for the % UO, in
the original UOaxH,0: 93.93, 93.85, 93.93, and 93.41(?)%. The value obtained by ignition to U30, was
Locating the U+3 U+4 end-point by differential titration
A number of investigators have demonstrated by direct potentiometric titration the sudden increase
in potential occurring when U+' is completely oxidized to U+4. In attempting to determine the precise
location of this end-point by the method of differential titration, we have come upon unexpected diffi-
culties. If a 0.1N solution of oxidizing reagent is used with a considerable quantity of over-reduced
uranium the end-point may be found. However, when the titrating increments are decreased by using
a more dilute reagent in order to obtain greater accuracy the drifts in potential characteristic of the
U+.-U+4 system become more and more dominant until they completely obscure the end-point.
If an over-reduced solution is taken in the usual apparatus for differential titration,4 using platinum
electrodes and an inert gas, the following typical behavior is noted:
The usual method of igniting at the higher temperature of a Meeker burner, but for the much
shorter.period of 5 minutes, gave 95.06% UOa as the composition of the hydrated oxide.
a) Erratic potential differences appear, increasing as the rate of stirring is increased.
b) When all circulation is stopped, the electrodes come to the same potential.
c) If one electrode is isolated and the solution about the other stirred, a large increasing potential
difference appears which persists when stirring is stopped.
Therefore, if a drop of titrating reagent is added to the bulk of the solution with one electrode iso-
lated and both electrodes at the same potential, as after (b), the resulting natural difference in potential
will be more or less covered up by that caused by necessary stirring (c).
In our investigations, we have taken precautions to exclude all traces of oxygen from the reduced
solution and from the titrating reagent. Nitrogen was passed through two electric comoustion furnaces
containing reduced copper, heated to about 4500C,and into the titrating vessel through all glass connections
and shown to be oxygen-free by the sensitive phosphorous streamer test. If a persistent trace of oxygen
was the disturbing impurity, its effect should have been decreased by greatly increasing the acidity of
the solution. Adding acid, however, had no apparent stabilizing action. After extensive tests with several
forms of apparatus, we were finally led to conclude that the difficulty was not caused by impure gas but
was due to the inherent instability of the over-reduced system.
It is noteworthy, that as soon as the first end-point is passed, the erratic behavior (a) suddenly
completely disappears and the electrodes come together, whether or not stirring is continued. The U"3
ion therefore appears to be the cause of the instability. Solutions of U's salts are known to liberate
hydrogen readily; crystalline UH(SO4),, for example, on contact with water, produces a strong evolution
of hydrogen. This tendency of an over-reduced uranium solution to evolve hydrogen is probably cat-
alyzed at the electrode surface. The potential difference between the electrodes would then depend on
the relative rates at which the hydrogen liberated is swept away from the electrodes.
With a view to finding a metal with a favorable overvoltage under the conditions of titration, we
have tested gold, silver, tungsten, and mercury as electrode materials in addition to the platinum. All
these metals behaved alike, except the mercury which was rather insensitive to all effects.
Observations on the reduction of uranium by silver
Using a micro-silver reductor with small amounts (6 mg) of uranium, we discovered a decided
dependence of the extent of reduction upon the temperature and rate at which the uranyl solution was
passed over the silver. This reductor was completely jacketed so that the temperature could be con-
trolled accurately. Since it contained only a small amount of silver, the solution to be reduced was
passed through very slowly. At a given low rate of passage the extent of reduction was found to be a
function of the temperature: at 50* reduction to the uranous stage was 55% of quantitative, at 80 it was
105 to 110%, indicating over-reduction. Blank corrections were small and uniform. The reductor gave
precise results with 5 mg of iron.
The uranium content of a preparation of hydrated UO, was determined by three independent proce-
dures: prolonged successive ignitions in air at 700, 820*, and 930 to lower oxides of the approximate
composition U300, further reduction of the same samples to UO, by dry hydrogen at 900, and titration
with ceric sulfate of the U(SO4,) obtained by passing a sulfuric acid solution of the UO3 through a Jones
reductor and aerating the somewhat over-reduced solution. The results of these three methods showed
an extreme difference of 0.08%.
MDDC 910 [ 5
A modified titration procedure was adapted to 5 mg of UO,, in which the over-reduced uranium
sulfate solution was aerated to the uranous stage and oxidized to the uranyl condition by excess per-
manganate, the excess being determined by differential potentiometric titration with ferrous sulfate.
The results generally agreed closely but failed to give the true quantity of uranium.
Extensive application of the differential titration method to over-reduced solutions failed to locate
the precise end-point of the UI3 -U4 change. The difficulty is believed to reside with the excessively
strong reducing power of the U+s ion.
The reduction of uranyl solutions with silver has been generally assumed to proceed quantilatively
to the uranous stage. However, some over-reduction appears to be possible.
1. Furman and Schoonover, J. Am. Chem. Soc. 53:2570 (1931).
2. Lundell and Knowles, J. Am. Chem. Soc. 47:2637 (1925).
3. Blitz and Muller, Z. anorg. Chem. 163:260 (1927).
4. MacInnes and Cowperthwaite, J. Am. Chem. Soc. 53:555 (1931).
5. Gmelin Handbuch der anorg. Ch.,55:146 (1936).
UNIVERSITY OF FLORIDA
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