The effect of dissolved air on the reduction of tracer level plutonium-IV by uranium-IV

The effect of dissolved air on the reduction of tracer level plutonium-IV by uranium-IV

J. Inorg. Nucl. Chem.. 1960, Vol. 13, pp. 323 to 333. PergamonPressLtd. Printed in Northern Ireland NOTES The effect of dissolved air on the reduct...

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J. Inorg. Nucl. Chem.. 1960, Vol. 13, pp. 323 to 333. PergamonPressLtd. Printed in Northern

Ireland

NOTES

The effect of dissolved air on the reduction of tracer level plutonium-IV by uranium-IV (Received 26 October 1959) RYDBERGtll has shown that 0.05 M solutions of uranous salts in 0.7 M HNOa will remove 95 ± 3 % of tracer level plutonium from a solution of hexavalent and tetravalent plutonium in hexone, on shaking for 5 min in the absence of dissolved oxygen. In the presence of air, only 70 4- 10% of the plutonium is removed. This note shows that dissolved oxygen in the presence of excess uranium-IV causes a partial re-oxidation of plutonium-III to plutonium-IV, and offers an explanation.

Experimental Materials 10-s M Pu(IV) in 10 M HNOa was prepared from a concentrated stock and was shown by TTA assay to be at least 98"5 % tetravalent. Uranous nitrate was prepared by electrolysis at a mercury cathode, in the presence of sulphamic acid to suppress nitrite. Provided that nitrite is suppressed and that nitric acid is not less than 1 M, uranous nitrate solutions are much more stable than reported by RYDBERG,~1)e.g. 2 × 10 - z M U(IV) in 3 M HNO3, 0"1 M sulphamic acid without removal of dissolved air lost only 14% of the U(IV) during 5 days storage in a clear glass flask in diffuse daylight.

Analysis Residual Pu(IV) in reaction mixtures (3 M HNOa) was extracted into an equal volume of 20% tributyl phosphate in odourless kerosene, centrifuged, and an aliquot was mounted for alpha counting. The results were corrected for the small extraction of Pu(IIl). ") Pu(IV) in solutions of lower acidity was separated by TTA extractionJ S)

Results Rate of reduction of Pu(IV) in the presence of air The results in Tables 1 and 2 were obtained in 3 M HNO3, 0"I M sulphamic acid at room temperature. TABLE

Initial concentrations (M) Pu(lV)

U(IV)

i0-;

10-~

10 -5 10-*

IO-S

i0-~

Time for 50% reduction (min)

25-30 3-5 Less than I

(1~ j. RYDBERG,J. lnorg. Nucl. Chem. 5, 79 (1957). (e) G. F. BEST, H. A. C. MCKAY and P. R. WOODGAT£,J. Inorg. Nucl. Chem. 4, 315 (1957). (s) j. G. CUNtNGHAMEand G. L. MILES,J. Inorg. Nucl. Chem. 3, 54 (1956). 323

324

Notes

The half-lives reported in Table 1 are an order of magnitude less than those calculated for a perchloric acid solution, probably due to catalysis of the reaction by sulphamic acid or sulphurie acid. c''

Residual levels of Pu(IV) in the presence and absence of air The results in Table 2 were obtained after 2 hr reaction. Similar residual levels persisted on standing overnight. The higher amounts of U(IV) gave green solutions which were obviously stable over the period of the reaction. Some results with ferrous sulphamate are included for comparison. TABLE 2 Initial concentrations (M) Pu(IV)

U(IV)

lO-: 10-;

lO-S

I0-;

10-2 10-2

Residual PuflV)(%)

Fe(lI)

10-4 10-* 10-* 10-~

In air

m

m

lO-S

2 × 10-2 2 × 10-* 2 × 10-s

10-s

2 × 10 -2

I0-;

10-r

Under N~

15 7.5 7 1.5 3 3 1.5 1.2 0.8

m

1 I

In the presence of dissolved air, the reduction with ferrous sulphamate is more complete than with an equivalent concentration of U(IV), but the latter reagent is practically quantitative in the absence of air.

Re-oxidation of Pu(lll) by air in the presence of U(IV) An air-free solution of 10-* M U(IV), 10-' M Pu in 3 M HNO3, 0.02 M sulphamic acid was stored under nitrogen for three days. Excess U(IV) was still present (green colour) and analysis under nitrogen showed only 0-5 % Pu(IV). On bubbling air through the solution for 10 min it remained green but showed 4 % re-oxidation to Pu(IV). In a further experiment, 10-5 M Pu(IV) was reduced under nitrogen (99% complete) and then aliquots were diluted with various concentrations of nitric acid, containing dissolved air. The dilutions were l0 -e M Pu, 10-2 M U(IV), 2 ><. 10-2 M sulphamic acid. The extent of re-oxidation to PuflV) after 14 hr is shown (Table 3) to increase at lower acidities. TABLE 3

HNO3 (M) Re-oxidation (%)

0'3 36

0"9 14

2"1 10

3"0 5

Application of counter-current solvent extraction to the reduction of Pu(IV) and (VI) by U(IV) The residual Pu(IV) which escapes reduction by U(IV) in the presence of air may be extracted into 20% tributylphosphate and subsequently removed by stripping with a fresh portion of U(IV). In a laboratory mixer-settler with five extraction stages and five stripping stages (strip solution = 10-* M U(IV) in 3 M HNO~), >99"5 ~o reduction was achieved of 10-8 M Pu(IV) or Pu(VI) in 0.25 M uranyl nitrate, 2 M HNOs. The plutonium (mixed with some U(IV) was recovered quantitatively in the aqueous raffinate, and the solvent product contained uranium, free from plutonium. No precautions were iaken to exclude air. ~o T. W. N£WTON,J. Phys. Chem. 62, 943 (1958); 63, 1493 (1959).

Notes

325

Discussion

The re-oxidation of Pu(lIl) in 3 M HNO,, 0-1 M sulphamic acid at room temperature requires dissolved oxygen and U(IV), and under certain conditions the yield of Pu(IV) actually increases w h e n more U(IV) is added Gable 2, results for 10-* M plutonium). It then appears that U(IV) has a dual role. It rapidly reduces Pu(IV) to Pu(III) in 3 M HNOe (and also Pu(VI) to Pu(III)---see the section on counter-current experiments) but it also acts as a catalyst in promoting the re-oxidation of Pu(III) by dissolved oxygen, the latter reaction occurring more readily at lower acidity. The active agent in the re-oxidation must be an intermediate product in the slow reaction between U(IV) and oxygen, studied by HAt.PERN and SMITH.Ce~ They have shown that this proceeds by a chain mechanism, involving UO, + and HO~ as chain carriers. UO~ + would not oxidise Pu(lll), according to the known redox potentials, but the peroxide radical HOs is a likely agent. The rate of the reaction between U(IV) and O, is inversely proportional to the acidity: the higher stationary state concentration of HOI at lower acidities would explain the results in Table 3. E. N. JENKINS A.E.R.E., Harwell Didcot, Berks A c k n o w l e d g e m e n t - - T h e author wishes to thank Mr. R. J. W. ST~EETONfor his assistance, particularly with the mixer-settler experiments.

~=~J. HALP~RNand J. G. SMSTH,Canad. J. Chem. 34, 1419 (1956).

C o m l m a n d s in the s y s t e m T i O 2 - C r 2 0 3 - F e 2 0 3 (Received 28 October 1959)

RECENTLYa number of investigations into the system TiO2-CrmOs have been reported. ANDERSON, et ai. c1~ found the compounds Tia-sCrsOsn-i with n = 6, 7, 8 and 9, which are isomorphous with

members of the homologous series TiaOsn-1 and VnOsa_l, and also a compound TisCraO7 which has a different structure. The latter is in agreement with earlier investigations by HAMEt.IN.~sj A compound TiCrIO5 which could be expected to be isomorphous with TiFe2Os and TiAIIO5 (pseudobrookite structure) was not found. ASaRrNK et ai. 's~ found small crystals with the composition TiCr,O6, after melting TiOs and CrsOs together in an electric furnace. The structure, however, was of the VaOe type, which is essentially different from the pseudobrookite structure. HAMEL~~4~reported the presence of an unknown compound W in the system TiOi-Al2Os-CrsOs; the composition of W is in the neighbourhood of TiCrsOi. HAMEUN'Ssamples, however, do not seem to have reached a state of equilibrium and therefore the composition of this compound is not exactly known. This paper describes some investigations into the system TiO,-CrtO,-Fe,Os. It is to be expected that samples of this system attain their equilibrium-states more easily than samples of the TiO2-AIIOs--Cr~O3 system. The results are given in Fig. 1. All the samples studied reached their equilibrium states by firing at 1300~C in O1. In the twophase system TiOs--Cr2OI, the Andersson structures Ti,Cr.Ou, TisCrtOla , TieCr2015, Ti7Cr2017 and a homogeneous area were found. On the other hand prolonged heating of Ti,CriOn and TisCr,O~s resulted in disproportionation into the stable compounds TieCrsO~5 and TiiCriOT. TiCr~Os was not found. From a study of the X-ray diagrams after a few per cent of FesOa were added to the Andersson structures, it followed that Ti,Cr,Ou and TisCrsO~, had completely disappeared whereas the other phases were now clearly distinguishable. In Ti,CrsO7 about [ of the Cr s+ ions can he replaced by Fe s+ ions, without the structure undergoing any change. In the part of the diagrams with 33~ttool % TiOs, FesTiO6 forms a region of solid solutions. In this compound 15 tool ~ Fe,Os may be replaced by CrsO,. Another solid solution area at 33t tool ~o TiOi is found from 10 tool ~o FeOl.~ to about 45 tool ~o FeO~.5. This phase has the VsO~ structure and is Ill S. ANDERSSON,A. SUNDHOLMand A. MAGNI~LI,Acta Chem. Scand. 13, 989 (1959). i,~ M. HAMELIN,Bull. Soc. Chim. Fr. 1421 (1957). ~s~S. ASeRINK,S. FRIn~nG,A. MAGNI~LIand G. ANOERSSO~,Acta Chem. Scand. 13, 603 (1959). ~*~M. HAUELIN.Bull. Soc. Chim. Fr. 1431 (1957). 1(~--(12pp.)