Working
Report
2008-75
Solubility of UO2 in the High pH Range
in 0.01 to 0.1 M NaCl Solution
Under Reducing Conditions
Kaija Ollila
November
POSIVA
OY
Olkiluoto
FI-27160 EURAJOKI, FINLAND
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+358-2-8372 31
Fax +358-2-8372 3709
2008
Working
Report
2008-75
Solubility of UO2 in the High pH Range
in 0.01 to 0.1 M NaCl Solution
Under Reducing Conditions
Kaija Ollila
VTT
November
2008
Working Reports contain information on work in progress
or pending completion.
The conclusions and viewpoints presented in the report
are those of author(s) and do not necessarily
coincide with those of Posiva.
Solubility of UO2 in the High pH Range in 0.01 to 0.1 M NaCl Solution
Under Reducing Conditions
ABSTRACT
The objective of this experimental study was to measure the solubility of UO2 (s) in
high-pH 0.01 M NaCl solution under reducing conditions. The pH was varied from
slightly alkaline pH 9 to strongly alkaline pH 13. The solubility was measured from undersaturation. The solid phase was crushed polycrystalline UO2 pellet material. A series
of preliminary tests was conducted in a higher ionic strength 0.1 M NaCl solution. The
experiments were performed under nitrogen atmosphere in a glove box. The reducing
conditions were maintained by metallic iron in the solution. Parallel experiments were
performed with sulfide (1 ppm S2-) in the solution. Before the start of the solubility experiments, the solid UO2 samples were pretreated conducting predissolution periods in
order to remove any surface material that may contain higher oxidation states of U.
The solubility of U was observed to increase with pH in 0.01 M NaCl solution. The
measured concentrations in the aqueous phase with an iron strip in solution increased by
one order of magnitude from pH 9 to pH 13. The effect of pH was similar in the tests
with sulfide in the solution. The U concentrations in the solution were at the level of the
theoretical solubility of amorphous UO2 or higher. The increase in the ionic strength of
the solution from 0.01 M to 0.1 M did not seem to increase solubility. The measured
concentrations in 0.1 M NaCl solution were at the same level with the concentrations
measured in 0.01 M NaCl solution in the presence of iron (pH 9). The results in the
higher ionic strength solution were close to the detection limit of the analytical method,
which included the separation of U from saline solution with anionic exchange in
strongly acidic HCl solution.
Keywords: uranium dioxide, UO2, solubility, high pH, NaCl solution, reducing conditions, metallic iron
UO2:n liukoisuus korkealla pH-alueella 0.01 - 0.1 M NaCl liuoksessa
pelkistävissä olosuhteissa
TIIVISTELMÄ
Tämän kokeellisen tutkimuksen tarkoituksena oli mitata UO2 :n liukoisuus korkealla
pH-alueella 0.01 M NaCl - liuoksessa pelkistävissä olosuhteissa. Liuoksen pH vaihteli
välillä pH 9 - 13. Liukoisuutta mitattiin alikyllästystilasta lähtien. Kiinteä faasi oli
murskattu monikiteinen UO2 - pellettimateriaali. Sarja alustavia mittauksia tehtiin korkeamman ionivahvuuden 0.1 M NaCl-liuoksessa. Kokeet tehtiin hapettomassa kaapissa,
jossa oli typpiatmosfääri. Pelkistävät vesiolosuhteet saatiin aikaan lisäämällä metallirautalevy liuokseen. Rinnakkaisissa kokeissa liuokseen lisättiin sulfidia (1 ppm S2-).
Sulfidilisäys alentaa veden redox-potentiaalia (Eh ~ -0.2 V). Ennen varsinaisten liukoisuuskokeiden aloitusta tehtiin esiliuotuskokeita, joiden tarkoitus oli poistaa UO2:n pinnalta mahdollinen hapettunut aines, jossa uraanin hapetustila on korkeampi kuin IV.
Liukoisuuskokeiden tulosten mukaan U:n liukoisuus kasvoi pH:n funktiona 0.01 M
NaCl - liuoksessa pelkistävissä olosuhteissa. Mitatut U konsentraatiot metalliraudan
läsnäollessa kasvoivat yhden kertaluokan, kun pH muuttui lievästi alkaalisesta (pH 9)
voimakkaasti alkaaliseksi (pH 13). Vaikutus oli samansuuntainen, kun liuokseen oli
lisätty sulfidia. Mitatut konsentraatiot liuoksessa olivat amorfisen UO2:n teoreettisen
liukoisuuden tasolla tai jonkin verran korkeampia. Ionivahvuuden kasvu 0.01 M:sta
0.1 M:ksi ei vaikuttanut liukoisuuteen raudan läsnäollessa, vaan U-konsentraatiot liuoksessa olivat samalla tasolla (pH 9). Korkeamman ionivahvuuden liukoisuuskokeissa
mittaustulokset olivat lähellä käytetyn analyysimenetelmän määritysrajaa.
Avainsanat: uraanidioksidi, UO2, liukoisuus, korkea pH, NaCl, pelkistävät olosuhteet,
metallirauta
1
TABLE OF CONTENTS
ABSTRACT
TIIVISTELMÄ
1
INTRODUCTION ............................................................................................... 2
2
PREVIOUS STUDIES ........................................................................................ 3
3
EXPERIMENTAL PROCEDURES ..................................................................... 8
4
5
3.1
Solid phase ............................................................................................ 8
3.2
Aqueous phase ...................................................................................... 8
3.3
Redox conditions ..................................................................................... 8
3.4
Pretreatment of the UO2 solid samples ................................................... 9
3.5
U separation from high-ionic-strength solution ........................................ 9
SOLUBILITY EXPERIMENTS .......................................................................... 11
4.1
U solubilities in 0.01 M NaCl .................................................................. 11
4.2
U solubilities in 0.1 M NaCl .................................................................... 16
CONCLUSIONS .............................................................................................. 19
REFERENCES ......................................................................................................... 20
ACKNOWLEDGEMENTS .......................................................................................... 22
APPENDICES 1 - 2 .......................................................................................... 23 - 24
2
1
INTRODUCTION
The UO2 matrix solubility is a critical parameter for predicting the stability of spent nuclear fuel under disposal conditions. A large number of measurements concerning the
solubility of U(IV) oxides under anoxic or reducing conditions have been performed.
The solubilities reported in the literature for hydrous UO2 xH2O(s) (Rai et al. 1995),
amorphous or crystalline UO2(s) (Bruno et al. 1987, Rai et al. 1990, Yajima et al. 1995,
Rai et al. 2003) are very scattered. The extent of dissolution is known to be dependent
on the morphology of the solid phase, however, the reported data varies over orders of
magnitude for similar morphology.
In practice, in addition to the uncertainty involved with the solid phase, there are difficulties in excluding the effects of trace oxygen in the experiments. The equilibrium potential of the U(IV)/U(VI) couple is very low and hence U(IV) is very easily oxidized to
U(VI) in the presence of trace oxygen content. Very low O2 partial pressures
( 10-65 atm) are needed to ensure the thermodynamic stability of UO2. The trace oxygen concentration in the atmosphere of an anaerobic glove box normally is 1 ppm. It
is on the order of 0.1 ppm in the best inert atmosphere glove boxes. It corresponds to the
O2 partial pressure of 10-7 atm. An additional reducing agent is needed to lower the
oxygen concentration.
The ease of oxidation of U(IV) increases with increasing pH. The modelling calculations suggest, that the stability of UO2 in high pH area is very redox sensitive and very
low redox potentials are required to stabilize UO2, lower than in the neutral pH region
(OECD/NEA Report 2001). At high pH, the solubility of uranium starts to increase at
relatively low redox values. Conditions with high pH may occur due to the use of cementitious materials in the repository.
The objective of this experimental study was to measure the solubility of UO2 in
slightly to strongly alkaline pH in 0.01 M NaCl. The pH varied from 9 to 13. The pH
value as high as pH 13 is an extreme value and is not probable under disposal conditions. A series of preliminary tests was conducted at pH 8 - 9 in a higher ionic strength
solution in 0.1 M NaCl. The solid phase was crushed polycrystalline UO2 pellet material. The solubility is based on the direct measurement of the solution concentration in
the experiment vs. time. The experiments were performed under reducing conditions,
which were achieved through the inclusion of an iron strip in the solution and with nitrogen atmosphere. This simulates the conditions inside the disposal canister caused by
corrosion of iron. Parallel experiments were performed under anoxic conditions, in
which sulfide (1 ppm S2-) was added to the solution in order to decrease the amount of
trace oxygen under glove box conditions. One test series was performed in deaerated
0.1 M NaCl solution without sulfide.
These solubility experiments were conducted during the years 2005 and 2006. The redox conditions were similar to those used in the dissolution rate measurements with the
isotope dilution method in the NF-PRO project ( Ollila 2008).
3
2
PREVIOUS STUDIES
There are only a few studies available from the literature in which the solubility of uranium dioxide has been measured at high pH under reducing conditions in the presence
of iron or other reducing agents.
Rai et al. (1990) measured the solubility of amorphous UO2 x H2O in deoxygenated
deionized water at room temperature in the pH range from 2 to 12. They conducted the
experiments in a glove box (few ppm O2) in the presence of Eu2+ or Fe powder to maintain very low O2 fugacities. Solubility was approached from oversaturation. The studies
were confined to short time periods (days) to avoid the precipitates becoming crystalline
in order to obtain the solubility of amorphous UO2·x H2O. The solubilities observed at
higher pH values were about 3-4 orders of magnitude lower than those measured in earlier studies, see Figure 2.1. The solubility was about 10-8 M in the pH range 6-10 and
seems to increase in the pH range 10-12.
Yajima et al. (1995) measured the solubility of solid UO2, UO2(s), in 0.1 M NaClO4 under reducing conditions at room temperature. The pH ranged between 2 and 12. The experiments were performed in a glove box (Ar, O2 < 1 ppm). They used sodium hydrosulfite (Na2S2O4) to keep the valence state of uranium tetravalent. Dissolution equilibrium was approached from over- and undersaturation. In the undersaturation experiments, the UO2 (s) sample was prepared by reduction of ammonium diuranate
(NH4)2U2O7 at 650 C. The presence of crystalline UO2 in the solid sample was shown
by the analyses with X-ray diffraction. The measured aqueous uranium concentrations
after different contact times (7, 14, 28 days) are given in Figure 2.2.
Figure 2.1. Comparison of aqueous U concentrations in equilibrium with
UO2·xH2O(am) at pH 2-12 obtained in the study by Rai et al. 1990 (black symbols) with
earlier results (Rai et al. 1990 and the references therein).
4
Figure 2.2. U concentrations measured in the oversaturation and undersaturation experiments as a function of pH (Yajima et al. 1995).
In the pH range 3 to 12, the U concentrations seem to be independent of pH. They range
between about 10-9 M and 10-8 M. The XRD results suggested that the crystallization of
precipitates was progressing in the experiments between 7 and 28 days. This had little
effect on the U concentrations in the solutions (Yajima et al. 1995). On the other hand,
the kinetics of precipitation of crystalline tetravalent actinide oxides, including UO2, is
known to be very slow at room temperature. Rai et al. (2003) discuss that the solubility
cannot be approached from the oversaturation direction in a reasonable amount of time.
The aging periods of 7 to 28 days in the experiments of Yajima et al. (1995) are relatively short.
Rai et al. (2003) conducted the solubility study by precipitating UO2 (s) at 90 C. The
kinetics of precipitation is expected to be relatively rapid at elevated temperatures. They
precipitated UO2 (s) samples from low-pH U(IV) solutions (0.02 - 0.65 M Cl-) and
equilibrated the samples for 24 days at 90 C and 1 day at 22 C before analyses. The
pH ranged from 0.2 to 5.3. All experiments were conducted in atmospheric control
chamber (Ar, < 10 ppm O2). 0.01 M EuCl2 was added to the suspensions to maintain
reducing conditions. The measured U in solution at pH 5.3 was 8.6 · 10-10 M. The XRD
analyses of the solid phases identified amorphous UO2 as the dominant phase at pH 5.3.
In this study, no measurements at higher pH were performed.
Neck and Kim (2001) have summarized reported solubilities for hydrous
UO2 · x H2O (s), amorphous or microcrystalline UO2 (s), see Figure 2.3. There is a large
scatter in the solubility values. They conclude that the solubilities probably do not refer
to a unique material, but rather to a range of solids with different thermodynamic properties. The solubilities in 0.03 – 0.2 M solution (Figure 2.3, a) would seem to increase at
higher pH values. In addition to the uncertainty in the nature of the solid phase, the differences in the redox conditions of the solubility experiments may explain the large
scatter in the solubility results.
5
Figure 2.3. Solubility of UO2(s) as a function of the H+ concentration at 20-25 C, a)
at I= 0.03 – 0.2 M (see above), b) in 1 M NaCl. The solid lines were calculated for
0.1M (a) and 1 M NaCl (b). The dotted lines show the range of uncertainty (Neck and
Kim 2001 and the references therein).
Ollila (2002) measured the solubility of U(IV) in 0.01 M NaCl solution in the pH range
2 to 12 under anoxic conditions in the glove box (N2, O2 < 1 ppm). No additional reducing agents were used, as the objective of the tests was to test the experimental procedures. The solubility was measured from oversaturation. An aliquot of U(IV) solution
(0.1 M HCl) was added to 0.01 M NaCl. Upon pH adjustment dark-green precipitates
were immediately formed in all solutions. After this, the precipitates were allowed to
equilibrate. The U in solution was measured during the ageing of the precipitates.
6
10-3
10-4
U (mol/l)
10-5
10-6
10-7
10-8
10-9
10-10
10-11
0
2
4
6
8
10
12
14
pH
Figure 2.4. Solubility of U in 0.01 M NaCl under anoxic conditions (N2). The U concentrations were measured for ultrafiltered samples (~ 3.2 nm pore size) in three parallel experiments.
Figure 2.4 gives the measured U concentrations after the equilibration period of
14 months. The value for U in solution was (2.6 0.6) · 10-10 M at pH 8. The U concentration in solution was higher at pH 11 and 12. The nature of the precipitated solid phase
was not studied in these experiments. Werme et al. (2004) have concluded that the solubilities at this level are representative for pure UO2 (am). The solubility of crystalline
UO2 based on thermodynamic data (Guillamount et al. 2003) is several orders of magnitude lower (10-15 M).
The lowest U concentrations have been measured in the presence of metallic iron in the
dissolution studies with unirradiated UO2 and UO2 doped with 233U in synthetic Allard
groundwater (Ollila and Oversby 2005). The pH was between 9.1 and 9.2. The U concentrations in the solution samples at the end of the contact period (52 days) were below
the detection limit of the analytical method (8.4 · 10-11 M). By using the isotope dilution
method the sensitivity of the method could be increased by a factor of about 10. The
estimate for U solubility under these conditions was (1.3 0.7) 10-11 M.
Theoretical calculations suggest that the solubility of uranium at high pH critically depends on the system redox state (Eh). The solubility of uranium was modelled in hyperalkaline cement pore water (pH 12.6), see Figure 2.5. A series of conceivable iron
and/or sulphur bearing redox-controlling solid phase combinations together with their
resulting equilibrium Eh have been presented. They cover a broad Eh range from -780 to
-35 mV. A very low Eh is probable in the presence of metallic iron and its corrosion
products under disposal conditions. The solubility curve for U (Figure 2.5) shows the
redox sensitivity of the solubility under cement pore water conditions.
7
Figure 2.5. Theoretical solubility of uranium in cementitious environments as a function of redox potential, pH 12.6. Conceivable redox controlling phase assemblies were:
Magnetite/Fe(0)/Pyrite, Magnetite/Goethite, Magnetite/Troilite/Pyrite, Magnetite/
Troilite, Magnetite/Pyrite, Magnetite alone, Magnetite/Hematite, Magnetite/Fe(OH)3
(microcrystalline) (OECD/NEA Report 2001).
8
3
EXPERIMENTAL PROCEDURES
3.1 Solid phase
The UO2 solid phase was polycrystalline UO2 pellets. The pellets were unirradiated fuel
pellets supplied by a commercial nuclear fuel manufacturer (ABB). They contained
2.8 % 235U. The pellets were gently crushed to fragments under air atmosphere. The material in the size range of 2 to 4 mm was selected for the experiments. After this, the
fragments were immediately transferred into the anaerobic glove box. Tests used
3 grams of fragments in 100 ml of NaCl solution. The UO2 samples (3 g) were allowed
to stay in loosely closed experimental vessels for a couple weeks prior to the start of the
predissolution periods in order to remove loosely bound oxygen from the material.
3.2 Aqueous phase
The aqueous phase was 0.01 or 0.1 M NaCl solution. A concentrated stock solution
(1 - 2 M NaCl) was prepared from suprapure NaCl under atmospheric conditions. Subsequently, the stock solution was flushed with nitrogen, transferred to the glove box,
and allowed to equilibrate with nitrogen atmosphere in the glove box for a few days
before the dilution. The deionized water that was used for dilution had been deaerated
as follows: The water (2 l) was flushed with nitrogen over night and transferred to the
glove box. Flushing with nitrogen was continued in the glove box for 4 - 5 hours. After
this, the water was allowed to equilibrate with nitrogen atmosphere for 2 - 3 weeks (the
cap of the vessel loosely closed). The water was ready for use in the tests, as the oxygen
content was checked to be at zero with an Orbisphere oxygen analyzer.
The pH adjustment and the stability of the 0.01 M NaCl solutions in the presence of iron
were tested before the start of the solubility tests. pH measurements during the solubility
experiments may disturb the redox conditions of the tests. The test solutions with pH 8,
9, 10, 11, 12 and 13 (100 ml) were prepared in the glove box and an iron strip was immersed in each solution. The pH was adjusted with NaOH and the solutions were allowed to stay in the glove box in the closed vessels. The adjustment of pH affected the
ionic strength of the solutions at pH 12 and 13 ( Table App. 1-1). The pH was monitored for three months, see Figure App. 1-1. The pH stability of the solutions was good
at pH 9 … 13. At pH 8, the pH of the solution increased rapidly to 9.2 and was stable
afterwards. It turned out to be impossible to stabilize the pH below 9 in the presence of
metallic iron.
3.3 Redox conditions
The solubilities were measured under reducing conditions. The reducing conditions
were achieved through the inclusion of an actively corroding iron strip (1.5 x 3 cm) in
the solution (100 ml) under anaerobic conditions in the glove box (N2, with < 1 ppm
O2). After the immersion of the iron strip in solution, the Eh of the solution drops rapidly to a low value (~ -0.4 V, relative to the standard hydrogen electrode scale, SHE).
Parallel experiments were performed under anoxic conditions with sulfide (1 ppm S2-)
in solution. An aliquot of S(II) stock solution was added. The stock solution was prepared by dissolving Na2S· 9H2O in deionized water in the glove box. The addition of
9
sulfide reduced the Eh of the solution to about - 0.2 V. All the experiments were conducted at the ambient temperature of the glove box (24 – 25 C).
3.4 Pretreatment of the UO2 solid samples
Before the start of the solubility tests, the solid UO2 samples were pretreated with the
help of predissolution periods under anoxic conditions in order to remove any surface
material that may contain higher oxidation states of U. The method was as follows.
First, eight sequential predissolution periods (4, 2, 2, 6, 4, 4, 7 and 7 days) were conducted in deaerated 0.01 M NaCl solution under nitrogen atmosphere in the glove box.
Each UO2 sample (3 g) was placed on a silica saucer and the saucer with the solid sample was immersed in 30 ml of 0.01 M NaCl in a Nalgene jar. The solutions were
changed for fresh solutions between the dissolution periods. In all, 22 UO2 samples
were pretreated. The U concentrations in the solutions were measured after the termination of the last predissolution period of 7 days, see Table 3.1.
Thereafter, two additional predissolution periods (12 and 17 days) were conducted consecutively in 0.01 M NaCl with an iron strip in solution. The solutions were changed for
fresh solutions between these periods. Fresh iron strips were also used for both periods.
The U concentrations in the solutions were measured after the termination of both predissolution periods with Fe, see Table 3.1. The presence of iron in solution decreased
the U concentration by two orders of magnitude. It was very similar at the end of the
sequential dissolution periods with iron.
Table 3.1. The U concentrations in the solutions after the 7 d predissolution periods in
deaerated 0.01 M NaCl solution, and after the 12 d and 17 d predissolution periods
with Fe. The concentrations are the average values of the predissolution periods with
22 UO2 samples. The average of the absolute deviations of the measured data from
their mean value is also given.
U concentration
(ppb)
7 d predissolution period
in deaerated solution (N2)
12 day predissolution period
with Fe in solution
17 d predissolution period
with Fe in solution
U concentration
(mol/l)
17.4
2.5
(7.3
1.1) · 10-8
0.11
0.04
(4.8
1.5) · 10-10
0.11
0.04
(4.5
1.7) · 10-10
3.5 U separation from high-ionic-strength solution
The separation of U(VI) was performed with anion-exchange chromatography in HCl
medium. The separation is based on the fixation of U(VI) chloride complexes on the
anionic resin in HCl medium. The method is basically similar to the method developed
and tested for the determination of U(VI) and U(IV) oxidation states in anoxic aqueous
solutions (Hussonnois et al. 1989, Ollila 1996).
10
The method included the following steps:
1.
The anaerobic sample (3-7 ml) was brought out of the glove box and acidified
(1 M HCl). An aliquot of 233U solution was added to the sample for yield determination.
The sample was allowed to oxidize at least 3 hours.
2.
Concentrated HCl was added to the sample to form U chloride complexes. The
resulting solution should have 4.5 M Cl-. This breaks up colloids and/or unknown hydrolyzed and complexed species present in the sample. The strongly acidic solution was
allowed to stay for a couple of hours.
3.
Anionic resin (Dowex 1 x 8, 200-400 mesh Cl) was added to an empty column
(2 ml, Eichrom). The resin was washed with 4.5 and 0.1 M HCl to remove trace uranium from the resin. Finally, 4.5 M HCl was allowed to flow through the column.
4.
The U was fixed to the anionic resin by allowing the sample solution to flow
through the column.
5.
The U was recovered by elution with 0.1 M HCl.
6.
The concentration of U in 0.1 M HCl was determined with ICP-MS.
The acid solutions (0.1 M, 4.5 M HCl) were prepared using suprapure HCl. The flow
rate through the column was kept low (0.2 ml/min). An aliquot of 233U solution was
added to the samples for the yield determination of the separation.
The method was tested with 0.1 – 0.5 M NaCl solutions having known U concentrations
(0.2 … 1 ng U/sample) (Zilliacus et al. 2005). The yield was typically between 80 and
90 %. The tests showed that the purity of the resin must be checked. U contamination
from the resin was observed occasionally. The detection limit for U in suprapure 0.1 M
NaCl solution was tested to be 0.06 ng in the sample of 1 ml. If there is a sample of
10 ml available for the analysis, the detection limit would be 0.006 ppb U in solution
(2.5 10-11 mol/l).
11
4
SOLUBILITY EXPERIMENTS
4.1 Solubilities in 0.01 M NaCl
The solubility of UO2 was approached from the undersaturation direction in all experiments. The experiments were started immediately after the predissolution periods.
The 0.01 M NaCl solutions for duplicate tests with iron at each pH (pH 9, 10, 11, 12
and 13) were prepared in the glove box. Iron strips were added to the solutions. The solutions were allowed to stay with iron strips in the test vessels for three days, until the
UO2 solid samples (3 g) were immersed in the solutions (100 ml). The 0.01 M NaCl solutions with 1 ppm S2- were allowed to equilibrate also for three days after the addition
of sulfide and pH adjustment.
After the immersion of the UO2 solid samples in test solutions, the evolution of the U
concentration in solution was followed by taking small samples (1 ml) for analyses with
ICP-MS (Inductively Coupled Plasma – Mass Spectrometry). The vessels were kept
tightly closed during the experiments and opened for short periods for samplings and
pH measurements. The pH was checked every three months and readjusted with NaOH
or HCl if necessary. The samples for pH measurements were returned to the test vessels.
The samples for U analyses were acidified with nitric acid prior to the measurements
(1 M HNO3).
The following Figures 4.1…4.3 give the U concentrations, which were measured in
small aliquots from test solutions vs. time at different pH. The duration of the experiments was 14 months. Duplicate tests were performed under both redox conditions at
pH 9, 10, 11 and 12. At pH 13, duplicate tests were conducted under reducing conditions with iron and a single test under anoxic conditions with 1 ppm S2- in solution.
10-8
10-8
2-
U (mol/l)
pH 9, Fe
10-9
pH 9, 1 ppm S
10-9
10-10
10-10
UO2 (am), Eh= - 0.4 V
10-11
UO2(am), Eh= - 0.2 V
10-11
0
100
200
300
Time, days
400
500
0
100
200
300
400
500
Time, days
Figure 4.1. U concentrations vs. time in 0.01 M NaCl solution at pH 9. Duplicate tests
(filled and open symbols) were conducted under both redox conditions (metallic Fe or
1 ppm S2- in solution). The dashed line presents the theoretical solubility of UO2(am).
12
10-8
10-8
U (mol/l)
pH 10, Fe
pH 10, 1 ppm S
(
10-9
10-9
10-10
UO2 (am), Eh - 0.4 V
10-11
(
)
10-10
UO2(am), Eh= - 0.2 V
100
200
300
400
500
0
100
Time, days
200
300
400
500
Time, days
10-8
10-8
pH 11, Fe
U (mol/l)
)
10-11
0
10-9
pH 11, 1 ppm S
2-
10-9
10-10
UO2 (am), Eh = - 0.4 V
10-10
UO2(am), Eh= - 0.2 V
10-11
10-11
0
100
200
300
400
0
500
100
200
300
400
500
Time, days
Time, days
10-8
10-8
pH 12, Fe
U (mol/l)
2-
10-9
pH 12, 1 ppm S
2-
10-9
10-10
UO2(am), Eh= - 0.4 V
10-11
10-10
UO2(am), Eh= - 0.2 V
10-11
0
100
200
300
Time, days
400
500
0
100
200
300
400
500
Time, days
Figure 4.2. U concentrations vs. time in 0.01 M NaCl solution at pH 10, 11 and 12.
Duplicate tests (filled and open symbols) were conducted under both redox conditions
(metallic Fe or 1 ppm S2- in solution) at each pH. The dashed line presents the theoretical solubility of UO2(am).
13
10-7
pH 13
2black: Fe, white: 1 ppm S
U (mol/l)
10-8
10-9
10-10
UO2(am), Eh= - 0.2 V
10-11
0
100
200
300
400
500
Time, days
Figure 4.3. U concentrations vs. time in 0.01 M NaCl solution at pH 13. Duplicate tests
were conducted under reducing conditions (Fe) and a single test under anoxic conditions (1 ppm S2-). The dashed line presents the theoretical solubility of UO2(am).
At pH 9, the concentration of U in solution increased in the early stages during the first
100 days under both redox conditions (Figure 4.1). The increase was sharper in the
presence of iron. Afterwards, the U concentration decreased and was at the same level
under both redox conditions. The decrease in U concentration suggests the precipitation
or coprecipitation of U due to the strongly reducing conditions in the aqueous phase. It
is also possible that there was some U(VI) left on the UO2 surface at the beginning of
the experiments that is dissolved and precipitated after reduction to U(IV). The results
of the dissolution rate measurements of UO2 with the isotope dilution method showed
the precipitation or sorption of U to occur under similar conditions (Ollila 2008). U was
found in precipitates (colloids) in solution, as well as on the vessel walls. If the precipitation of U occurs in the tests of Figure 4.1, then the observed solubility is probably that
of amorphous UO2 rather than crystalline UO2 . The kinetics of precipitation of crystalline tetravalent actinide oxides, including UO2, at room temperature is known to be very
slow (Rai et al. 2003). The theoretical solubility for UO2 (am), that is presented in the
Figure was calculated with the EQ3NR Speciation-Solubility Code (Version 8.0, Wolery and Jarek 2003). The Data0.com.V8.R6 thermodynamic database by Lawrence Livermore National Laboratory was used with one modification. This composite datafile
includes the compilation of thermodynamic data for U minerals, aqueous species and
gases by the Nuclear Energy Agency (Grenthe et al. 1992). The data for U(OH)4 (aq)
was modified. The hydrolysis constant, log 14 = -10.00, for the reaction:
U4+ + 4 H2O = U(OH)4(aq) + 4 H+
was used (Guillamount et al. 2003). The U concentrations measured in the solutions at
pH 9 at the end of the experimental time are at the level of the calculated solubility of
UO2(am) under strongly reducing conditions, see Figure 4.1. The calculated solubility
for a crystalline solid, UO2 (c), is many orders of magnitude lower (1.45 · 10-15 M). The
measured U is at the same level as in the previous solubility experiments from oversaturation in 0.01 M NaCl (pH 8, Figure 2.4).
14
At pH 10, the concentration of U in solution in the presence of iron increased in a similar manner to the pH 9 tests in the early stages of the experiments, Figure 4.2. The decrease in U concentration was not observed afterwards. The [U] in solution seemed to
stabilize and remained at a higher level than the theoretical solubility of UO2(am) under
reducing conditions. The reason for this behaviour is not known. Generally, the oxidation of U(IV) to U(VI) is known to occur at lower redox potential when pH increases. If
oxidation occurs, the formation of U(VI) hydrolysis complexes may explain the higher
U concentrations. Oversby (2008) has suggested that the presence of large U(IV)
polynuclear complexes in solution may increase solubility under reducing conditions.
Some scatter in the results of the analyses was observed in the data for the tests in
0.01 M NaCl with 1 ppm S2- in the solution at pH 10 - 13. This is probably related to the
use of Bi as internal standard in the ICP-MS measurements. 10 ppb Bi was added to the
acidified sample (1 M HNO3) before the measurements. Bi apparently precipitates with
sulphide if the anaerobic sample was not allowed to stay long enough under aerobic
conditions with acid to oxidize S(II) before the addition of Bi. The inhomogeneity of the
sample solution may have disturbed U measurements in some samples. After these tests,
In replaced Bi as internal standard in U analyses.
The U concentrations seemed to stabilize at the end of the tests at each pH. The [U] in
solution at the end of the tests increased slowly as a function of pH under both redox
conditions, being 10-9 M at pH 13. It was close to or above the theoretical solubility of
U(IV)O2 (am). In the last sampling at 432 days, triplicate samples were taken from each
test solution for U analyses. One of the samples from each test was filtered using a
membrane with 0.22 μm pore size. The data for 432 days in Figures 4.1 - 4.3 are the
average values of the unfiltered samples. Microfiltering did not change the results. Figures 4.4 and 4.5 give the average values of all results of analysis at each pH. The range
bars give the upper and lower limits for the measured values. There is roughly an increase of one order of magnitude in U solubility between pH 9 and 13. The difference in
redox conditions had only a small effect on the results. The pH values were measured
for the test solutions at the end of the tests, see App. 2. A decrease in pH was observed
due to the dissolution of U in the tests at pH 9, 10 and 11. There was a slight decrease at
pH 12 and at pH 13. The changes in the [OH-] contents suggest the formation of uranium hydroxide complexes (U(OH)x).
15
10-7
N2, Fe
U (mol/l)
10-8
10-9
10-10
10-11
8
9
10
11
12
13
14
pH
Figure 4.4. U concentrations in 0.01 M NaCl at the end of the tests (432 days) vs. pH
under reducing conditions with Fe. The data are the average values for parallel tests.
Triplicate samples were taken in each test. The range bars give the upper and lower
limits for the measured values.
10-7
N2, 1 ppm S2-
U (mol/l)
10-8
10-9
10-10
10-11
8
9
10
11
12
13
14
pH
Figure 4.5. U concentrations in 0.01 M NaCl at the end of the tests (432 days) vs. pH
under reducing conditions with 1 ppm S2- in solution. The data are the average values
for parallel tests. Triplicate samples were taken in each test. The range bars give the
upper and lower limits for the measured values.
16
4.2 U solubilities in 0.1 M NaCl
Finally, preliminary solubility measurements were performed in 0.1 M NaCl solution.
The redox conditions were chosen to be deaerated solution under N2 atmosphere and
reducing conditions produced by metallic iron. Duplicate tests were conducted under
both redox conditions. Test series in deaerated solution were chosen to conduct because
the U concentrations in these redox conditions were known to be higher than the ones
under reducing conditions, which are close to the detection limit. It was not known if it
is possible to measure the U concentration in the solution under reducing conditions because it was necessary to use a separation method for U for samples in order to measure
the U contents in 0.1 M NaCl solution with ICP-MS. A single test was performed for
comparison with 1 ppm S2- in solution. The pH was between 8 and 9 (see Table A2-2,
App. 2.)
The solubility tests were conducted similarly with the previous tests in 0.01 M NaCl.
After the immersion of the UO2 solid samples (3 g) in test solutions (100 ml), the evolution of the U concentration in solution was followed by taking small samples (2-5 ml)
for ICP-MS analyses. The samples were treated with anion-exchange in 4.5 M HCl to
separate the U from 0.1 M NaCl solution (for the method, see paragraph 3.5).
The measured U concentrations in dearated 0.1 M NaCl solution are given in Figure 4.6.
The data are the average values of the parallel tests. The U concentrations in the samples were from 0.2 to 1.5 ppb (0.2 to 1.5 ng U/1 ml sample), being clearly above the estimated detection limit of the separation method, which was 0.06 ng U in the 1 ml sample. The sample volume of 2 ml was used in these tests. The yields of the separations
varied between 65 and 95 %. The U concentration increased slowly in the solution until
the end of the experimental time.
U (mol/l)
10-8
10-9
UO2 (am), Eh= -0.1 V
10-10
10-11
0
100
200
300
400
Time (days)
Figure 4.6. Average U concentration vs. time in deaerated 0.1 M NaCl solution
(pH 8.3) under N2 atmosphere. The range bars give the upper and lower limits for the
analysed concentrations in parallel tests. The dashed line presents the theoretical solubility of UO2(am).
17
There is a break in the samplings between 180 and the last sampling of 322 days. This
was due to problems with ICP-MS, which was out of order during that period. The U
concentration in solution is at higher level than the theoretical solubility of UO2(am).
The solubility was calculated using the Eh value of -0.1 V. It is probable that the formation of U(VI) hydroxide complexes increases solubility under these conditions, because
no reducing agent was added to solution.
Under reducing conditions with an iron strip in solution, the U concentrations in the
samples were varying from 0.02 to 0.2 ppb (0.02 to 0.2 ng U/1 ml sample). They were
close to the detection limit of the separation method except in the beginning. The sample volume of 4 - 5 ml was used. The analyses gave similar results in the last sampling
at 320 days for parallel tests, 0.02 0.01 ppb U. Triplicate samples from each test were
treated in this sampling with anion exchange. The yield of the separation varied from 84
to 94 %. The U concentration decreased slowly as the test proceeded, see Figure 4.7.
This suggests the precipitation or coprecipitation of uranium solid phase due to the
strongly reducing conditions in the aqueous phase. The U concentration in solution is at
the level of the theoretical solubility of amorphous UO2. This is in agreement with the
results of the solubility tests in 0.01 M NaCl (see Figure 4.1).
Figure 4.8 compares the measured concentrations in 0.1 M NaCl under different redox
conditions. A single test was performed under anoxic conditions with 1 ppm S2- in the
solution. The U concentrations were halfway between the concentrations measured in
deaerated solution and in the presence of an iron strip in solution. They varied from 0.1
to 0.3 ppb (0.1 to 0.3 ng U/1 ml sample), being above the detection limit of the separation method. The sample volume of 2 - 3 ml was used. The yields of the separations
were lower in the case of these samples containing sulphur. They varied from 30 to 70
%, in most cases they were near 50 %. The reason for this is not known.
U (mol/l)
10-8
10-9
UO2 (am), Eh= -0.4V
10
-10
10-11
0
100
200
300
400
Time (days)
Figure 4.7. Average U concentration vs. time in 0.1 M NaCl solution (pH 9.0) under
reducing conditions (Fe). A test in duplicate was conducted. The range bars give the
upper and lower limits for the analysed concentrations. The dashed line presents the
theoretical solubility of UO2(am).
18
10-8
U (mol/l)
deaerated solution, N2
10-9
N2,1 ppm S2-
10-10
N2, Fe
10-11
0
100
200
300
400
Time, days
Figure 4.8. Comparison of U concentrations measured vs. time in 0.1 M NaCl under
different redox conditions.
It seems that the presence of sulphur in the sample disturbs the fixation of U on the anionic resin. The results given in Figure 4.8 include the yield corrections. The effect of the
redox conditions is seen clearly. The solubility is lowest in the presence of an iron strip
and highest in deaerated solution. There is a difference between the measured concentrations in 0.1 M NaCl in the presence of sulphide and metallic iron, which was not observed in 0.01 M NaCl (Figure 4.1). The observed difference in 0.1 M NaCl is in
agreement with the measured Eh of the solutions. The measured redox potential in the
solution which contains an iron strip, is extremely low (- 0.4 V). In the sulphide containing solution, the corresponding value is around - 0.2 V. The Eh measurements were performed in the beginning of the tests for fresh solutions.
19
5
SUMMARY AND CONCLUSIONS
The solubility of UO2 was measured under reducing conditions in the slightly to
strongly alkaline pH range (pH 9 - 13). The aqueous phase was 0.01 M NaCl. A series
of preliminary solubility tests was conducted in 0.1 M NaCl (pH 8 - 9). The solid phase
was crushed polycrystalline UO2 pellet material. The reducing conditions were achieved
with the help of an iron strip in solution. Parallel experiments were performed under
anaerobic conditions, in which sulfide (1 ppm S2-) was added to the solution to decrease
the effect of trace oxygen in glove box conditions.
Some theoretical calculations found from literature suggested that the solubility of uranium at high pH area is very redox sensitive, and very low redox potentials are required
to stabilize UO2. Previous experimental studies under anoxic conditions showed some
increase at higher pH. However, there was a large scatter in the solubility values found
from literature.
In this study, a pH effect was observed in the high pH range in 0.01 M NaCl under
strongly reducing conditions produced by metallic iron. The solubility of U increased by
one order of magnitude from pH 9 to pH 13. The U concentrations in 0.01 M NaCl containing 1 ppm S2- were very similar at the end of the experiments and showed a similar
trend with pH. The U concentrations in the solution were at the level of the theoretical
solubility of amorphous UO2 or higher. The theoretical solubility of crystalline UO2 is
many orders of magnitude lower (10-15 M). At pH 9, the U concentration in solution increased in the early stages and decreased afterwards in both redox conditions suggesting
the precipitation of U. At pH 10 - 12, the decrease of U concentration was not observed
in the later stages. The U seemed to stabilize at higher level than the calculated solubility of amorphous UO2. The reason is not known. It is possible that the formation of
U(VI) hydrolysis complexes may explain higher U concentrations in solution because
the ease of oxidation increases with pH. The formation of large polynuclear U(IV) complexes has been suggested. A speciation analysis of the solution would be needed to
clarify mechanisms. A method which is sensitive enough was not available for these
tests. New tests using larger amounts of solid and aqueous phases would be needed.
The measured solubility of U in 0.1 M NaCl solution (pH 8-9) under reducing conditions with iron was at the same level with the solubility in 0.01 M NaCl. The increase in
ionic strength did not seem to increase solubility. The analysis results in these conditions were close to the detection limit of the analytical method, which included the separation of U from saline solution with anionic exchange in strongly acidic HCl solution.
The U concentrations in 0.1 M NaCl under anaerobic conditions with 1 ppm S2- in solution were somewhat higher.
20
REFERENCES
Bruno J., Casas I., Lagerman B. and Munoz M. 1987. The determination of the solubility of amorphous UO2(s) and the mononuclear hydrolysis constants of uranium(IV) at
25 C. In: Scientific Basis for Nuclear Waste Management X (Eds. J. K. Bates and W.
B. Seefeldt), Pittsburgh, Pennsylvania, USA: Mat. Res. Soc. Symp. Proc. Vol. 84, pp.
153-160.
Grenthe I., Fuger J., Konings R. J. M., Lemire R. J., Muller A. B., Nguyen-Trung Cregu
C. & Wanner H. 1992. Chemical Thermodynamics of Uranium (eds. H. Wanner and I.
Forest), OECD/Nuclear Energy Agency, The netherlands: North-Holland Elsevier Science Publishers B. V. 715 p.
Guillamount R., Fanghänel T., Fuger J., Grenthe I., Neck V., Palmer D. A. and Rand M.
H. 2003. Update on the chemical thermodynamics of uranium, neptunium, plutonium,
americium and technetium (Eds. F. J. Mompean, M. Illemassene, C. Domenech-Orti
and K. Ben Said), OECD/Nuclear Energy Agency, The Netherlands: Elsevier B. V.
919 p.
Hussonnois M., Guillaumont R., Brillard L. and Fattahi M. 1989. A method for determining the oxidation state of uranium at concentration as low as 10-10 M. In: Scientific
Basis for Nuclear Waste Management XII (eds. W. Lutze and R.C. Ewing), Pittsburgh,
Pennsylvania, USA: Mat. Res. Soc. Symp. Proc. Vol. 127, pp. 979-985.
Neck V. and Kim J. I. 2001. Solubility and hydrolysis of tetravalent actinides. Radiochim. Acta 89, pp. 1-16.
OECD/NEA Report 2001. The use of thermodynamic databases in performance assessment. Workshop Proceedings, Barcelona, Spain, 29-30 May 2001.
Ollila K. 2002. Uranium solubility. In: Nuclear Waste Managements in Finland. Final
report of Public Sector’s Research Programme, JYT 2001 (Ed. K. Rasilainen), Ministry
of Trade and Industry, Helsinki, Finland, pp. 77-79.
Ollila K. and Oversby V. 2005. Dissolution of unirradiated UO2 and UO2 doped with
U under reducing conditions. POSIVA Report 2005-05, Olkiluoto, Finland: Posiva
Oy. 28 p.
233
Ollila K. 1996. Determination of U oxidation state in anoxic (N2) aqueous solutions –
method development and testing. Helsinki, Finland: Posiva Oy. POSIVA Report 96-01.
33 p.
Ollila K. 2008. Dissolution of unirradiated UO2 and UO2 doped with 233U in low- and
high-ionic-strength NaCl under anoxic and reducing conditions. To be published as
POSIVA Report 2008.
Oversby V. VMO Konsult, private communication 4.4.2008.
21
Rai D., Felmy A. R. and Ryan J. L. 1990. Uranium(IV) hydrolysis constants and solubility product of UO2 x H2O(am). Inorg. Chem. 29, pp. 260-264.
Rai D., Felmy A. R., Moore D. A. and Mason M. J. 1995. The solubility of Th(IV) and
U(IV) hydrous oxides in concentrated NaHCO3 and Na2CO3 solutions. In: Scientific
Basis for Nuclear Waste Management XVIII (Eds. T. Murakami and R.C. Ewing),
Pittsburgh, Pennsylvania, USA: Mat. Res. Soc. Symp. Proc. Vol. 353, pp. 1143-1150.
Rai D., Yui M. Z. and Moore D. 2003. A. Solubility and solubility product at 22 C of
UO2(c) precipitated from aqueous U(IV) solutions. J. Solution Chem. 32, pp. 1-17.
Werme L. O., Johnson L. H., Oversby V. M., King F., Spahiu K., Grambow B. and
Shoesmith D. W. 2004. Spent fuel performance under repository conditions: A model
for use in SR-Can. SKB Technical Report TR-04-19, Stockholm, Sweden: Swedish Nuclear Fuel Waste Management Co. 34 p.
Wolery T. W. and Jarek R. L. 2003. Software user’s manual. EQ3/6, Version 8.0. Sandia National Laboratories. January 2003.
Yajima T., Kawamura Y., Ueta S. 1995. Uranium(IV) solubility and hydrolysis constants under reduced conditions. In: In: Scientific Basis for Nuclear Waste Management
XVIII (Eds. T. Murakami and R.C. Ewing), Pittsburgh, Pennsylvania, USA: Mat. Res.
Soc. Symp. Proc. Vol. 353, pp. 1137-1142.
Zilliacus R., Lipponen M. and Ollila K. 2005. Development of analytical methods for
saline waters. Posiva Working Report 2005-42 (In Finnish). Olkiluoto, Finland: Posiva
Oy. 32 p.
22
ACKNOWLEDGEMENTS
The author takes pleasure in thanking Ms. R. Zilliacus, Ms. M. Lipponen and Ms. J.
Rantanen for performing the ICP-MS analyses. The contribution of K. Helosuo is also
gratefully acknowledged. The author gives thanks to Dr. V. Oversby for valuable comments.
23
Appendix 1
14
pH0 8.3
pH0 9.0
13
pH0 10.0
pH0 11.0
12
pH0 12.0
pH
pH0 13.0
11
10
9
8
0
20
40
60
80
100
Time (days)
Figure A1-1. pH stability of the 0.01 M NaCl test solutions in the presence of iron
strips under N2 atmosphere (without UO2 solid phases).
Table A1-1. The ionic strength of the test solutions after pH adjustment.
pH
8.3
9.0
10.0
11.0
12.0
13.0
Ionic strength (M)
0.01
0.01
0.01
0.011
0.03
0.11
24
Appendix 2
Table A2-1. Measured initial and final pH values for the test solutions in the solubility
tests with UO2 in 0.01 M NaCl. Changes in the [OH-] are also given (M).
Redox
Final
pH
8.7
8.6
8.8
8.7
Initial
[OH-]
10-5
1.3 10-5
1.3 10-5
7.9 10-6
[OH-]
pH 9, Test 1
pH 9, Test 2
pH 9, Test 1
pH 9, Test 2
N2, Fe
“
N2, 1 ppm S2“
Initial
pH
9.0
9.1
9.1
8.9
Final
[OH-]
5.0 10-6
4.0 10-6
6.3 10-6
5.0 10-6
5.0
9.0
6.7
2.9
10-6
10-6
10-6
10-6
pH 10, Test 1
pH 10, Test 2
pH 10, Test 1
pH 10, Test 2
N2, Fe
“
N2, 1 ppm S2“
10.1
10.1
10.0
10.0
9.4
9.4
9.6
9.6
10-4
10-4
10-4
10-4
2.5
2.5
4.0
4.0
10-5
10-5
10-5
10-5
7.5
7.5
6.0
6.0
10-5
10-5
10-5
10-5
pH 11, Test 1
pH 11, Test 2
pH 11, Test 1
pH 11, Test 2
N2, Fe
“
N2, 1 ppm S2“
11.0
11.0
11.0
11.0
10.4
10.4
10.5
10.5
10-3
10-3
10-3
10-3
2.5
2.5
3.2
3.2
10-4
10-4
10-4
10-4
7.5
7.5
6.8
6.8
10-4
10-4
10-4
10-4
pH 12, Test 1
pH 12, Test 2
pH 12, Test 1
pH 12, Test 2
N2, Fe
“
N2, 1 ppm S2“
12.0
12.0
12.0
12.0
11.9
11.9
11.9
11.8
10-2
10-2
10-2
10-2
7.9
7.9
7.9
6.3
10-3
10-3
10-3
10-3
2.1
2.1
2.1
3.7
10-3
10-3
10-3
10-3
pH 13, Test 1
pH 13, Test 2
pH 13, Test 1
N2, Fe
“
N2, 1 ppm S2-
13.0
13.0
13.0
12.9
13.0
12.9
10-1
10-1
10-1
7.9 10-2
10-1
7.9 10-2
2.1 10-2
2.1 10-2
Table A2-2. Measured initial and final pH values in test solutions in the solubility tests
with UO2 in 0.1 M NaCl.
Redox
Test 1
Test 2
Test 3
Test 4
Test 5
deaerated solution, N2
“
N2, Fe
“
N2, 1 ppm S2-
Initial
pH
8.3
8.3
9.0
9.0
8.4
Final
pH
8.4
8.3
8.4
8.4
8.2