thermodynamic model for the solubility of tco2⋅xh2o in aqueous oxalate systems

J Solution Chem (2008) 37: 1471–1487DOI 10.1007/s10953-008-9328-5

Thermodynamic Model for the Solubility of TcO2·xH2Oin Aqueous Oxalate Systems

Nancy J. Hess · Odeta Qafoku · Yuanxian Xia ·Dean A. Moore · Andrew R. Felmy

Received: 11 April 2008 / Accepted: 4 June 2008 / Published online: 17 September 2008© Pacific Northwest National Laboratory 2008

Abstract The room temperature solubility of amorphous, hydrous technetium(IV) oxide(TcO2·xH2O) was studied across a broad range of pH values extending from 1.5 to 12 andin oxalate concentrations from dilute (10−6 mol·kg−1) to complete saturation with respectto sodium bioxalate at lower pH values, and to saturation with respect to sodium oxalateat higher pH values. The solubility was measured to very long equilibration times (i.e., aslong a 1000 days or longer). The thermodynamic modeling results show that the dominantspecies in solution must have at least one more hydroxyl moiety present in the complex thanproposed by previous investigators (e.g., TcO(OH)Ox− rather than TcO(Ox)(aq)). Inclusionof the single previously unidentified species TcO(OH)Ox− in our aqueous thermodynamicmodel explains a wider range of observed solubility data for TcO2·xH2O(am) in the presenceof oxalate and over a broad range of pH values. Inclusion of this species is also supportedby the recently proposed thermodynamic data for the TcO(OH)+ hydrolysis species thatindicates that this species is stable at pH values as low as one.

Keywords Technetium(IV) · Oxalate · Complexation · Stability constants · Solubility ·Speciation

1 Introduction

Technetium-99 is a long-lived fission product (t1/2 = 2.13×105 years and fission yield about6%) produced in nuclear reactors from the fission of uranium 235 and plutonium 239, andtherefore is present in large quantities in high-level nuclear waste (HLW). In total, approxi-mately 33 kCi (1940 kg) of 99Tc are currently stored in the underground HLW tanks at the

N.J. Hess (�) · O. Qafoku · Y. Xia · D.A. MooreEnvironmental Dynamics and Simulation, Pacific Northwest National Laboratory, Richland, WA 99354,USAe-mail: [email protected]

A.R. FelmyW.R. Wiley Environmental Molecular Science Laboratory, Richland, WA 99354, USA

mailto:[email protected]

1472 J Solution Chem (2008) 37: 1471–1487

U.S Department of Energy (DOE) Hanford site [1]. Several of the tanks at Hanford alsocontain significant concentrations of organic chelating agents including ethylenediaminete-traacetic acid (EDTA), citric acid, and oxalic acid. Chemical speciation studies of HLWsupernatant in tanks indicate that Tc is not only present as TcO−

4 , but also can exist as or-ganic complexes of reduced technetium species, most likely Tc(IV) [2–4]. Currently, HLWsludge is being retrieved from selected tanks both at the Hanford site and at the SavannahRiver site (SRS) for processing into glass for eventual repository disposal. Saturated oxalicacid is the primary waste retrieval solution at the SRS and has also been used to facilitatesludge dissolution for tank C-106 at the Hanford site. The solution chemistry of technetiumin tank waste is therefore of significant interest especially since some of these tanks havebeen found to leak into the subsurface.

It is well known that technetium can exist in multiple oxidation states (such as III, IV, Vand VII). Among these oxidation states, Tc(IV) and Tc(VII) are the most stable oxidationstates under environmental and tank wastes conditions. The chemical behavior of technetiumin these two oxidation states differs drastically. In the heptavalent state, technetium formshighly soluble solids. In aqueous solution, Tc(VII) exists as pertechnetate, TcO−

4 , which ishighly mobile in the environment. In contrast, tetravalent technetium forms fairly insolublehydrous oxides (TcO2(am)), and is relatively immobile in the absence of strongly chelatingagents. Oxalic acid could enhance the solubility of Tc(IV) by forming soluble complexes [5,6], thus affecting the migration behavior of Tc(IV) in the environment. Unfortunately, therehas been only one previous study on the impacts of oxalate on the aqueous complexation ofTc(IV) by oxalate [7] and these results were limited strictly to acidic conditions (pH < 2).

In this study we examine the complexes of Tc(IV) with oxalic acid across a range ofpH values from 1 to 12 and across the entire range of oxalate or bioxalate concentrationsextending from dilute to complete saturation with respect to sodium bioxalate or sodiumoxalate. The studies also extend to very long equilibration times (as long as 1000 days).

2 Experimental

All experiments and sample preparations were conducted at ambient temperature (23±1 ◦C)in an atmosphere-controlled chamber under an Ar(g) atmosphere. Deionized distilled water,degassed by boiling and cooling in an Ar atmosphere, was used in all cases. Steady-stateTc concentrations in the solubility experiments were approached from the undersaturationdirection.

A 0.29 mol·L−1 NH4TcO4 stock solution was prepared, following purification, from ir-radiated MoO3 obtained from Oak Ridge National Laboratory. The purity of the ammo-nium pertechnetate was checked spectrophotometrically, and its content was establishedby comparative measurement of the β-activity of actual and standard 99Tc solutions. TheTcO2·xH2O(am) precipitate was prepared individually for each sample starting from thestock solution and under basic conditions. Specifically, a small amount (0.18 mL) of theTc stock solution was added to a 25-mL glass centrifuge tube that contained 4.12 mL wa-ter and 0.195 mol·L−1 freshly prepared Na2S2O4 solution and the pH of this solution wasthen adjusted to about 12 using NaOH. A black precipitate formed quickly and was allowedto mature in the mother liquid for 72 hrs. Approximately 5 mg of the precipitate was thenwashed three times with 20 mL of freshly prepared 0.01 mol·L−1 Na2S2O4 in near neutralaqueous solution. The precipitate was equilibrated with 20 mL of pH-adjusted solution forat least three days prior to the initial sampling.

Two types of experiments were conducted using the washed precipitates: experiments inwhich the concentration of oxalate was fixed at 0.01, 0.025, and 0.25 mol·kg−1 while the pH

J Solution Chem (2008) 37: 1471–1487 1473

varied from 0.1 to 12, and experiments where the pH was maintained at fixed values of ca.1.5, 3.5, and 5.5 and the oxalate concentration was varied from 10−7 to 10−2 mol·kg−1. Un-der an Ar atmosphere, 20 mL of the appropriate solution containing 0.02 mol·L−1 hydrazine(used to maintain reducing conditions) was added to the washed precipitates. The pH val-ues of the samples were readjusted using NaOH or HCl. The tubes were tightly cappedand placed on a shaker in an argon atmospheric chamber. After 3 to 15 days, the pH wasmeasured using a Orion-Ross combination glass electrode calibrated against pH buffers.1

Although the solutions have varying ionic strength no correction to the pH values was ap-plied, thus the reported pH values are measured values. After different equilibration periods(4 to 17 days, 25 to 44 days, 65 to 73 days, 245 to 263 days, and 584 to 1137 days), 1.5 mLof the solution was withdrawn for analyses. The sampling procedure involved centrifugingthe aliquot at 2000 rpm for 10 minutes followed by filtering through a Centricon-30 filter(Amicon, Inc.) with an approximate 0.0036 μm pore size. The color of the filtered solutionsvaried from clear and colorless to clear grey at higher Tc concentrations. The total Tc con-centration in solution was determined by β-scintillation counting using a 0.5 mL subsampleof the filtrate and 10 mL of scintillation cocktail. The detection limit for this technique isapproximately 10−8 mol·L−1 Tc. The concentrations of Tc(VII) and reduced Tc in the sam-ple were determined using the solvent extraction technique in which TcO−

4 is extracted intoCHCl3 with tetraphenylphosphonium chloride (TPPC), leaving reduced Tc in the aqueousphase. For this oxidation state analysis, the pH of a second 0.5 mL subsample of the fil-trate was adjusted to 1 mol·L−1 HCl using 6 mol·L−1 HCl or 5 mol·L−1 NaOH and dilutedto 1.0 mL with water. The diluted and pH-adjusted filtrates were shaken vigorously with1 mL of the 0.05 mol·L−1 TPPC in the chloroform solvent mixture for five minutes. Aftercentrifugation, the Tc concentration in the organic and aqueous fractions was subsequentlydetermined by β-scintillation counting. The concentrations of chloride and sodium in theequilibrated samples were calculated from the input molalities.

A limited number of solid samples were selected for powder X-ray diffraction (XRD)measurements on a Scintag PAD V X-ray diffractometer using Cu Kα radiation at 45 kVand 40 mA. Approximately 0.5 mg of the solid phase was washed with Fe-equilibratedwater to prevent Tc oxidation and then suspended in alcohol. Using a disposable pipette thesuspended solids were deposited onto a glass slide and the residual alcohol was allowed toevaporate. Diffraction data were collected from 5 to 65◦ 2� at a rate of 0.04◦·min−1.

Additional solid samples were isolated for Raman spectroscopic measurements. Approx-imately 0.2 mg of washed solid phase was placed on a convex glass slide and sealed with aCoverWell imaging chamber under anaerobic conditions. The amorphous solid was imagedusing a 100× objective of a confocal Raman microscope. The Raman signal was excited us-ing 5 mW of 532 nm excitation from a doubled CW Nd:YAG diode laser. The scattered lightwas collected in backscattering geometry and passed through a DILOR 800 XY triplespec-trometer and finally dispersed onto a liquid-nitrogen cooled, charge couple device detector.

1For solutions with pH less than 7 we used Fischer Chemical buffers at pH = 7.00 ± 0.01, 4.00 ± 0.01 and1.00 ± 0.02. For solutions with pH greater than 7, we used buffers pH = 4.00 ± 0.01 and 7.00 ± 0.01. Formost experiments in our work the pH = 1.00±0.02 buffer was used as a check after electrode calibration. TheNernst slopes obtained were between 99.6% and 98.8% of the theoretical values. The estimated experimentalaccuracy of the final pH measurements was ±0.4% based on measuring the pH value of a buffer standardsolution after calibrating the pH electrode with respective buffer solutions. For measurement of hydrogenconcentrations, usually the equation pcH (= − log[H+]) = pH (meter reading) + constant can be used whereconstant is the correction factor depending on the solution composition, especially at the ionic strength ofthe solution. In our study, for most experiments the concentrations of oxalic acid used were low. Thus theliquid-junction potential was not accounted for in our work.

1474 J Solution Chem (2008) 37: 1471–1487

The spectrometer was calibrated using Hg emission lines. Data were collected from 200to 1400 cm−1 in three segments and averaged for approximately 600 seconds per segment.Raman spectra were collected after 1200 days of equilibration.

3 Thermodynamic Model

The ion-interaction model of Pitzer and co-workers [8, 9] was used to interpret the solubilitydata. This aqueous thermodynamic model emphasizes a detailed description of the specificion interactions in the solution. The effects of specific ion interactions on the solution ex-cess Gibbs energy are contained in the expressions for the activity coefficients. The activitycoefficients can be expressed in a virial-type expansion as

lnγi = lnγ DHi +

∑

j

βij (I )mj +∑

j

∑

k

Cijkmjmk + · · · (1)

where m is the molality, γ DHi is a modified Debye-Hückel activity coefficient that is a uni-

versal function of ionic strength, and βij (I ) and Cijk are specific for each ion interactionand are functions of the ionic strength. The third virial coefficient, C, is understood to be in-dependent of ionic strength. A detailed description of the exact form of Eq. 1 was publishedin Felmy and Weare [10] and Felmy et al. [11].

4 Results and Discussion

4.1 Characterization of Amorphous TcO2·xH2O

Powder X-ray diffraction and Raman spectral analyses were used to characterize the solidphases. These studies indicate the presence of amorphous TcO2·xH2O as the solubility-limiting phase under the pH conditions investigated at the beginning of the equilibrations.X-ray diffraction analysis of the solids after approximately 1200 days of equilibrationshowed the solids to be amorphous under all solution conditions examined as shown inFig. 1. Similarly, Raman analysis (data not shown) of the solid phases revealed only broadfeatures indicative of amorphous solids with no significant change as a function of equili-bration time, oxalate concentration, or pH.

4.2 Technetium Oxidation States in Solution

Oxidation state analyses using the solvent extraction technique indicate that under mostsolution conditions the percentage of reduced Tc species in solution exceeded 80% of thetotal measured Tc even after 1137 days of equilibration. Lower percentages occur where themeasured Tc concentration approaches the detection limit of the beta-scintillation countingtechnique, approximately 10−8 mol·L−1 Tc, where it is difficult to reliably determine Tcconcentrations in solution. The measured pH, the measured Tc(IV) concentration based onsolvent extraction analyses, and the percentage of Tc(IV) of the total measured Tc in solutionfor all data sets are reported in Tables 3–13 of the Appendix.

J Solution Chem (2008) 37: 1471–1487 1475

Fig. 1 Diffractogram of theamorphous TcO2·xH2O solidphase over a range pH conditionsand under (a) low and (b) highoxalate concentrations

4.3 Analysis of TcO2·xH2O Solubility Data

The solubility of TcO2·xH2O as a function of added oxalate (Fig. 2) shows a trend of rel-atively constant concentrations at low total added oxalate (<10−4 mol·kg−1) followed bya rapid increase at higher added oxalate indicating the onset of significant Tc(IV)-oxalatecomplex formation. In general the solutions appear to come into equilibrium with the solidsat relatively early equilibration times (9 to 44 days), although a slower equilibration processis observed at higher oxalate concentration. An exception to the rather rapid solid/solutionequilibration occurs at low total added oxalate at the lowest pH range (Fig. 2a). A remarkabledecrease in sample concentration from 23 to 584 days might be an indication of crystalliza-

1476 J Solution Chem (2008) 37: 1471–1487

Fig. 2 The solubility of TcO2·xH2O as a function of oxalate concentration and equilibration time:(a) pH = 1.56 to 1.65; (b) pH = 3.21 to 5.11; and (c) 5.05 to 6.49. A dashed line represents solubilitypredictions using existing thermodynamic data (see text). Solid lines represent predictions using the thermo-dynamic model developed as part of this study (see Tables 1 and 2). Smaller, open symbols represent samplescontaining less than 80% Tc(IV)

tion of a solid phase, although no evidence for this could be obtained from either diffractionor Raman analysis.

The data for TcO2·xH2O solubility in the presence of high amounts of added oxalate(Fig. 3) show a different trend of relatively slower equilibration between solids and solutions.In fact, the solubilities increase by approximately an order of magnitude in concentrationbetween 14 to 629 days equilibration in the pH range 1.3 to 1.5 (Fig. 3a) and an equalincrease is observed between 18 to 263 days in the pH range 3.4 to 4.2 (Fig. 3b). Theseresults indicate that exceptionally longer equilibration times maybe required when the netmass transfer between the solid and solution phase is large.

The results for the solubility of TcO2·xH2O at constant oxalate concentration and vari-able pH (Fig. 4) also show a similar trend of a more rapid approach to steady-state concentra-tions at pH values where the solubility is low and a relatively slow approach to equilibriumwhen the solubility is high. Hence the measured solubilities at lower pH values approachequilibrium only very slowly. There is also a consistent overall trend in the solubility datashowing a near constant solubility at a fixed oxalate concentration in the pH range 1 to 4,

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Fig. 3 The solubility of TcO2·xH2O as a function of oxalate concentration and equilibration time:(a) pH = 1.30 to 1.49; and (b) pH = 3.40 to 4.15. A dashed line represents solubility predictions usingexisting thermodynamic data (see text). Solid lines represent predictions using the thermodynamic model de-veloped as part of this study (see Tables 1 and 2). Smaller, open symbols represent samples containing lessthan 80% Tc(IV)

followed by a rapid decrease in solubility as a function of increasing pH at higher pHs. Thisoverall trend of near constant solubility at lower pH values and rapidly decreasing solubilityat higher pH values becomes a critical factor in terms of developing improved thermody-namic models as described below.

4.4 Thermodynamic Modeling

As mentioned in the Introduction, there is only one previous study of the complexing ofTc(IV) species by oxalate [7]. In that study the complexation constants for the Tc(IV)-oxalate complexes {proposed as TcO(Ox)(aq) and TcO(Ox)2−

2 } were determined by thesolvent extraction method as a function of ionic strength at high acidity (pH = 1.4 to 2.0).However, the question remains whether such results can be extrapolated to higher pH valueswhere the hydrolysis of Tc(IV) species predominates [12–14].

In order to evaluate whether the existing thermodynamic models for Tc(IV) complexa-tion are adequate to describe our current solubility data for TcO2·xH2O across a broad rangeof pH values and oxalate concentrations, we developed a thermodynamic model based uponthe most recently available data in the literature. The initial aqueous thermodynamic modelincluded: i) the solubility product for TcO2·1.6H2O proposed by Rard [12] and later ver-ified [14] to be applicable to hydrous TcO2·xH2O prepared by the same method as thesolid phase used in this study; ii) the recently determined stability constants for the ox-alate complexes TcO(Ox)(aq) and TcO(Ox)2−

2 ; [7] iii) new values for the Tc(IV) hydrolysisspecies determined by Hess et al. [14] and later critically reviewed by Rard [13]; and iv) thePitzer ion-interaction parameters for oxalate and bioxalate species determined by Qafokuand Felmy [15] and for Tc(IV) species by Hess et al. [14]. This initial model uses all ofthe stability constants and Pitzer ion interaction parameters given in Tables 1 and 2 with theexception of those involving the TcO(OH)Ox− species.

Application of the initial model to the solubility data determined in this study (see dashedlines in Figs. 2 to 4) significantly under predicts the observed solubilities with the possible

1478 J Solution Chem (2008) 37: 1471–1487

Fig. 4 The solubility of TcO2 · xH2O as a function of pH, oxalate concentration and equilibration time:(a) 0.01 mol·kg−1 oxalate; (b) 0.025 mol·kg−1 oxalate; and (c) 0.25 mol·kg−1 oxalate. The dashed line rep-resents solubility predictions using existing thermodynamic data (see text). The data represented by triangleand square symbols are from separate experimental sets. Smaller, open symbols represent samples containingless than 80% Tc(IV).

exception of the lowest pH values and lowest oxalate concentrations (Fig. 2a).2 In orderto improve the existing thermodynamic model and predict the measured solubility results,numerous combinations of aqueous species were evaluated. However, in all cases it wasnecessary to include a previously unidentified species formulated here as TcO(OH)Ox−. Thereason why this species is required becomes most obvious when examining the solubilitydata as a function of pH at different added oxalate concentrations (Fig. 4). These data setsshow a consistent trend of constant solubility as a function of pH below about pH = 4followed by decreasing solubilities above pH = 4. The solubilities also increase with addedoxalate. These results are can be explained with the formation of the TcO(OH)Ox− complex

2In this region the model gives excellent predictions of the solubility data at early time frames but doesnot predict the much lower values at very long equilibration times. These solutions also exhibit very littlepredicted oxalate complexation and these solubility data are significantly lower than previously reportedvalues in the absence of oxalate [14]. Hence this single set of solubility data at 584 days of equilibration willnot be included in our thermodynamic analysis.

J Solution Chem (2008) 37: 1471–1487 1479

Table 1 Equilibrium constants for Tc(IV) aqueous species and solid phases used in new model

Reaction log10 K Reference

TcO2+ + H2O � TcO(OH)+ + H+ 0.01 [14]

TcO2+ + 2H2O � TcO(OH)2(aq) + 2H+ −4.00 [12]

TcO(OH)2(aq) + H2O � TcO(OH)−3 + H+ −10.89 [12]

TcO(OH)+ + Ox2− � TcO(OH)Ox− 6.66 This study

TcO2+ + Ox2− � TcO(Ox)(aq) 7.13 [7]

H2Ox(aq) � HOx− + H+ −1.39 [15]

HOx− � Ox2− + H+ −4.26 [15]

TcO2·1.6H2O(am) � TcO(OH)2(aq) + 0.6H2O −8.40 [12]

Table 2 Pitzer ion-interaction parameters for technetium (from reference [14]) and oxalate species (fromreference [15]) used in this study

Binary parameters

Species i Species j β(0)ij

β(1)ij

β(2)ij

Cφij

Na+ HC2O−4 −0.206 0.468 0 0.074

Na+ C2O2−4 −0.172 1.991 0 0.131

TcO2+ Cl− 0.305 1.709 0 0.002

TcO(OH)+ Cl− −0.053 0.040 0 0.008

Cation–cation and anion–anion ion interaction parameters

Species i Species j θij

TcO2+ H+ 0.092

TcO(OH)+ H+ 0.036

Neutral species ion interaction parameters

Species i Species j θij

H2C2O04 Cl− 0.122

H2C2O04 Na+ –0.061

Triple ion interaction parameters

Species i Species j Species k ψijk

TcO2+ H+ Cl− –0.015

TcO(OH)+ H+ Cl− –0.004

according to the following reactions:

TcO2·xH2O(am) + HOx− � TcO(OH)Ox− + xH2O (pH < 4) (1)

TcO2·xH2O(am) + Ox2− � TcO(OH)Ox− + OH− + (x − 1)H2O (pH > 4). (2)

The pKa for the acid dissociation reaction between bioxalate (HOx−) and oxalate (Ox2−) is4.26, (see Table 1), which closely corresponds to the pH where the TcO2·xH2O solubilityis observed to rapidly decrease as shown in Fig. 3. Both reactions (1) and (2) also predict

1480 J Solution Chem (2008) 37: 1471–1487

a nearly unit wise increase in solubility with a unit increase in total oxalate concentration.Therefore the formation of this single species explains the majority of the solubility data atboth low and high pH. The formation of a species with an additional hydroxyl moiety overthat of the species proposed by Xia et al. [7] is also consistent with the findings of Hess etal. [14], which revealed a much stronger hydrolysis behavior for Tc(IV) than originally pro-posed by Rard [13]. In fact, the hydrolysis species TcO(OH)+ proposed by Hess et al. [14]is the dominant species in solution even below pH one. Hence it is not surprising that theTcO(OH)Ox− complex containing the hydroxyl moiety could form, even at low pH.

In order to evaluate the equilibrium constant for formation of this species we fit only thedata at low pH and high oxalate concentrations, determined at the longest equilibration times(the long time frame data shown in Fig. 3a, b). All other solubility data were predicted fromthis relatively simple model formulation. The log10K value for the TcO(OH)Ox− specieswas determined to be 6.66 ± 0.27. The predictions of this new thermodynamic model areshown as the solid lines in Figs. 2, 3 and 5a. This proposed simple model explains essentiallyall of the TcO2·xH2O solubility data determined in this study including the 0.01 mol·kg−1

total oxalate concentration data across the entire pH range of 1 to 10. The only exceptionsare the 0.025 mol·kg−1 and 0.25 mol·kg−1 total oxalate data above pH = 7 (enclosed in thedashed oval in Fig. 5a). The increased solubility at high oxalate concentrations and higherpH does not correlate with the presence of sharp diffraction peaks or notable changes in theamorphous peaks in the XRD scans and thus the increased solubility cannot be explained bya change in the solid phase. Other possible contributions to the increased solubility includethe inadvertent introduction of an additional complexing component (e.g., CO2−

3 ) or theformation of additional Tc(IV)-oxalate aqueous species not currently accounted for in ourthermodynamic model.

To explore the possible formation of additional Tc(IV)-oxalate aqueous species, severaldifferent aqueous Tc(IV)-oxalate complexes were tested with increasing numbers of hy-droxyl groups and/or oxalate ligands added to the TcO(OH)Ox− base complex in an attemptto model the higher observed solubility above pH = 7 for 0.025 and 0.25 mol·kg−1 oxalateconcentrations. The results showed that the addition of oxalate ligands to the TcO(OH)Ox−complex to form a complex like (TcO(OH)(Ox)3−

2 ) cannot explain the experimental data.For example, refitting the solubility data with a TcO(OH)(Ox)3−

2 species included in ourmodel resulted in an over prediction of the 0.01 mol·kg−1 oxalate solubility data (whichpreviously showed excellent agreement), a good fit to the 0.025 mol·kg−1 data, and a dra-matic over prediction of the 0.25 mol·kg−1 solubility data (see Fig. 5b). The formation ofspecies containing additional oxalate ligands was therefore removed from consideration.The species that seemed to improve the prediction of the data involved the addition of asingle hydroxyl moiety to the TcO(OH)Ox− to form a complex like (TcO(OH)2(Ox)2−),Fig. 5c. Although the fit is improved over the single species TcO(OH)Ox− model, the pre-dicted concentrations at very high pH (>9) and high ligand concentration (0.25 mol·kg−1)

do not match the solubility trend. The reason for this disagreement at very high pH is re-lated to the solubility reaction of TcO2·xH2O as a function of pH with this species in-cluded,

TcO2·xH2O(am) + Ox2− � TcO(OH)2Ox2− + (x − 1)H2O. (3)

Reaction (3) shows that the predicted solubility of TcO2·xH2O would become inde-pendent of pH at higher pH values, where the TcO(OH)2Ox2− complex is the dom-inant species in solution. Hence, although our high pH and high oxalate TcO2·xH2Osolubility data do indicate that additional Tc(IV)-oxalate-hydroxyl complexes proba-bly form in solution, the agreement between our models and the experimental data

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Fig. 5 The solubility of TcO2·xH2O as a function of pH and oxalate concentration. Lines represent predic-tions using the thermodynamic model developed as part of this study (see Tables 1 and 2): (a) inclusion ofthe TcO(OH)Ox− complex, note that inclusion of the TcO(OH)Ox− complex does not model the enhancedsolubility of data points within the dashed oval; (b) inclusion of TcO(OH)Ox− and TcO(OH)(OX)3−

2 com-

plexes; and (c) inclusion of TcO(OH)Ox− and TcO(OH)2Ox2− complexes. The data points representedthe longest-term equilibration data available for each oxalate concentration. Smaller, open symbols representsamples containing less than 80% Tc(IV)

is not precise enough for us to confidently estimate thermodynamic values for suchspecies.

Finally, although the experimental studies were conducted in a controlled atmosphereglove box, it is always possible that over such a long-term equilibration the introduction ofa small level of CO2 could have occurred. However, the analyzed total inorganic carbon inthese long-term samples was always extremely low (<10−4 mol·kg−1), and based upon thestability constants for Tc(IV)-carbonate complexes [12], such low carbonate levels couldnot have resulted in the higher observed TcO2·xH2O solubilities.

5 Conclusions

In summary, the inclusion of the single previously unidentified species TcO(OH)Ox− inour aqueous thermodynamic model explains a wider range of observed solubility data for

1482 J Solution Chem (2008) 37: 1471–1487

TcO2·xH2O in the presence of oxalate and over a broad range of pH values. Inclusion of thisspecies is also supported by the recently proposed thermodynamic data for the TcO(OH)+hydrolysis species which shows an enhanced stability for this species at pH values as lowas one. Hence inclusion of this species in any thermodynamic model for determining thecomplexation of Tc(IV) with oxalate is required. The only significant limitation of this sin-gle species model occurs at high base concentration (pH > 7) and in high oxalate con-centration (>0.025 mol·kg−1). In this region our current solubility data indicate that otherpossible Tc(IV)-oxalate complexes form, most likely involving the addition of hydroxylcomplexes to the TcO(OH)Ox− moiety rather than oxalate groups. The currently proposedthermodynamic model works exceptionally well in the acidic concentration range extendingto high oxalate concentration and over the entire pH range at lower oxalate concentration(<0.025 mol·kg−1).

Acknowledgements This work was supported by the Environmental Management Science Program(EMSP), U.S. Department of Energy. Pacific Northwest National Laboratory is operated by Battelle MemorialInstitute for the U.S. Department of Energy under Contract DE-AC06-76RLO 1830.

Appendix

Solubility measurements for data used in the figures. The % Tc(IV) is calculated as the Tcconcentration remaining in the aqueous fraction after solvent extraction, divided by the totalTc concentration measured in the aqueous solution prior to extraction. Values in italics havemeasured Tc(IV) values either 80% < T c(IV ) > 120%.

Table 3 Set 1: Tc072903. TcO2·xH2O(am) solubility at different equilibration periods as a function of thepH in aqueous solution containing 0.01 mol·kg−1 oxalate and 0.02 mol·kg−1 hydrazine

Sample 15 days 34 days 1134 days

number pH log10 Tc(IV) %Tc(IV) pH log10 Tc(IV) %Tc(IV) pH log10 Tc(IV) % Tc(IV)

1 4.80 −6.00 93 4.84 −5.80 94 5.01 −4.99 93

2 6.47 −6.55 93 6.50 −6.41 99 6.68 −6.39 96

3 7.26 −6.95 87 7.20 −6.84 95 7.08 −7.22 103

4 7.98 −7.18 85 7.89 −7.15 82 7.63 −7.55 107

5 8.99 −7.26 83 8.93 −7.28 80 8.64 −7.69 113

6 9.88 −6.87 81 9.86 −6.91 81 9.57 −7.12 100

7 10.82 −6.35 73 10.77 −6.68 73 10.15 −6.95 96

8 11.71 −6.59 54 11.72 −7.06 51 11.12 −7.63 74

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Table 4 Set 2: Tc102704. TcO2·xH2O(am) solubility at different equilibration periods as a function of pHin aqueous solution containing 0.01 mol·kg−1 oxalate and 0.02 mol·kg−1 hydrazine

Sample 4 days 14 days 711 days

number pH log10 Tc(IV) %Tc(IV) pH log10 Tc(IV) %Tc(IV) pH log10 Tc(IV) %Tc(IV)

1 0.99 −4.94 95 1.07 −4.80 98 0.61 −3.80 97

2 1.45 −5.24 97 1.53 −5.03 98 1.14 −4.05 99

3 1.90 −5.17 94 2.01 −4.96 99 1.79 −4.06 97

4 2.46 −5.11 96 2.84 −4.91 98 3.13 −4.13 96

5 3.13 −5.03 94 3.78 −4.85 97 3.82 −4.18 94

6 4.36 −5.28 94 5.72 −5.22 93 6.38 −5.35 94

7 7.78 −5.90 97 7.77 −6.02 95 7.42 −7.01 98

8 7.76 −6.24 97 7.76 −6.33 93 7.53 −7.11 98

Table 5 Set 3: Tc102804. TcO2·xH2O(am) solubility at different equilibration periods in aqueous solutionas a function of oxalate concentration from 0.003 to 0.9 mol·kg−1 at ca. pH = 1.5 and containing 0.02mol·kg−1 hydrazine

SN [ox] 4 days 14 days 629–640 days

pH log10 Tc(IV) %Tc(IV) pH log10 Tc(IV) % Tc(IV) pH log10 Tc(IV) % Tc(IV)

1 0.94 1.49 −4.51 95 1.57 −4.22 97 1.30 −2.79 97

2 0.51 1.49 −4.79 58 1.59 −4.25 100 1.34 −2.82 99

3 0.25 1.47 −4.63 96 1.58 −4.31 99 1.34 −2.84 96

4 0.1 1.48 −4.80 96 1.60 −4.49 99 1.39 −3.05 95

5 0.025 1.47 −5.01 96 1.61 −4.69 98 1.44 −3.46 96

6 0.0125 1.46 −5.14 97 1.61 −4.83 98 1.49 −3.70 96

7 6.3 × 10−3 1.46 −5.18 97 1.63 −4.90 99 1.33 −3.90 97

8 3.2 × 10−3 1.48 −5.33 98 1.64 −5.08 98 1.36 −4.13 93

Table 6 Set 4: Tc121304. TcO2·xH2O(am) solubility at different equilibration periods in aqueous solutionas a function of oxalate concentration from 5 × 10−7 to 1 × 10−4 mol·kg−1 at ca. pH = 1.5 and containing0.02 mol·kg−1 hydrazine

SN [ox] 4 days 25 days 584 days

pH log10 Tc(IV) %Tc(IV) pH log10 Tc(IV) %Tc(IV) pH log10 Tc(IV) %Tc(IV)

1 1 × 10−2 0.07 −4.28 92 0.07 −3.78 94 0.07 −3.21 84

2 1 × 10−2 0.19 −4.56 92 0.19 −4.07 93 0.19 −3.22 87

3 1 × 10−4 1.64 −5.97 98 1.63 −5.99 97 1.53 −7.14 91

4 5 × 10−5 1.63 −5.93 98 1.62 −6.00 97 1.54 −7.78 79

5 1 × 10−5 1.63 −5.97 100 1.63 −6.08 96 1.56 −7.77 82

6 5 × 10−6 1.62 −6.05 97 1.63 −6.13 58 1.59 −7.82 76

7 1 × 10−6 1.62 −6.01 96 1.63 −6.09 83 1.61 −7.90 72

8 5 × 10−7 1.62 −6.00 85 1.63 −6.08 91 1.65 −7.79 84

1484 J Solution Chem (2008) 37: 1471–1487

Table 7 Set 5: Tc100405. TcO2·xH2O(am) solubility at different equilibration periods in aqueous solutionas a function of oxalate concentration from 1 × 10−6 to 0.05 mol·kg−1 at ca. pH = 3.0 and containing 0.02mol·kg−1 hydrazine

SN [ox] 9 days 17 days 44 days 288 days

log10 % log10 % log10 % log10 Tc(IV)%

pH Tc(IV) Tc(IV) pH Tc(IV) Tc(IV) pH Tc(IV) Tc(IV) pH Tc(IV)

1 5 × 10−2 3.16 −3.98 98 3.22 −3.94 93 3.29 −3.79 96 3.21 −3.53 97

2 1 × 10−3 3.93 −5.39 93 3.50 −5.46 94 3.80 −5.39 97 4.56 −5.28 92

3 5 × 10−4 3.52 −6.02 93 3.41 −6.00 95 3.63 −5.97 96 4.45 −6.20 91

4 1 × 10−4 3.40 −7.44 77 3.50 −7.30 118 3.83 −7.61 98 4.46 −8.77 1

5 5 × 10−5 3.40 −7.03 145 3.46 −7.18 102 4.98 −8.41 257

6 1 × 10−5 3.40 −7.33 85 3.46 −7.24 98 4.90 −9.97 29

7 5 × 10−6 3.33 −7.41 3.37 −6.75 102 4.51 −8.55 306

8 1 × 10−6 3.35 −7.49 88 3.38 −7.38 100 5.11 −9.19 86

Table 8 Set 6: Tc100505. TcO2·xH2O(am) solubility at different equilibration periods in aqueous solutionas a function of oxalate concentration from 1 × 10−6 to 0.05 mol·kg−1 at ca. pH = 5.0 and containing0.02 mol·kg−1 hydrazine

SN [ox] 7 days 17 days 44 days 288 days

log10 % log10 % log10 % log10 %

pH Tc(IV) Tc(IV) pH Tc(IV) Tc(IV) pH Tc(IV) Tc(IV) pH Tc(IV) Tc(IV)

1 5 × 10−2 4.74 −4.43 94 4.90 −4.36 95 5.04 −4.31 95 5.05 −4.24 90

2 1 × 10−3 5.10 −6.48 95 5.43 −6.48 84 5.85 −6.41 112 6.30 −6.92 84

3 5 × 10−4 5.20 −6.91 105 5.53 −6.99 92 5.96 −7.10 93 6.34 −7.81 87

4 1 × 10−4 5.34 −7.91 123 5.66 −8.15 3 6.08 −8.51 94 6.41 −9.27 7

5 5 × 10−5 5.27 −7.78 175 5.52 −8.20 130 6.46 −9.12 117

6 1 × 10−5 5.20 −7.57 155 5.56 −8.03 114 6.41 −8.63 2200

7 5 × 10−6 5.24 −7.82 184 5.63 −8.23 106 6.45 −8.77 213

8 1 × 10−6 5.19 −7.67 236 5.66 −7.68 390 6.49 −8.74 3400

J Solution Chem (2008) 37: 1471–1487 1485

Table 9 Set 7: Tc110805. TcO2·xH2O(am) solubility at different equilibration periods in aqueous solution asa function of oxalate concentration from 0.003 to 0.9 mol·kg−1 at ca. pH = 3.5 and containing 0.02 mol·kg−1

hydrazine

SN [ox] 8 days 18 days 60 days


1 0.94 3.31 −4.50 73 3.41 −4.37 79 3.51 −4.08 89

2 0.51 3.33 −4.45 77 3.42 −4.33 82 3.54 −3.86 90

3 0.25 3.35 −4.28 79 3.42 −4.05 86 3.51 −3.60 93

4 0.1 3.40 −4.66 95 3.47 −4.39 94 3.56 −4.00 94

5 0.025 3.59 −4.93 95 3.66 −4.65 95 3.74 −4.36 96

6 0.0125 3.78 −5.11 95 3.86 −4.82 96 3.95 −4.44 94

7 6.3 × 10−3 3.85 −5.32 95 3.97 −5.06 95 4.00 −4.56 93

8 3.2 × 10−3 3.76 −5.52 93 3.89 −5.25 95 3.99 −4.78 95

SN [ox] 70 days 263 days

pH log10 Tc(IV) %Tc(IV) pH log10 Tc(IV) %Tc(IV)

1 0.94 3.49 −4.01 89 3.49 −4.01 89

2 0.51 3.47 −3.82 92 3.47 −3.82 92

3 0.25 3.46 −3.55 92 3.46 −3.55 92

4 0.1 3.53 −3.95 94 3.53 −3.95 94

5 0.025 3.75 −4.33 93 3.75 −4.33 93

6 0.0125 3.92 −4.39 94 3.92 −4.39 94

7 6.3 × 10−3 4.04 −4.51 93 4.04 −4.51 93

8 3.2 × 10−3 4.03 −4.71 92 4.03 −4.71 92

Table 10 Set 9: Tc012406. TcO2·xH2O(am) solubility at different equilibration periods as a function of pHin aqueous solution containing 0.25 mol·kg−1 oxalate and 0.02 mol·kg−1 hydrazine

SN 7 days 17 days 34 days 259 days

log10 % log10 % log10 % log10 %


1 0.97 −4.15 95 0.98 −3.94 96 0.98 −3.77 96 0.70 −3.00 98

2 1.52 −4.14 96 1.53 −3.95 96 1.54 −3.79 97 1.27 −3.02 102

3 1.97 −4.09 95 2.03 −3.89 89 2.03 −3.76 96 1.78 −2.97 98

4 2.50 −4.08 95 2.55 −3.90 95 2.56 −3.75 96 2.33 −2.98 101

5 3.02 −4.14 93 3.04 −3.94 94 3.05 −3.80 95 2.78 −3.14 98

6 3.98 −4.26 95 3.99 −4.08 94 4.00 −3.96 96 3.82 −3.44 99

7 5.03 −4.44 93 5.12 −4.32 93 5.26 −4.22 96 5.31 −3.99 99

8 7.92 −5.54 96 7.28 −5.20 93 7.64 −5.22 95 7.65 −5.13 95

1486 J Solution Chem (2008) 37: 1471–1487

Table 11 Set 10: Tc020606. TcO2·xH2O(am) solubility at different equilibration periods as a function ofpH in aqueous solution containing 0.25 mol·kg−1 oxalate and 0.02 mol·kg−1 hydrazine


log10 % log10 % log10 % log10 %


1 3.47 −4.44 94 3.49 −4.22 92 3.46 −3.97 94 3.29 −3.11 96

2 4.56 −4.61 95 4.53 −4.43 94 4.56 −4.24 94 4.40 −3.28 93

3 5.60 −4.91 95 5.65 −4.81 96 5.65 −4.69 96 5.74 −4.13 100

4 8.07 −5.46 96 8.12 −5.41 97 8.17 −5.38 97 7.93 −5.33 95

5 9.01 −5.75 95 9.06 −5.79 97 9.07 −5.84 103 8.76 −6.07 94

6 9.99 −6.22 96 10.07 −6.40 98 10.06 −6.67 97 9.69 −7.50 87

7 11.09 −6.83 92 11.12 −7.24 92 11.09 −7.88 62

8 12.01 −7.33 38 12.05 −7.57 39 12.05 −8.23 17 11.63 −7.74 51


SN 7 days 28 days 245 days


1 3.57 −4.40 94 3.72 −4.16 95 3.73 −4.00 97

2 4.85 −4.73 94 7.47 −5.09 96 7.22 −5.20 97

3 8.23 −5.79 97 8.43 −5.94 97 8.11 −6.35 93

4 8.94 −6.38 96 8.97 −6.58 95 8.71 −7.34 97

5 9.50 −6.58 52 9.46 −7.04 90 9.22 −7.49 89

6 10.22 −6.89 126 10.13 −7.55 79 9.79 −7.38 81

7 11.04 −7.36 92 10.87 −7.46 72 10.49 −6.95 74

8 11.95 −6.82 87 11.86 −7.08 61 11.45 −6.78 52



log10 % log10 % log10 % log10 %


1 0.89 −4.65 94 1.00 −4.51 94 0.88 −4.38 96 0.98 −3.85 99

2 1.41 −4.63 95 1.50 -4.52 89 1.38 −4.37 95 1.44 −3.79 98

3 1.94 −4.62 95 2.03 −4.51 90 1.97 −4.36 96 2.21 −3.87 98

4 2.63 −4.48 96 2.77 −4.37 94 2.81 −4.25 95 3.16 −3.91 98

5 3.11 −4.46 95 3.23 −4.35 93 3.26 −4.24 93 3.46 −3.93 95

6 4.09 −4.48 94 4.24 −4.41 90 4.27 −4.30 94 4.46 −4.09 97

7 7.31 −5.59 95 8.39 −5.81 96 8.34 −5.87 96 8.16 −6.15 96

8 8.56 −6.23 99 8.72 −6.30 92 8.61 −6.31 96 8.45 −6.63 134

J Solution Chem (2008) 37: 1471–1487 1487

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thermodynamic model for the solubility of tco2⋅xh2o in aqueous oxalate systems

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