Thorium ionic strength

Thorium. Experimental and theoretical studies of thorium speciation, solubility, and sorption in low-ionic-strength waters are described by Langmuir and Herman (1980), Laflamme and Murray (1987), Osthols et ai (1994), Osthols (1995), and Quigley et al. (1996). Langmuir and Herman (1980) provide a critically evaluated thermodynamic database for natural waters at low temperature that is widely used. However, it does not contain information about important thorium carbonate complexes, and the stability of phosphate complexes may be overestimated (US EPA, 1999b). 

The release of uranium and thorium from minerals into natural waters will depend upon the formation of stable soluble complexes. In aqueous media only Th is known but uranium may exist in one of several oxidation states. The standard potential for the oxidation of U in water according to equation (2) has been re-evaluated as E° - 0.273 0.005 V and a potential diagram for uranium in water at pH 8 is given in Scheme 3. This indicates that will reduce water, while U is unstable with respect to disproportionation to U and U Since the Earth s atmosphere prior to about 2 x 10 y ago was anoxic, and mildly reducing, U " would remain the preferred oxidation state in natural waters at this time. A consequence of this was that uranium and thorium would have exhibited similar chemistry in natural waters, and have been subject to broadly similar redistribution processes early in the Earth s history. Both U " and Th are readily hydrolyzed in aqueous solutions of low acidity. A semiquantitative summary of the equilibrium constants for the hydrolysis of actinide ions in dilute solutions of zero ionic strength has been... 

Simple thermodynamic calculations based on literature data (5-12) support the choice of phosphates as the optimum mineral phases for actinide immobilization. The calculations considered every relevant species reported (5-72) that contained protons, hydroxide, or the ligand in question for each metal ion. Where necessary, equilibrium constants were corrected to 0.1 M ionic strength using the Davies equation. As an example, the calculated solubility of europium, thorium, and uranium in various media at p[H] 7.0 (p[H] = - log of the hydrogen ion concentration), 0.001 M total ligand concentration, 0.1 M ionic strength, and 25 °C are shown in Table I. Within the constraints of the calculation, the solubility of thorium is limited by Th(OH)4, but the lowest europium and uranyl solubilities are observed for phosphates. 

In spite of its theoretical limitations, this hnear extrapolation of the thorium data was adopted by [1975RAN] and [1976FUG/OET], As pointed out by Wagman [1976WAG], a Debye-Htickel correction of the solution data, followed by an extrapolation against the ionic strength, yields an extrapolated value of (ThCLt,... 

Experimental studies on the hydrolysis of the Th" ion, its complexes with strong inorganic hgands, and the solubility of thorium oxides or hydroxides ate usually performed with low concentrations of thorium in perchlorate, chloride, and nitrate media. There is no evidence for complex formation between Th" and CIO4 however, chloride and nitrate form weak Th(lV) complexes as discussed in Sections Vlll.2.2.1 and X. 1.3.3, respectively. For the evaluation of equihbrium constants at zero ionic strength from data in chloride and nitrate media we have therefore the general problem to decide if the activity of Th", , should be calculated using a... 

Amorphous thorium oxyhydroxides with varying chemical composition, water content and particle size have also different thermodynamic properties. This may reflect the differences between reported solubility data. Typical examples of discrepancies between solubility data measured at the same ionic strength are shown in Figure VII-14. 

Figure VIII-12 Experimental and predicted (NONLINT-SIT) values of HD vs. total [ BrOj ] concentration or thorium concentrations vv total [ BrO ] concentration at the ionic strength 0.51 m from [1950DAY/STO], The predicted lines based on the ion-interaction parameters reported in Table VIII-27, and A(.G°/RT (Th" ) = -284.305 (this review), AfG°/RT(Br03) = 7.693 [2003GUI/FAN], and...

The available thermodynamic data are of two types stabihty constants, enthalpy and entropy of reaction for the formation of soluble complexes Th(S04) " " and solubihty data for various solid phases. The two sources are linked because the solubility of the solid phases depends on the chemical speciation, i.e., the sulphate complexes present in the aqueous phase. The analysis of the experimental stability constants has been made using the SIT model however, this method cannot be used to describe the often very high solubility of the solid sulphate phases. In order to describe these data the present review has selected a set of equilibrium constants for the formation of Th(S04) and Th(S04)2(aq) at zero ionic strength based on the SIT model and then used these as constants in a Gibbs energy minimisation code (NONLINT-SIT) for modelling experimental data to determine equilibrium constants for the formation of Th(S04)3 and the solubility products of different thorium sulphate solids phases. 

The calculations are accounting for the continuous variation of logj,, log,(, K2 (protonation constant of COj ) and the formation constants of the thorium complexes when ionic strength varies from I = at... 

Figure XI-8 Solubility and speciation of thorium at a total carbonate concentration of Ctot = 0.2 M. Experimental data for Na6[Th(C03)5]-12H20(s) in 0.2 M NaHCOs /1.8 M NaNOs and in 0.2 M NazCOs /1.6 M NaNOs [1973DER/FAU3] and calculated solubility for Th02(am, hyd) in 0.2 MNa2C03 /1.6 MNaNOs containing 0.01-1.0 M NaOH. The calculations are based on the equilibrium constants and SIT coefficients selected in the present review taking into accounting the variation of the solution composition and ionic strength.

Table XI-6 Equilibrium constants at zero ionic strength and molar standard Gibbs energies of formation selected for solid and aqueous thorium carbonate compounds at 25°C.

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