chemical thermodynamics of neptunium and plutonium - U.S. ...

12.1 Neptunium carbon compounds and complexes 265used to deduce any thermodynamic information or, indeed, any information at all onthe solution species. Eriksen et al. [93ERI/NDA] also proposed two mixed complexesbased on a single series of experiments. This review accepts the conclusions of RaiandRyan[85RAI/RYA] and of Vitorge [95VIT] that solubility studies published upto and including 1990 [71MOS2, 85RAI/RYA, 89MOR/PRA, 90PRA/MOR] shownofirm evidence of any soluble carbonate or mixed hydroxide Np(IV) carbonate complex.When such species were proposed in literature, they are the result of incorrectexperimental methodology and treatment of data. An efficient reducing agent must beused for Np(IV) solubility measurements in neutral solutions or in solutions containinglow concentrations of carbonate/bicarbonate [85RAI/RYA], to avoid oxidation ofcomplexes and changes in the solid phase during solubility measurements. Accordingto the following equilibriaNp(OH) 4 (am) + xCO 2−3Å Np(CO 3 ) (4−2x)x + 4OH − (12.28)the plot of experimental log 10 (solubility) as a function of (log 10 [CO 2−3 ] −(4/x) log 10 [OH − ]) should be linear with the slope x, when Np(CO 3 ) (4−2x)x solublecomplex predominate. In the present review, x = 5 and 4 are used and the corresponding4/x values are 0.8 and 1. It is noted, however, that a line can be drawn through datareported in [89MOR/PRA, 90PRA/MOR] with a slope that might indicate formationof Np(CO 3 ) 4−4(or Np(CO 3 ) 2 (OH) 2−2[99RAI/HES]). The scatter in the experimentaldata could be due to the formation of colloidal NpO 2 (am, hyd). However, such specieswould be important only under a very limited range of conditions. Possible formationof other species cannot be excluded [93ERI/NDA] (alsocf. the Appendix A entry for[99RAI/HES]).12.1.2.1.4.dTemperature influenceDelmau, Vitorge and Capdevila [96DEL/VIT] measured the influence of temperatureon the Np(V)/Np(IV) couple. By varying temperature up and down over severalweeks, they obtained reasonable precision (± 10 or 20 mV), when compared to previouswork [79FED/PER, 84VAR/BEG]. As discussed above (Section 12.1.2.1.4.a), themain uncertainty was lack of correct pH measurements and this could induce shifts ofas much as 330 mV. Since reproducible data were obtained during temperature cycleson a single solution, it is clear that eventually each solution was buffered. Hence themain source of uncertainty cancels when calculating the enthalpy of reaction. Theauthors calculated, by linear regression, ∂ E ◦′ /∂T (12.21) =−2.0584, −2.1676 and−2.2409 mV·K −1 in 0.6, 1 and 1.5 M Na 2 CO 3 respectively. This data treatment neglectedheat capacity, hence assuming entropy and enthalpy changes were independentof temperature. This review calculates r H m (12.21) =−82.40, −85.86, −86.91kJ·mol −1 respectively. Extrapolation to zero ionic strength using the SIT and linear regression[94GIF/VIT]gives r H ◦ m (12.21, 298.15 K) =−(84.39±4.38) kJ·mol−1 andT ◦ (dε(12.21))/dT) = (0.50 ± 0.34) kg·mol −1 ,whereT ◦ is the standard temperature(298.15 K).