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

768 A. Discussion of selected referencesof magnitude as the NaClO 4 concentration. Therefore, the ionic medium wasnot constant over the series of solubility measurements, and the concentration ofNa + varied from 0.15 to 0.43 M (at least as recalculated in the present review-seebelow). The pH measurements reported by the authors (9.61 to 10.5) are, for somesamples, considerably greater than the theoretical values calculated assuming theabove procedure. The authors also may have added a strong base (possibly NaOH)to vary the pH. It is not possible to accurately calculate the amount of NaOH thatmight have been added for pH adjustment, as equilibration with air during the pHadjustment would have also added an acidic reactant (carbon dioxide gas). Thus, theconcentrations of Na + and the ionic strength corrections to the equilibrium constantsfor the carbonate protonation reactions cannot be estimated with good accuracy (andthe ones used by the authors are then probably in error).The authors reported the absorption spectrum of the solution most concentrated inNp(VI) and total carbonate. Although of poor quality compared with the previouslypublished spectra [81WES/SUL, 86GRE/RIG] cited by the authors, the spectrum (contraryto the contention of the authors) still clearly corresponds to the spectrum of thepure limiting complex, NpO 2 (CO 3 ) 4−3. The analysis of the absorbance measurementsstrongly relies on the values of the molar absorbances for the pure species. Such valueswere not proposed by the authors, nor can the values be deduced because there was noevidence for a second absorbing species.The authors plotted their experimental solubility data as a function of the logarithmof the aqueous concentration of CO 2−3, and proposed a slope analysis. In the presentreview it is concluded that the range of experimental conditions was too small, thechemical conditions too poorly defined, and the scatter of the experimental results toolarge, to obtain any useful result from such an analysis. The authors’ proposed initialcomplex, final complex, and the number of OH ligands they claim were exchanged arecompletely in error.[93PRA/YAM]See [91YAM/PRA].[93SAW/MAH]The authors carried out potentiometric titrations using a fluoride electrode at (23±1) ◦ C(“room temperature”) and I = 1M(NaClO 4 ). This method is well established and hasa high precision. This study supersedes the one published earlier by the same group[86MAH/SAW], in which log 10 β 1 (A.69, q = 1, I = 1M) = (3.07 ± 0.04) was reported.This earlier result was called into question by these authors [93SAW/MAH], becausehigher complexes had not been considered in the evaluation in [86MAH/SAW].Pure starting solutions of Pu(III) were obtained by adding an excess of quinhydroneto a Pu stock solution before each experiment. The remaining quinhydrone servedas a holding reductant for Pu(III). The absorption spectra were recorded before eachexperiment to ensure absence of Pu(IV). The liquid junction potential was carefullyevaluated at each titration point as a function of the free proton concentration. The data